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14
Chemical Kinetics
CHAPTER OBJECTIVES
• To understand the factors that affect the rate of chemical
reactions
• To be able to determine a reaction rate
• To understand the meaning of a rate law
• To be able to determine the concentration dependence
of the rate of a chemical reaction
• To be able to predict the individual steps of a simple
reaction
• To begin to understand why chemical reactions occur
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Chemistry: Principles, Patterns,
and Applications, 1e
14.1 Factors That Affect
Reaction Rates
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14.1 Factors That Affect Reaction Rates
• Chemical kinetics
– Study of reaction rates, or the changes in the concentrations of
reactants and products with time
– By studying kinetics, insights are gained into how to control
reaction conditions to achieve a desired outcome
• Chemical kinetics of a reaction depend on
various factors
1. Reactant concentrations
2. Temperature
3. Physical states and surface areas of reactants
4. Solvent and catalyst properties
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Concentration Effects
• Two substances cannot react with each other unless
their constituent particles come into contact; if there is no
contact, the rate of reaction will be zero.
• The more reactant particles that collide per unit time, the
more often a reaction between them can occur.
• The rate of reaction usually increases as the
concentration of the reactants increases.
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Temperature Effects
• Increasing the temperature of a system increases the
average kinetic energy of its constituent particles.
• As the average kinetic energy increases, the particles
move faster, so they collide more frequently per unit time
and possess greater energy when they collide, causing
increases in the rate of the reaction.
• Rate of all reactions increases with increasing
temperature and decreases with decreasing temperature.
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Phase and Surface Area Effects
• If reactants are uniformly dispersed in a single homogeneous
solution, the number of collisions per unit time depends on
concentration and temperature.
• If the reaction is heterogeneous, the
reactants are in two different phases,
and collisions between the reactants can
occur only at interfaces between phases;
therefore, the number of collisions between
the reactants per unit time is reduced, as
is the reaction rate. The rate of a
heterogeneous reaction depends on the
surface area of the more condensed phase.
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Solvent Effects
• The nature of the solvent can affect the reaction rates of
solute particles.
• Solvent viscosity is also important in determining
reaction rates.
1. In highly viscous solvents, dissolved particles diffuse much more
slowly than in less viscous solvents and collide less frequently per
unit time.
2. Rates of most reactions decrease rapidly with increasing solvent
viscosity.
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Catalyst Effects
• Catalyst is a substance that participates in a chemical
reaction and increases the rate of the reaction without
undergoing a net chemical change itself.
• Catalysts are highly selective and often determine the
product of a reaction by accelerating only one of several
possible reactions that could occur.
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Chemistry: Principles, Patterns,
and Applications, 1e
14.2 Reaction Rates and
Rate Laws
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Reaction Rates
• Reaction rates
– Expressed as the concentration of reactant consumed or the
concentration of product formed per unit time
– Units are moles per liter per unit time (M/s, M/min or M/h)
– To measure reaction rates
1. initiate the reaction;
2. measure the concentration of the reactant or product at
different times as the reaction progresses;
3. plot the concentration as a function of time on a graph;
4. calculate the change in the concentration per unit time.
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Reaction Rates
• Reaction Rates
– The change in the concentration of either the reactant or the product
over a period of time.
– For a simple reaction (A  B),
rate = [B] = – [A]
t
t
– Square brackets indicate concentration; and  means “change in.”
– Concentration of A decreases with time; and the concentration of B
increases with time.
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Reaction Rates
• Determining the rate of hydrolysis of aspirin
– One can calculate the average reaction rate for a given time
interval from the concentrations of either the reactant or one of
the products at the beginning of the interval (time = t0) and at the
end of the interval (tf).
– Using salicylic acid, one can find the rate of the reaction for the
interval between t = 0 and t = 2h; one can also calculate the rate
of the reaction from the concentrations of aspirin at the beginning
and the end of the same interval.
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Reaction Rates
• Calculating the Rate of Fermentation of Sucrose
– A reaction in which coefficients are not all the same
–The coefficients show that the reaction produces four
molecules of ethanol and four molecules of carbon dioxide for
every one molecule of sucrose that is consumed
–The coefficients in the balanced equation show that the rate at
which ethanol is formed is four times faster than the rate at which
sucrose is consumed: [ethanol] = – 4[sucrose]
t
t
This can also be expressed in terms of the reactant or product
with the smallest coefficient in the balanced equation
rate = – [sucrose] / t = ¼([ethanol] /t).
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Reaction Rates
• Instantaneous rate of reaction
– The rate at any given point in time
– As the period of time used to calculate an average
rate of a reaction becomes shorter and shorter, the
average rate approaches the instantaneous rate
– In chemical kinetics, focus is on one particular
instantaneous rate, t = 0, which is the initial rate of the
reaction; initial rates are determined by measuring the
rate of the reaction at various times and then
extrapolating a plot of rate versus time to t = 0
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Rate Laws
• Describes the relationships between reactant rates and
reactant concentrations
• May be written from either of two different, but related,
perspectives:
1. Differential rate law
– Expresses the rate of a reaction in terms of changes in the
concentration of one or more reactants, [R], over a specific time
interval, t
– Describes what is occurring on a molecular level during a
reaction
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Rate Laws
2. Integrated rate law
– Describes the rate of a reaction in terms of the initial
concentration, [R]0, and the measured concentration of one or more
reactants, [R], after a given amount of time, t
– Used for determining the reaction order and the value of the rate
constant from experimental measurements
• Rate law must give the proper units for the rate, M/s
• Rate law must be determined experimentally
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Rate Laws
• Reaction orders
– For a reaction with the general equation
aA + bB  cC + dD,
the experimentally determined rate law has the form
rate = k[A]m [B]n.
– The proportionality constant, k, is called the rate constant.
1. Value is characteristic of the reaction and reaction conditions
2. A given reaction has a particular value of the rate constant
under a given set of conditions, such as temperature,
pressure, and solvent
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Rate Laws
– Rate of a reaction depends on the rate constant for the given set of
reaction conditions and on the concentration of each reactant, raised
to the powers m and n
– Values of m and n are derived from experimental measurements of
the changes in reactant concentrations over time and indicate the
reaction order, the degree to which the rate of the reaction
depends on the concentration of each reactant
– m and n are not related to the stoichiometric coefficients a and b in
the balanced chemical equation but must be determined
experimentally
– Overall reaction order is the sum of all the exponents in the rate
law, or m + n
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Chemistry: Principles, Patterns,
and Applications, 1e
14.3 Methods of
Determining Reaction
Orders
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Zeroth-Order Reactions
• Zeroth-order reaction
– Reaction whose rate is independent of concentration
– Its differential rate law is rate = k
– One can write their rate in a form such that the exponent of the
reactant in the rate law is 0
rate = – [A] = k[reactant]0 = k(1) = k
t
– Since rate is independent of reactant concentration, a graph of
the concentration of any reactant as a function of time is a
straight line with a slope of –k (concentration decreases with
time); a graph of the concentration of any product as a function
of time is a straight line with a slope of +k
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Zeroth-Order Reactions
– Integrated rate law for a zeroth-order reaction produces a
straight line and has the general formula
[A] = [A]0 – kt,
where [A]0 is the initial concentration of reactant A; the rate
constant must have the same units as the rate of the reaction,
M/s, in a zeroth-order reaction
– Equation has the form of the equation for a straight line (y = mx
+ b); y = [A], mx = – kt, and b = [A]0
– Occur most often when the reaction rate is determined by
available surface area
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First-Order Reactions
• First-order reaction
– Reaction rate is directly proportional to the concentration of
one of the reactants
– Have the general form A  products
– Differential rate for a first-order reaction is
rate = – [A] = k[A]
t
– If the concentration of A is doubled, the rate of the reaction
doubles; if the concentration of A is increased by a factor of 10,
the rate increases by a factor of 10
– Units of a first-order rate constant are inverse seconds, s–1
– First-order reactions are very common
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First-Order Reactions
– Integrated rate law for a first-order reaction can be written in
two different ways, one using exponentials and one using
logarithms
1. Exponential form, [A] = [A]0e–kt, where [A]0 is the initial
concentration of reactant A at t = 0; k is the rate constant, and e is
the base of the natural logarithms, which has the value 2.718.
Concentration of A will decrease in a smooth exponential curve over
time
2. The logarithmic expression of the relationship between the
concentration of A and t is obtained by taking the natural logarithm
of each side of the preceding equation and rearranging: ln[A] =
ln[A]0 – kt; the equation has the form of the equation for a straight
line; y = ln[A] and b = ln[A]0; and a plot of ln[A] vs. t for a first-order
reaction gives a straight line with a slope of –k and an intercept of
ln[A]0
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Second-Order Reactions
• Two kinds of second-order reactions
1. The simplest kind of second-order reaction is one whose rate
is proportional to the square of the concentration of the reactant
and has the form 2A  products.
– Differential rate law is rate = – [A] = k[A]2
2t
– Doubling the concentration of A quadruples the rate of the
reaction
– Units of rate constant is M–1s–1 or L/mols
– Concentration of the reactant at a given time is described by the
following integrated rate law: 1/ [A] =
1/ [A]0 + kt, which has the form of an equation of a straight line; y =
1/ [A], b = 1/ [A]0; and a plot of 1/ [A] vs t is a straight line with a
slope of k and an intercept of 1/[A]0
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Second-Order Reactions
2. The second kind has a rate that is proportional to the product
of the concentrations of two reactants and has the form A + B 
products.
– Reaction is first order in A and first order in B
– Differential rate law for the reaction is
rate = – [A] = – [B] = k[A] [B]
t
t
– Reaction is first order both in A and in B and has an overall
reaction order of 2
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Determining the Rate Law of a
Reaction
• Understanding reaction mechanisms (sets of steps in a reaction) simplifies
chemical reactions.
• The first step in discovering the mechanism of a reaction is to determine the
reaction’s rate law, which can be done by designing experiments that
measure the concentration(s) of one or more of the reactants or products as
a function of time.
• For the reaction A + B  products, one needs to determine the value of k
and the exponents m and n in the equation
rate = k[A]m [B]n.
• Rate data is given in the following table:
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Determining the Rate Law of a
Reaction
• One can determine values of k and exponents m and n
in several ways:
1. Keep the initial concentration of B constant while varying the
initial concentration of A and calculating the initial rate of the
reaction and then deducing the order of the reaction with respect
to A
2. Determine the order of reaction with respect to B by studying
the reaction rates when the initial concentration of A is kept
constant while the concentration of B is varied
3. Determine order of reaction with respect to a given reactant by
comparing the different rates obtained when only the
concentration of the reactant in question was changed
4. Determine reaction orders by taking the quotient of the rate
laws for two different experiments
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Determinin the Rate Law of a
Reaction
5. Obtain the value of m and n directly by finding the ratio of the
rate laws for two experiments in which the concentration of one
of the reactants is the same such as Experiments 1 and 3 in the
table 14.4.
rate1 = k[A1]m [B1]n
rate3 = k[A3]m [B3]n
6. By selecting two experiments in which the concentration of B
is the same, one can solve for the value of m; by selecting two
experiments in which the concentration of A is the same, one can
solve for n
7. Calculate the rate constant by inserting data from any line
from the table into the experimentally determined rate law and
solve for k
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Chemistry: Principles, Patterns,
and Applications, 1e
14.4 Using Graphs to
Determine Rate Laws, Rate
Constants, and Reaction Orders
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14.4 Using Graphs to Determine Rate Laws,
Rate Constants, and Reaction Orders
• For a zeroth-order reaction, a plot of the concentration of
any reactant versus time is a straight line with a slope
of – k.
• For a first-order reaction, a plot of the logarithm of the
concentration of a reactant versus time is a straight line
with a slope of – k.
• For a second-order reaction, a plot of the inverse of the
concentration of a reactant versus time is a straight line
with a slope of k.
• Properties of reactions that obey zeroth-, first-, and
second-order rate laws are summarized in the following
table.
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14.4 Using Graphs to Determine Rate Laws,
Rate Constants, and Reaction Orders
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Chemistry: Principles, Patterns,
and Applications, 1e
14.5 Half-Lives and
Radioactive Decay
Kinetics
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Half-Lives
• Another approach to describe reaction rates is based on
the time required for the concentration of a reactant to
decrease to one-half its initial value.
• The period of time is called the half-life of the reaction,
written as t½ .
• The half-life of a reaction is the time required for the
reactant concentration to decrease from [A]0 to [A]0 /2.
• If two reactions have the same order, the faster reaction
will have a shorter half-life and the slower reaction will
have a longer half-life.
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Half-Lives
•
The half-life of a first-order reaction under a given set of reaction conditions
is a constant; this is not true for zeroth- or second-order reactions.
•
The half-life of a first-order reaction is independent of the concentration of
the reactants.
•
Rearranging the integrated rate law for a first-order reaction produces the
equation
[A]0
ln
[A] = kt
•
Substituting [A]0 /2 for [A] and t½ for t (to indicate a half-life) into the above
equation gives
[A]0
ln [A] /2 = ln 2 = kt½
0
• Solving for t½: t½ = 0.693/k
• For a first-order reaction, each successive
half-life is the same length of time and is
independent of [A].
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Radioactive Decay Rates
• Radioactivity, or radioactive decay, is the emission of a
particle or a photon that results from the spontaneous
decomposition of the unstable nucleus of an atom.
• The rate of radioactive decay is an intrinsic property of
each radioactive isotope, independent of the chemical
and physical form of the radioactive isotope.
• Rate is also independent of temperature.
• In a sample of a given radioactive substance, the
number of atoms of the radioactive isotope must
decrease with time as their nuclei decay to nuclei of a
more stable isotope.
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Radioactive Decay Rates
• Using N to represent the number of atoms of the
radioactive isotope, the rate of decay of the sample
(also called its activity, A) can be defined as the
decrease in the number of the radioisotope’s nuclei per
unit time A = – N/t
• Activity is measured in disintegrations per second (dps)
or disintegrations per minute (dpm)
• Activity of a sample is directly proportional to the number
of atoms of the radioactive isotope in the sample A = kN
• k is the radioactive decay constant and has units of
inverse time (s–1, yr –1) and has a characteristic value for
each radioactive isotope
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Radioactive Decay Rates
• Combining equations, the relationship between the
number of decays per unit time and the number of atoms
of the isotope in a sample is obtained
– N/t = kN
• The equation is the same as the equation for the rate of
a first-order reaction; except that it uses number of
atoms instead of concentrations
• Radioactive decay is a first-order process and can be
described in terms of either the differential rate law as
above or the integrated rate law
N = N0e–kt or ln N/N0 = – kt
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Radioactive Decay Rates
• Because radioactive decay is a first-order
process, the time required for half of the nuclei
in any sample of a radioactive isotope to decay
is a constant, called the half-life of the isotope
• Half-life tells how radioactive an isotope is (the
number of decays per unit time) and is the most
commonly cited property of any isotope
• Isotopes with a short half-life decay more rapidly,
undergoing a greater number of radioactive
decays per unit time than do isotopes with a
long half-life
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Radioisotope Dating Techniques
• Using the half-lives of isotopes, one can estimate the
ages of geological and archaeological artifacts.
• Techniques that have been developed for this application
are known as radioisotope dating techniques.
• The most common method for measuring the age of
ancient objects is carbon-14 dating; carbon-14 isotope,
created in the upper regions of Earth’s atmosphere,
reacts with atmospheric oxygen or ozone to form 14CO2.
• The CO2 that plants use as a carbon source include a
proportion of 14CO2 molecules as well as nonradioactive
12CO and 13CO ; animals that eat plants ingest these
2
2
carbon isotopes.
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Radioisotope Dating Techniques
• When the animals or plants die, the carbon-14 nuclei in its
tissue decay to nitrogen-14 nuclei by a radioactive
process known as beta decay, which releases low-energy
electrons ( particles) that can be detected and measured:
14C  14N + – with a half-life of 5700 ± 30 yr.
• The 14C/ 12C ratio in living organisms is 1.3 x 10–12 with a
decay rate of 15 dpm per gram of carbon (dpm/g carbon).
• Comparing the disintegrations per minute per gram of
carbon from an archaeological sample with those from a
recently living sample enables scientists to estimate the
age of the artifact.
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Chemistry: Principles, Patterns,
and Applications, 1e
14.6 Reaction Rates—A
Microscopic View
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14.6 Reaction Rates—A Microscopic
View
• One of the major reasons for studying chemical kinetics
is to use measurements of the macroscopic properties of
a system to discover the sequence of events that occur
at the molecular level during a reaction.
• Molecular description is the mechanism of the reaction; it
describes how individual atoms, ions, or molecules
interact to form particular products.
• Stepwise changes are called the reaction mechanism,
the microscopic path by which reactants are transformed
into products by a complex series of reactions that take
place in a stepwise fashion.
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14.6 Reaction Rates—A Microscopic
View
• Each step or individual reaction is called an elementary
reaction, and the overall sequence of elementary
reactions is the mechanism of the reaction.
• The sum of the individual steps, or elementary reactions,
in the mechanism must give the balanced chemical
equation for the overall reaction.
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Molecularity and the Rate-Determining
Step
• Species that are formed in one step and consumed in
another are intermediates; and they do not appear in
the balanced chemical equation for the reaction.
• Following the two-step mechanism is an example:
Step 1 NO2 + NO2  NO3 + NO
Elementary reaction
Step 2 NO3 + CO  NO2 + CO2
Elementary reaction
Sum NO2 + CO  NO + CO2
Overall reaction
• NO3 is an intermediate in the reaction.
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Molecularity and the Rate-Determining
Step
• Using molecularity to describe a rate law
– Each elementary step can be described in terms of its
molecularity, the number of molecules that collide in that step.
– If there is only a single reactant molecule in an elementary
reaction, that step is designated as unimolecular.
– If there are two reactant molecules, it is bimolecular; if there are
three reactant molecules, it is termolecular.
– Order of the elementary reaction is the same as its molecularity,
but the rate law for the reaction cannot be determined from the
balanced equation for the overall reaction.
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Molecularity and the Rate-Determining
Step
– The general rate law for a unimolecular elementary reaction (A
 products) is rate = k[A].
– For bimolecular reactions, the reaction rate depends on the
number of collisions per unit time, which is proportional to the
product of the concentrations of the reactants.
– For a bimolecular elementary reaction of the form 2A 
products, the general rate law is rate = k[A]2.
– Common types of elementary reactions and their rate laws are
summarized in the following table:
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Molecularity and the Rate-Determining
Step
• Identifying the rate-determining step
– The balanced chemical equation does not
necessarily reveal the individual elementary reactions
by which the reaction occurs.
– One cannot obtain the rate law for a reaction from
the overall balanced equation alone.
– The rate law for the overall reaction is the same as
the rate law for the slowest step in the reaction
mechanism, the rate-determining step, because any
process that occurs through a sequence of steps can
take place no faster than the slowest step in the
sequence.
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Chain Reactions
• Many reaction mechanisms consist of long series of elementary
reactions called chain reactions, in which one or more elementary
reactions that contain a highly reactive species repeat again and
again during the reaction process.
• Chain reactions have three stages:
1. initiation, a step that produces one or more reactive intermediates;
often these intermediates are radicals, species that have an unpaired
valence electron;
2. propagation, reactive intermediates are continuously consumed and
regenerated while products are formed;
3. termination, intermediates are also consumed, usually by forming
stable products.
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Chemistry: Principles, Patterns,
and Applications, 1e
14.7 The Collision Model
of Chemical Kinetics
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14.7 The Collision Model of
Chemical Kinetics
• A useful tool for understanding the behavior of reacting
chemical species
• The collision model explains the following:
– A chemical reaction can occur only when the reactant
molecules, atoms, or ions collide with more than a certain
amount of kinetic energy and in the proper orientation
– Explains why most collisions between molecules do not result
in a chemical reaction, because in most collisions, the molecules
simply bounce off one another without reacting
– Why such chemical reactions occur more rapidly at higher
temperatures
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Activation Energy
• A minimum energy (activation energy, Ea) is required for a collision
between molecules to result in a chemical reaction.
• Reacting molecules must have enough energy to overcome
electrostatic repulsion and a minimum amount of energy to break
chemical bonds so that new ones may be formed.
• Molecules that collide with less than the activation energy bounce off
one another chemically unchanged, with only their direction of travel
and their speed altered by the collision.
• Molecules that are able to overcome the energy barrier react and
form an arrangement of atoms called the activation complex or the
transition state.
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Graphing the Energy Changes during a
Reaction
• One can graph the energy of a reaction by plotting the potential
energy of the system as the reaction progresses with time.
• Plots show an energy barrier that must be overcome for the reaction
to occur, which means that the activation energy is always positive.
• Ea provides information about the rate of a reaction and how rapidly
the rate changes with temperature.
• For two similar reactions under comparable conditions, the reaction
with the smallest Ea will occur more rapidly.
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Graphing the Energy Changes during a
Reaction
• Even when the energy of collisions between two reactant species is
greater than Ea, most collisions do not produce a reaction; the
probability of a reaction occurring depends not only on the collision
energy, but also on the spatial orientation of the molecules when
they collide.
• The fracture of orientations that result in a reaction is called the
steric factor, p, and its value can range from 0 (no orientations of
molecules result in reaction) to 1 (all orientations result in reaction).
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The Arrhenius Equation
• In the following figure, both the kinetic energy
distributions and a potential energy diagram for a
reaction are shown
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The Arrhenius Equation
•
– Shaded areas show that at the lower temperature, only a small
fraction of molecules collide with kinetic energy greater than Ea;
at the higher temperature, a much larger fraction of molecules
collide with kinetic energy greater than Ea
– The rate of the reaction is much slower at the lower
temperature because only a few molecules collide with enough
energy to overcome the potential energy barrier
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The Arrhenius Equation
• For an A + B elementary reaction, all the factors that affect the
reaction rate can be summarized in a single series of relationships:
rate = (collision frequency) (steric factor) (fraction of collisions with E > Ea)
where rate = k[A] [B]
• Arrhenius used these relationships to arrive at an equation that
relates the magnitude of the rate constant of a reaction to the
temperature; the activation energy; and the constant, A, called the
frequency factor, which converts concentrations to collisions per
second
k = Ae–Ea /RT (Arrhenius equation)
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The Arrhenius Equation
• The Arrhenius equation summarizes the collision
model of chemical kinetics, where T is the
absolute temperature (in K) and R is the ideal
gas constant [8.314 J/(K•mol)].
• The value of Ea indicates the sensitivity of the
reaction to changes in temperature.
• The rate of a reaction with a large Ea increases
rapidly with increasing temperature, and the rate
of a reaction with a smaller Ea increases much
more slowly with increasing temperature.
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The Arrhenius Equation
• If the rate of a reaction at various temperatures is known,
the Arrhenius equation can be used to calculate the
activation energy by taking the natural logarithm of both
sides of the Arrhenius equation.
ln k = ln A + (–Ea /RT) = ln A + [(–Ea /R) (1/T )]
• The preceding equation is the equation for a straight line,
where y = ln k and x = 1/T; a plot of ln k versus 1/T is a
straight line with a slope of –Ea /R and an intercept of ln
A.
• Knowing the value of Ea at one temperature predicts the
rate of a reaction at other temperatures.
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Chemistry: Principles, Patterns,
and Applications, 1e
14.8 Catalysis
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14.8 Catalysis
• Catalysts
– Substances that increase the rate of a chemical reaction without being
consumed in the process
– A catalyst does not appear in the overall stoichiometry of the reaction it
catalyzes, but it must appear in at least one of the elementary steps in
the mechanism for the catalyzed reaction
– Catalyzed pathway has a lower Ea, but the net change in energy that
results from the reaction (the difference between the energy of the
reactants and the energy of the products) is not affected by the
presence of a catalyst
– Because of its lower Ea, the rate of a catalyzed reaction is faster than
the rate of the uncatalyzed reaction at the same temperature
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14.8 Catalysis
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14.8 Catalysis
– A catalyst decreases the height of the energy barrier, and its
presence increases the rates of both the forward and the reverse
reactions by the same amount
– There are three major classes of catalysts
1. Heterogeneous catalysts
2. Homogeneous catalysts
3. Enzymes
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Heterogeneous Catalysis
• In heterogeneous catalysis, the catalyst is in a different
phase from the reactants.
• At least one of the reactants interacts with the solid
surface (in a physical process called adsorption) in
such a way that a chemical bond in the reactant
becomes weak and then breaks.
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Homogeneous Catalysis
• In homogeneous catalysis, the catalyst is in the same
phase as the reactant(s); the number of collisions
between reactants and catalyst is at a maximum
because the catalyst is uniformly dispersed throughout
the reaction mixture.
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Enzymes
• Enzymes are catalysts that occur naturally in living organisms and
are almost all protein molecules with typical molecular masses of
20,000–100,000 amu.
• Some are homogeneous catalysts that react in aqueous solution
within a cellular compartment of an organism.
• Some are heterogeneous catalysts embedded within the membranes
that separate cells and cellular compartments from their
surroundings.
• A reactant in an enzyme-catalyzed reaction is called a substrate.
• Enzymes can increase reaction rates by enormous factors and tend
to be very specific, typically producing only a single product in
quantitative yield.
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Enzymes
• Enzymes are expensive, and often cease functioning at
temperatures higher than 37ºC, and have limited stability
in solution.
• Enzyme inhibitors cause a decrease in the rate of an
enzyme-catalyzed reaction by binding to a specific
portion of an enzyme and thus slowing or preventing a
reaction from occurring.
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15
Chemical
Equilibrium
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
CHAPTER OBJECTIVES
• To understand what is meant by chemical equilibrium
• To know the relationship between the equilibrium constant and the
rate constants for the forward and reverse reactions
• To be able to write an equilibrium constant expression for any
reaction
• To be able to solve quantitative problems involving chemical
equilibrium
• To predict the direction of reaction
• To predict the effects of stresses on a system at equilibrium
• To understand different ways to control the products of a reaction
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Chemistry: Principles, Patterns,
and Applications, 1e
15.1 The Concept of
Chemical Equilibrium
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15.1 The Concept of Chemical Equilibrium
• Chemical equilibrium
– A dynamic process
– Consists of a forward reaction, in which reactants are
converted to products, and a reverse reaction, in which products
are converted to reactants
– At equilibrium, the forward and reverse reactions proceed at
equal rates
– Double arrow (⇋) indicates that both the forward and reverse
reactions are occurring simultaneously and is read “is in
equilibrium with”
– At equilibrium, the composition of the system no longer
changes with time
– Composition of an equilibrium mixture is independent of the
direction from which equilibrium is approached
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Chemistry: Principles, Patterns,
and Applications, 1e
15.2 The Equilibrium
Constant
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15.2 The Equilibrium Constant
• Equilibrium state is achieved when the rate of the forward reaction
equals the rate of the reverse reaction.
• Under a given set of conditions, there must be a relationship
between the composition of the system at equilibrium and the
kinetics of a reaction represented by rate constants.
• The ratio of the rate constants yields a new constant, the
equilibrium constant (K), a unitless quantity and is defined as K =
kf /kr.
• The fundamental relationship between chemical kinetics and
chemical equilibrium states that, under a given set of conditions, the
composition of the equilibrium mixture is determined by the
magnitudes of the rate constants for the forward and reverse
reactions.
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Developing an Equilibrium Constant
Expression
• In1864, Guldberg and Waage discovered that for any reversible
reaction of the general form
aA + bB ⇋ cC + dD,
where A and B are reactants, C and D are products, and a, b, c, and
d are the stoichiometric coefficients in the balanced equation for the
reaction, the ratio of the product of the equilibrium concentrations of
the products (raised to their coefficients in the balanced equation) to
the product of the equilibrium concentrations of the reactants (raised
to their coefficients in the balanced equation) is always a constant
under a given set of conditions.
This equation is called the equilibrium equation.
• This relationship is known as the law of mass action
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Developing an Equilibrium Constant
Expression
• Law of mass action is stated as
K = [C]c [D]d
[A]a [B]b
– K is the equilibrium constant for the reaction
– Right side of the equation is called the equilibrium constant
expression
– Relationship is true for any pair of opposing reactions
regardless of the mechanism of the reaction or of the number of
steps in the mechanism
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Developing an Equilibrium Constant
Expression
• Equilibrium constant (K)
– Can vary over a wide range of values and is unitless
– Values of K greater than 103 indicate a strong tendency for
reactants to form products, so equilibrium lies to the right, favoring
the formation of products (kf >>kr)
– Values of K less than 10–3 indicate that the ratio of products to
reactants at equilibrium is very small; reactants do not tend to form
products readily, and equilibrium lies to the left, favoring the
formation of reactants (kf<<kr)
– Values of K between 103 and 10–3 are not very large or small, so
there is no strong tendency to form either products or reactants; at
equilibrium, there are significant amounts of both products and
reactants (kf  kr)
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Developing an Equilibrium Constant
Expression
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Variations in the Form of the Equilibrium
Constant Expression
• Equilibrium can be approached from either direction in a
chemical reaction, so the equilibrium constant
expression and the magnitude of the equilibrium
constant depend on the form in which the chemical
reaction is written.
• When a reaction is written in the reverse direction,
cC + dD ⇋ aA + bB
K and the equilibrium constant expression are inverted:
K´= [A]a [B]b
[C]c [D]d
so K´ = 1/K
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Equilibrium Constant Expressions for
Systems that Contain Gases
• For reactions that involve nongaseous substances, the
concentrations used in equilibrium calculations are expressed in
moles/liter.
• For gases, the concentrations are expressed in terms of partial
pressures where the standard state is 1 atm of pressure.
• Symbol Kp is used to denote equilibrium constants calculated from
partial pressures.
• For the general reaction aA + bB ⇋ cC + dD in which all the
components are gases, the equilibrium constant is the ratio of the
partial pressures of the products and reactants, each raised to its
coefficient in the chemical equation.
(PA)c (PC)d
Kp =
(PB)a (PD)b
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Equilibrium Constant Expressions for
Systems that Contain Gases
• Kp is unitless.
• Partial pressures are expressed in atmospheres or mmHg, so the molar
concentration of a gas and its partial pressure do not have the same
numerical value but are related by the ideal gas constant R and the
temperature
Kp = K(RT)n
where K is the equilibrium constant expressed in units of concentration and
n is the difference between the number of moles of gaseous products and
gaseous reactants; temperature is expressed in kelvins.
• If n = 0, Kp = K
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Homogeneous and Heterogeneous
Equilibria
• Homogeneous equilibrium
– When the products and reactants of an equilibrium reaction
form a single phase, whether gas or liquid
– Concentrations of the reactants and products can vary over a
wide range
• Heterogeneous equilibrium
– A system whose reactants, products, or both are in more than
one phase
– An example is the reaction of a gas with a solid or liquid
• Molar concentrations of pure liquids and solids do not
vary with temperature, so they are treated as constants,
this simplifies their equilibrium constant expressions
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Equilibrium Constant Expressions for the
Sums of Reactions
• When a reaction can be expressed as the sum of two or
more reactions, its equilibrium constant is equal to the
product of the equilibrium constants for the individual
reactions.
• H for the sum of two or more reactions is the sum of
the H values for the individual reactions.
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Chemistry: Principles, Patterns,
and Applications, 1e
15.3 Solving Equilibrium
Problems
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15.3 Solving Equilibrium Problems
• Two fundamental kinds of equilibrium problems
1. Those in which the concentrations of the reactants and
products at equilibrium are given and the equilibrium constant for
the reaction needs to be calculated
2. Those in which the equilibrium constant and the initial
concentrations of reactants are known and the concentration of
one or more substances at equilibrium needs to be calculated
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Calculating an Equilibrium Constant from
Equilibrium Concentrations
• An equilibrium constant can be calculated when
equilibrium concentrations, molar concentrations, or
partial pressures are substituted into the equilibrium
constant expression for the reaction.
• Sometimes the concentrations of all the substances are
not given or the equilibrium concentrations of all the
relevant substances for a particular system are not
measured. In these cases, the equilibrium
concentrations can be obtained from the initial
concentrations of the reactants and the balanced
equation for the reaction, as long as the equilibrium
concentration of one of the substances is known.
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Calculating Equilibrium Concentrations
from the Equilibrium Constant
• Equilibrium constants can be used to calculate the
equilibrium concentrations of reactants and products by
using the quantities or concentrations of the reactants,
the stoichiometry of the balanced equation for the
reaction, and a tabular format to obtain the final
concentrations of all species at equilibrium.
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Chemistry: Principles, Patterns,
and Applications, 1e
15.4 Nonequilibrium
Conditions
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15.4 Nonequilibrium Conditions
• One must often decide whether a system has reached
equilibrium or the composition of the mixture will
continue to change with time
• To make this determination, one needs to know how to
analyze the composition of a reaction mixture
quantitatively
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The Reaction Quotient (Q)
• Reaction quotient (Q)
– A quantity used to determine whether a system has reached
equilibrium
– Expression for the reaction quotient has the same form as the
equilibrium constant expression
– Q may be derived from a set of values measured at any time
during the reaction of any mixture of the reactants and products,
regardless of whether the system is at equilibrium
– For the general reaction aA + bB ⇋ cC + dD, reaction quotient
is defined as Q = [C]c [D]d
[A]a [B]b
– Qp can be written for any reaction that involves gases by using
the partial pressures of the components
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The Reaction Quotient (Q)
• Comparing the magnitudes of Q and K allows the determination of
whether a reaction mixture is already at equilibrium and, if it is not,
how to predict whether its composition will change with time
(whether the reaction will proceed to the right or to the left)
1. If Q = K, the system is at equilibrium, no further change in the
composition of the system will occur unless the conditions are changed
2. If Q < K, then the ratio of the concentrations of products to the
concentration of reactants is less than the ratio at equilibrium; reaction
will proceed to the right, forming products at the expense of reactants
3. If Q > K, then the ratio of the concentrations of products to the
concentrations of reactants is greater than at equilibrium; reaction will
proceed to the left, forming reactants at the expense of products
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The Reaction Quotient (Q)
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Predicting the Direction of Reaction
Using a Graph
• Graphs derived by plotting a few equilibrium
concentrations for a system at a given temperature and
pressure can be used to predict the direction in which a
reaction will proceed
• Points that do not lie on the line represent
nonequilibrium states and the system will adjust to
achieve equilibrium
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Chemistry: Principles, Patterns,
and Applications, 1e
15.5 Factors That Affect
Equilibrium
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15.5 Factors That Affect Equilibrium
• Strategies are used to increase the yield of the desired
products of reactions
• Reaction conditions are controlled to obtain the
maximum amount of the desired product
• Changes in reaction conditions affect the equilibrium
composition of a system
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Le Châtelier’s Principle
•
When a system at equilibrium is perturbed in some
way, the effects of the perturbation can be predicted
qualitatively using Le Châtelier’s principle.
•
This principle states that if a stress is applied to a
system at equilibrium, the composition of the system
will change to relieve the applied stress.
• Stress occurs when any change in the system affects
the magnitude of Q or K.
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Le Châtelier’s Principle
•
Three types of stresses can change the
composition of an equilibrium mixture
1. A change in the concentrations (or partial pressures) of the
components by the addition or removal of reactants or
products
2. A change in the total pressure or volume
3. A change in the temperature of the system
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Changes in Concentration
• An equilibrium is disturbed by adding or removing a
reactant or product
1. Stress of an added reactant or product is relieved by
reaction in the direction that consumes the added substance
a. Add reactant—reaction shifts right toward product
b. Add product—reaction shifts left toward reactant
2. Stress of removing reactant or product is relieved by
reaction in the direction that replenishes the removed
substance
a. Remove reactant—reaction shifts left
b. Remove product—reaction shifts right
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Changes in Concentration
• Changes that occur due to changes in the
value of Qc
1. Add reactant—denominator in Qc expression becomes larger
A. Qc < Kc
B. To return to equilibrium, Qc must increase
I. Numerator of Qc expression must increase and
the denominator must decrease
II. Implies net conversion of reactants to
products; reaction shifts right
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Changes in Concentration
2. Remove reactant—denominator in Qc expression becomes
smaller
A. Qc > Kc
B. To return to equilibrium, Qc must decrease
I. Numerator of Qc expression must decrease and
the denominator must increase
II. Implies net conversion of products to reactants;
reaction shifts left
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Changes in Total Pressure or Volume
• If a balanced reaction contains different numbers of
gaseous reactant and product molecules, the equilibrium
will be sensitive to changes in volume or pressure;
change in pressure (due to changing volume) changes
the composition of the equilibrium mixture
• Increase in pressure (due to decrease in volume) results
in a reaction in the direction of a fewer number of moles
of gas
• Decrease in pressure (due to increase in volume) results
in a reaction in the direction of a greater number of
moles of gas
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Changes in Total Pressure or Volume
• Changes occur due to changes in the value of Qc
1. Decrease volume—molarity increases
2. If reactant side has more moles of gas
a. Increase in denominator is greater than increase in numerator
and Qc < Kc
b. To return to equilibrium, Qc must increase; the numerator of the
Qc expression must increase and denominator must
decrease—it shifts toward fewer moles of gas (reactants to
products)
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Changes in Total Pressure or Volume
3. If product side has more moles of gas
a. Increase in numerator is greater than increase in denominator
and Qc > Kc
b. To return to equilibrium, Qc must decrease; the denominator of
the Qc expression must decrease and the numerator must
increase—it shifts toward fewer moles of gas (products to
reactants)
• If the reaction involves no change in the number of
moles of gas, there is no effect on the composition of the
equilibrium mixture
• Effect of pressure changes on solids and liquids can be
ignored
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Changes in Temperature
• Changes in temperature can change the value of the
equilibrium constant without affecting the reaction
quotient (Q  K)
• System is no longer at equilibrium, and the composition
of the system will change until Q equals K at the new
temperature
• To predict how an equilibrium system will respond to a
change in temperature, one must know the enthalpy
change of the reaction, Hrxn
1. Exothermic (heat is released, H < 0): reactants ⇋ products +
heat
2. Endothermic (heat is absorbed, H > 0): reactants + heat ⇋
products
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Changes in Temperature
• Heat is a product in an exothermic reaction and a
reactant in an endothermic reaction; increasing the
temperature of a system corresponds to adding heat
• Le Châtelier’s principle predicts
1. that an exothermic reaction will shift to the left (toward
reactants) if the temperature of the system is increased (heat is
added);
2. that an endothermic reaction will shift to the right (toward the
products) if the temperature of the system is increased;
3. that if Hrxn = 0, then a change in temperature will not affect
the equilibrium composition.
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Changes in Temperature
• Value of Kc
1. Increasing the temperature increases the magnitude of the
equilibrium constant for an endothermic reaction
2. Increasing the temperature decreases the equilibrium
constant for an exothermic reaction
3. Increasing the temperature has no effect on the
equilibrium constant for a thermally neutral reaction
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Chemistry: Principles, Patterns,
and Applications, 1e
15.6 Controlling the
Products of Reactions
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15.6 Controlling the Products of
Reactions
• One of the primary goals of modern chemistry is to
control the identity and quantity of the products of
chemical reactions.
• Two approaches
1. To get a high yield of a desired compound, make the rate of
the desired reaction much faster than the rate of any of the other
possible reactions that might occur in the system; altering
reaction conditions to control reaction rates, thereby obtaining a
single product or set of products is called kinetic control.
2. Thermodynamic control—consists of adjusting conditions so
that at equilibrium only the desired products are present in
significant quantities.
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16
Aqueous
Acid-Base
Equilibria
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CHAPTER OBJECTIVES
•
To understand the autoionization reaction of liquid water
•
To know the relationship among pH, pOH, and pKw
•
To understand the concept of conjugate acid-base pairs
•
To know the relationship between acid or base strength and the magnitude
of Ka, Kb, pKa, and pKb
•
To understand the leveling effect
•
To be able to predict whether reactants or products are favored in an
acid-base equilibrium
•
To understand how molecular structure determines acid and base strengths
•
To be able to use Ka and Kb values to calculate the percent ionization and
pH of a solution of an acid or a base
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CHAPTER OBJECTIVES
•
To be able to calculate the pH at any point in an acid-base titration
•
To understand how the addition of a common ion affects the position of an
acid-base equilibrium
•
To understand how a buffer works and how to use the
Henderson-Hasselbalch equation to calculate the pH of a buffer
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Chemistry: Principles, Patterns,
and Applications, 1e
16.1 The Autoionization of
Water
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16.1 The Autoionization of Water
• Acids and bases can be defined in different
ways:
1. Arrhenius definition: An acid is a substance that dissociates in
water to produce H+ ions (protons), and a base is a substance
that dissociates in water to produce OH– ions (hydroxide); an
acid-base reaction involves the reaction of a proton with the
hydroxide ion to form water
2. Brønsted–Lowry definition: An acid is any substance that can
donate a proton, and a base is any substance that can accept a
proton; acid-base reactions involve two conjugate acid-base
pairs and the transfer of a proton from one substance (the acid)
to another (the base)
3. Lewis definition: A Lewis acid is an electron-pair acceptor, and
a Lewis base is an electron-pair donor
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16.1 The Autoionization of Water
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Acid-Base Properties of Water
• Water is amphiprotic: it can act as an acid by donating
a proton to a base to form the hydroxide ion, or as a
base by accepting a proton from an acid to form the
hydronium ion, H3O+
• Structure of the water molecule
1. Polar O–H bonds and two lone pairs of electrons on
the oxygen atom
2. Liquid water has a highly polar structure
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The Ion-Product Constant of Liquid Water
• Because water is amphiprotic, one water molecule can react with
another to form an OH– ion and an H3O+ ion in an autoionization
process:
2H2O(l)⇋H3O+ (aq) + OH– (aq)
• Equilibrium constant K for this reaction can be written as
K = [H3O+] [OH–]
[H2O]2
• When pure liquid water is in equilibrium with hydronium and
hydroxide ions at 25ºC, the concentrations of hydronium ion and
hydroxide ion are equal: [H3O+] = [OH–] = 1.003 x 10–7 M
• At 25ºC, the density of liquid water is 0.0997 g/mL
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The Ion-Product Constant of Liquid Water
• The concentration of liquid water at 25ºC is
[H2O] = mol/L = (0.997 g/mL) (1 mol/18.02 g) (1000 mL/L) = 55.3 M
• Because the number of dissociated water molecules is very small,
the equilibrium of the autoionization reaction lies far to the left, so
the concentration of water is unchanged by the autoionization
reaction and can be treated as a constant
• By treating [H2O] as a constant, a new equilibrium constant, the
ion-product constant of liquid water (Kw), can be defined:
K[H2O]2 = [H3O+] [OH–] or Kw = [H3O+] [OH–]
• Substituting the values for [H3O+] and [OH–] at 25ºC gives
Kw = (1.003 x 10–7) (1.003 x 10–7) = 1.006 x 10–14
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The Ion-Product Constant of Liquid Water
• Kw varies with temperature, ranging from 1.15 x 10–15 at
0ºC to 4.99 x 10–13 at 100ºC
• In pure water, the concentrations of the hydronium ion
and the hydroxide ion are the same, so the solution is
neutral
• If [H3O+] > [OH–], the solution is acidic
• If [H3O+] < [OH–], the solution is basic
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The Relationship among pH, pOH, and pKw
• The pH scale is a concise way of describing the H3O+ concentration
and the acidity or basicity of a solution
• pH and H+ concentration are related as follows:
pH = –log10[H+] or [H+] = 10–pH
• pH of a neutral solution ([H3O+] = 1.00 x 10–7 M) is 7.00
• pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 10–7
• pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 10–7
• The pH scale is logarithmic, so a pH difference of 1 between two
solutions corresponds to a difference of a factor of 10 in their
hydronium ion concentrations
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The Relationship among pH, pOH, and pKw
• There is an analogous pOH scale to describe the hydroxide ion
concentration of a solution; pOH and [OH–] are related as follows:
pH = –log10[OH–] or [OH–] = 10–pOH
• A neutral solution has [OH–] = 1.00 x 10–7, so the pOH of a neutral
solution is 7.00
• The sum of the pH and the pOH for a neutral solution at 25ºC is 7.00
+ 7.00 = 14.00
pKw = –log Kw = –log([H3O+] [OH–]) =
(–log[H3O+]) + (–log[OH–]) = pH + pOH
• At any temperature, pH + pOH = pKw, and at 25ºC, where Kw = 1.01
x 10–14, pH + pOH = 14.00; pH of any neutral solution is just half the
value of pKw at that temperature
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Chemistry: Principles, Patterns,
and Applications, 1e
16.2 A Qualitative
Description of Acid-Base
Equilibria
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Conjugate Acid-Base Pairs
• Two species that differ by only a proton constitute a conjugate
acid-base pair.
1. Conjugate base has one less proton than its acid; A– is the conjugate
base of HA
2. Conjugate acid has one more proton than its base; BH+ is the
conjugate acid of B
• In the reaction of HCl with water, HCl, the parent acid, donates a
proton to a water molecule, the parent base, forming Cl–; HCl and
Cl– constitute a conjugate acid-base pair.
• In the reverse reaction, the Cl– ion in solution acts as a base to
accept a proton from H3O+, forming H2O and HCl; H3O+ and H2O
constitute a second conjugate acid-base pair.
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Conjugate Acid-Base Pairs
• Any acid-base reaction must contain two conjugate
acid-base pairs, which in this example are HCl/Cl– and
H3O+/H2O
• HCl (aq) + H2O (l)  H3O+ (aq) + Cl– (aq)
parent acid
parent base
conjugate acid conjugate base
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Acid-Base Equilibrium Constants:
Ka, Kb, pKa, and pKb
• The magnitude of the equilibrium constant for an ionization reaction can be
used to determine the relative strengths of acids and bases
• The general equation for the ionization of a weak acid in water, where HA is
the parent acid and A– is its conjugate base, is
HA(aq) + H2O(l) ⇋ H3O+(aq) + A–(aq)
• The equilibrium constant for this dissociation is
K = [H3O+] [A–]
[H2O] [HA]
• The concentration of water is constant for all reactions in aqueous solution,
so [H2O] can be incorporated into a new quantity, the acid ionization
constant (Ka):
Ka = K[H2O] = [H3O+] [A–]
[HA]
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Acid-Base Equilibrium Constants:
Ka, Kb, pKa, and pKb
• Numerical values of K and Ka differ by the concentration of water
(55.3 M); the larger the value Ka, the stronger the acid and the
higher the H+ concentration at equilibrium
• Weak bases react with water to produce the hydroxide ion, B(aq) +
H2O(l) ⇋ BH+(aq) + OH–(aq), where B is the parent base and BH+ is
its conjugate acid
• Equilibrium constant for this reaction is the base ionization
constant (Kb); concentration of water is constant and does not
appear in the equilibrium constant expression but is included in the
value of Kb
• The larger the value of Kb, the stronger the base and the higher the
OH– concentration at equilibrium
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Acid-Base Equilibrium Constants:
Ka, Kb, pKa, and pKb
• The sum of the reactions described by Ka and Kb is the
equation for the autoionization of water, and the product
of the two equilibrium constants is Kw
• For any conjugate acid-base pair, KaKb = Kw
• pKa = –log10Ka and pKb = –log10Kb
• Smaller values of pKa correspond to larger acid
ionization constants and stronger acids
• Smaller values of pKb correspond to larger base
ionization constants and stronger bases
• At 25ºC, pKa + pKb = 14.00
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Acid-Base Equilibrium Constants:
Ka, Kb, pKa, and pKb
• There is an inverse relationship between the strength of
the parent acid and the strength of the conjugate base;
the conjugate base of a strong acid is a weak base, and
the conjugate base of a weak acid is a strong base
• One can use the relative strengths of acids and bases to
predict the direction of an acid-base reaction by following
a simple rule: An acid-base equilibrium always favors
the side with the weaker acid and base
stronger acid + stronger base
weaker acid + weaker base
• In an acid-base reaction, the proton always reacts with
the stronger base
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Solutions of Strong Acids and Bases:
The Leveling Effect
• No acid stronger than H3O+ and no base stronger than OH– can exist
in aqueous solution, leading to the phenomenon known as the
leveling effect.
• Any species that is a stronger acid than the conjugate acid of water
(H3O+) is leveled to the strength of H3O+ in aqueous solution
because H3O+ is the strongest acid that can exist in equilibrium with
water.
• In aqueous solution, any base stronger than OH– is leveled to the
strength of OH– because OH– is the strongest base that can exist in
equilibrium with water
• Any substance whose anion is the conjugate base of a compound
that is a weaker acid than water is a strong base that reacts
quantitatively with water to form hydroxide ion
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Polyprotic Acids and Bases
• Polyprotic acids contain more than one ionizable proton,
and the protons are lost in a stepwise manner.
• The fully protonated species is always the strongest acid
because it is easier to remove a proton from a neutral
molecule than from a negatively charged ion; the fully
deprotonated species is the strongest base.
• Acid strength decreases with the loss of subsequent
protons, and the pKa increases.
• The strengths of the conjugate acids and bases are
related by pKa + pKb = pKw, and equilibrium favors
formation of the weaker acid-base pair.
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Acid-Base Properties of Solutions
of Salts
• A salt can dissolve in water to produce a neutral, basic, or acidic
solution, depending on whether it contains the conjugate base of a
weak acid as the anion (A–) or the conjugate acid of a weak base as
the cation (BH+), or both.
• Salts that contain small, highly charged metal ions produce acidic
solutions in water.
• The most important parameter for predicting the effect of a metal ion
on the acidity of coordinated water molecules is the charge-to-radius
ratio of the metal ion.
• The reaction of a salt with water to produce an acidic or basic
solution is called a hydrolysis reaction, which is just an acid-base
reaction in which the acid is a cation or the base is an anion.
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Chemistry: Principles, Patterns,
and Applications, 1e
16.3 Molecular Structure
and Acid-Base Strength
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Bond Strengths
• The acid-base strength of a molecule depends strongly
on its structure.
• The stronger the A–H or B–H+ bond, the less likely the
bond is to break to form H+ ions, and thus the less acidic
the substance.
• The larger the atom to which H is bonded, the weaker
the bond.
• Acid strengths of binary hydrides increase as we go
down a column of the periodic table.
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Stability of the Conjugate Base
• The conjugate base (A– or B) contains one more lone
pair of electrons than the parent acid (AH or BH+).
• Any factor that stabilizes the lone pair on the conjugate
base favors dissociation of H+ and makes the parent acid
a stronger acid.
• Acid strengths of binary hydrides increase as we go from
left to right across a row of the periodic table.
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Inductive Effects
• Atoms or groups in a molecule other than those to which
H is bonded can induce a change in the distribution of
electrons within the molecule, called an inductive effect;
this can have a major effect on the acidity or basicity of
the molecule.
• Inductive effect can weaken an O–H bond and allow
hydrogen to be more easily lost as H+ ions.
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Chemistry: Principles, Patterns,
and Applications, 1e
16.4 Quantitative
Aspects of Acid-Base
Equilibria
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Determining Ka and Kb
• The ionization constants Ka and Kb are equilibrium
constants that are calculated from experimentally
measured concentrations.
• What does the concentration of an aqueous solution of a
weak acid or base exactly mean?
– A 1 M solution is prepared by dissolving 1 mol of acid or base in
water and adding enough water to give a final volume of exactly 1 L.
– If the actual concentrations of all species present in the solution
were listed, it would be determined that none of the values is exactly
1 M because a weak acid or a weak base always reacts with water
to some extent.
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Determining Ka and Kb
– The extent of the reaction depends on the value of Ka or Kb, the
concentration of the acid or base, and the temperature.
– Only the total concentration of both the ionized and unionized species
is equal to 1 M.
– The analytical concentration (C) is defined as the total concentration of
all forms of an acid or base that are present in solution, regardless of
their state of protonation.
– Thus; a 1 M solution has an analytical concentration of 1 M, which is
the sum of the actual concentrations of unionized acid or base and the
ionized form.
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Determining Ka and Kb
– In addition to the analytical concentration of the acid or base,
one must be able to measure the concentration of a least one of
the species in the equilibrium constant expression in order to
determine the value of Ka or Kb.
– Two common ways to obtain the concentrations
1. By measuring the electrical conductivity of the solution, which is
related to the total concentration of ions present
2. By measuring the pH of the solution, which gives [H+] or [OH–]
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Determining Ka and Kb
• Procedure for determining Ka for a weak acid
and Kb for a weak base
1. The analytical concentration of the acid or base is the initial
concentration
2. The stoichiometry of the reaction with water determines the
change in concentrations
3. The final concentrations of all species are calculated from the
initial concentrations and the changes in the concentrations
4. Inserting the final concentration into the equilibrium constant
expression enables the value of Ka or Kb to be calculated
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Calculating Percent Ionization from
Ka or Kb
• Need to know the concentrations of all species in solution
• The reactivity of a weak acid or a weak base is very different from the
reactivity of its conjugate base or acid; need to know the percent
ionization of a solution of an acid or base in order to understand a
chemical reaction
• Percent ionization is defined as
percent ionization of acid = [H+]
CHA
x 100
percent ionization of base = [OH–]
CB
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x 100
Calculating Percent Ionization from
Ka or Kb
• To determine the concentrations of species in
solutions of weak acids and bases, use a tabular
method
1. Make a table that lists the following values for each of the
species involved in the reaction
a. Initial concentration
b. The change in concentration on preceding to equilibrium
(–x or +x)
c. The final concentration—sum of the initial concentration and
the change in concentration
d. In constructing the table, define x as the concentration of the
acid that dissociates
2. Solve for x by substituting the final concentrations from the
table into the equilibrium constant expression
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Calculating Percent Ionization from
Ka or Kb
3. Calculate the concentrations of the species
present in the solution by inserting the value of x into
the expressions in the last line of the table (final
concentration)
4. Calculate the pH = –log[H3O+]
5. Use the concentrations to calculate the fraction of
the original acid that is ionized (the concentration of
the acid that is ionized divided by the analytical or
initial concentration of the acid times 100%
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Calculating Percent Ionization from
Ka or Kb
• Strong acids and bases ionize essentially completely in
water; the percent ionization is always approximately
100%, regardless of the concentration
• The percent ionization in solutions of weak acids and
bases is small and depends on the analytical
concentration of the weak acid or
base; percent ionization of a weak
acid or a weak base actually
increases as its analytical
concentration decreases and
percent ionization increases as the
magnitude of the ionization
constants Ka and Kb increases
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Determining Keq from Ka and Kb
• The value of the equilibrium constant for the reaction of a
weak acid with a weak base can be calculated from Ka
(or pKa), Kb (or pKb), and Kw
• One can quantitatively determine the direction and
extent of reaction for a weak acid and a weak base by
calculating the value of K for the reaction
• The equilibrium constant for the reaction of a weak acid
with a weak base is the product of the ionization
constants of the acid and the base divided by Kw
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Determining Keq from Ka and Kb
• Calculations
1. Write the dissociation reactions for a weak acid and a weak base and
then sum them:
Acid
Base
HA ⇋ H+ + A–
B + H2O ⇋ HB+ + OH–
Sum HA + B + H2O ⇋ H+ + A– + HB+ + OH–
Ka
Kb
Ksum = KaKb
2. Obtain an equation that includes only the acid-base reaction by simply
adding the equation for the reverse of the autoionization of water
(H+ + OH– ⇋ H2O), for which K = 1/Kw, to the overall reaction and
canceling
HA + B + H2O ⇋ H+ + A– + HB+ + OH–
⇋ H2O
⇋ A– + HB+
Ksum = KaKb
H+ + OH–
1/Kw
HA + B
K = (KaKb)/Kw
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Chemistry: Principles, Patterns,
and Applications, 1e
16.5 Acid-Base Titrations
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16.5 Acid-Base Titrations
• In acid-base titrations, a buret is used to deliver
measured volumes of an acid or base solution of known
titration (the titrant) to a flask that contains a solution of a
base or an acid, respectively, of unknown concentration
(the unknown).
• If the concentration of the titrant is known, then the
concentration of the unknown can be determined.
• Plotting the pH changes that occur during an acid-base
titration against the amount of acid or base added
produces a titration curve; the shape of the curve
provides important information about what is occurring in
solution during the titration.
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Titrations of Strong Acids and Bases
• Before addition of any strong base, the initial [H3O+]
equals the concentration of the strong acid.
• Addition of strong base before the equivalence point, the
point at which the number of moles of base (or acid)
added equals the number of moles of acid (or base)
originally present in the solution, decreases the [H3O+]
because added base neutralizes some of the H3O+
present.
• Addition of strong base at the equivalence point
neutralizes all the acid initially present and pH = 7.00;
the solution contains water and a salt derived from a
strong base and a strong acid.
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Titrations of Strong Acids and Bases
• Addition of a strong base after the equivalence causes an excess of
OH– and produces a rapid increase in pH.
• A pH titration curve shows a sharp increase in pH in the region near
the equivalence point and produces an S-shaped curve; the shape
depends only on the concentration of the acid and base, not on their
identity.
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Titrations of Strong Acids and Bases
• For the titration of a monoprotic strong acid with a monobasic strong
base, the volume of base needed to reach the equivalence point can
be calculated from the following relationship:
moles of base = moles of acid
(volume)b (molarity)b = (volume)a (molarity)a
VbMb = VaMa
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Titrations of Weak Acids and Bases
• The shape of the titration curve for a weak acid or a weak base
depends dramatically on the identity of the acid or base and the
corresponding value of Ka or Kb.
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Titrations of Weak Acids and Bases
• The pH changes much more gradually around the equivalence point
in the titration of a weak acid or a weak base.
• [H+] of a solution of a weak acid (HA) is not equal to the
concentration of the acid but depends on both its pKa and its
concentration.
• Only a fraction of a weak acid dissociates, so [H+] is less than [HA];
therefore, the pH of a solution of a weak acid is higher than the pH
of a solution of a strong acid of the same concentration.
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Titrations of Weak Acids and Bases
• Comparing the titration curve of a strong acid with a
strong base with the titration curve of a weak acid and a
strong base
1. Below the equivalence point, the two curves are very different;
before any base is added, the pH of the weak acid is higher than
the pH of the strong acid
2. pH changes more rapidly during the first part of the titration in
a weak acid and strong base titration
3. Due to the higher starting pH, the pH of the weak acid at the
equivalence point is greater than 7.00, so solution is basic
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Titrations of Weak Acids and Bases
4. Change in pH for the weak acid/strong base titration around
the equivalence point is about half as large as for the strong acid
titration; the magnitude of the change at the equivalence point
depends on the pKa of the acid being titrated
5. Above the equivalence point, the two curves are identical; once
acid has been neutralized, the pH of the solution is controlled
only by the amount of excess of OH– present, regardless of
whether the acid is weak or strong
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Titrations of Weak Acids and Bases
• Calculating the pH of a solution of a
weak acid or base
– If Ka or Kb and the initial concentration of a weak acid or base
are known, one can calculate the pH of a solution of a weak acid
or base by setting up a table of initial concentrations, changes in
concentrations, and final concentrations
– Define x as [H+] due to the dissociation of the acid
– Insert values for final concentrations into the equilibrium
equation and solve for x and then pH (pH = –log[H+])
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Titrations of Weak Acids and Bases
• Calculating the pH during titration of a
weak acid or base
– Solved in two steps: a stoichiometric calculation followed by
an equilibrium calculation
1. Use stoichiometry of the neutralization reaction to calculate the
amounts of acid and conjugate base present in solution after the
neutralization reaction has occurred
2. Use the equilibrium equation K = [H3O+] [A–] / [H2O] [HA] to
determine [H+] of the resulting solution
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Titrations of Weak Acids and Bases
• Identity of the weak acid or base being titrated strongly
affects the shape of the titration curve.
• The shape of titration curves as a function of the pKa or
pKb shows that as the acid or base being titrated
becomes weaker (its pKa or pKb becomes larger), the
pH change around the equivalence point decreases
significantly.
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Titrations of Weak Acids and Bases
• The midpoint of a titration is defined as the point at which exactly
enough acid (or base) has been added to neutralize one-half of the
acid (or base) originally present and occurs halfway to the
equivalence point.
• At the midpoint of the titration of an acid, [HA] = [A–].
• The pH at the midpoint of the titration of a weak acid is equal to the
pKa of the weak acid.
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Titrations of Polyprotic Acids or Bases
• When a strong base is added to a solution of a polyprotic
acid, the neutralization reaction occurs in stages.
1. The most acidic group is titrated first, followed by the next
most acidic, and so forth
2. If the pKa values are separated by at least three pKa units,
then the overall titration curve shows well-resolved “steps”
corresponding to the titration of each proton
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Titrations of Polyprotic Acids or Bases
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Indicators
• Most acid-base titrations are not monitored by recording
the pH as a function of the amount of the strong acid or
base solution used as a titrant
• Instead, an acid-base indicator is used, and they are
compounds that change color at a particular pH and if
carefully selected, undergo a dramatic color change at
the pH corresponding to the equivalence point of the
titration
• Acid-base indicators are typically weak acids or bases
whose changes in color correspond to deprotonation or
protonation of the indicator itself
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Indicators
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Indicators
• The chemistry of indicators are described by the general
equation Hn(aq) ⇋ H+ (aq) + n–(aq), where the protonated
form is designated by Hn and the conjugate base by n–
• The ionization constant for the deprotonation of indicator
Hn is Kin = [H+] [n–] / [Hn]
• The value of pKin determines the pH at which the
indicator changes color
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Indicators
• A good indicator must have the following properties:
1. Color change must be easily detected
2. Color change must be rapid
3. Indicator molecule must not react with the substance being
titrated
4. The indicator should have a pKin that is within one pH unit of
the expected pH at the equivalence point of the titration
• Synthetic indicators have been developed that meet the
above criteria and cover the entire pH range
• An indicator does not change color abruptly at a
particular pH but undergoes a pH titration like any other
acid or base
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Indicators
• Choosing the correct indicator for an acid-base
titration
1. For titrations of strong acids and strong bases (and vice versa),
any indicator with a pKin between 4 and 10 will do
2. For the titration of a weak acid, the pH at the equivalence
point is greater than 7, and an indicator such as phenolphthalein
or thymol blue, with pKin > 7, should be used
3. For the titration of a weak base, where the pH at the
equivalence point is less than 7, an indicator such as methyl red
or bromcresol blue, with pKin < 7, should be used
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Indicators
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Indicators
• Paper or plastic strips that
contain combinations of
indicators estimate the pH of a
solution by simply dipping a
piece of pH paper into it and
comparing the resulting color
with standards printed on the
container
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Chemistry: Principles, Patterns,
and Applications, 1e
16.6 Buffers
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16.6 Buffers
• Buffers are solutions that maintain a relatively constant
pH when an acid or a base is added; they protect or
“buffer” other molecules in solution from the effects of the
added acid or base
• Buffers contain either a weak acid (HA) and its conjugate
base (A–) or a weak base (B) and its conjugate acid (BH+)
• Buffers are critically important for the proper functioning
of biological systems; every biological fluid is buffered to
maintain its physiological pH
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The Common Ion Effect
• The ionization equilibrium of a weak acid (HA) is affected
by the addition of either the conjugate base of the acid
(A–) or a strong acid (a source of H+); LeChâtelier’s
principle is used to predict the effect on the equilibrium
position of the solution
• Common-ion effect—the shift in the position of an
equilibrium on addition of a substance that provides an
ion in common with one of the ions already involved in
the equilibrium; equilibrium is shifted in the direction that
reduces the concentration of the common ion
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The Common Ion Effect
• Buffers are characterized by the following:
1. the pH range over which they can maintain a constant pH—
depends strongly on the chemical properties of the weak acid or
base used to prepare the buffer (on K)
2. their buffer capacity, the amount of strong acid or base that
can be absorbed before the pH changes significantly—depends
solely on the concentration of the species in the buffered solution
(the more concentrated the buffer solution, the greater its buffer
capacity)
3. observed change in the pH of the buffer is inversely
proportional to the concentration of the buffer
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The Common Ion Effect
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Calculating the pH of a Buffer
• The pH of a buffer can be calculated from the
concentrations of the weak acid or the weak base used
to prepare it, the concentration of the conjugate base or
acid, and the pKa or pKb of the weak acid or base
• An alternative method used to calculate the pH of a
buffer solution is based on a rearrangement of the
equilibrium equation for the dissociation of a weak acid
• Ionization reaction is HA⇋H+ + A– and the equilibrium
constant expression is
Ka = [H+] [A–]
[HA]
or
[H+] = Ka[HA]
[A–]
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Calculating the pH of a Buffer
• Taking the logarithm of both sides and multiplying both
sides by –1 gives
–log[H+] = –logKa – log([HA]/[A–]) = – logKa + log([A–]/[HA])
• Replacing the negative logarithms gives
pH = pKa + log([A–]/[HA]) or pH = pka + log([base]/[acid])
Both forms of the Henderson-Hasselbalch equation
• Henderson-Hasselbalch equation is valid for solutions
whose concentrations are at least a hundred times
greater than the value of their Ka’s
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Calculating the pH of a Buffer
• Three special cases where the Henderson-Hasselbalch
equation is interpreted without the need for calculations
1. [base] = [acid]. Under these conditions, [base]/[acid] = 1.
Because log 1 = 0, pH = pKa, regardless of the actual
concentrations of the acid and base (corresponds to the midpoint in
the titration of a weak acid or base)
2. [base]/[acid] = 10. Because log 10 = 1, pH = pKa + 1
3. [base]/[acid] = 100. Because log 100 = 2, pH = pKa+ 2
• Each time the [base]/[acid] ratio is increased by 10, the
pH of the solution increases by one unit; if the
[base]/[acid] ratio is 0.1, then pH = pKa – 1, so each
additional factor-of-10 decrease in the [base]/[acid] ratio
causes the pH to decrease by one pH unit
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Calculating the pH of a Buffer
• The Henderson-Hasselbalch equation can also be used
to calculate the pH of a buffer solution after the addition
of a given amount of strong acid or base
• A buffer that contains equal amounts of the weak acid (or
weak base) and its conjugate base (or acid) in solution is
equally effective at neutralizing either added base or
added acid
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The Relationship between Titrations
and Buffers
• A strong correlation exists between the effectiveness of a
buffer solution and the titration curves
• In a titration of a weak acid with a strong base;
– the region around pKa corresponds to the midpoint of the titration,
when half the weak acid has been neutralized; this portion of the
titration curve corresponds to a buffer because it exhibits the smallest
change in pH per increment of added strong base (horizontal nature of
the curve in this region);
– the flat portion of the curve extends only from a pH value of one unit
less than the pKa to a pH value of one unit greater than the pKa ; that is
why buffer solutions have a pH that is within ±1 pH units of the pKa of
the acid component of the buffer;
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The Relationship between Titrations
and Buffers
– in the region of the titration
curve at the lower left, before
the midpoint, the acid-base
properties of the solution are
dominated by the equilibrium
for dissociation of the weak
acid, corresponding to Ka;
– in the region of the titration
curve at the upper right, after
the midpoint, the acid-base
properties of the solution are
dominated by the equilibrium
for reaction of the conjugate
base of the weak acid with
water, corresponding to Kb.
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Blood: A Most Important Buffer
• Metabolic processes produce large amounts of acids and bases, but
organisms are able to maintain a constant internal pH because their
fluids contain buffers.
• pH is not uniform throughout all cells and tissues of a mammal; even
within a cell, different compartments can have very different pH
values.
• Because no single buffer system can effectively maintain a constant
pH value over the physiological range of 5 to 7.4, biochemical
systems use a set of buffers with overlapping ranges; most
important of these is the CO2/HCO3– system, which dominates the
buffering action of blood plasma.
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17
Solubility and
Complexation
Equilibria
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CHAPTER OBJECTIVES
• To be able to calculate the solubility of an ionic
compound from its Ksp
• To understand the factors that determine the solubility of
ionic compounds
• To be able to describe complex-ion formation
quantitatively
• To understand why the solubility of many compounds
depends on pH
• To know how to separate metal ions by selective
precipitation
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Chemistry: Principles, Patterns,
and Applications, 1e
17.1 Determining the
Solubility of Ionic
Compounds
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The Solubility Product, Ksp
• When a slightly soluble ionic compound is added to water, some of it
dissolves to form a solution, establishing an equilibrium between the
pure solid and a solution of its ions.
• The equilibrium constant for the dissolution of a sparingly soluble
salt is the solubility product of the salt, Ksp.
• The concentration of a pure solid is a constant and does not appear
in the equilibrium constant expression.
• Solubility products are determined experimentally by either directly
measuring the concentration of one of the component ions or by
measuring the solubility of the compound in a given amount of water.
• Ksp is defined in terms of molar concentrations of the component
ions.
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The Ion Product
• The ion product (Q) of a salt is the product of the
concentrations of the ions in solution raised to the same
powers as in the solubility product expression.
• The ion product describes concentrations that are not
necessarily equilibrium concentrations, whereas Ksp
describes equilibrium concentrations.
• The process of calculating the value of the ion product
and comparing it with the magnitude of the solubility
product is a way to determine if a solution is unsaturated,
saturated, or supersaturated and whether a precipitate
will form when solutions of two soluble salts are mixed.
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The Ion Product
• Three possible conditions for an aqueous
solution of an ionic solid
1. Q < Ksp: the solution is unsaturated, and more of the ionic
solid will dissolve
2. Q = Ksp: the solution is saturated and at equilibrium
3. Q > Ksp: the solution is supersaturated, and ionic solid will
precipitate
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The Common Ion Effect and Solubility
• Solubility product expression
– Equilibrium concentrations of cation and anion are inversely
related
– As the concentration of the anion increases, the maximum
concentration of the cation needed for precipitation to occur
decreases, and vice versa
– Ksp is constant
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The Common Ion Effect and Solubility
• Common ion effect
– The solubility of an ionic compound depends on the
concentrations of other salts that contain the same ions.
– This dependency is an example of the common ion effect;
adding a common cation or anion shifts a solubility equilibrium
in the direction predicted by LeChâtelier’s principle.
– The solubility of any sparingly soluble salt is almost always
decreased by the presence of a soluble salt that contains a
common ion.
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Chemistry: Principles, Patterns,
and Applications, 1e
17.2 Factors That Affect
Solubility
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17.2 Factors That Affect Solubility
• The solubility product of an ionic compound describes
the concentrations of ions in equilibrium with a solid.
• There are four reasons that the solubility of a compound
may be other than expected:
1. Ion-pair formation
2. Incomplete dissociation of molecular solutes
3. Formation of complex ions
4. Changes in pH
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Ion-Pair Formation
•
An ion pair consists of a cation and anion that are in
intimate contact in solution, rather than separated by
solvent.
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Ion-Pair Formation
• Ions in an ion pair are held together by the same
attractive electrostatic forces as for ionic solids.
– Ions in an ion pair migrate as a single unit, whose net
charge is the sum of the charges on the ions.
– The ion pair is a species intermediate between the ionic
solid (in which each ion participates in many cation-anion
interactions that hold the ions in a rigid array) and the
completely dissociated ions in solution (where each is fully
surrounded by water molecules and free to migrate
independently).
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Ion-Pair Formation
• Ion pairs
– A second equilibrium must be included to describe the
solubility of salts that form ion pairs.
– An ion pair is represented by the symbols of the individual ions
separated by a dot, to indicate that they are associated in
solution (Ca2+ ·SO42–).
– The formation of an ion pair is a dynamic process, so a
particular ion pair may exist only briefly before dissociating into
the free ions, each of which may later associate briefly with
other ions.
– Ion-pair formation has a major effect on the measured solubility
of a salt and is most important for salts that contain M2+ and
M3+ ions; it is unimportant for salts that contain monocations.
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Incomplete Dissociation
• A molecular solute may be more soluble than
predicted by the measured concentrations of
ions in solution due to incomplete dissociation.
– Common for weak organic acids
– Weak acids do not dissociate completely into their
constituent ions (H+ and A–) in water
– The molecular (undissociated) form of a weak acid
(HA) is quite soluble in water
– Many carboxylic acids have a limited solubility in
water
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Incomplete Dissociation
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Chemistry: Principles, Patterns,
and Applications, 1e
17.3 Complex-Ion
Formation
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17.3 Complex-Ion Formation
• Metal ions in aqueous solution are hydrated—surrounded by a shell
of four to six water molecules.
• A hydrated ion is one kind of complex ion, a species formed
between a central metal ion and one or more surrounding ligands,
molecules or ions that contain at least one lone pair of electrons.
• A complex ion forms from a metal ion and a ligand because of a
Lewis acid-base interaction.
– The positively charged metal ion acts as a Lewis acid, and the ligand,
with one or more lone pairs of electrons, acts as a Lewis base.
– Small, highly charged metal ions have the greatest tendency to
act as Lewis acids and to form complex ions.
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The Formation Constant
• The equilibrium constant for the formation of the complex
ion from the hydrated ion is called the formation
constant (Kf).
• Equilibrium constant expression for Kf has the same
general form as any other equilibrium constant
expression.
• The larger the value of Kf, the more stable the product.
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The Effect of Complex-Ion Formation
on Solubility
• The solubility of a sparingly soluble salt increases if a
ligand that forms a stable complex ion is added to the
solution.
• The formation of a complex ion by the addition of a
complexing agent increases the solubility of a
compound.
• Complexing agents are molecules or ions that increase
the solubility of metal salts by forming soluble metal
complexes.
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Chemistry: Principles, Patterns,
and Applications, 1e
17.4 Solubility and pH
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17.4 Solubility and pH
• Solubility of many compounds depend strongly
on the pH of the solution
– The anion in many sparingly soluble salts is the conjugate
base of a weak acid that may become protonated in solution.
– The solubility of simple binary compounds such as oxides and
sulfides are dependent on pH.
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The Effect of Acid-Base Equilibria on the
Solubility of Salts
• Examining the effect of pH on the solubility of a
representative salt, M+A–, where A– is the
conjugate base of the weak acid HA
– When the salt dissolves in water, this reaction occurs:
M A(s) ⇋ M+(aq) + A–(aq)
Ksp = [M+] [A–]
– The anion can also react with water in a hydrolysis reaction:
A–(aq) + H2O(l)⇋OH–(aq) + HA(aq)
– If a strong acid is added to the solution, the added H+ will react
essentially completely with A– to form HA, which decreases
[A–] and which in turn decreases the magnitude of the ion
product, Q = [M+] [A–]
– More MA will dissolve until Q = Ksp
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The Effect of Acid-Base Equilibria on the
Solubility of Salts
– An acidic pH dramatically increases the solubility of sparingly
soluble salts whose anion is the conjugate base of a weak acid;
pH has little to no effect on the solubility of salts whose anion is
the conjugate base of a strong acid
– caves and their associated pinnacles and spires of stone
provide one of the most impressive examples of pH-dependent
solubility equilibria
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Acidic, Basic, and Amphoteric Oxides
and Hydroxides
• Oxides and hydroxides can be classified as
either basic or acidic.
1. Basic oxides and hydroxides either react with water to
produce a basic solution or dissolve readily in aqueous acid.
2. Acidic oxides or hydroxides either react with water to
produce an acidic solution or are soluble in aqueous base.
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Acidic, Basic, and Amphoteric Oxides
and Hydroxides
• There is a clear correlation between the acidic or basic character of
an oxide and the position of the element combined with oxygen in
the periodic table.
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Acidic, Basic, and Amphoteric Oxides
and Hydroxides
– Oxides of metallic elements are generally basic oxides.
– Oxides of nonmetallic elements are acidic oxides.
– There is a gradual transition from basic metal oxides to acidic
nonmetal oxides from the lower left to the upper right of the
periodic table; a broad diagonal band of oxides of intermediate
character separates the two extremes.
– Oxides of the elements in the diagonal region are soluble in both
acidic and basic solutions and are called amphoteric oxides;
which dissolve in acid to produce water or dissolve in base to
produce a soluble complex.
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Acidic, Basic, and Amphoteric Oxides
and Hydroxides
• The difference in reactivity is due to the difference in
bonding in the two kinds of oxides.
– Metals at the far left of the periodic table have low
electronegativities; their oxides contain discrete Mn+ cations
and O2– anions.
– Nonmetal oxides have high electronegativities and form oxides
that contain covalent bonds to oxygen.
– Because of the high electronegativity of oxygen, the covalent
bond between oxygen and the other atom is polarized E+–O– .
– These oxides act as Lewis acids that react with water to
produce an oxoacid.
– Oxides of metals in high oxidation states tend to be acidic
oxides; they contain covalent bonds to oxygen.
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Selective Precipitation Using pH
• Many dissolved metal ions can be separated by
selective precipitation of the cations from solution under
specific conditions.
• pH is used to control the concentration of the anion in
solution; this in turn controls which cations precipitate.
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Chemistry: Principles, Patterns,
and Applications, 1e
17.5 Qualitative Analysis
Using Selective
Precipitation
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17.5 Qualitative Analysis Using
Selective Precipitation
• The composition of relatively complex mixtures of metal
ions can be determined using qualitative analysis, a
procedure for discovering the identity of metal ions
present in the mixture.
• The technique consists of selectively precipitating only
a few kinds of metal ions at a time under given sets of
conditions; consecutive precipitation steps become
progressively less selective until almost all the metal ions
are precipitated.
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17.5 Qualitative Analysis Using
Selective Precipitation
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17.5 Qualitative Analysis Using
Selective Precipitation
• The traditional scheme of analysis for
metal cations involves the separation of
cations into five groups by selective
precipitation
1. Group 1: Insoluble chlorides
– Metal chloride salts are soluble in water; only Ag+, Pb2+,
and Hg22+ form chlorides that precipitate from water
– First step in qualitative analysis is to add 6 M HCl,
causing AgCl, PbCl2, and/or Hg2Cl2 to precipitate;
precipitate is collected by filtration or centrifugation
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17.5 Qualitative Analysis Using
Selective Precipitation
2. Group 2: Acid-insoluble sulfides
– Acidic solution is saturated with H2S
– Only those metal ions that form very insoluble
sulfides precipitate as their sulfide salts under
these acidic conditions; all others remain in
solution
– Precipitates collected by filtration or
centrifugation
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17.5 Qualitative Analysis Using
Selective Precipitation
3. Group 3: Base-insoluble sulfides (and
hydroxides)
– Ammonia or NaOH is added to the solution
until it is basic, and then (NH4)2 S is added
– Treatment removes any remaining cations that
form insoluble hydroxides or sulfides
4. Group 4: Insoluble carbonates or
phosphates
– The next metal ions to be removed are those
that form insoluble carbonates and phosphates
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17.5 Qualitative Analysis Using
Selective Precipitation
5. Group 5: Alkali metals
– All the metal ions that form water-insoluble
chlorides, sulfides, carbonates, or phosphates
have been removed
– Only common ions that remain are the alkali
metals and ammonium
– Take a second sample from the original solution
and add a small amount of NaOH to neutralize the
ammonium ion and produce NH3; any ammonia
produced can be detected by odor or by litmus
paper test, and alkali metal ions, which produce
characteristic colors in flame tests, allows them to
be identified
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17.5 Qualitative Analysis Using
Selective Precipitation
• Metal ions that precipitate together are separated by
various techniques, such as forming complex ions,
changing the pH of the solution, or increasing the
temperature to redissolve some of the solids
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18
Chemical
Thermodynamics
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CHAPTER OBJECTIVES
• To understand the connections among work, heat, and energy
• To be familiar with the concept of PV work
• To be able to calculate changes in internal energy
• To understand the relationship between internal energy and entropy
• To be able to use a thermodynamic cycle to calculate changes in
entropy
• To understand the relationship between Gibbs free energy and work
• To know the difference between the information that
thermodynamics and kinetics provide about a system
• To understand the importance of thermodynamics in biochemical
systems
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Chemical Thermodynamics
• Chemical reactions obey two fundamental laws:
1. The law of conservation of mass
– States that matter can be neither created nor destroyed
– Explains why equations must balance and is the basis for
stoichiometry and equilibrium calculations
2. The law of conservation of energy
– States that energy can be neither created nor destroyed
– Energy takes various forms that can be converted from one to
the other
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Chemical Thermodynamics
• Thermodynamics
– The study of the interrelationships among heat, work, and the
energy content of a system at equilibrium
– Tells whether a particular reaction is energetically possible in
the direction in which it is written and the composition of the
reaction system at equilibrium
– Does not say anything about whether an energetically feasible
reaction will actually occur as written
– Tells nothing about the rate of the reaction or the pathway by
which it will occur
– Provides a bridge between the macroscopic properties of a
substance and the individual properties of its constituent
molecules and atoms
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Chemistry: Principles, Patterns,
and Applications, 1e
18.1 Thermodynamics
and Work
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18.1 Thermodynamics and Work
• A system is that part of the universe in which we are interested; the
surroundings are everything else—the rest of the universe.
• System + surroundings = universe.
• A closed system cannot exchange matter with its surroundings; an
open system can.
• State function—the property of a system that depends only on the
present state of the system and not on its history.
• A change in state function depends only on the difference between
the initial and final states, not on the pathway used to go from one to
the other.
• Thermodynamics is concerned with state functions and does not
deal with how the change between the initial and final state occurs.
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The Connections among Work, Heat,
and Energy
• The internal energy (E) of a system is the sum of the
potential energy and the kinetic energy of all the
components; internal energy is a state function.
• A closed system cannot exchange matter with its
surroundings, but it can exchange energy with its
surroundings in two ways:
1. By doing work
2. By releasing or absorbing heat, the flow of thermal energy
• Work and heat are two distinct ways of changing the
internal energy of a system.
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The Connections among Work, Heat,
and Energy
• Work (w) is defined as a force (F) acting through a distance (d): w =
Fd.
• Work occurs only when an object moves against an opposing force;
work requires that the system and its surroundings be physically
connected.
• The flow of heat, the transfer of energy due to differences in
temperature between two objects, represents a thermal connection
between the system and its surroundings.
• Work causes a physical displacement, and flow of heat causes a
temperature change.
• Units of work and heat must be the same because both processes
result in the transfer of energy; units are joules (J).
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PV Work
• In chemistry, most work is expansion work (PV work)
done as the result of a volume change during a reaction
when air molecules are pushed aside.
• Amount of work done by an expanding gas is given by
the equation w = – PV, where P is the pressure against
which the system must push and V is the change in
volume of the system.
• When the pressure or volume of a gas is
changed, any mechanical work done is
PV work.
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PV Work
• Work done by the system on the surroundings has a
negative value, and work done on the system by the surroundings
has a positive value.
• Work is not a state function; it depends on the path taken.
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Chemistry: Principles, Patterns,
and Applications, 1e
18.2 The First Law of
Thermodynamics
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18.2 The First Law of Thermodynamics
• The Relationship between the energy change of the system and that of
the surroundings is given by the first law of thermodynamics, which
states that the energy of the universe is constant.
• This law can be expressed mathematically as
Euniv = Esys + Esurr = 0
Esys = – Esurr
The change in energy of the system is identical in magnitude but
opposite in sign to the change in energy of the surroundings.
• One of the most important factors that determine the outcome of a
chemical reaction is the tendency of all systems to move toward the
lowest possible energy state.
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18.2 The First Law of Thermodynamics
• Heat and work are the only two ways in which energy can be
transferred between a system and the surroundings.
• Any change in the internal energy of the system is the sum of the
heat transferred, q, and the work done, w:
Esys = q + w
• q and w are not state functions on their own; their sum, Esys, is
independent of the path taken and is therefore a state function.
• Any machine that converts energy to work is designed to want to
maximize the amount of work obtained and to minimize the amount
of energy released to the environment as heat
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Enthalpy
• To understand the relationship between heat flow, q, and
the resulting change in internal energy E, one must
look at two sets of limiting conditions:
1. Reactions that occur at constant volume
2. Reactions that occur at constant pressure
• Assume that PV work is the only kind of work possible
for the system, so E = q – PV
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Enthalpy
• Constant volume
– If reaction occurs in a closed vessel, the volume of the system is
fixed and V is zero
– Heat flow (qv) must equal E
qv = E
– No PV work can be done, and the change in the internal energy
of the system is equal to the amount of heat transferred from the
system to the surroundings, or vice versa
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Enthalpy
• Constant pressure
– If reaction occurs in an open container at a constant pressure of
1 atm, heat flow is given the symbol qp
– Replacing q with qp gives the equation qp = E + PV
– At constant pressure, the heat flow for any process is equal to
the change in the internal energy of the system plus the PV work
done
– A new state function called enthalpy (H) is defined as
H = E + PV
– At constant pressure, the change in the enthalpy of a system is
H = E + (PV) = E + PV
– At constant pressure, the change in the enthalpy of a system is
equal to the heat flow: H = qp
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The Relationship between H and E
• If H for a reaction is known, one can use the change in the enthalpy
of the system to calculate its change in internal energy.
• When a reaction involves only solids, liquids, liquid solutions, or any
combination of these, the volume does not change (V = 0), so H =
E.
• If gases are involved, H and E can differ significantly; one can
calculate E from the measured value of H by using the equation
H = E + PV and the ideal gas law, PV = nRT.
• (PV) = (nRT), so H = E + (PV) = E + (nRT).
• At constant temperature, (nRT) = RTn, where n is the difference
between the final and initial moles of gas, so E = H – RTn.
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The Relationship between H and E
• For reactions that result in a net production of gas, n >
0 and E < H.
• Endothermic reactions that result in a net consumption
of gas have n < 0 and E > H.
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Chemistry: Principles, Patterns,
and Applications, 1e
18.3 The Second Law of
Thermodynamics
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18.3 The Second Law of
Thermodynamics
• How to predict whether a particular process or
reaction would occur spontaneously
– Most spontaneous reactions are exothermic, but there are
many that are not exothermic.
– Reactions can be both spontaneous and highly endothermic.
– Enthalpy changes are not the only factors that determine
whether a process is spontaneous.
– An additional state function called entropy (S) can help explain
why some processes proceed spontaneously in only one
direction.
– Entropy is a thermodynamic property of all substances that is
proportional to their degree of disorder.
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Entropy
• Chemical and physical changes in a system are accompanied by
either an increase or decrease in the disorder of the system,
corresponding to an increase in entropy (S > 0) or a decrease in
entropy (S < 0), respectively.
• A change in entropy is defined as the difference between the
entropies of the final and initial states: S = Sf – Si.
• When a gas expands into a vacuum, its entropy increases because
the increased volume allows for greater atomic or molecular disorder;
the greater the number of atoms or molecules in the gas, the greater
the disorder.
• The magnitude of the entropy of a system depends on the number
of microscopic states, or microstates, associated with it; the greater
the number of microstates, the greater the entropy.
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Entropy
• A disordered system has a greater number of possible
microstates than does an ordered system, so it has a
higher entropy.
• Liquids that have highly ordered structures due to
hydrogen bonding or other intermolecular interactions
tend to have significantly higher values of Svap.
• The formation of a solution is a process that is
accompanied by entropy changes; formation of a liquid
solution from a crystalline solid and a liquid solvent
results in an increase in the disorder of the system and
in its entropy.
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Reversible and Irreversible Changes
• In a reversible process, every intermediate state between the
extremes is an equilibrium state, regardless of the direction of the
change.
• An irreversible process is one in which the intermediate states are
not equilibrium states, and change occurs spontaneously in only one
direction.
• A reversible process can change direction at any time, whereas an
irreversible process cannot.
• Work done during the expansion of a gas depends on the opposing
external pressure (w = PextV), so work done in a reversible process
is always equal to or greater than work done in a corresponding
irreversible process: wrev  wirrev.
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Reversible and Irreversible Changes
• Whether a process is reversible or irreversible,
E = q + w.
• E is a state function, so the magnitude of E does not
depend on reversibility and is independent of the path
taken: E = qrev + wrev = qirrev + wirrev.
• E for a process is the same whether that process is
carried out in a reversible or an irreversible manner.
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The Relationship between Internal
Energy and Entropy
• The quantity of heat transferred, qrev, is directly proportional to the
absolute temperature of an object, T (qrev  T), so the hotter the
object, the greater amount of heat transferred.
• Adding heat to a system increases the kinetic energy of the
component atoms and molecules and their disorder (S  qrev).
• For any reversible process S = qrev/T or qrev = TS;
units of S are joules/kelvin (J/K).
• Work done in a reversible process at constant
pressure is wrev = –PV, so E = qrev + wrev
= TS – PV; change in the internal energy of
the system is related to the change in entropy,
the absolute temperature, and the PV work done.
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The Relationship between Internal
Energy and Entropy
• Entropy of the universe is unchanged in reversible
processes and constitutes part of the second law of
thermodynamics: the entropy of the universe remains
constant in a reversible process, whereas the entropy of
the universe increases in an irreversible (spontaneous)
process.
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Chemistry: Principles, Patterns,
and Applications, 1e
18.4 Entropy Changes
and the Third Law of
Thermodynamics
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
18.4 Entropy Changes and the Third Law
of Thermodynamics
•
The atoms, molecules, or ions that make up a chemical system can
undergo several types of molecular motion: translation, rotation, and
vibration.
•
The greater the molecular motion of a system, the greater the number of
possible microstates and the higher the entropy.
•
A perfectly ordered system, a perfect crystal at a temperature of absolute
zero (0 K) that exhibits no motion, with only a single microstate available
would have an entropy of zero.
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
18.4 Entropy Changes and the Third Law
of Thermodynamics
• Absolute zero is an ideal temperature that is
unobtainable, and a perfect single crystal is an ideal that
cannot be achieved, however, the combination of these
two ideals constitutes the basis for the third law of
thermodynamics: The entropy of any perfectly ordered,
crystalline substance at absolute zero is zero.
• The third law of thermodynamics has two important
consequences:
1.
2.
It defines as positive the sign of the entropy of any
substance at temperatures above absolute zero
It provides a fixed reference point that allows the
measurement of the absolute entropy of any substance at
any temperature
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18.4 Entropy Changes and the Third Law
of Thermodynamics
• Two different ways to calculate S for a reaction
or physical change
1. Uses tabulated values of absolute entropies of substances,
based on the definition of absolute entropy provided by the third
law
2. Uses a thermodynamic cycle, based on the fact that entropy is a
state function
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Calculating S from Standard Molar
Entropy Values
• One way of calculating S for a reaction is to use tabulated values of the
standard molar entropy (Sº), which is the entropy of 1 mol of a substance
at a standard temperature of 298 K
• Units of Sº are J/(molK)
• It is possible to obtain absolute entropy values by measuring the entropy
change that occurs between the reference point of 0 K, corresponding to S
= 0, and 298 K
• For substances with the same molar mass and number of atoms, Sº values
fall in the order Sº(gas) > Sº(liquid) > Sº(solid)
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Calculating S from Standard Molar
Entropy Values
• Substances with similar molecular structures
have similar Sº values
1.
2.
Those with the lowest entropies tend to be rigid crystals
composed of small atoms linked by strong, highly directional
bonds
Those with higher entropies are soft crystalline substances
that contain larger atoms and increased molecular motion and
disorder
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Calculating S from Standard Molar
Entropy Values
• Absolute entropy of a substance tends to increase with
increasing molecular complexity because the number of
available microstates increases with molecular
complexity
• Substances with strong hydrogen bonds have lower
values of Sº, reflecting a more ordered structure
• To calculate Sº for a chemical reaction from standard
molar entropies, the “products minus reactants” rule is
used; here the absolute entropy of each reactant and
product is multiplied by its stoichiometric coefficient in
the balanced chemical equation
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Calculating S from Thermodynamic
Cycles
• A change in entropy can also be calculated
using a thermodynamic cycle
– The molar heat capacity Cp is the amount of heat needed to raise
the temperature of 1 mol of a substance by 1ºC at constant
pressure
– Cv is the amount of heat needed to raise the temperature of 1
mol of a substance by 1ºC at constant volume
– Increase in entropy with increasing temperature is proportional to
the heat capacity of the substance
– Entropy change S is related to heat flow, qrev, by S = qrev/T
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Calculating S from Thermodynamic
Cycles
– qrev = nCpT at constant pressure or nCvT at constant volume,
where n is the number of moles of substance present
– The change in entropy for a substance whose temperature
changes from T1 to T2 is S =nCplnT2/T1 (constant pressure) or
S = nCvlnT2/T1 (constant volume)
– A combination of heat capacity easurements and experimentally
measured values of enthalpies of fusion or vaporization can be
used to calculate the entropy change corresponding to a change
in the temperature of the sample
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Calculating S from Thermodynamic
Cycles
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Chemistry: Principles, Patterns,
and Applications, 1e
18.5 Free Energy
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18.5 Free Energy
• One major goal of chemical thermodynamics is to establish criteria
for predicting whether a particular reaction or process will occur
spontaneously
• The sign of Suniv is a universally applicable and infallible indicator
of the spontaneity of a reaction; if Suniv > 0, the process will occur
spontaneously as written; if Suniv < 0, a process cannot occur
spontaneously; and if Suniv = 0, the system is at equilibrium
• Using Suniv requires the calculation of S for both the system and
the surroundings. This is not useful because we are much more
interested in the system than in the surroundings; it is also difficult to
make quantitative measurements of the surroundings
• A criterion of spontaneity that is based solely on state functions of
the system is more convenient and is provided by a new state
function called the Gibbs free energy
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Gibbs Free Energy and the Direction of
Spontaneous Reactions
•
The Gibbs free energy (G), or free energy, is defined in terms of
three other state functions—enthalpy, temperature, and entropy—
and is a state function itself:
G = H – TS
•
The criterion for predicting spontaneity is based on G, the change
in G at constant temperature and pressure
G = H – TS,
where all thermodynamic quantities are those of the system
•
At constant pressure, H = q whether a process is reversible or
irreversible, and TS = qrev, so G = q – qrev
•
So G is the difference between the heat released during a process
(via a reversible or an irreversible path) and the heat released for
the same process occurring in a reversible manner
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Gibbs Free Energy and the Direction of
Spontaneous Reactions
• To understand how the sign of G for the system
determines the direction in which change is spontaneous,
the equation S = qrev/T and q = H is rewritten to give
Ssurr = – Hsys/T
• The entropy change of the surroundings is related to the
enthalpy change of the system, and since for a
spontaneous reaction, Suniv > 0, the equation becomes
Suniv = Ssys + Ssurr > 0 or Suniv = Ssys – Hsys/T > 0
• Multiplying both sides of the inequality by –T reverses its
sign and rearranges the equation to Hsys– TS < 0,
which is equal to G; therefore, for a spontaneous
process, G < 0
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Gibbs Free Energy and the Direction of
Spontaneous Reactions
• Relationship between the entropy change of the surroundings and
the heat gained or lost by the system provides the key connection
between the thermodynamic properties of the system and the
change in entropy of the universe
• This relationship allows one to predict spontaneity by focusing
exclusively on the thermodynamic properties and temperature of the
system. Highly exothermic processes (H << 0) that increase the
disorder of the system (Ssys >> 0) would occur spontaneously
• For a system at constant temperature and pressure,
1. if G < 0, the process occurs spontaneously
2. if G = 0, the system is at equilibrium
3. if G > 0, the process is not spontaneous as written but occurs
spontaneously in the reverse direction
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Gibbs Free Energy and the Direction of
Spontaneous Reactions
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The Relationship between G and Work
• Change in free energy, G, is equal to the maximum
amount of work that a system can perform on the
surroundings while undergoing a spontaneous change
(at constant temperature and pressure): G = wmax
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Standard Free-Energy Change
• Absolute free energies cannot be measured. However,
changes in free energy can.
• The standard free-energy change (Gº) is the change
in free energy when one substance or set of substances
in their standard states is converted to one or more other
substances also in their standard states
• The standard free-energy change can be calculated from
the definition of free energy if the standard enthalpy and
entropy changes are known: Gº = Hº – TSº
• If Sº and Hº for a reaction have the same sign, then the
sign of Gº depends on the relative magnitudes of the
Hº and TSº terms
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Standard Free-Energy Change
• The standard free energy of formation (Gºf) of a
compound is the change in free energy that occurs when
1 mol of a substance in its standard state is formed from
the elements in their standard states
• Standard free energy of formation of an element in its
standard state is zero at 298.15 K
• Standard free energy of formation of a compound can be
calculated from the standard enthalpy of formation, Hºf,
and from the standard entropy of formation, Sºf, using
the definition of free energy: Gºf = Hºf – TSºf
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Standard Free-Energy Change
• Using standard free energies of formation, the standard
free energy of a reaction can be calculated by employing
the “products minus reactants” rule:
Gºrxn = mGºf (products) – nGºf (reactants),
where m and n are the stoichiometric coefficients of each
product and reactant in the balanced chemical equation
• The effect of temperature on the spontaneity of a
reaction depends on the sign and magnitude of both Hº
and Sº; the temperature at which a given reaction is at
equilibrium can be calculated by setting Gº = 0
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Chemistry: Principles, Patterns,
and Applications, 1e
18.6 Spontaneity and
Equilibrium
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18.6 Spontaneity and Equilibrium
•
Three criteria have been identified for whether a given
reaction will occur spontaneously
1.
2.
3.
Suniv > 0
Gsys < 0
The relative magnitude of the reaction quotient Q versus the
equilibrium constant K
a. Q < K, reaction proceeds spontaneously to the right as written,
resulting in the net conversion of reactants to products
b. Q > K, reaction proceeds spontaneously to the left as written,
resulting in the net conversion of products to reactants
c. Q = K, system at equilibrium, no net reaction occurs
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Free Energy and the Equilibrium
Constant
• Because H º and S º determine the magnitude of Gº
and because the equilibrium constant K is a measure of
the ratio of the concentrations of products to the
concentrations of reactants, K can be expressed in
terms of Gº and vice versa
• For a reversible process that does not involve external
work, the change in free energy can be expressed in
terms of volume, pressure, entropy, and temperature,
eliminating H from the equation for G: G = VP –
ST and G = VP at constant temperature (T = 0)
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Free Energy and the Equilibrium
Constant
• Assuming ideal gas behavior in the equation G = VP, V can be
replaced by nRT/P (where n is the number of moles of gas and R is
the ideal gas constant). G can be expressed in terms of the initial
and final pressures (P1 and P2, respectively).
G = (nRT/P)P = nRT(P/P) = nRTln(P2 /P1)
• If the initial state is the standard state with P1 = 1 atm, then the
change in free energy upon going from the standard state to any
other state with a pressure P can be written as G = Gº + nRTlnP
• Using the hypothetical reaction aA + bB ⇋ cC + dD, in which all
reactants and products are ideal gases and the lowercase letters
correspond to the stoichiometric coefficients for the various species,
the expression for G can be written as G = Gproducts – Greactants
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Free Energy and the Equilibrium
Constant
• Substituting G = Gº + nRTlnP into the equation:
G = [(cGºC + cRTlnPC) + (dGºD + dRTlnPD)] – [(aGºA + aRTlnPA) + (bGºB + bRTlnPB)]
• Combining terms gives the following relationship
between G and the reaction quotient Q
G = Gº + RTln(PcCPdD/PaAPbB) = Gº + RTlnQ
where Gº indicates that all reactants and products are in
their standard states
• For gases Q = Kp at equilibrium, and for a system G = 0
at equilibrium, so the relationship between Gº and Kp for
gases can be described as Gº = –RTlnKp
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Free Energy and the Equilibrium
Constant
• If the products and reactants are in their standard states and Gº < 0,
then Kp > 1 and products are favored over reactants; if Gº > 0, then
Kp < 1 and reactants are favored over products; if Gº = 0, then Kp =
1 and neither reactants nor products are favored and the system is
at equilibrium
• Kp is defined in terms of the partial pressures of reactants and
products, and the equilibrium constant K is defined in terms of the
concentrations of reactants and products; therefore, Kp = K(RT)n,
where n is the number of moles of gaseous product minus the
number of moles of gaseous reactant
• For reactions where n = 0, Kp = K, so Gº = –RT ln K
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Free Energy and the Equilibrium
Constant
• The following equation provides insight into how
the components of Gº influence the magnitude
of the equilibrium constant: Gº = Hº – TSº = –
RT ln K
– K becomes larger as Sº becomes more positive,
indicating that the magnitude of the equilibrium
constant is directly influenced by the tendency of the
system to move toward maximum disorder
– K increases as Hº decreases, so the magnitude of
the equilibrium constant is also directly influenced by
the tendency of the system to seek the lowest energy
state possible
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Temperature Dependence of the
Equilibrium Constant
• The fact that Gº and K are related explains why equilibrium
constants are temperature-dependent
ln K = – Hº/RT + Sº/R
• Assuming Hº and Sº are temperature-dependent, for an exothermic
reaction (Hº < 0), the magnitude of K decreases with increasing
temperature and for an endothermic reaction (Hº > 0), the
magnitude of K increases with increasing temperature
• The magnitude of Hº dictates how rapidly K changes as a function
of temperature and the magnitude and sign of Sº affects the
magnitude of K, but not its temperature dependence
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Temperature Dependence of the
Equilibrium Constant
• If the value of K at a given temperature and the value of
Hº for a reaction are known, the value of K can be
estimated at any other temperature, even in the absence
of information on Sº
• If K1 and K2 are the equilibrium constants for a reaction
at temperatures T1 and T2, respectively, then
ln K2 – ln K1 = ln K2/K1 = Hº/R(1/T1 – 1/T2)
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Chemistry: Principles, Patterns,
and Applications, 1e
18.7 Comparing
Thermodynamics and
Kinetics
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18.7 Comparing Thermodynamics
and Kinetics
• Thermodynamics
– Deals with state functions and can be used to describe the
overall properties, behavior, and equilibrium composition of a
system
– Provides a significant constraint on what can occur during a
reaction process
• Kinetics
– Concerned with the particular pathway by which physical or
chemical changes occur, so it can address the rate at which a
particular process will occur
– Describes the detailed steps of what actually occurs on an
atomic or molecular level
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18.7 Comparing Thermodynamics
and Kinetics
• The following table gives the numerical values of the
equilibrium constant K that correspond to various values
of Gº
– If Gº  + 10 kJ/mol or Gº  –10 kJ/mol, an equilibrium is ensured to lie
all the way to the left or to the right, respectively
– If Gº is quite small (10 kJ/mol), significant amounts of both products
and reactants are present at equilibrium
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18.7 Comparing Thermodynamics
and Kinetics
• Most reactions have equilibrium constants greater than 1, with the
equilibrium strongly favoring either products or reactants
• In many cases, reactions that are strongly favored by
thermodynamics do not occur at a measurable rate, and reactions
that are not thermodynamically favored do occur under certain
nonstandard conditions
• A reaction that is not thermodynamically spontaneous under
standard conditions can be made to occur spontaneously by varying
reaction conditions, by using a different reaction to obtain the same
product, by supplying external energy, or by coupling the
unfavorable reaction to another reaction for which
Gº<< 0
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Chemistry: Principles, Patterns,
and Applications, 1e
18.8 Thermodynamics
and Life
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18.8 Thermodynamics and Life
• A living cell can be viewed as a low-entropy
system that is not in equilibrium with its
surroundings and is capable of replicating itself
• A constant input of energy is needed to maintain
the cell’s highly organized structure and its
intricate system of chemical reactions
• A cell needs energy to synthesize complex
molecules from simple precursors, to create and
maintain differences in the concentrations of
various substances inside and outside of the cell,
and to do mechanical work
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Energy Flow between the Cell and
Its Surroundings
• A cell is an open system that can exchange matter with its
surroundings as well as absorb energy from its environment in the
form of heat or light
• Cells utilize the energy obtained to maintain the nonequilibrium state
that is essential for life
• Nonequilibrium thermodynamics have
been developed to quantitatively describe
open systems such as living cells
• The only way a cell can maintain a
low-entropy, nonequilibrium state
characterized by a high degree of
structural organization is to increase the
entropy of the surroundings; a cell
releases some of the energy it obtains
from its environment as heat that is
transferred to its surroundings, resulting in
an increase in Ssurr
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Extracting Energy from the Environment
• Organisms can be divided into two categories
1. Phototrophs, whose energy source is light
2. Chemotrophs, whose energy source is chemical compounds,
obtained by consuming or breaking down other organisms
• All organisms utilize oxidation-reduction, or redox,
reactions to drive the synthesis of complex biomolecules
– Organisms that can use only O2 as the oxidant are aerobic
organisms that can’t survive in the absence of O2
– Many organisms that use other oxidants or oxidized organic
compounds can live only in the absence of O2 and are called
anaerobic organisms
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Extracting Energy from the Environment
• Fundamental reaction by which all green plants and
algae (phototrophs) obtain energy from sunlight is
photosynthesis, the photochemical reduction of CO2 to
a reduced carbon compound
• One of the main processes chemotrophs use to obtain
energy is respiration, which is the reverse of
photosynthesis
• Some chemotrophs obtain energy by fermentation, in
which both the oxidant and reductant are organic
compounds
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The Role of NADH and ATP in
Metabolism
• Regardless of the identity of the substances from which an organism
obtains energy, the energy must be released in very small
increments if it is to be useful to the cell
• Cells store part of the energy that is released as ATP (adenosine
triphosphate)
• Most organisms use a number of intermediate species to shuttle
electrons between the terminal reductant and the terminal oxidant;
intermediate species oxidizes the energy-rich reduced compound,
and the now-reduced intermediate migrates to another site, where it
is oxidized
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The Role of NADH and ATP in
Metabolism
• The most important of these electron-carrying intermediates is NAD+,
whose reduced form is NADH
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The Role of NADH and ATP in
Metabolism
• Energy from the oxidation of nutrients is made available to cells
through the synthesis of ATP; its energy is used by the cell to
synthesize substances through coupled reactions and to perform
work
• For biochemical reactions, a new standard state has been defined in
which the H+ concentration is 1 x 10–7 M (pH 7) and all other
reactants and products are present in standard-state conditions (1 M
or 1 atm)
• The free-energy change and equilibrium constant for a reaction
under these new standard conditions are denoted by the addition of
a prime sign (´) to the conventional symbol
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Energy Storage in Cells
• All organisms use ATP as the immediate free-energy
source in biochemical reactions, but ATP is not an
efficient form in which to store energy on a long-term
basis
• The body stores energy as sugars, proteins, and fats
before using it to produce ATP
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19
Electrochemistry
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CHAPTER OBJECTIVES
• To distinguish between galvanic and electrolytic cells
• To predict spontaneous reactions using redox potentials
• To balance redox reactions using half-reactions
• To understand the relationship between cell potential and the
equilibrium constant
• To be able to measure solution concentrations using cell potentials
• To describe how commercial galvanic cells operate
• To describe the process of corrosion
• To understand electrolysis and to be able to describe it quantitatively
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Electrochemistry
• In oxidation-reduction (redox) reactions, electrons are
transferred from one species (the reductant) to another
(the oxidant).
• Transfer of electrons provides a means for converting
chemical energy to electrical energy, or vice versa.
• The study of the relationship between electricity and
chemical reactions is called electrochemistry.
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Chemistry: Principles, Patterns,
and Applications, 1e
19.1 Describing
Electrochemical Cells
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19.1 Describing Electrochemical Cells
• Electrochemical process—electrons flow from one
chemical substance to another, driven by an
oxidation-reduction (redox) reaction
• Redox reaction
– Occurs when electrons are transferred from a substance that is
oxidized to one that is being reduced
– Reductant is the substance that loses electrons and is
oxidized in the process
– Oxidant is the species that gains electrons and is reduced in
the process
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19.1 Describing Electrochemical Cells
– Described as two half-reactions, one representing the
oxidation process and one the reduction process
Reductive half-reaction: Br2 (aq) + 2e–  2Br –(aq)
Oxidative half-reaction: Zn (s)  Zn2+(aq) + 2e–
– Adding the two half-reactions gives the overall chemical
reaction
Zn (s) + Br2 (aq)  ZnBr2 (aq)
– A redox reaction is balanced when the number of electrons lost
by the reductant is equal to the number of electrons gained by
the oxidant; overall process is electrically neutral
– An electric current is produced from the flow of electrons from
reductant to oxidant
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19.1 Describing Electrochemical Cells
• Electrochemical cell
– An apparatus that is used to generate electricity
from a spontaneous redox or that uses electricity to
drive a nonspontaneous redox reaction
– There are two types of electrochemical cells
1. Galvanic cell (voltaic cell)—uses the energy released
during a spontaneous redox reaction (G < 0) to generate
electricity
2. Electrolytic cell—consumes electrical energy from an
external source, using it to cause a nonspontaneous redox
reaction to occur (G > 0)
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19.1 Describing Electrochemical Cells
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19.1 Describing Electrochemical Cells
– Both types of cells contain two electrodes, which are solid
metals connected to an external circuit that provides an
electrical connection between systems
– Oxidation half-reaction occurs at one electrode, the anode,
and the reduction half-reaction occurs at the other, the
cathode
– When circuit is closed, electrons flow from the anode to the
cathode; electrodes are connected by an electrolyte, which is
an ionic substance or solution that allows ions to transfer
between the electrodes, thereby maintaining the system’s
electrical neutrality
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Galvanic (Voltaic) Cells
• To illustrate the basic principles of a galvanic cell
– look at the reaction of metallic zinc with cupric ion (Cu2+) to
give copper metal and Zn2+ ion
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
– a copper strip is inserted into a beaker that contains a 1 M
solution of Cu2+ ions, and a zinc strip is inserted into a
different beaker that contains a 1 M solution of Zn2+ ions
– two metal strips serve as electrodes and are connected by a
wire that allows electricity to flow; the compartments are
connected by a salt bridge, a U-shaped tube inserted into
both solutions and contains a concentrated liquid or gelled
electrolyte to maintain electrical neutrality
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Galvanic (Voltaic) Cells
– when the circuit is closed, a spontaneous reaction occurs:
zinc metal is oxidized to Zn2+ ions at the zinc electrode (the
anode), and Cu2+ ions are reduced to Cu metal at the
copper electrode (the cathode)
– as reaction progresses, the zinc strip dissolves and the
concentration of Zn2+ ions in the Zn2+ solution increases; the
copper strip gains mass and the concentration of Cu2+ ions in
the Cu2+ solution decreases
– electrons that are released at the anode flow through the wire,
producing an electric current; galvanic cells transform
chemical energy into electrical energy that can be used to do
work
– the electrolyte in the salt bridge serves two purposes: to
complete the circuit by carrying electrical charge and to
maintain electrical neutrality in both solutions by allowing
ions to migrate between the two solutions
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Galvanic (Voltaic) Cells
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Galvanic (Voltaic) Cells
– a voltmeter is used to measure the difference in electrical
potential between the two compartments
– the potential of the cell, measured in volts, is the difference in
electrical potential between the two half-reactions; electrical
potential is related to the energy needed to move a charged
particle in an electric field
– electrons from the oxidative half-reaction are released at the
anode, so the anode in a galvanic cell is negatively charged; the
cathode, which attracts electrons, is positively charged
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Constructing a Cell Diagram
• Because galvanic cells are cumbersome to describe in
words, a line notation called a cell diagram has been
developed
• In a cell diagram
– the identity of the electrodes and the chemical contents of the
compartments are indicated by their chemical formulas, with the
anode written on the far left and the cathode on the far right;
– phase boundaries are shown by single vertical lines;
– the salt bridge, which has two phase boundaries, is shown by a
double vertical line;
Here’s a cell diagram for Zn/Cu cell:
Zn(s)| Zn2+(aq, 1M) || Cu2+(aq, 1 M) | Cu(s)
Anode
Salt Bridge
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Cathode
Constructing a Cell Diagram
• In a single-compartment galvanic cell, the voltage
produced by a redox reaction can be measured more
accurately using two electrodes immersed in a single
beaker containing an electrolyte that completes the
circuit; arrangement reduces errors caused by resistance
to the flow of charge at a boundary, called the junction
potential
– Cell diagram does not include a double vertical line for the salt
bridge (no salt bridge) and does not include solution
concentrations
Pt(s) | H2(g) | HCl(aq) | AgCl(s) | Ag(s)
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Chemistry: Principles, Patterns,
and Applications, 1e
19.2 Standard Potentials
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19.2 Standard Potentials
• In a galvanic cell, current is produced when electrons
flow externally through the circuit from the anode to the
cathode because of a difference in potential energy
between two electrodes in the electrochemical cell
• The flow of electrons in an
electrochemical cell depends on the
identity of the reacting substances,
the difference in the potential
energy of their valence electrons,
the concentrations of the reacting
species, and the temperature of the
system
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
19.2 Standard Potentials
• The potential of the cell under standard conditions (1 M
for solutions and 1 atm for gases, pure solids, or liquids
for other substances) and at a fixed temperature (25ºC)
is called the standard cell potential, Eºcell
– Used in order to develop a scale of relative potentials that will
allow the prediction of the direction of an electrochemical
reaction and the magnitude of the driving force for the reaction
– Used to measure the potentials for oxidations and reductions of
different substances under comparable conditions
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Measuring Standard Electrode Potentials
• Impossible to measure the potential of a single electrode;
only the difference between the potentials of two
electrodes can be measured
• Can compare the standard cell potentials for two
different galvanic cells that have one kind of electrode in
common—allows the measurement of the potential
difference between two dissimilar electrodes
• All tabulated values of standard electrode potentials by
convention are listed as standard reduction potentials
in order to compare standard potentials for different
substances
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Measuring Standard Electrode Potentials
• The standard cell potential is the reduction potential of the reductive
half-reaction minus the reduction potential of the oxidative
half-reaction (Eºcell = Eºcathode – Eºanode).
• The potential of the standard hydrogen electrode (SHE) is defined
as 0 V under standard conditions.
• The potential of a half-reaction measured against the SHE under
standard conditions is called the standard electrode potential for
that half-reaction.
• The standard cell potential is a measure of the driving force for a
given redox reaction.
• All Eº values are independent of the stoichiometric coefficients for the
half-reactions.
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Balancing Redox Reactions Using the
Half-Reaction Method
• Redox reactions can be balanced using the half-reaction
method, where the overall redox reaction is divided into
an oxidation half-reaction and a reduction half-reaction,
each one balanced for mass and charge
• The half-reactions selected from tabulated lists must
exactly reflect reaction conditions
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Balancing Redox Reactions Using the
Half-Reaction Method
• In an alternative method, the atoms in each half-reaction
are balanced, and then the charges are balanced; one
does not need to use the tabulated half-reactions;
instead, focus on the atoms whose oxidation states
change by using the following steps:
1. Write the reduction half-reaction and the oxidation
half-reaction
2. Balance the atoms by balancing elements other than O and
H; then balance O atoms by adding H2O, and balance H
atoms by adding H+
3. Balance the charges in each half-reaction by adding
electrons
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Balancing Redox Reactions Using the
Half-Reaction Method
4.
5.
6.
Multiply the reductive and oxidative half-reactions by
appropriate integers to obtain the same number of electrons in
both half-reactions
Add the two half-reactions and cancel substances that appear
on both sides of the equation
Check to make sure that all atoms and charges are balanced
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Calculating Standard Cell Potentials
• The standard cell potential for a redox reaction, Eºcell, is a
measure of the tendency of the reactants in their standard
states to form the products in their standard states—it is a
measure of the driving force for the reaction (voltage)
• Calculations for the standard potential for the Zn/Cu cell
represented by the cell diagram:
Zn(s) | Zn2+(aq, 1 M || Cu2+(aq, 1M) | Cu(s)
Cathode: Cu2+(aq) + 2e–  Cu(s)
Eºcathode = 0.34 V
Anode: Zn(s)  Zn2+(aq, 1M) + 2e– Eºanode = –0.76 V
Overall: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Eºcell = Eºcathode – Eºanode = 1.10 V
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Calculating Standard Cell Potentials
• If the value of Eºcell (the standard cell potential) is positive,
the reaction will occur spontaneously as written
• If the value of Eºcell is negative, then the reaction is not
spontaneous and it will not occur as written under
standard conditions; it will proceed spontaneously in the
opposite direction
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Reference Electrodes and Measuring
Concentrations
• When using a galvanic cell to measure the concentration
of a substance, we are interested in the potential of only
one of the electrodes of the cell, the indicator electrode,
whose potential is related to the concentration of the
substance being measured
• To ensure that any change in the measured potential of
the cell is due to only the substance being analyzed, the
potential of the other electrode, the reference electrode,
must be constant
• Whether oxidation or reduction occurs depends on the
potential of the sample versus the potential of the
reference electrode
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Reference Electrodes and Measuring
Concentrations
• Types of reference electrodes
1. Standard hydrogen electrode (SHE)—consists of a strip of
platinum wire in contact with an aqueous solution containing 1
M H+, which is in equilibrium with H2 gas at a pressure of 1 atm at
the Pt-solution interface
2. Silver-silver chloride electrode—consists of a silver wire
coated with a thin layer of AgCl that is dipped into a chloride ion
solution with a fixed concentration
3. Saturated calomel electrode (SCE)—consists of a platinum
wire inserted into a moist paste of liquid mercury, Hg2Cl2, and
KCl
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Reference Electrodes and Measuring
Concentrations
• One of the most common uses of
electrochemistry is to measure the H+ ion
concentration of a solution
– A glass electrode is used for this purpose, in which an internal
Ag/AgCl electrode is immersed in a 0.10 M HCl solution that is
separated from the solution by a very thin glass membrane that
absorbs protons
– The extent of the adsorption on the inner side is fixed but the
adsorption of protons on the outer surface depends on the pH of
the solution
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Reference Electrodes and Measuring
Concentrations
• Ion-selective electrodes
– Used to measure the concentration of a particular species in
solution; designed so that their potential depends on only the
concentration of the desired species
– Contains an internal reference electrode that is connected by a
solution of an electrolyte to a crystalline inorganic material or
membrane, which acts as the sensor
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Chemistry: Principles, Patterns,
and Applications, 1e
19.3 Comparing
Strengths of Oxidants
and Reductants
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19.3 Comparing Strengths of Oxidants
and Reductants
• The following table lists the
standard potentials for a wide
variety of chemical substances that
allow us to compare the oxidative
and reductive strengths of a variety
of substances
• The half-reaction for the standard
hydrogen electrode lies halfway
down on the table
– All species that lie above it in the table
are stronger oxidants than H+, and all
those that lie below it are weaker
– All species that lie below H2 are
stronger reductants than H2, and those
that lie above H2 are weaker
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19.3 Comparing Strengths of Oxidants
and Reductants
• Because the half-reactions in the table are arranged in
order of their Eº values, the table can be used to predict
the relative strengths of various oxidants and reductants
– Any species on the left side of a half-reaction will spontaneously
oxidize any species on the right side of another half-reaction that
lies below it in the table
– Any species on the right side of one half-reaction will
spontaneously reduce any species on the left side of another
half-reaction that lies above it in the table
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Chemistry: Principles, Patterns,
and Applications, 1e
19.4 Electrochemical
Cells and
Thermodynamics
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The Relationship between Cell Potential
and Free Energy
• Electrochemical cells convert chemical to electrical energy, and vice
versa
• Total amount of energy produced by an electrochemical cell and the
amount of energy available to do electrical work depends on both
the cell potential and the total number of electrons that are
transferred from the reductant to the oxidant during the course of the
reaction
• Resulting electric current measured in coulombs (C), an S unit that
measures the number of electrons passing a given point in 1 s;
coulomb defined as 6.25 x 1018 e–/s and relates electrical potential
(in volts) to energy (in joules)
1J/1V = 1 coulomb = 6.25 x 1018 e–/s
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
The Relationship between Cell Potential
and Free Energy
• Electric current is measured in amperes (A); 1 ampere is defined as
the flow of 1 coulomb per second past a given point (1C = 1A/s)
• In chemical reactions one must relate the coulomb to the charge on
a mole of electrons; multiplying the charge on the electron by
Avogadro’s number gives the charge on 1 mol of electrons, the
faraday (F)
F = (1.60218 x 10–19C) (6.02214 x 1023/1 mol e–) =
9.64855 x 104C/mol e– = 96,485.5 J(V·mol e–)
• Total charge transferred from the reductant to the oxidant is nF,
where n is the number of moles of electrons
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The Relationship between Cell Potential
and Free Energy
• The maximum amount of work that can be produced by an
electrochemical cell, wmax, is equal to the product of the cell potential,
Ecell, and the total charge transferred during the reaction, nF: wmax
= –nFEcell
• Work is expressed as a negative number because work is being
done by the system on the surroundings
• G is also a measure of the maximum amount of work that can be
performed during a chemical reaction (G = wmax; there must be a
relationship between the potential of an electrochemical cell and the
change in free energy, G
G = –nFEcell
• A spontaneous redox reaction is characterized by a negative value of
G and a positive value of Ecell
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Potentials for the Sums of Half-Reactions
• When the standard potential for a half-reaction is not available, the
relationships between standard potentials and free energy can be
used to obtain the potential of any other half-reaction that can be
written as the sum of two or more half-reactions whose standard
potentials are available
• Values of Eº for half-reactions cannot be added to give Eº for the sum
of the half-reactions because Eº is not a state function
• Because Gº is a state function, the sum of the G values for the
individual reactions gives Gº for the overall reaction, which is
proportional to both the potential and the number of electrons (n)
transferred
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Potentials for the Sums of Half-Reactions
• To obtain the value of Eº for the overall half-reaction, the
values of Gº (= –nFEº) must be added for each
half-reaction to obtain Gº for the overall half-reaction
• Substitute values into equation Gº = nFEºcell and solve
for Eº
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The Relationship between Cell Potential
and the Equilibrium Constant
• Can use the relationship between Gº and the
equilibrium constant K to obtain a relationship between
Eºcell and K
• For a general reaction of the type aA + bB  cC + dD,
the standard free-energy change and the equilibrium
constant are related by the equation Gº = –RTlnK
• Given the relationship between the standard free-energy
change and the standard cell potential, the following
equation can be written: –nFEºcell = –RTlnK
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The Relationship between Cell Potential
and the Equilibrium Constant
• Rearranging the equation gives
Eºcell = (RT/nF)lnK
• For T = 298 K, the equation is simplified to
Eºcell = (RT/nF)lnK = [8.314 J/(mol·K)] (298 K) 2.303 log K
n[96,486 J/(V·mol)]
= (0.0591 V/n)log K
• The standard cell potential, Eºcell, is directly proportional
to the logarithm of the equilibrium constant
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The Relationship between Cell Potential
and the Equilibrium Constant
• The following figure summarizes the relationships
developed based on properties of the system (based on
the equilibrium constant, standard free-energy change,
and standard cell potential) and the criteria for
spontaneity (Gº < 0)
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The Effect of Concentration on Cell
Potential: The Nernst Equation
• The actual free-energy change for a reaction under nonstandard conditions, G, is given by
G = Gº + RTln Q
We also know that G = –nFEcell and Gº = –nFEºcell, so
substituting these expressions in the preceding equation
gives
–nFEcell = –nFEºcell + RT lnQ
• Dividing both sides of this equation by –nF gives the
Nernst equation: Ecell = Eºcell – (RT/nF)lnQ
• The Nernst equation determines the spontaneous
direction of any redox reaction under any reaction
conditions from values of the relevant standard reduction
potentials
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
The Effect of Concentration on Cell
Potential: The Nernst Equation
• When a redox reaction is at equilibrium (G = 0), the
Nernst equation reduces to
Eºcell = (RT/nF)ln K because Q = K and there is no net
transfer of electrons (Ecell = 0)
• Substituting the values of the constants into the Nernst
equation with T = 298 K and converting to base-10
logarithms gives the relationship of the actual cell
potential (Ecell), the standard cell potential (Eºcell), and the
reactant and product concentrations at room
temperature (contained in Q):
Ecell = Eºcell – (0.0591 V/n)logQ
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The Effect of Concentration on Cell
Potential: The Nernst Equation
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Concentration Cells
• A voltage can be generated by constructing an
electrochemical cell in which each compartment contains
the same redox active solution but at different
concentrations; voltage is produced as the
concentrations equilibrate
• An electrochemical cell in which the anode and cathode
compartments are identical except for the concentration
of a reactant is called a concentration cell
• Because G = 0 at equilibrium, the measured potential
of a concentration cell is zero at equilibrium
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Using Cell Potentials to Measure
Solubility Products
• A galvanic cell can be used to measure the solubility
product of a sparingly soluble substance
• Measure the solubility product of AgCl, Ksp = [ Ag+][Cl–]
– One compartment contains a silver wire inserted into a 1.0 M
solution of Ag+
– The other compartment
contains a silver wire inserted
into a 1.0 M Cl– solution
saturated with AgCl
– The potential due to the
difference in [Ag+] between
the two cells can be used to
determine Ksp
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Using Cell Potentials to Measure
Concentrations
• A galvanic cell can be used to calculate the
concentration of a species given a measured potential
and the concentrations of all the other species by using
the Nernst equation
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Chemistry: Principles, Patterns,
and Applications, 1e
19.5 Commercial Galvanic
Cells
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19.5 Commercial Galvanic Cells
• Galvanic cells can be self-contained and portable and
can be used as batteries and fuel cells
1. A battery (storage cell) is a galvanic cell (or a series of
galvanic cells) that contains all the reactants needed to
produce electricity.
2. A fuel cell is a galvanic cell that requires a constant
external supply of one or more reactants in order to
generate electricity.
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Batteries
• Two basic kinds of batteries
1. Disposable, or primary, batteries in which the electrode
reactions are effectively irreversible and which cannot
be recharged
2. Rechargeable, or secondary, batteries, which form an
insoluble product that adheres to the electrodes; can
be recharged by applying an electrical potential in the
reverse direction, which temporarily converts a
rechargeable battery from a galvanic cell to an
electrolytic cell
• Major difference between batteries and galvanic cells
is that commercial batteries use solids or pastes
rather than solutions as reactants to maximize the
electrical output per unit mass
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Batteries
• When a battery consists of more than one galvanic
cell, the cells are connected in series—that is, with
the positive (+) terminal of one cell connected to the
negative (–) terminal of the next, and so on
• The overall voltage of the battery is the sum of the
voltages of the individual cells
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Batteries
• Leclanché dry cell
– A “wet cell” in which the electrolyte is an acidic
water-based paste containing MnO2, NH4Cl, ZnCl2,
graphite, and starch
– Used in flashlights, Walkmen, and GameBoys and is
disposable
– Cell not very efficient in producing electrical energy and
has a limited shelf life
– The alkaline battery is a Leclanché cell adapted to
operate under alkaline, or basic, conditions; has a longer
shelf life and more constant output voltage than the
Leclanché dry cell
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Batteries
• “Button” batteries
– The anode is a zinc-mercury amalgam, and the cathode
can be either HgO or Ag2O as the oxidant
– Are reliable and have a high output-to-mass ratio, which
allows them to be used in calculators, cameras, hearing
aids, and watches, where their small size is crucial
– Disposable
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Batteries
• Lithium-iodine battery
– Water-free battery
– Consists of two cells separated by a metallic nickel mesh
that collects charge from the anodes
– The anode is lithium metal, and the cathode is a solid
complex of 2
– Electrolyte is a layer of solid Li that allows Li+ ions to
diffuse from the cathode to the anode
– Highly reliable and long-lived
– Used in cardiac pacemakers, medical implants, smoke
alarms, and in computers
– Disposable
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Batteries
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Batteries
• Nickel-cadmium (nicad) battery
– Used in small electrical appliances and in devices like
drills and portable vacuum cleaners
– A water-based cell with a cadmium anode and a highly
oxidized nickel cathode
– This design maximizes the surface area of the electrodes
and minimizes the distance between them, which gives
the battery both a high discharge current and a high
capacity
– Lightweight, rechargeable, and high capacity but tend to
lose capacity quickly and do not store well; also presents
disposal problems because of the toxicity of cadmium
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Batteries
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Batteries
• Lead-acid (lead storage) battery
– Provides the starting power in automobiles and boats;
can be discharged and recharged many times
– The anodes in each cell of this rechargeable battery are
plates or grids of lead containing spongy lead metal,
while the cathodes are similar grids containing powdered
lead dioxide, PbO2
– The electrolyte is an aqueous solution of sulfuric acid
– The value of Eº for such a cell is 2 V; connecting three
cells in series produces a 6-V battery, and a typical 12-V
car battery contains six of these cells connected in series
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Batteries
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Fuel Cells
• A galvanic cell that requires an external supply of
reactants because the products of the reaction are
continuously removed
• Does not store electrical energy but allows electrical
energy to be extracted directly from a chemical
reaction
• Have reliability problems and are costly
• Used in space vehicles
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Chemistry: Principles, Patterns,
and Applications, 1e
19.6 Corrosion
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19.6 Corrosion
• Corrosion is a galvanic process by which metals
deteriorate through oxidation, usually but not always
to their oxides
• One of the most common techniques used to prevent
corrosion is to apply a protective coating of another
metal that is more difficult to oxidize
• Alternatively, a more easily oxidized metal can be
applied to a metal surface, thus providing cathodic
protection of the surface; galvanized steel is protected
by a thin layer of zinc
• Sacrificial electrodes can also be attached to an
object to protect it
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Chemistry: Principles, Patterns,
and Applications, 1e
19.7 Electrolysis
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19.7 Electrolysis
• Galvanic cells—a spontaneous chemical reaction is
used to generate electrical energy
• In an electrolytic cell, the opposite process, called
electrolysis, occurs: an external voltage is applied to
drive a nonspontaneous reaction
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Electrolytic Cells
• An electrochemical cell in which one electrode is copper
metal immersed in a 1 M Cu2+ solution and the other
electrode is cadmium metal immersed in a 1 M Cd2+
solution is set up, and then the circuit is closed
– Cadmium electrode begins to dissolve (Cd is oxidized to Cd2+)
and thus is the anode, while metallic copper will be
deposited on the copper electrode (Cu2+ is reduced to Cu),
which is the cathode
– Overall reaction is thermodynamically spontaneous as written
(Gº < 0); in this direction, the system is acting as a galvanic
cell
– The reverse reaction, the reduction of Cd2+ by Cu, is
thermodynamically nonspontaneous and will only occur with
an input of an applied voltage
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Electrolytic Cells
– Can force the reaction to proceed in the reverse direction by
applying an electrical potential from an external power supply;
applied voltage forces electrons through the circuit in the reverse
direction, converting a galvanic cell to an electrolytic cell
– The copper electrode is now the anode (Cu is oxidized) and the
cadmium electrode is now the cathode (Cd2+ is reduced)
– Signs of the electrodes have changed to reflect the flow of
electrons in the circuit
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Electrolytic Cells
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Electrolytic Cells
• Differences between galvanic and
electrolytic cells
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Electrolytic Reactions
• At sufficiently high temperatures, ionic solids melt to form liquids that
conduct electricity extremely well due to the high concentrations of
ions
• Sodium metal is produced commercially by electrolysis of molten
mixture of NaCl and CaCl2 in a Downs cell
• The Hall-Heroult process is used to produce aluminum commercially
by electrolysis of a molten mixture of aluminum oxide and cryolite
• Electrolysis can also be used to drive the thermodynamically
nonspontaneous decomposition of water into its constituent
elements, H2 and O2, by adding a small amount of an ionic solute to
the water to make it electrically conductive
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Electrolytic Reactions
• Overvoltages are needed in all electrolytic processes
• An overvoltage is an added voltage and represents the
additional driving force required to overcome barriers
such as a large activation energy and a junction potential
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Electroplating
• In a process called electroplating, a layer of a second
metal is deposited on the metal electrode that acts as
the cathode during electrolysis
• Electroplating is used to enhance the appearance of
metal objects and protect them from corrosion
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Quantitative Considerations
• The amount of material consumed or produced in a
reaction can be calculated from the stoichiometry of an
electrolysis reaction, the amount of current passed, and
the duration of the electrolytic reaction
Charge (C) = current (A) X times(s)
moles e– =
charge (C)
96,486 C/mol (1 faraday)
• The stoichiometry can be used to determine the
combination of current and time needed to produce a
given amount of material
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20
Nuclear Chemistry
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CHAPTER OBJECTIVES
• To understand the factors that affect nuclear stability
• To know the different kinds of radioactive decay
• To be able to balance a nuclear reaction
• To be able to interpret a radioactive decay series
• To know the differences between ionizing and nonionizing radiation and their
effects on matter
• To be able to identify natural and artificial sources of radiation
• To be able to calculate a mass-energy balance and a nuclear binding energy
• To understand the differences between nuclear fission and fusion
• To understand how nuclear reactors operate
• To understand how nuclear transmutation reactions led to the formation of the
elements in the stars and how they can be used to synthesize transuranium
elements
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Nuclear Reactions
•
Differences between nuclear reactions and
chemical processes
1. In a nuclear reaction, the identities of the elements change
2. Nuclear reactions are accompanied by the release of
enormous amounts of energy
3. The yields and rates of a nuclear reaction are unaffected
by changes in temperature, pressure, or the presence of
catalysts
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Chemistry: Principles, Patterns,
and Applications, 1e
20.1 The Components of
the Nucleus
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20.1 The Components of the Nucleus
• Most of the known elements have at least one isotope
whose atomic nucleus is stable indefinitely
• A great majority of elements also have isotopes that are
unstable and disintegrate, or decay, at measurable rates
by emitting radiation
• Some elements have no stable isotopes and eventually
decay to other elements
• The process of nuclear decay is a nuclear reaction that
results in changes inside the atomic nucleus
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The Atomic Nucleus
• Each element can be represented by the
notation AZ X
– A is the mass number, the sum of the numbers of
protons and neutrons
– Z is the atomic number, the number of protons
– The protons and neutrons that make up the nucleus
of an atom are called nucleons
– An atom with a particular number of protons and
neutrons is called a nuclide
– Nuclides that have the same number of protons but
different numbers of neutrons are called isotopes
– The number of neutrons is equal to A – Z
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The Atomic Nucleus
• Isotopes of oxygen can be represented in any of these
ways:
A
Z X:
16O
8
17O
8
18O
8
AX:
16O
17O
18O
Element-A:
Oxygen-16
Oxygen-17
Oxygen-18
• Isotopes of naturally occurring elements on Earth are
present in nearly fixed proportions with each proportion
constituting an isotope’s natural abundance
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The Atomic Nucleus
• Any nucleus that is unstable and decays spontaneously
is said to be radioactive, emitting subatomic particles
and electromagnetic radiation
• The emissions are collectively called radioactivity and
can be measured
• Isotopes that emit radiation are called radioisotopes
• The rate at which radioactive decay occurs is
characteristic of the isotope and is reported as a half-life
(t½), the amount of time required for half the initial
number of nuclei present to decay in a first-order
reaction
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Nuclear Stability
• The nucleus of an atom occupies a tiny fraction of the volume of the
atom and contains the number of protons and neutrons that is
characteristic of a given isotope
• Electrostatic repulsions would cause the positively charged protons
to repel each other, but the nucleus does not fly apart because of
the strong nuclear force, an extremely powerful but very shortrange attractive force between nucleons
• All stable nuclei except the hydrogen-1 nucleus contain at least one
neutron to overcome the electrostatic repulsion between protons
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Nuclear Stability
• As the number of protons in the nucleus increases, the number of neutrons needed for
a stable nucleus increases even more rapidly; too many protons (or too few neutrons)
in the nucleus result in an imbalance between forces, which leads to nuclear instability
• Relationship between the numbers of protons and neutrons in stable nuclei is shown in
the following figure
– The stable isotopes form a “peninsula of stability” in a “sea of instability”
– Only three stable isotopes, 1H, 3He, and 4He, have a neutron-to-proton ratio less
than or equal to 1; all other stable nuclei have a higher neutron-to-proton ratio,
which increases steadily to about 1.5 for the heaviest nuclei
– All elements with Z > 83 are unstable and radioactive
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Nuclear Stability
• More than half of the stable nuclei have even numbers of
both neutrons and protons; only 6 of the 279 stable
nuclei do not have odd numbers of both
• Certain numbers of neutrons or protons result in
especially stable nuclei; these are the so-called magic
numbers 2, 8, 20, 50, 82, and 126
• Nuclei with magic numbers of both protons and neutrons
are said to be “doubly magic” and are even more stable
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Superheavy Elements
• In addition to the “peninsula of stability,” the preceding
figure shows a small “island of stability” that exists in the
upper right corner
• The island corresponds to the superheavy elements,
with atomic numbers near the magic number of 126, and
may be stable enough to exist in nature
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Chemistry: Principles, Patterns,
and Applications, 1e
20.2 Nuclear Reactions
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20.2 Nuclear Reactions
• Two general kinds of nuclear reactions
1. Nuclear decay reaction (or radioactive decay)
– An unstable nucleus emits radiation and is transformed into the
nucleus of one or more other elements
– Resulting daughter nuclei have a lower mass and are lower in
energy (more stable) than the parent nucleus that decayed
– Occur spontaneously under all conditions
2. Nuclear transmutation reaction
– A nucleus reacts with a subatomic particle or another nucleus to
form a product nucleus that is more massive than the starting
material
– Occur spontaneously only under special conditions
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Classes of Radioactive Nuclei
• Each of the three general classes of radioactive
nuclei is characterized by a different decay
process or set of processes
1. Neutron-rich nuclei
– Have too many neutrons and have a neutron-to-proton ratio that is
too high to give a stable nucleus
– These nuclei decay by a process that converts a neutron to a
proton, thereby decreasing the neutron-to-proton ratio
2. Neutron-poor nuclei
– Have too few neutrons and have a neutron-to-proton ratio that is
too low to give a stable nucleus
– These nuclei decay by processes that convert a proton to a
neutron, thereby increasing the neutron-to-proton ratio
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Classes of Radioactive Nuclei
3. Heavy nuclei
– Heavy nuclei (with A  200) are intrinsically unstable, regardless of
the neutron-to-proton ratio
– All nuclei with Z > 83 are unstable
– Decay by emitting an  particle, which decreases the number of
protons and neutrons in the original nucleus by 2
– Since the neutron-to-proton ratio in an  particle is 1, the net result
of alpha emission is an increase in the neutron-to-proton ratio
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Nuclear Decay Reactions
• Can use the number and type of nucleons present to
write a balanced equation for a nuclear decay reaction
– Procedure allows us to predict the identity of either the parent or
daughter nucleus if the identity of only one is known
– Regardless of the mode of decay, the total number of nucleons
is conserved in all nuclear reactions, as is the total positive
charge
• To describe nuclear decay reactions, the AX notation for
nuclides has been extended to include radioactive
emissions
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Nuclear Decay Reactions
• The following table lists the name and symbol for each
type of emitted radiation
1. The left superscript in the symbol for a particle gives the mass
number, which is the total number of protons and neutrons
– For a proton or a neutron, A = 1
– Because neither an electron nor a positron contains protons or
neutrons, its mass number is 0
2. The left subscript gives the charge of the particle
–
–
–
–
Protons carry a positive charge, so Z = +1 for a proton
A neutron contains no protons and is electrically neutral, so Z = 0
For an electron, Z = –1, and for a positron, Z = +1
Because  rays are high-energy photons, both A and Z are 0
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Nuclear Decay Reactions
3. In some cases, two different symbols are used for particles that are
identical but produced in different ways
0
– Symbol -1e, simplified to e– represents a free electron or an electron
associated with an atom
– Symbol 0, simplified to – denotes an electron that originates from
-1
within the nucleus, which is a  particle
– 42He refers to the nucleus of a helium atom, and 24 is an identical
particle ejected from a heavier nucleus
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Nuclear Decay Reactions
• There are six fundamentally different kinds of nuclear decay
reactions, each of which releases a different kind of particle or
energy (see table)
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Nuclear Decay Reactions
1. Alpha decay
– Nuclei with mass numbers greater than 200 undergo alpha
decay, which results in the emission of a helium-4 nucleus
as an  particle, 24
– The daughter nuclide contains two fewer protons and two
fewer neutrons than the parent, thus -particle emission
produces a daughter nucleus with a mass number A that is
lower by 4 and a nuclear charge Z that is lower by 2 than
the parent nucleus
A
ZX
→
parent
A-4
Z-2X′
+
4

2
daughter αparticle
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Nuclear Decay Reactions
2. Beta decay
– Nuclei that contain too many neutrons undergo beta decay,
in which a neutron is converted to a proton and a highenergy electron that is ejected from the nucleus as a 
particle
1
n
0
→
unstable neutron
in nucleus
1p
1
0
+
proton retained
by nucleus
β
-1
beta particle emitted
by nucleus
– Beta decay does not change the mass number of the
nucleus but results in an increase of +1 in the atomic
number due to the addition of a proton in the daughter
nucleus; beta decay decreases the neutron-to-proton ratio,
moving the nucleus toward the band of stable nuclei
A
X
Z
parent
→
Z+1X′ +
A
daughter
0
β
-1
 particle
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Nuclear Decay Reactions
3. Positron emission
– A positron has the same mass as an electron but opposite
charge
– Positron emission is the opposite of beta decay and is
characteristic of neutron-poor nuclei which decay by the
transformation of a proton to a neutron and a high-energy
positron that is emitted
1

p
1
1
n
0
+
0

+1
– Positron emission does not change the mass number of the
nucleus, however the atomic number of the daughter nucleus is
lower by 1 than that of the parent. The neutron-to-proton ratio
increases, moving nucleus closer to the band of stable nuclei
A
X
Z
parent

A
X′
Z-1
+
daughter
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0

+1
positron
Nuclear Decay Reactions
4. Electron capture
– A neutron-poor nucleus can decay by either positron
emission or electron capture (EC), in which an electron in
an inner shell reacts with a proton to produce a neutron
1
p
1
+
0
e
-1

1
n
0
– When a second electron moves from an outer shell to take
the place of the lower-energy electron that was absorbed by
the nucleus, an X-ray is emitted. The overall reaction for
electron capture is
0
X
+
e 
Z
-1
A
parent
A
X’
Z-1
electron
+ X-ray
daughter
– The mass number does not change, but the atomic number
of the daughter nucleus is lower by 1 than that of the parent;
neutron-to-proton ratio increases, moving the nucleus
toward the band of stable nuclei
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Nuclear Decay Reactions
5. Gamma emission
– Many nuclear decay reactions produce daughter nuclei that are
in a nuclear excited state
– A nucleus in an excited state releases energy in the form of a
photon when it returns to the ground state
– These high-energy photons are  rays
– Gamma emission can occur instantaneously or after a
significant delay
– General formula
A
X*
Z

A
X +
Z

0
0
– Because  rays are energy, their emission does not affect either
the mass number or the atomic number of the daughter nuclide;
gamma-ray emission is the only kind of radiation that does not
involve the conversion of one element to another but is
observed in conjunction with some other nuclear decay reaction
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Nuclear Decay Reactions
6. Spontaneous fission
– Only very massive nuclei with high neutron-to-proton ratios can
undergo spontaneous fission, in which the nucleus breaks into
two pieces that have different atomic numbers and atomic
masses
– Process most important for trans-actinide elements with Z  104
– Spontaneous fission is accompanied by the release of large
amounts of energy and is accompanied by the emission of
several neutrons
– The number of nucleons is conserved; the sum of the mass
numbers of the products equals the mass number of the
reactant; the sum of the atomic numbers of the products is the
same as the atomic number of the parent nuclide
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Radioactive Decay Series
• Impossible for any nuclide with Z > 85 to decay to a stable daughter
nuclide in a single step, except via nuclear fission
• Radioactive isotopes with Z > 85 usually decay to a daughter
nucleus that is radioactive, which in turn decays to a second
radioactive daughter nucleus, and so forth, until a stable nucleus
finally results
• This series of sequential alpha- and beta-decay reactions is called a
radioactive decay series
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Induced Nuclear Reactions
• Some nuclei spontaneously transform into nuclei with a
different number of protons, producing a different
element
• These naturally occurring radioactive isotopes decay by
emitting subatomic particles
• Should be possible to carry out the reverse reaction,
converting a stable nucleus to another more massive
nucleus by bombarding it with subatomic particles in a
nuclear transmutation reaction
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Synthesis of Transuranium Elements
• Uranium (Z = 92) is the heaviest naturally occurring
element; all the elements with Z > 92, the transuranium
elements, are artificial and have been prepared by
bombardment of suitable target nuclei with smaller
particles
• Bombarding the target with more massive nuclei creates
elements that have atomic numbers greater than that of
the target nucleus
• Accelerating positively charged particles to the speeds
needed to overcome the electrostatic repulsions
between them and the target nuclei requires a device
called a particle accelerator, which uses electrical and
magnetic fields to accelerate the particles
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Synthesis of Transuranium Elements
• Types of particle accelerators
1.
2.
The linear accelerator is the simplest particle accelerator in
which a beam of particles is injected at one end of a long
evacuated tube; rapid alternation of the polarity of the
electrodes along the tube causes the particles to be alternately
accelerated toward a region of opposite charge and repelled
by a region with the same charge, resulting in a tremendous
acceleration as the particle travels down the tube.
A cyclotron achieves the same outcome in less space and
forces the charged particles to travel in a circular path;
particles are injected into the center of a ring and accelerated
by rapidly alternating the polarity of two large D-shaped
electrodes above and below the ring, which accelerates the
particles outward along a spiral path toward the target.
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Synthesis of Transuranium Elements
3.
The synchrotron is a hybrid of the previous two designs and
contains an evacuated tube similar to that of the linear
accelerator, but the tube is circular and can be more than a
mile in diameter; charged particles are accelerated around the
circle by a series of magnets whose polarities rapidly alternate.
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Chemistry: Principles, Patterns,
and Applications, 1e
20.3 The Interaction of
Nuclear Radiation with
Matter
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20.3 The Interaction of Nuclear Radiation
with Matter
• Nuclear reactions do not cause chemical reactions
directly.
• The particles and photons emitted during nuclear decay
are very energetic, and they can indirectly produce
chemical changes in the matter surrounding the nucleus
that has decayed.
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Ionizing versus Nonionizing Radiation
• Effects of radiation on matter are determined by the
energy of the radiation, which depends on the nuclear
decay reaction that produced it.
1.
Nonionizing radiation
–
–
–
Low in energy; when it collides with an atom in a molecule or
ion, most of its energy can be absorbed without causing a
structural or chemical change
The kinetic energy of the radiation is transferred to the atom or
molecule with which it collides, causing it to rotate, vibrate, or
move more rapidly
This energy can be transferred to adjacent molecules or ions
in the form of heat, so many radioactive substances are warm
to the touch
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Ionizing versus Nonionizing Radiation
2.
Ionizing radiation
–
–
–
Higher in energy and some of its energy can be transferred to
one or more atoms with which it collides as it passes through
matter
If enough energy is transferred, electrons can be excited to
very high energy levels, resulting in the formation of positively
charged ions
Molecules ionized in this way are highly reactive and can
decompose or undergo other chemical changes that create a
cascade of reactive molecules that can damage biological
tissues and other materials
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The Effects of Ionizing Radiation
on Matter
• The effects of ionizing radiation depend on
four factors
1. The type of radiation, which dictates how far it can
penetrate into matter
2. The energy of the individual particles or photons
3. The number of particles or photons that strike a
given area per unit time
4. The chemical nature of the substance exposed to
the radiation
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The Effects of Ionizing Radiation
on Matter
• The relative abilities of the various forms of
ionizing radiation to penetrate tissues are:
1.  radiation
– Reacts strongly with matter because of its high charge and
mass
– Does not penetrate deeply into an object and can be stopped
by clothing or skin
– Alpha particles most damaging if their source is inside the body
because their energy is absorbed by internal tissues
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The Effects of Ionizing Radiation
on Matter
2.
3.
 radiation
–  rays, with no charge and no mass, do not interact strongly with matter
and penetrate deeply into most objects, including the human body
– Lead or concrete needed to completely stop  rays
– The most dangerous type when they come from a source outside the
body
 radiation
– Intermediate in mass and charge between  particles and  rays, so
interaction with matter is intermediate
– Beta particles penetrate paper or skin but can be stopped by wood or a
thin sheet of metal
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The Effects of Ionizing Radiation
on Matter
• There are many ways to measure radiation
exposure, or the dose
– The roentgen (R) is used to measure the amount of energy
absorbed by dry air and is used to describe exposure
quantitatively
– Damage to biological tissues is proportional to the amount of
energy absorbed by tissues, not air
– The most common unit to measure the effects of radiation on
biological tissue is the rad (radiation absorbed dose); S unit is
the gray (Gy)
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The Effects of Ionizing Radiation
on Matter
– Rad is defined as the amount of radiation that causes 0.01 J of
energy to be absorbed by 1 kg of matter, and the gray is defined
as the amount of radiation that causes 1 J of energy to be
absorbed per kilogram
1 rad = 0.010 J/kg
1 Gy = 1J/kg
– The amount of tissue damage caused by 1 rad of  particles is
much greater than the damage caused by 1 rad of  particles or
 rays because  particles have higher masses and charge
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The Effects of Ionizing Radiation
on Matter
– A unit called the rem (roentgen equivalent in man)
describes the actual amount of tissue damage caused
by a given amount of radiation
– The number of rems of radiation is equal to the
number of rads multiplied by the RBE (relative
biological effectiveness) factor, which is 1 for 
particles,  rays, and X-rays, and 20 for  particles
– Most measurements are reported in millirems
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Natural Sources of Radiation
• We are continuously exposed to measurable
background radiation from a variety of natural
sources, which is equal to about 150–600
mrem/yr
1. Cosmic rays, high-energy particles, and  rays emitted by
the sun and other stars that bombard Earth continuously
2. Cosmogenic radiation, produced by the interaction of cosmic
rays with gases in the upper atmosphere
3. Terrestrial radiation, due to the remnants of radioactive
elements that were present on the primordial Earth and their
decay products
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Natural Sources of Radiation
4. Tissues also absorb radiation (40 mrem/yr) from
naturally occurring radioactive elements present
in our bodies
5. Radon is the most important source of
background radiation
– The heaviest of the noble gases and tends to
accumulate in enclosed spaces
– Radon exposure can cause lung damage or cancer
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Artificial Sources of Radiation
• In addition to naturally occurring background radiation,
humans are exposed to small amounts of radiation from
a variety of artificial sources
1. X-rays used for diagnostic purposes in medicine and dentistry;
X-rays are photons with much lower energy than  rays
2. Television screens and computer monitors with cathode-ray
tubes that produce X-rays
3. Luminescent dials
4. Residual fallout from atmospheric nuclear-weapons testing
5. Nuclear power industry
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Assessing the Impact of Radiation
Exposure
• The radiation exposure from
artificial sources, when
combined with the exposure
from natural sources, poses a
significant risk to human health
• The effects of single radiation
doses of different magnitudes
on humans are listed in the
following table
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Assessing the Impact of Radiation
Exposure
• A large dose of radiation spread over time is less harmful
than the same total amount of radiation administered
over a short time
• Tissues most affected by large, whole-body exposures
are bone marrow, intestinal tissue, hair follicles, and
reproductive organs
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Chemistry: Principles, Patterns,
and Applications, 1e
20.4 Thermodynamic
Stability of the Atomic
Nucleus
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20.4 Thermodynamic Stability of the
Atomic Nucleus
• Nuclear reactions are accompanied by changes in
energy
• Energy changes in nuclear reactions are enormous
compared with those of even the most energetic
chemical reactions
• Energy changes in a typical nuclear reaction are so large
that they result in a measurable change of mass
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Mass-Energy Balance
• Nuclear reactions are accompanied by large changes in energy,
which result in detectable changes in mass.
• The relationship between mass, m, and energy, E, is expressed in
the equation E = mc2, where c is the speed of light (2.998 x 108 m/s),
and energy and mass are expressed in units of joules and kilograms,
respectively.
• Every mass has an associated energy, and any reaction that
involves a change in energy must be accompanied by a change in
mass.
• Large changes in energy in nuclear reactions are reported in units of
keV or MeV; a change in energy that accompanies a nuclear
reaction can be calculated from the change in mass (1 amu = 931
MeV).
• Chemical reactions are accompanied by changes in mass, but these
changes are too small to be detected
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Nuclear Binding Energies
• The mass of an atom is always less than the sum of the
masses of its component particles; the only exception is
hydrogen-1.
• The difference between the sum of the masses of the
components and the measured atomic mass is called the
mass defect of the nucleus.
• The amount of energy released when a nucleus forms
from its component nucleons is the nuclear binding
energy.
• The magnitude of the mass defect is proportional to the
nuclear binding energy, so both values indicate the
stability of the nucleus.
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Nuclear Binding Energies
• Not all nuclei are equally stable; the relative stability of
different nuclei are described by comparing the binding
energy per nucleon, which is obtained by dividing the
nuclear binding energy by the mass number A of the
nucleus.
• The binding energy per nucleon increases rapidly with
increasing atomic number until Z = 26, where it levels off
and then decreases slowly.
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Nuclear Fission and Fusion
• Nuclear fission
– The splitting of a heavy nucleus into two lighter ones
– Nucleus usually divides asymmetrically rather than into equal
parts, and the fission of a given nuclide does not give the same
products every time
– In a typical nuclear fission reaction, more than one neutron is
released by each dividing nucleus; when these neutrons collide
with and induce fission in other neighboring nuclei, a
self-sustaining series of nuclear fission reactions known as a
nuclear chain reaction can result
– Each series of events is called a generation
– The minimum mass capable of supporting sustained fission is
called the critical mass
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Nuclear Fission and Fusion
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Nuclear Fission and Fusion
– If the mass of the fissile isotope is greater than the critical
mass, then under the right conditions, the resulting
supercritical mass can release energy explosively
• Nuclear fusion
– Two light nuclei combine to produce a heavier, more stable
nucleus and is the opposite of a nuclear fission reaction
– The positive charge on both nuclei results in a large
electrostatic energy barrier to fusion; barrier can be
overcome if one or both particles have sufficient kinetic
energy to overcome the electrostatic repulsions, allowing the
two nuclei to approach close enough for a fusion reaction to
occur
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Nuclear Fission and Fusion
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Chemistry: Principles, Patterns,
and Applications, 1e
20.5 Applied Nuclear
Chemistry
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Nuclear Reactors
• When a critical mass of a fissile isotope has been achieved, the
resulting flux of neutrons can lead to a self-sustaining reaction; a
variety of techniques can be used to control the flow of neutrons,
which allows nuclear fission reactions to be maintained at safe
levels
• Many levels of control are required, along with a fail-safe design;
otherwise, the chain reaction can accelerate so rapidly that it
releases enough heat to melt or vaporize the fuel and the container,
causing the release of enough radiation to contaminate the
surrounding area
• If the neutron flow in a reactor is carefully regulated so that only
enough heat is released to boil water, then the resulting steam can
be used to produce electricity
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Nuclear Reactors
• Light-water reactors
– Used to produce electricity
– Fuel rods containing a fissile isotope in a structurally
stabilized form (uranium oxide pellets encased in a corrosionresistant zirconium alloy) are suspended in a cooling bath
that transfers the heat generated by the fission reaction to a
secondary cooling system
– Heat is used to generate steam for the production of
electricity
– Control rods are utilized to absorb neutrons and control the
rate of the nuclear chain reaction
– Pulling the control rods out increases the neutron flow,
allowing the reactor to generate more heat; inserting the rods
completely stops the reaction
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Nuclear Reactors
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Nuclear Reactors
• Heavy-water reactors
– Deuterium (2H) absorbs neutrons less effectively than does
hydrogen (1H), but it is about twice as effective at scattering
neutrons
– A nuclear reactor that uses D2O instead of H2O as the
moderator is so efficient that it can use unenriched uranium
as fuel, which reduces the operating costs and eliminates the
need for plants that produce enriched uranium
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Nuclear Reactors
• Breeder reactors
– A nuclear fission reactor that produces more fissionable fuel
than it consumes; the fuel produced is not the same as the
fuel consumed
– Overall reaction is the conversion of nonfissile 238U to fissile
239Pu, which can be isolated chemically and used to fuel a
new reactor
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Nuclear Reactors
• Nuclear fusion reactors
– Nuclear fusion reactions are thermodynamically spontaneous,
but the positive charge on both nuclei results in a large
electrostatic energy barrier to the reaction; high temperatures
are required to overcome the electrostatic repulsions and
initiate a fusion reaction
– Achieving these temperatures, controlling the materials to be
fused, and extracting the energy released by the fusion
reaction are difficult problems
– These nuclear reactions are called thermonuclear
reactions because a great deal of thermal energy must be
invested to initiate the reaction; amount of heat released by
the reaction is several orders of magnitude greater than the
energy needed to initiate it
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Uses of Radioisotopes
• Radiation is destructive to rapidly dividing cells such as
tumor cells and bacteria, so it has been used medically
to treat cancer; many radioisotopes are available for
medical use, and each has specific advantages for
certain applications
• Radiation therapy
– Radiation is delivered by a source planted inside the body, or in
some cases, physicians take advantage of the body’s own
chemistry to deliver a radioisotope to the desired location
– In cases where a tumor is surgically inaccessible, an external
radiation source is used to aim a tightly focused beam of  rays at it
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Uses of Radioisotopes
• Medical imaging
– A radioisotope is temporarily
localized in a particular tissue or
organ where its emissions provide a
map of the tissue or organ
– Positron emission tomography (PET)
is an imaging technique that
produces remarkably detailed
three-dimensional images;
biological molecules that have been
tagged with a positron-emitting
isotope can be used to probe the
functions of organs
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Uses of Radioisotopes
• Ionizing radiation is used in the irradiation of food to kill
bacteria
• In addition to the medical uses of radioisotopes,
radioisotopes have hundreds of other uses: smoke
alarms, dentistry, detectors, and gauges
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Chemistry: Principles, Patterns,
and Applications, 1e
20.6 The Origin of the
Elements
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Relative Abundances of the Elements on
Earth and in the Universe
• The relative abundances of the elements in the universe
and on Earth relative to silicon are shown in the
following figure
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Relative Abundances of the Elements on
Earth and in the Universe
• Data are estimates based on the characteristic emission
spectra of the elements in stars, the absorption spectra
of matter in clouds of interstellar dust, and the
approximate composition of Earth as measured by
geologists
• Data illustrate two points
1. Except for hydrogen, the most abundant elements have
even atomic numbers
2. The relative abundances of the elements in the universe
and on Earth are very different
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Relative Abundances of the Elements on
Earth and in the Universe
• All the elements originally present on Earth were
synthesized from hydrogen and helium nuclei in the
interiors of the stars that have long since exploded and
disappeared
• Six of the most abundant elements in the universe
(carbon, oxygen, neon, magnesium, silicon, and iron)
have nuclei that are integral multiples of the helium-4
nucleus, which suggests that helium-4 is the primary
building block for heavier nuclei
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Synthesis of the Elements in Stars
• Elements are synthesized in discrete stages
during the lifetime of a star, and some steps
occur only in the most massive stars known
– All stars are formed by the aggregation of interstellar dust,
which is mostly hydrogen
– As the cloud of dust slowly contracts due to gravitational
attraction, its density reaches 100g/cm3 and the temperature
increases to 1.5 x 107 K, forming a dense plasma of ionized
hydrogen nuclei
– Self-sustaining nuclear reactions begin and the star ignites,
creating a yellow star
– In the first stages of life, the star is powered by a series of
nuclear fusion reactions that convert hydrogen to helium
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Synthesis of the Elements in Stars
– Overall reaction is the conversion of four hydrogen nuclei to a
helium-4 nucleus, accompanied by the release of two positrons,
two  rays, and a great deal of energy
– When large amounts of helium-4 have been formed, they
become concentrated in the core of the star, which slowly
becomes denser and hotter
– At a temperature of 2 x 108 K, the helium-4 nuclei begin to fuse,
producing beryllium-8, which is unstable and decomposes in 10–
16 s, long enough for it to react with a third helium-4 nucleus to
form the stable carbon-12
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Synthesis of the Elements in Stars
– Sequential reactions of carbon-12 with helium-4 produce the
elements with even numbers of protons and neutrons up to
magnesium-24
– So much energy is released by these reactions that it causes the
surrounding mass of hydrogen to expand, producing a red giant
that is 100 times larger than the original yellow star
– As the star expands, the heavier nuclei accumulate in its core,
which contracts to a density of 50,000 g/cm3 and becomes hotter
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Synthesis of the Elements in Stars
– At a temperature of 7 x 108 K, carbon and oxygen nuclei undergo
nuclear fusion reactions to produce sodium and silicon nuclei
– At these temperatures, carbon-12 reacts with helium-4 to initiate
a series of reactions that produce more oxygen-16, neon-20,
magnesium-24, and silicon-28, as well as heavier nuclides such
as sulfur-32, argon-36, and calcium-40
– Energy released by these reactions causes a further expansion
of the star to form a red supergiant
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Synthesis of the Elements in Stars
– Core temperature increases steadily, at a temperature of 3 x 109
K, the nuclei that have been formed exchange protons and
neutrons freely
– This equilibration process forms heavier elements up to iron-56
and nickel-58, which have the most stable nuclei known
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Formation of Heavier Elements
in Supernovas
• All naturally occurring elements heavier than
nickel are formed in the rare but spectacular
cataclysmic explosions called supernovas
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Formation of Heavier Elements
in Supernovas
– Fuel in the core of a massive star is consumed, so its gravity causes
it to collapse in about 1 s
– As the core is compressed, the iron and nickel nuclei within it
disintegrate to protons and neutrons, and many of the protons
capture electrons to form neutrons
– The resulting neutron star is so dense that atoms no longer exist
– The energy released by the collapse of the core causes the
supernova to explode in a violent event
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Formation of Heavier Elements
in Supernovas
– Force of the explosion blows the star’s matter into space,
creating a gigantic and rapidly expanding dust cloud called a
nebula
– The concentration of neutrons is so great that multiple
neutron-capture events occur, leading to the production of
the heaviest elements and many of the less-stable nuclides
– Force of the explosion distributes these elements throughout
the galaxy surrounding the supernova and are eventually
captured in the dust that condenses to form new stars
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21
Periodic Trends
and the s-Block
Elements
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CHAPTER OBJECTIVES
• To be able to describe the physical and chemical
properties of hydrogen and to predict its reactivity
• To be able to describe how the alkali and alkaline earth
metals are isolated
• To become familiar with the reactions, compounds, and
complexes of the alkali and alkaline earth metals
• To know some of the uses of the alkali and alkaline earth
metals
• To become familiar with the role of the s-block elements
in biology
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Chemistry: Principles, Patterns,
and Applications, 1e
21.1 Overview of
Periodic Trends
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
21.1 Overview of Periodic Trends
• The single most important unifying principle in
understanding the chemistry of the elements is the
systematic increase in atomic number, accompanied by
the orderly filling of atomic orbitals by electrons, which
leads to periodicity in such properties as atomic and
ionic size, ionization energy, electronegativity, and
electron affinity
• Same factors lead to periodicity in valence-electron
configurations, which for each group results in
similarities in oxidation states and the formation of
compounds with common stoichiometries
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21.1 Overview of Periodic Trends
• The most important periodic trends in atomic
properties are summarized in the following table:
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21.1 Overview of Periodic Trends
–
–
These trends are based on periodic variations in a
single fundamental property, the effective nuclear
charge (Zeff), which increases from left to right and
from bottom to top in the periodic table
The diagonal line separates the metals (to the left of
the line) from the nonmetals (to the right)
1. Metals have low electronegativities; they tend to lose
electrons in chemical reactions to form compounds in
which they have positive oxidation states
2. Nonmetals have high electronegativities; they tend to
gain electrons in chemical reactions to form
compounds in which they have negative oxidation
states
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21.1 Overview of Periodic Trends
– The semimetals lie along the diagonal line dividing metals and
nonmetals; they exhibit properties and reactivities intermediate
between those of metals and nonmetals
– Elements of Groups 13, 14, and 15 span the diagonal line
separating metals and nonmetals, and their chemistry is more
complex than predicted based solely on their valence-electron
configuration
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Unique Chemistry of the Lightest Elements
• The chemistry of the second-period element of each
group (n = 2; Li, Be, B, C, N, O, and F) differs in many
important aspects from that of the heavier members, or
congeners, of the group
• The elements of the third period (n = 3; Na, Mg, Al, Si, P,
S, and Cl) are generally more representative of the
group to which they belong
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Unique Chemistry of the Lightest Elements
• The anomalous chemistry of second-period
elements results from three important
characteristics
1. Small radii
– Due to their small radii, second-period elements have
lower electron affinities than would be predicted from
general periodic trends
– When an electron is added to such a small atom,
increased electron-electron repulsions substantially
destabilize the anion
– The small sizes of these elements prevent them from
forming compounds in which they have more than four
nearest neighbors
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Unique Chemistry of the Lightest Elements
– Simple binary ionic compounds of second-period
elements have more covalent character than the
corresponding compounds formed from their heavier
congeners
– The very small cations derived from second-period
elements have a high charge-to-radius ratio and can
polarize the filled valence shell of an anion
– Bonding in these compounds has a significant covalent
component, giving the compounds properties that can
differ significantly from those expected for simple ionic
compounds
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Unique Chemistry of the Lightest Elements
2. Energetically unavailable d orbitals
– Because d orbitals are never occupied for principal
quantum numbers less than 3, the valence electrons of
second-period elements occupy 2s and 2p orbitals only
– Electron configurations with more than four electron pairs
around a central, second-period element are not
observed
3. Tendency to form  bonds with other atoms
– One of the most dramatic differences between the
lightest main group elements and their heavier
congeners is the tendency of the second-period elements
to form species that contain multiple bonds
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Unique Chemistry of the Lightest Elements
• Another important trend to note in main-group chemistry
is the chemical similarity between the lightest element of
one group and the element immediately below and to the
right of it in the next group, a phenomenon known as the
diagonal effect
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The Inert-Pair Effect
• The inert-pair effect refers to the empirical observation
that the heavier elements of Groups 13—17 have
oxidation states that are lower by 2 than the maximum
predicted for their group
• Two major reasons for the inert-pair effect
1. Increasing ionization energies
– Filled (n –1)d or (n –2)f subshells are poor at shielding
electrons in ns orbitals
– The two electrons in the ns subshell experience a high
effective nuclear charge Zeff and are strongly attracted to
the nucleus, reducing their participation in bonding
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The Inert-Pair Effect
2. Decreasing bond strengths
– Going down a group, the atoms become larger and the
overlap between the valence orbitals of the bonded
atoms decreases
– Bond strengths tend to decrease down a column
• The net effect of these two factors is that as you
go down a group in the p block, the additional
energy released by forming two additional bonds
eventually is not great enough to compensate
for the additional energy required to remove the
two ns2 electrons
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Chemistry: Principles, Patterns,
and Applications, 1e
21.2 The Chemistry of
Hydrogen
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Isotopes of Hydrogen
• Hydrogen
– The most abundant element in the universe
– The ultimate source of all the other elements by the process of
nuclear fusion
• There are three isotopes of hydrogen, all of which
contain one proton and one electron per atom
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Isotopes of Hydrogen
1. Protium (1H or H)—most common isotope
2. Deuterium (2H or D)—has an additional neutron
3. Tritium (3H or T)—the rarest isotope of hydrogen
– Produced in the upper atmosphere by a nuclear reaction
when cosmic rays strike nitrogen and other atoms; it is
then washed into the oceans by rainfall
– Radioactive, decaying to 3He with a half-life of 12.3 years
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Isotopes of Hydrogen
• Different masses of the three isotopes of hydrogen
cause them to have different physical properties; H2, D2,
and T2 differ in their melting points, boiling points,
densities, and heats of fusion and vaporization
• Deuterium and tritium are important research tools
– By incorporating these isotopes into specific positions in selected
molecules, they act as labels, or tracers
– Tracers are substances that enable biochemists to follow the path of
a molecule through an organism or cell and to provide information
about the mechanism of enzymatic reactions
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Bonding in Hydrogen and
Hydrogen-Containing Compounds
• The 1s1 electron configuration of hydrogen indicates a
single valence electron; the 1s orbital has a maximum
capacity of two electrons, so hydrogen can form
compounds with other elements in three ways:
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Bonding in Hydrogen and
Hydrogen-Containing Compounds
1. By losing its electron to form a proton (H+) with an
empty 1s orbital
– The proton is a Lewis acid that can accept a pair of
electrons from another atom to form an electron-pair
bond
2. By accepting an electron to form a hydride ion (H–),
which has a filled 1s2 orbital
– Hydrogen reacts with electropositive metals, such as
the alkali metals and alkaline earths, to form ionic
hydrides, which contain metal cations and H– ions
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Bonding in Hydrogen and
Hydrogen-Containing Compounds
3. By sharing its electron with an electron on another
atom to form an electron-pair bond
– With a half-filled 1s1 orbital, the hydrogen atom can
interact with singly occupied orbitals on other atoms to
form either a covalent or polar covalent electron-pair
bond, depending on the electronegativity of the other
atom
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Bonding in Hydrogen and HydrogenContaining Compounds
• Hydrogen can also act as a bridge between two or
more atoms. Two example of this are:
1.
The hydrogen bond, an electrostatic interaction between a
hydrogen bonded to an electronegative atom and an atom that
has one or more lone pairs of electrons
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Bonding in Hydrogen and HydrogenContaining Compounds
2. The three-center bond, in which a hydride ion
bridges two electropositive atoms
−
In these bonds, only two bonding electrons are used to
hold three atoms together and are called electron-deficient
bonds
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Synthesis, Reactions, and Compounds
of Hydrogen
• The first known preparation of elemental hydrogen was
in 1671, when Boyle dissolved iron in dilute acid and
obtained a colorless, odorless, gaseous product
• Hydrogen was identified as an element in 1766, when
Cavendish showed that water was the sole product of
the reaction of the gas with oxygen
• Explosive properties of mixtures of hydrogen with air
were discovered in the eighteenth century
• Due to its low molecular mass, hydrogen gas is difficult
to condense to a liquid and solid hydrogen has one of
the lowest melting points known
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Synthesis, Reactions, and Compounds
of Hydrogen
• Hydrogen gas
– Most common way to produce small amounts of highly pure
hydrogen gas is to react an active metal with dilute acid
– Hydrogen gas can also be generated by the reaction of metals
with a strong base
– Also produced by the reaction of ionic hydrides with water;
expensive—used for specialized purposes
– On an industrial scale, H2 is produced from methane by means
of catalytic steam reforming, a method used to convert
hydrocarbons to a mixture of CO and H2 known as synthesis gas,
or syngas
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Synthesis, Reactions, and Compounds
of Hydrogen
• Most of the elements in the periodic table form binary
compounds with hydrogen, called hydrides, which can
be divided into three classifications, each with its own
characteristic properties
1. Covalent hydrides contain hydrogen bonded to another
atom via a covalent or polar covalent bond and are
molecular substances that are volatile and have low melting
points
2. Ionic hydrides contain the hydride ion as the anion and
cations derived from electropositive metals; they and are
nonvolatile solids that contain three-dimensional lattices of
cations and anions, which decompose upon heating to H2
gas and the parent metal
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Synthesis, Reactions, and Compounds
of Hydrogen
3. Metallic hydrides are formed by hydrogen and less
electropositive metals such as the transition metals
– Have properties similar to those of the parent metal
– Best viewed as metals that contain many hydrogen atoms
present as interstitial impurities
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Chemistry: Principles, Patterns,
and Applications, 1e
21.3 The Alkali Metals
(Group 1)
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Preparation of Alkali Metals
• Alkali metals are so reactive that they are never found in
nature in elemental form
• Because the alkali metals are among the most potent
reductants known, obtaining them in pure form requires
a considerable input of energy
– Pure Li, Na, and K are prepared by electrolytic reduction of
the molten chlorides
– Metallic sodium and potassium are produced by the
electrolysis of molten mixtures of NaCl and CaCl2 or KCl and
CaCl2
– Rubidium and cesium are obtained by reaction of their
hydroxide salts with a reductant such as Mg
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Preparation of Alkali Metals
– Massive deposits of pure NaCl and KCl are found in
nature and are the major sources of sodium and
potassium
– Other alkali metals found in low concentrations in a
wide variety of minerals
– Alkali metals are obtained from silicate ores in a
multistep process that uses the pH-dependent
solubility of selected salts of each metal ion
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Preparation of Alkali Metals
• The steps in this process are:
1. leaching—use sulfuric acid to dissolve the
desired alkali metal and Al3+ from the ore
2. basic precipitation—to remove Al3+ from the
mixture as Al(OH)3
3. selective precipitation of the insoluble alkali metal
salt
4. dissolution of the salt in hydrochloric acid
5. isolation of the metal by evaporation and
electrolysis
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General Properties of Alkali Metals
• Various properties of the Group-1 elements are
summarized in the following table
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General Properties of Alkali Metals
– The atomic and ionic radii increase smoothly from Li to Cs, and
the first ionization energies decrease as the atoms become
larger
– Due to their low first ionization energies, the alkali metals have a
tendency to form an ion with a +1 charge—they have high
electron affinities because the addition of an electron produces
an anion (M–) with an ns2 electron configuration
– Densities of the elements increase from Li to Cs
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General Properties of Alkali Metals
– Melting and boiling points decrease from Li to Cs
– Standard reduction potentials, Eº, of the alkali metals do not
follow the trend based on ionization energies; lithium is the
strongest reductant and sodium is the weakest
– Li+ is much smaller than the other alkali metal cations, and its
hydration energy is the highest; lithium metal is the strongest
reductant in aqueous solution
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Reactions and Compounds of
Alkali Metals
• All alkali metals are electropositive elements with an ns1
valence-electron configuration, forming the monocation
(M+) by the loss of the single valence electron
• Chemistry of the alkali metals is largely that of ionic
compounds that contain the M+ ions
• The lighter Group-1 elements form a series of
organometallic compounds that contain polar covalent
M–C bonds
• All of the alkali metals react vigorously with the halogens
(Group 17) to form the corresponding ionic halides
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Reactions and Compounds of
Alkali Metals
• Alkali metals react with the heavier chalcogens (sulfur,
selenium, and tellurium in Group 16) to produce metal
chalcogenides
• The reaction of the alkali metals with oxygen is complex,
and the stoichiometry of the product depends on both
the metal:oxygen ratio and the size of the metal
• Alkali metal peroxides and superoxides are potent
oxidants that react with a wide variety of reducing agents
• Lithium, the lightest alkali metal, is the only one that
reacts with atmospheric nitrogen, forming lithium nitride
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Reactions and Compounds of
Alkali Metals
• All the alkali metals react with the larger Group-15
elements phosphorus and arsenic to form metal
phosphides and arsenides
• Alkali metals react with all the Group-14 elements, but
the compositions and properties of the products vary
significantly
• The heavier alkali metals ( K, Rb, and Cs) react with
carbon in the form of graphite, where the metals insert
themselves between the sheets of carbon atoms to give
new substances called graphite intercalation
compounds
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Reactions and Compounds of
Alkali Metals
• All the alkali metals react directly with gaseous hydrogen
at elevated temperatures to produce ionic hydrides
(M+H–) and all are capable of reducing water to produce
hydrogen gas
• Alkali metal cations are found in a wide variety of ionic
compounds; any alkali metal salt can be prepared by
reacting the alkali metal hydroxide with an acid and then
evaporating the water
• Hydroxides of alkali metals can react with organic
compounds that contain an acidic hydrogen to produce a
salt
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Complexes of Alkali Metals
• Because of their low positive charge (+1) and large ionic
radii, alkali metal cations have only a weak tendency to
react with simple Lewis bases to form metal complexes
• Complex formation is most significant for the smallest
cation, Li+, and decreases with increasing radius
• Complex formation is due to the electrostatic interaction
of the metal cation with the polar water molecules;
anhydrous salts containing Li+ and Na+ ions are used as
drying agents due to their high affinity for water
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Complexes of Alkali Metals
• Electrostatic interactions allow alkali metal ions to form
complexes with certain cyclic polyethers and related
compounds, such as crown ethers and cryptands
– Crown ethers are cyclic polyethers that contain four or more
oxygen atoms separated by two or three carbon atoms; they
also have a central cavity that can accommodate a metal ion
coordinated to the ring of oxygen atoms
– Cryptands are spherical analogues of crown ethers and are
more powerful and selective complexing agents; they consist
of three (-CH2CH2O-)n chains connected by two nitrogen
atoms and can completely surround a metal ion of the
appropriate size, coordinating to the metal by a lone pair of
electrons on each O atom and the two N atoms
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Liquid Ammonia Solutions
• Alkali metals dissolve reversibly in liquid ammonia which
produces hydrogen gas and the metal salt of the
conjugate base of the solvent (the amide ion NH2–)
M(s) + NH3 (l)  1/2H2 (g) + M+(am)NH2–(am)
– (am) designation refers to an ammonia solution
– Reaction needs a catalyst
– In many cases the alkali metal amide salt (MNH2) is not very
soluble in liquid ammonia and precipitates, but when dissolved,
solutions of the alkali metal are produced that can be very
concentrated
• Solutions of alkali metals in liquid ammonia are intensely
colored and are good conductors of electricity due to the
presence of solvated electrons, e–(NH3), which are not
attached to single atoms
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Liquid Ammonia Solutions
• In addition to solvated electrons, solutions of
alkali metals in liquid ammonia contain the metal
cation (M+), the neutral metal atom (M), metal
dimers (M2), and the metal anion (M–)
– Anion is formed by the addition of an electron to
the singly occupied ns valence orbital of the metal
atom
– These solutions are not stable and decompose to
the thermodynamically favored products M+NH2–
and hydrogen gas
– Solvated electron is a potent reductant
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Organometallic Compounds of the
Group-1 Elements
• Organometallic compounds contain a metal covalently
bonded to a carbon atom of an organic species
• Properties and reactivities of organometallic compounds
differ greatly from those of either the metallic or organic
components
– Lithium forms an extensive series of covalent organolithium
compounds that are volatile, low-melting-point solids or
liquids that can be sublimed or distilled at low temperatures
and are soluble in nonpolar solvents; the molten solids do not
conduct electricity and have a tendency to form oligomers
– Organosodium and organopotassium compounds are more
ionic than organolithium compounds, contain discrete M+ and
R– ions, and are insoluble or only sparingly soluble in
nonpolar solvents
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Uses of the Alkali Metals
• Sodium remains liquid over a wide temperature range
and is used as a coolant in specialized high-temperature
applications
• Cesium, because of its low ionization energy, is used in
photosensors in automatic doors, toilets, burglar alarms,
and other electronic devices
• Compounds of sodium and potassium are produced on a
huge scale in industry: NaOH, used in a wide variety of
industrial processes; Na2CO3, used in the manufacture
of glass; K2O, used in porcelain glazes; and Na4SiO4,
used in detergents
• Li2CO3 is an effective treatment for manic depression, or
bipolar disorder
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Chemistry: Principles, Patterns,
and Applications, 1e
21.4 The Alkaline Earth
Metals (Group 2)
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21.4 The Alkaline Earth Metals (Group 2)
• The alkaline earth metals are so reactive that they are
never found in elemental form in nature; they form +2
ions that have very negative reduction potentials so large
amounts of energy are needed to isolate them from their
ores
• Four of the six Group-2 elements—magnesium (Mg),
calcium (Ca), strontium (Sr), and barium (Ba)—were
isolated in the early nineteenth century by Davy
• Beryllium (Be), the lightest alkaline earth metal, was
obtained in 1828 by Wöhler and Bussy
• Radium was discovered in 1898 by the Curies
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Preparation of Alkaline Earth Metals
• Alkaline earths are produced for industrial use by
electrolytic reduction of their molten chlorides
• Group-2 metal chlorides are obtained from a variety of
sources
• Chemical reductants can also be used to obtain the
Group-2 elements
• Alkaline earths are somewhat easier than the alkali
metals to isolate from their ores because their carbonate,
sulfate, and some hydroxide salts are insoluble
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General Properties of Alkaline Earth
Metals
• Several important properties of the alkaline
earths are summarized in the following table
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General Properties of Alkaline Earth
Metals
– Many of these properties are similar to those of the alkali
metals, but certain key differences are attributable to the
differences in the valence-electron configuration of the two
groups (ns2 for the alkaline earth metals versus ns1 for the
alkali metals)
– Atomic and ionic radii of the alkaline earth metals increase
smoothly from Be to Ba, and the ionization energies
decrease
– The first ionization energy of an alkaline earth, with an ns2
valence-electron configuration, is always higher than that of
the alkali metal immediately preceding it
– The density of Ca is lower than that of Be and Mg, the two
lightest members of the group, and Mg has the lowest
melting and boiling points
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General Properties of Alkaline Earth
Metals
– The heaviest alkaline earth, Ba, is the strongest reductant, and
the lightest, Be, is the weakest
– Reduction potentials of Ca and Sr are not very different from that
of Ba, indicating that the opposing trends in ionization energies
and hydration energies are of equal importance
– One major difference between the Group-1 and Group-2
elements is their electron affinities; alkaline earths have little or
no tendency to accept an additional electron because their ns
valence orbitals are already full, while the alkali metals have a
significant affinity for an additional electron due to their half-filled
ns orbitals
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Reactions and Compounds of Alkaline
Earth Metals
• With their low first and second ionization energies, the
Group-2 elements exclusively form ionic compounds that
contain the M+2 ions
• The lightest element, Be, with its higher ionization
energy and small size, forms compounds that are largely
covalent; some compounds of Mg2+ have significant
covalent character; so organometallic compounds are
important for Be and Mg in Group 2
• All the alkaline earth metals react vigorously with the
halogens (Group 17) to form the corresponding halides
(MX2)
– these compounds are ionic in nature, containing the M2 cation
and two X– anions (with the exception of the beryllium halides)
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Reactions and Compounds of Alkaline
Earth Metals
• Beryllium halides, with properties more typical of
covalent compounds, have a polymeric halide-bridged
structure in the solid state
– These compounds are volatile, producing vapors that contain
the linear X-Be-X molecules
– They have four valence electrons around the central atom
– They are potent Lewis acids and react readily with Lewis
bases to form tetrahedral adducts in which the central
beryllium is surrounded by an octet of electrons
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Reactions and Compounds of Alkaline
Earth Metals
• Reactions of the alkaline earth metals with oxygen are
less complex than those of the alkali metals
– All the Group-2 elements, except barium, react directly with
oxygen to form the simple oxide MO
– barium forms barium peroxide (BaO2)
– BeO is prepared by direct reaction with oxygen; other
alkaline earth oxides are prepared by thermal decomposition
of carbonate salts
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Reactions and Compounds of Alkaline
Earth Metals
• The reactions of the alkaline earths with the heavier
chalcogens (Y) are similar to those of the alkali metals
– When the reactants are present in a 1:1 ratio, the binary
chalcogenides (MY) are formed
– at higher Y:M ratios, salts containing polychalcogenide ions
(Yn2–) are formed
• Oxides of Ca, Sr, and Ba react with CO2 to form
carbonate; the carbonates of the alkaline earths react
with aqueous acid to give CO2 and H2O
• Except for BeO, which has significant covalent character
and is amphoteric, all the alkaline earth oxides are basic
and react with water to form the hydroxides, M(OH)2, and
dissolve in aqueous acid
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Reactions and Compounds of Alkaline
Earth Metals
• Hydroxides of the lighter alkaline earths are insoluble in
water, but their solubility increases as the atomic number
of the metal increases; BeO and MgO are more inert
than the other Group-2 oxides
• Reactivities of the alkaline earth metals with nitrogen is
the opposite of that observed for the alkali metals
– Only the lightest element, Be, does not react readily with N2
to form the nitride (M3N2), although finely divided Be will react
at high temperatures
– The higher lattice energy due to the highly charged M2+ and
N3– ions is sufficient to overcome the chemical inertness of
the N2 molecule
• All alkaline earths react with the heavier Group-15
elements to form binary compounds such as phosphides
and arsenides with the general formula M3Z2
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Reactions and Compounds of Alkaline
Earth Metals
• When heated, all the alkaline earths, except for beryllium,
react directly with carbon to form ionic carbides with the
general formula MC2
– Most important is calcium carbide, which reacts with water to
produce acetylene
– Beryllium reacts with carbon to form Be2C, which reacts with water
or aqueous acid to produce methane
• Beryllium does not react with hydrogen, although BeH2
can be prepared by an indirect route
– All the heavier alkaline earths (Mg through Ba) react directly
with hydrogen to produce the binary hydrides (MH2), which
are ionic, but BeH2 and MgH2 have polymeric structures with
covalent character
– All alkaline earth hydrides are good reducing agents that
react rapidly with water or aqueous acid to produce
hydrogen gas
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Reactions and Compounds of Alkaline
Earth Metals
• Like the alkali metals, the heavier alkaline earths are
electropositive and dissolve in liquid ammonia
– In this case, two solvated electrons are formed per metal
atom, and no equilibria involving metal dimers or metal
anions are known
• Like the alkali metals, the alkaline earths form a wide
variety of simple ionic salts with oxoanions such as
carbonate, sulfate, and nitrate
– Nitrate salts tend to be soluble, but the carbonates and the
sulfates are quite insoluble because of the higher lattice
energy due to the doubly charged cation and anion
– Solubility of the carbonates and the sulfates decreases
rapidly down the group because hydration energies decrease
with increasing cation size
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Complexes of Alkaline Earth Metals
• Because of their higher positive charge (+2) and smaller ionic radii,
the alkaline earths have a greater tendency to form complexes with
Lewis bases than do the alkali metals. This tendency is most
important for the lightest cation, Be2+, and decreases rapidly with
increasing radius of the metal ion
• Chemistry of Be2+ dominated by its behavior as a Lewis acid,
forming complexes with Lewis bases that produce an octet of
electrons around the beryllium, which is amphoteric
• The heavier alkaline earths also form complexes, but with a
coordination number of 6 or higher; this behavior is most important
for the smaller cations Mg2+ and Ca2+
• Like the alkali metals, the alkaline earths form complexes with
neutral cyclic ligands like the crown ethers and cryptands
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Organometallic Compounds Containing
Group 2 Elements
• Like the alkali metals, the lightest alkaline earths (Be and
Mg) form the most covalent-like bonds with carbon, and
they form the most stable organometallic compounds
• Organometallic compounds of magnesium with the
formula RMgX, where R is an alkyl or aryl group and X is
a halogen, are called Grignard reagents, which are used
to synthesize various organic compounds
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Uses of the Alkaline Earth Metals
• Elemental magnesium is the only alkaline earth
metal that is produced on a large scale
– Its low density makes it an important component of
lightweight metal alloys used in aircraft frames and aircraft
and automobile engine parts
– Serves as a reductant for the production of a number of
metals
• Beryllium is widely used but is extremely toxic
– Increases the strength of copper and nickel alloys
– Its low atomic number gives it the lowest tendency to absorb
X-rays of all the metallic elements, which makes it suited for
applications involving radioactivity
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Uses of the Alkaline Earth Metals
• Tons of calcium compounds are used every year
– CaCl2 is used as road salt to lower the freezing point of water
on roads in cold temperatures
– CaCO3 is a major component of cement and an ingredient in
antacids
– “Quicklime” (CaO) is used in the steel industry to remove
oxide impurities, for making glass, and to neutralize acidic
soil
• BaSO4 used in milkshakes for identifying digestive
problems by X-rays
• Various alkaline earth compounds are used to produce
brilliant colors seen in fireworks
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Chemistry: Principles, Patterns,
and Applications, 1e
21.5 The s-Block
Elements in Biology
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21.5 The s-Block Elements in Biology
• The s-block elements play important roles in biological
systems
• Covalent hydrides are the building block of organic
compounds, and other compounds and ions containing
s-block elements are found in tissues and cellular fluids
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Covalent Hydrides
• There are three major classes of hydrides—covalent,
ionic, and metallic—but only covalent hydrides occur in
living cells and have any biochemical significance
–
–
–
Hydrogen is less electronegative than oxygen, nitrogen, or
sulfur (all symbolized by Z)
The H-Z bond in the hydrides of these elements is
polarized so the hydrogen atoms in H-Z bonds are acidic
Hydrides in which H is bonded to O, N, or S atoms are
polar, hydrophilic molecules that form hydrogen bonds and
undergo acid-base reactions by transferring a proton
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Covalent Hydrides
•
Hydrogen bonds are crucial in biochemistry because
they help hold proteins in their biologically active
folded structures
–
–
•
Hydrogen bonds connect the two intertwining strands of
DNA, the substance that contains the genetic code of all
organisms
Hydrogen bonds are easier to break than the covalent
bonds that form the individual DNA strands, so the two
intertwined strands can be separated to give intact single
strands, which is essential for the duplication of genetic
information
The extensive hydrogen-bonding network in water is
one of the keys to the existence of life on our planet
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Macrominerals
•
The members of Group 1 and Group 2 that are present
in the largest amounts in organisms are sodium,
potassium, magnesium, and calcium, all of which form
monatomic cations with a charge of +1 (Group 1, M+)
or +2 (Group 2, M2+); these are known as
macrominerals
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Macrominerals
• Ion transport
–
–
–
–
Na+ and Ca2+ are found in extracellular fluids, and K+ and Mg2+
are found in intracellular fluids
Energy is needed to transport each of these ions across the
cell membrane toward the side with the higher concentration
Biological machines responsible for the selective transport of
these metal ions are complex assemblies of proteins called
ion pumps, which recognize and discriminate between metal
ions with a high affinity for ions of a certain charge and radius
Defects in the ion pumps or their control mechanisms result in
major health problems
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Macrominerals
–
–
Some of the most important biological functions of the Group-1
and Group-2 metals are due to small changes in the cellular
concentrations of the metal ion
Within cells, K+ and Mg2+ activate particular enzymes by
binding to specific negatively charged sites in the enzyme
structure
• Ionophores
–
–
Molecules that facilitate the transport of metal ions across
membranes, because the health of cells depends on
maintaining the proper levels of cations in intracellular fluids
Potent antibiotics
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22
The p-Block
Elements
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CHAPTER OBJECTIVES
• To understand how valence-electron configurations and
periodic trends in atomic properties determine the
chemical properties of the p-block elements
• To use thermodynamics and kinetics to understand the
reactivity of the p-block elements
• To understand why the chemistry of the lightest member
of each group differs from that of the heavier elements of
the group
• To be able to predict the types of reactions the p-block
elements undergo
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The p-Block Elements
• p-block elements
– The line that divides metals from nonmetals in the periodic table
crosses the p block diagonally
– The differences between metallic and nonmetallic properties are
evident within each group, even though all members of the group
have the same valence-electron configuration
– The p-block only portion of the periodic table where the inert-pair
effect is seen
– The chemistry of the lightest member of each group in the p block
differs sharply from its heavier congeners but is similar to the
element immediately below and to the right of it in the next group;
diagonal similarities in chemistry are seen across the p block
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Chemistry: Principles, Patterns,
and Applications, 1e
22.1 The Elements of
Group 13
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
22.1 The Elements of Group 13
• Group 13, the first group to span the dividing line
between metals and nonmetals
• Its chemistry is more diverse than that of Groups 1 and 2,
which include only metallic elements
• Except for the lightest element, boron, the Group-13
elements are all electropositive; they tend to lose
electrons in chemical reactions rather than gain them
• None of these elements was known until the early
nineteenth century because they are never found in
nature in their free state
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Preparation and General Properties Of
Group-13 Elements
• Group-13 elements are not as powerful of reductants as
are the alkali metals and alkaline earths
• Their compounds with oxygen are thermodynamically
stable, and large amounts of energy are needed to
isolate them from their oxide ores
• The two most accessible elements are boron and
aluminum
• The other members of Group 13 are rather rare, and
these metals are usually obtained as by-products in the
processing of other metals
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Preparation and General Properties Of
Group-13 Elements
• The following table summarizes some important
properties of the Group-13 elements
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Preparation and General Properties Of
Group-13 Elements
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Preparation and General Properties Of
Group-13 Elements
– Large differences between boron and aluminum in size, ionization
energy, electronegativity, and reduction potential due to the fact
that boron behaves chemically like a nonmetal and aluminum like a
metal
– All the Group-13 elements have ns2np1 valence-electron
configurations, and all tend to lose their three valence electrons to
form compounds in the +3 oxidation state
– Heavier elements in the group also form compounds in the +1
oxidation state by loss of the single np valence electron
– Neutral compounds of the Group-13 elements contain only six
valence electrons and are all strong Lewis acids
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Preparation and General Properties Of
Group-13 Elements
– In contrast to Groups 1 and 2, the Group-13 elements show no
consistent trends in ionization energies, electron affinities, and
reduction potentials
– Electronegativity increases from aluminum to thallium
– Many of the inconsistencies observed in the properties of the
Group-13 elements can be explained by the increase in Zeff that
arises from poor shielding of the nuclear charge by the filled (n–
1)d10 and (n–2)f14 subshells; although actual nuclear charge
increases by 32 from indium to thallium, screening by the filled
5d and 4f subshells is so poor that Zeff increases significantly
from indium to thallium; therefore, the first ionization energy of
thallium is higher than that of indium
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Reactions and Compounds of Boron
• Elemental boron is a semimetal that is unreactive; the
other Group-13 elements all exhibit metallic properties
and reactivity and have fewer valence electrons than
valence orbitals, which results in delocalized, metallic
bonding
• Boron
– High ionization energy, low electron affinity, low electronegativity,
and small size
– Does not form a metallic lattice with delocalized valence
electrons; forms unique and intricate structures that contain
multicenter bonds, in which a pair of electrons holds together
three or more atoms
– Basic building block of elemental boron is not the individual
boron atom, as would be the case in a metal, but rather the B12
icosahedron, which do not pack together well and have voids,
resulting in low density
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Reactions and Compounds of Boron
– Elemental boron can be induced to react with many nonmetallic
elements to give binary compounds that have a variety of
applications
1. Boron carbide (B4C)—used in armor
2. Boron nitride (BN)—produced by heating boron with excess
nitrogen and is similar in many ways to elemental carbon
3. Boron oxide (B2O3)—formed when boron is heated with excess
oxygen, and it dissolves many metal and nonmetal oxides to
give a wide range of important borosilicate glasses
4. Boron trihalides (BX3)—formed by heating boron with excess
halogen
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Reactions and Compounds of Boron
− At high temperatures, boron also reacts with all metals to give
metal borides that contain regular three-dimensional networks,
or clusters, of boron atoms; metal borides are hard and
corrosion-resistant, and they are used in turbine blades and
rocket nozzles
− Binary hydrides were discovered in the early twentieth century
and have multicenter bonds
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Reactions and Compounds of the Heavier
Group-13 Elements
• Neutral compounds of the Group-13 elements are
electron deficient and behave like Lewis acids
• All four of the heavier elements (Al, Ga, In, and Tl) react
readily with the halogens to form compounds with the
stoichiometry MX3
– Of the halides, only the fluorides exhibit typical behavior of an
ionic compound with high melting points and low solubility in
nonpolar solvents
– The trichlorides, tribromides, and triiodides of aluminum, gallium,
and indium, as well as TlCl3, and TlBr3, are more covalent in
character and form halogen-bridged dimers, in which the
bonding is described as electron-pair bonds
– Group-13 trihalides are potent Lewis acids and react with Lewis
bases to form a Lewis acid-base adduct
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Reactions and Compounds of the Heavier
Group-13 Elements
– In water, the halides of the Group-13 metals hydrolyze to
produce the metal hydroxide [M(OH)3]; halides of the heavier
metals (In and Tl) are less reactive with water because of their
lower charge-to-radius ratio and dissolve to form the hydrated
metal ions [M(H2O)6]3+
• All the heavier Group-13 elements react with excess
oxygen at elevated temperatures to give the trivalent
oxide (M2O3)
– All the oxides dissolve in dilute acid, but aluminum and gallium
oxides are amphoteric and also dissolve in concentrated
aqueous base to form solutions that contain M(OH)4– ions
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Reactions and Compounds of the Heavier
Group-13 Elements
• Aluminum, gallium, and indium react with the other Group-16
elements (chalcogens) to form chalcogenides with the stoichiometry
M2Y3; thallium forms only the thallium() chalcogenides with the
stoichiometry Tl2Y
• Only aluminum reacts directly with N2 at high temperatures to give
AlN; GaN and InN are prepared by other methods
• All the metals, except Tl, react with the heavier Group-15 elements
(pnicogens) to form the -V compounds, which are semiconductors
• The heavier Group-13 elements do not react directly with hydrogen;
only aluminum and gallium hydrides are known and are prepared
indirectly
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Complexes of Group-13 Elements
• Boron has a limited tendency to form complexes
• Aluminum, gallium, indium, and thallium form a large
number of complexes; the simplest are the hydrated
metal ions, M(H2O)63+, that are strong Brønsted–Lowry
acids
• Group-13 metal ions also form stable complexes with
compounds that contain two or more negatively charged
groups; stability of such complexes increases as the
number of coordinating groups provided by the ligand
increases
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Chemistry: Principles, Patterns,
and Applications, 1e
22.2 The Elements of
Group 14
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
22.2 The Elements of Group 14
• The elements of Group 14 show a greater range of
chemical behavior than any other family in the periodic
table
• Three of the five elements—carbon, tin, and lead—have
been known since ancient times
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Preparation and General Properties of
Group-14 Elements
• The natural abundance of the Group-14
elements varies tremendously
– Elemental carbon ranks 17th on the list of constituents of
Earth’s crust
– After oxygen, the most abundant element on Earth is silicon,
the next member of Group 14; pure silicon is obtained by the
reaction of impure silicon with Cl2 to give SiCl4, followed by
reduction with H2; ultrapure silicon and germanium form the
basis of the modern electronics industry
– Concentrations of germanium and tin in the Earth’s crust are
1–2 ppm and lead is 13 ppm, which makes lead the most
abundant of the heavy Group-14 elements
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Preparation and General Properties of
Group-14 Elements
– No concentrated ores of germanium are known, so germanium
is recovered from flue dusts obtained by processing the ores of
metals and is used in optical devices
– Tin and lead are soft metals and are too weak for structural
applications; tin is used in food packaging, magnets, and lowmelting-point alloys, lead is used in lead-storage batteries
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Preparation and General Properties of
Group-14 Elements
• All the Group-14 elements form compounds in which they lose either
the two np and the two ns valence electrons or just the two np
valence electrons, giving compounds in the +4 or +2 oxidation state,
respectively
• The relative stability of the +2 oxidation state increases smoothly
from carbon to lead because covalent bonds decrease in strength
with increasing atomic size, and the ionization energies for the
heavier elements of the group are higher than expected due to a
higher Zeff
• Many carbon compounds contain multiple bonds formed by 
overlap of singly occupied 2p orbitals on adjacent atoms
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Preparation and General Properties of
Group-14 Elements
• Compounds of silicon, germanium, tin, and lead with the
same stoichiometry as those of carbon, tend to have
different structures and properties
• The tendency to catenate (to form chains of like atoms)
decreases rapidly going down Group 14 because bond
energies for both the E–E and E–H bonds decrease with
increasing atomic number (where E is any Group-14
element); the thermal stability of catenated compounds
decreases from carbon to lead
• Group-14 elements all have ns2np2 valence-electron
configurations
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Preparation and General Properties of
Group-14 Elements
• As shown in the following table, there is a
large difference between the lightest
element, C, and the others in size,
ionization energy, and electronegativity
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Preparation and General Properties of
Group-14 Elements
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Preparation and General Properties of
Group-14 Elements
– As in Group 13, the second and third elements (Si
and Ge) are similar
– There is a reversal in the trends of some properties,
such as ionization energy, between the fourth and
fifth elements (Sn and Pb), which can be explained
by the presence of filled (n–1)d and (n–2)f
subshells, whose electrons are poor at screening
the outermost electrons from the higher nuclear
charge
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Reactions and Compounds of Carbon
• Carbon
– The building block of all organic compounds; inorganic
compounds of carbon include metal carbonates
– Chemistry of carbon differs from that of its heavier congeners
– structures of the allotropes of carbon—diamond, graphite,
fullerenes, and nanotubes—are distinct, but they all contain
simple electron-pair bonds
– All the carbon tetrahalides (CX4) are known; they are not
obtained by the direct reaction of carbon with the elemental
halogens but by indirect methods; their stability decreases as
the halogen increases in size because of poor orbital overlap
and increased crowding
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Reactions and Compounds of Carbon
– Reacts with oxygen to form either CO or CO2, depending on
the stoichiometry
– CO can be prepared by dehydrating formic acid with
concentrated sulfuric acid; CO reacts with halogens to form the
oxohalides (COX2)
– CO2 can be prepared by the reaction of any metal carbonate or
bicarbonate salt with strong acid; CO2 reacts with water to form
acidic solutions that contain carbonic acid (H2CO3)
– Reaction of carbon with sulfur at high temperatures produces
carbon disulfide (CS2)
– Binary compounds of carbon with less electronegative
elements are called carbides
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Reactions and Compounds of Carbon
• The chemical and physical properties of
carbides depend on the identity of the second
element, resulting in three general classes
1. Ionic carbides
– Produced by the reaction of carbon at high temperatures with
electropositive metals such as those of Groups 1 and 2 and
aluminum
– Contain discrete metal cations and carbon anions (C4–, methide or
C22–, acetylide); identity of the anions depends on the size of
the second element
– Reaction of ionic carbides with dilute aqueous acid results in
protonation of the anions to give the parent hydrocarbons,
CH4 or C2H2
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Reactions and Compounds of Carbon
2. Interstitial carbides
– Produced by the reaction of carbon with most transition metals at
high temperatures
– Contain covalent metal-carbon interactions, which result in
different properties:
a. Good conductors of electricity
b. Have high melting points
c. Among the hardest substances known
– Exhibit a variety of nominal compositions, and they are
nonstoichiometric compounds whose carbon content can
vary over a wide range
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Reactions and Compounds of Carbon
3. Covalent carbides
– Formed with elements with an electronegativity similar to that
of carbon
– Extremely hard, high melting, and chemically inert
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Reactions and Compounds of the
Heavier Group-14 Elements
• Silicon, germanium, tin, and lead in their +4 oxidation states form
binary compounds with the same stoichiometry as carbon, but the
structures and properties of these compounds are different from
those of the carbon analogues
• Silicon is amphoteric; it dissolves in strong aqueous base to produce
hydrogen gas and solutions of silicates; the only aqueous acid it
reacts with is hydrofluoric acid
• Germanium is amphoteric and is more metallic in its behavior than
silicon; tin has an even more metallic character and is amphoteric;
shows that metallic behavior increases going down the group
• Lead is the only element in the group that behaves purely like a
metal
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Reactions and Compounds of the
Heavier Group-14 Elements
• All the Group-14 dichlorides are known, but their stability increases
as the atomic number of the central atom increases
• The first four elements of Group 14 form tetrahalides (MX4) with all
the halogens, but only fluorine is able to oxidize lead to the +4
oxidation state
– Tetrahalides of the semimetals silicon and germanium react rapidly
with water to give amphoteric hydroxides
– Tetrahalides of tin and lead react with water to give hydrated metal
ions
• The heavier Group-14 elements react with O2 or S8 at elevated
temperatures to form the corresponding dioxide or disulfide,
respectively; silicon has a tremendous affinity for oxygen because of
partial Si–O  bonding
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Reactions and Compounds of the
Heavier Group-14 Elements
• Dioxides of the Group-14 elements become increasingly
basic going down the group and their metallic character
increases
• Compounds with anions that contain only silicon and
oxygen are called silicates
– Basic building block of all silicates is the SiO44– unit
– The number of oxygen atoms shared between silicon atoms
and the way in which the units are linked vary in different
silicates
– In aluminosilicates, some of the Si atoms are replaced by Al
atoms and have three-dimensional framework structures with
large cavities connected by small tunnels
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Reactions and Compounds of the
Heavier Group-14 Elements
• Silicon and germanium react with nitrogen at high temperature to
form nitrides (M3N4) that are strong, hard, and chemically inert
• Silicides of active metals are ionic compounds that contain the Si4–
ion and react with aqueous acid to form silicon hydrides
• Unlike carbon, catenated silicon hydrides become
thermodynamically less stable as the chain lengthens and the
hydrides become thermodynamically less stable going down the
group
• As atomic size increases, multiple bonds between or to the Group-14
elements become weaker
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Reactions and Compounds of the
Heavier Group-14 Elements
• Only important organic derivatives of lead are
compounds such as tetraethyllead because Pb–C bonds
are weak
• Compounds that contain Si–C and Si–O bonds are
stable and important
– High-molecular-mass polymers called silicones contain an
(Si–O–)n backbone with organic groups attached to Si
– Properties of silicones are determined by the chain length, the
type of organic group, and the extent of cross-linking between
the chains
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Chemistry: Principles, Patterns,
and Applications, 1e
22.3 The Elements of
Group 15 (the Pnicogens)
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22.3 The Elements of Group 15
(the Pnicogens)
• The lightest member of Group 15, nitrogen, is found in
nature as the free element
• The heaviest elements have been known for centuries
because they are easily isolated from their ores
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Preparation and General Properties of
Group-15 Elements
• The atmosphere contains tons of elemental nitrogen with
a purity of 80%, so it is a huge source of nitrogen gas
– Distillation of liquefied air yields nitrogen gas that is about
100% pure
– Small amounts of very pure nitrogen gas can be obtained from
the thermal decomposition of sodium azide
– Earth’s crust is poor in nitrogen
• Phosphorus constitutes about 0.1% of Earth’s crust
– More abundant in ores than nitrogen
– Always found in combination with oxygen, and large inputs of
energy are required to isolate it
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Preparation and General Properties of
Group-15 Elements
• The other three pnicogens are much less
abundant
– Arsenic is found in Earth’s crust at a concentration of 2 ppm,
antimony is an order of magnitude less abundant, and
bismuth is rare
– All three elements have a high affinity for the chalcogens and
are found as the sulfide ores (M2S3), in combination with
sulfides of other heavy elements
– Major source of antimony and bismuth is flue dust obtained by
smelting the sulfide ores of the more abundant metals
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Preparation and General Properties of
Group-15 Elements
• In Group 15, there are large differences between the
lightest element, N, and its congeners in size, ionization
energy, electron affinity, and electronegativity as seen in
the following table
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Preparation and General Properties of
Group-15 Elements
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Preparation and General Properties of
Group-15 Elements
• The chemical behavior of the elements are:
– nitrogen and phosphorus behave chemically like nonmetals,
arsenic and antimony like semimetals, and bismuth like a
metal
– they have ns2np3 valence-electron configurations
– they all form compounds by losing the three np valence
electrons to form the +3 oxidation state, or by losing the three
np and two ns valence electrons to give the +5 oxidation state,
whose stability decreases smoothly from phosphorus to
bismuth
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Preparation and General Properties of
Group-15 Elements
– the high electron affinity of the lighter pnicogens
enables them to form compounds in the –3
oxidation state, in which three electrons are added
to the neutral atom to give a filled np subshell
– nitrogen has a high electron affinity and small size
and has the ability to form compounds in nine
different oxidation states, including –3, +3, and +5
– neutral covalent compounds of the trivalent
pnicogens contain a lone pair of electrons on the
central atom and tend to behave like Lewis bases
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Reactions and Compounds of Nitrogen
• Nitrogen
– Has four valence orbitals (one 2s and three 2p), so it can
participate in four electron-pair bonds by using sp3 hybrid
orbitals
– Nitrogen is smaller than carbon and has less tendency to
accommodate more than four nearest neighbors
– Unlike carbon, nitrogen does not form long chains due to
repulsive interactions between lone pairs of electrons on
adjacent atoms
– Stable compounds with N–N bonds are limited to chains of no
more than three N atoms
– Only pnicogen that normally forms multiple bonds with other
second-period atoms, using  overlap of adjacent np orbitals
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Reactions and Compounds of Nitrogen
– The stable form of elemental nitrogen is N2 with a strong triple
bond
– Compounds containing N–N single and NN double bonds are
thermodynamically unstable and potentially explosive; the
formation of the NN triple bond is thermodynamically favored
– In contrast to carbon, nitrogen undergoes only two important
chemical reactions at room temperature
1. Reacts with metallic lithium to form lithium nitride
2. It is reduced to ammonia by certain microorganisms
– At higher temperatures, N2 reacts with more
electropositive elements to give binary nitrides, which range
from covalent to ionic in character
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Reactions and Compounds of Nitrogen
– Binary compounds of nitrogen with oxygen, hydrogen, or other
nonmentals are covalent molecular substances
– Few binary molecular compounds of nitrogen are formed by
direct reaction of the elements: at elevated temperatures, N2
reacts with H2 to form ammonia, with O2 to form a mixture of
NO and NO2 , and with carbon to form cyanogen (NC–CN),
but elemental nitrogen does not react with the halogens or the
other chalcogens
– All the binary nitrogen halides (NX3) are known; except for NF3,
all are toxic, thermodynamically unstable, and explosive, and
all are prepared by the reaction of the halogen with NH3 rather
than N2
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Reactions and Compounds of Nitrogen
– The three oxides of nitrogen (NO, N2O, and NO2) are
thermodynamically unstable
– At elevated temperatures, nitrogen reacts with highly
electropositive metals to form ionic nitrides, which consist of
ionic lattices formed by Mn+ and N3– ions
– With less electropositive metals, nitrogen forms a range of
interstitial nitrides, in which nitrogen occupies holes in a closepacked metallic structure; these substances are hard, highmelting-point materials that have metallic luster and
conductivity
– Reacts with semimetals at very high temperatures to produce
covalent nitrides, which are solids with extended covalent
network structures and are high melting and chemically inert
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Reactions and Compounds of Nitrogen
– Ammonia (NH3) is one of the few thermodynamically stable
binary compounds of nitrogen with a nonmetal
– Forms two other important binary compounds with hydrogen
1. Hydrazoic acid (HN3), a colorless, highly toxic, and explosive
substance
2. Hydrazine (N2H4), also explosive
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Reactions and Compounds of the
Heavier Pnicogens
• The heavier pnicogens form catenated compounds that
contain only single bonds, whose stability decreases
rapidly going down the group
• Phosphorus forms multiple allotropes: white phosphorus,
an electrical insulator, and red phosphorus and black
phosphorus, which are semiconductors
• The three heaviest pnicogens—arsenic, antimony, and
bismuth—all have a metallic luster, are brittle, and are
poor electrical conductors
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Reactions and Compounds of the
Heavier Pnicogens
• Reactivity of the heavier pnicogens decreases
going down the column
– Phosphorus is the most reactive, forming binary compounds
with every element in the periodic table except antimony,
bismuth, and the noble gases
– Phosphorus reacts rapidly with O2, whereas arsenic burns in
pure O2, and antimony and bismuth react with O2 only when
heated
– None of the pnicogens reacts with nonoxidizing acids but all
dissolve in oxidizing acids
– Only bismuth behaves like a metal
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Reactions and Compounds of the
Heavier Pnicogens
• Heavier pnicogens use energetically accessible 3d, 4d, or 5d orbitals
to form dsp3, or d2sp3 hybrid orbitals for bonding; these elements
have coordination numbers of 5 or higher
• Phosphorus and arsenic form halides that are covalent molecular
species and behave like nonmetal halides, reacting with water to
form the corresponding oxoacids
• All the pentahalides are potent Lewis acids
• Bismuth halides have extended lattice structures and dissolve in
water to produce hydrated ions
• Except for BiF3, which is an ionic compound, the trihalides are
volatile covalent molecules with a lone pair of electrons in the
central atom and react rapidly with water
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Reactions and Compounds of the
Heavier Pnicogens
• With energetically accessible d orbitals, phosphorus and
arsenic are able to form  bonds with second period
atoms such as N and O
– Very strong P–O single bonds and even stronger PO double
bonds
– First four elements in Group 15 react with oxygen to produce
the corresponding oxide in the +3 oxidation state
– The two least metallic elements, phosphorus and arsenic, form
very stable oxides with the formula E4O10 in the +5 oxidation
state
• Heavier pnicogens form sulfides that range from
molecular species with three-dimensional cage
structures to layered or ribbon structures
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Reactions and Compounds of the
Heavier Pnicogens
• Reaction of the heavier pnicogens with metals produces
substances whose properties vary with the metal content
– Metal-rich phosphides (M4P) are hard, high-melting-point,
electrically conductive solids with a metallic luster
– Phosphorus-rich phosphides (MP15) are lower melting and less
thermally stable because they contain catenated Pn units
• Many of the organic or organometallic heavier pnicogens
that contain from one to five alkyl or aryl groups are
known, and their thermal stability decreases from
phosphorus to bismuth
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Chemistry: Principles, Patterns,
and Applications, 1e
22.4 The Elements of
Group 16 (the
Chalcogens)
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22.4 The Elements of Group 16 (the
Chalcogens)
• The chalcogens are the first group in the p block that
contain no stable metallic elements
– All isotopes of polonium (Po), the only metal in Group 16, are
radioactive
– Only one element in the group, tellurium (Te), can be described
as a semimetal
– The lightest element of Group 16, oxygen, is found in nature as
the free element
– Of the Group-16 elements, only sulfur was known in ancient
times; others were not discovered until the late eighteenth and
nineteenth centuries
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Preparation and General Properties of
Group-16 Elements
• Oxygen is the most abundant element in Earth’s crust
and in the hydrosphere
– Can be obtained by the electrolysis of water, the
decomposition of alkali metal or alkaline earth peroxides or
superoxides, or the thermal decomposition of simple inorganic
salts
• Sulfur is not very abundant, but it is found as elemental
sulfur in rock formations overlying salt domes; sulfur is
also recovered from H2S and organosulfur compounds in
crude oil and coal, and from metal sulfide ores
• Selenium and tellurium are found as minor contaminants
in metal sulfide ores and are recovered as by-products
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Preparation and General Properties of
Group-16 Elements
• Properties
– Have ns2np4 electron configurations
– Chalcogens are two electrons short of a filled valence shell;
they tend to acquire two additional electrons to form
compounds in the –2 oxidation state; and this tendency is
greatest for oxygen, the chalcogen with the highest
electronegativity
– Heavier, less electronegative chalcogens can lose either four
np electrons or four np and two ns electrons to form
compounds in the +4 and +6 oxidation state, respectively
– The lightest member in the group (oxygen) differs greatly from
the others in size, ionization energy, electronegativity, and
electron affinity
– Second and third members (sulfur and selenium) have similar
properties because of shielding effects; only polonium is
metallic as seen in the following table
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Preparation and General Properties of
Group-16 Elements
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Reactions and Compounds of Oxygen
• The lightest Group-16 member has the greatest
tendency to form multiple bonds, so elemental oxygen is
found in nature as a diatomic gas that contains a net
double bond, OO
• Electrostatic repulsion between lone pairs of electrons
on adjacent atoms prevents oxygen from forming stable
catenated compounds
– All compounds that contain O–O bonds are explosive
– Ozone, peroxides, and superoxides are all dangerous in pure
form
• Despite the strength of the OO bond, O2 is extremely
reactive, reacting directly with nearly all other elements
except the noble gases; properties of O2 and related
species are listed in the following table
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Reactions and Compounds of Oxygen
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Reactions and Compounds of Oxygen
• Chemistry of oxygen is restricted to negative oxidation
states because of its high electronegativity
– Oxygen does not form compounds in the +4 or +6 oxidation
state
– Second only to fluorine in its ability to stabilize high oxidation
states of metals in both ionic and covalent compounds
– Because oxygen is so electronegative, the O–H bond is highly
polar, creating a large bond dipole moment that makes
hydrogen bonding more important for compounds of oxygen
than for compounds of the other chalcogens
– Metal oxides are basic, nonmetal oxides are acidic, and oxides
of elements that lie on or near the diagonal band of emimetals
are amphoteric; nonmetal oxides are covalent compounds
where the bonds between oxygen and the nonmetal are
polarized and in water form oxoacids
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Reactions and Compounds of the
Heavier Chalcogens
• The tendency to catenate decreases going down the column
– Sulfur and selenium form an extensive series of catenated species
– There is only one stable allotrope of tellurium
– Polonium shows no tendency to form catenated compounds
• There is a striking decrease in structural complexity from sulfur to
polonium and a decrease in the strength of single bonds and an
increase in metallic character going down the group
• Reactivity of elements in Group 16 decreases from lightest to
heaviest
• Fluorine reacts directly with the chalcogens to produce hexafluorides
(YF6); four additional stable fluorides of sulfur are known, but only
two fluorides of selenium and three of tellurium are known
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Reactions and Compounds of the
Heavier Chalcogens
• Direct reaction of the heavier chalcogens with oxygen at
elevated temperatures gives the dioxides (YO2), which
exhibit a dramatic range of structures and properties
– Going down the column, the dioxides become increasingly
metallic in character and the coordination number of the
chalcogen increases
– Dioxides of sulfur, selenium, and tellurium react with water to
produce the weak, diprotic oxoacids (H2YO3)
– Sulfuric and selenic acids are strong acids, but telluric acid is
different because tellurium is larger than sulfur and selenium
and it forms weak  bonds to oxygen
– The most stable structure for telluric acid is Te(OH)6, with six
Te–OH single bonds rather than TeO double bonds
– Sulfuric, selenic, and telluric acids are oxidants, and the
stability of the highest oxidation state decreases with
increasing atomic number
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Reactions and Compounds of the
Heavier Chalcogens
• Sulfur and selenium react with carbon to form an
extensive series of compounds that are structurally
similar to their oxygen analogues
• Chalcogens react directly with nearly all metals to form
compounds with a wide range of stoichiometries and a
variety of structures
– Metal chalcogenides can contain either the simple
chalcogenide ion (Y2–) or polychalcogenide ions (Yn2–)
• Ionic chalcogenides react with aqueous acid to produce
binary hydrides
– Strength of the Y–H bond decreases with increasing atomic
radius, so the stability of the binary hydrides decreases rapidly
going down the column
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Chemistry: Principles, Patterns,
and Applications, 1e
22.5 The Elements of
Group 17 (the Halogens)
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22.5 The Elements of Group 17 (the
Halogens)
• Halogens are highly reactive, but none are found in
nature as the free element
• None of the halogens was recognized as an element
until the nineteenth century
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Preparation and General Properties of
Group-17 Elements
• All the halogens except iodine are found in nature as
salts of the halide ions (X–), so the methods used for
preparing F2, Cl2, and Br2 all involve oxidizing the
halide
•
Fluorine is a very powerful oxidant
– Reaction of CaF2 with concentrated sulfuric acid produces
gaseous hydrogen fluoride
– Fluorine is produced by the electrolysis of a 1:1 mixture of
HF and K+HF2–
– Both F2 and HF are highly corrosive
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Preparation and General Properties of
Group-17 Elements
•
Chlorine is less abundant than fluorine, but elemental
chlorine is produced on an enormous scale
– There are large subterranean deposits of rock salt around
the world; seawater consists of about 2% NaCl by mass,
and inland salt lakes are richer sources
– Chlorine is prepared by the chlor-alkali process
•
Bromine is much less abundant than fluorine or
chlorine but is easily recovered from seawater; salt
lakes and underground brines are richer sources
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Preparation and General Properties of
Group-17 Elements
•
Iodine is the least abundant of the nonradioactive
halogens and is a rare element
– Has low electronegativity and occurs in nature in an oxidized
form
– Most commercially important deposits of iodine are iodate
salts, so the production of iodine from these deposits
requires reduction rather than oxidation, which occurs in two
steps
1. Reduction of iodate to iodide with sodium hydrogen sulfite
2. Reaction of iodide with additional iodate
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Preparation and General Properties of
Group-17 Elements
• The heaviest halogen, astatine (At) is continuously
produced by natural radioactive decay
– All its isotopes are highly radioactive
– Most stable isotope has a half-life of 8 hours
– Least abundant naturally occurring element on Earth
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Preparation and General Properties of
Group-17 Elements
• Properties
– Halogens all have ns2np5 electron configurations, so their
chemistry is dominated by a tendency to accept an additional
electron to form the closed-shell ion (X–)
– Only the electron affinity and the bond dissociation energy of
fluorine differ significantly from the expected periodic trends as
shown in the following table
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Preparation and General Properties of
Group-17 Elements
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Preparation and General Properties of
Group-17 Elements
– Fluorine has a small atomic volume, so repulsive electronelectron interactions are important in the fluoride ion, making
the electron affinity of fluorine lower than that ofchlorine
– Repulsions between electron pairs on adjacent atoms are
responsible for the low F–F bond dissociation energy
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Preparation and General Properties of
Group-17 Elements
– Fluorine is the most electronegative element in the periodic
table, so it forms compounds in only the –1 oxidation state
– All the halogens except astatine have electronegativities
higher than 2.5, making their chemistry that of nonmetals
– Halogens all have high ionization energies, but the energy
required to remove electrons decreases going down the
column
– Heavier halogens form compounds in positive oxidation states
(+1, +3, +5, and +7), derived by the formal loss of ns and np
electrons
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Reactions and Compounds of the
Halogens
• Fluorine is the most reactive element in the periodic
table, forming compounds with every other element
except helium, neon, and argon
– Reactions of fluorine with most other elements range from
vigorous to explosive; only O2, N2, and Kr react slowly
– There are three reasons for the high reactivity of fluorine
1. Fluorine is so electronegative, it can remove or share the
valence electrons of any other element
2. Because of its small size, fluorine tends to form very strong
bonds to other elements, making its compounds
thermodynamically stable
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Reactions and Compounds of the
Halogens
3. The F–F bond is weak due to repulsion between lone pairs
of electrons on adjacent atoms, reducing both the
thermodynamic and kinetic barriers to reaction
– With highly electropositive elements, fluorine forms ionic
compounds that contain the closed-shell F– ion—with less
electropositive elements (or with metals in very high oxidation
states), fluorine forms covalent compounds that contain
terminal F atoms
– Because of its high electronegativity and 2s22p5 valenceelectron configuration, fluorine participates in only one
electron-pair bond; only a strong Lewis acid can share a lone
pair of electrons with a fluorine atom
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Reactions and Compounds of the
Halogens
• The halogens (X2) react with metals (M) according to
the general equation
M(s,l) + nX2(s,I,g)  MXn(s,l)
• For elements that exhibit multiple oxidation states,
fluorine tends to produce the highest possible
oxidation state and iodine the lowest
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Reactions and Compounds of the
Halogens
• Metal halides in the +1 or +2 oxidation state are ionic
halides, which have high melting points and are
soluble in water; as the oxidation state of the metal
increases, so does the covalent character of the halide
due to polarization of the M–X bond
– Fluoride, with its high electronegativity, is the least
polarizable, and iodide, with the lowest electronegativity, is
the most polarizable of the halogens
– Halides of small trivalent metal ions tend to be covalent, and
halides of larger trivalent metals are ionic
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Reactions and Compounds of the
Halogens
• All the halogens react vigorously with hydrogen to give
the hydrogen halides (HX)
– Because the H–F bond in HF is highly polarized, liquid HF is
extensively hydrogen bonded, has a high boiling point and
high dielectric constant, and is a polar solvent
• Except for fluorine, all the halogens react with water in
a disproportionation reaction:
X2(g,l,s) + H2O(l)  H+(aq) + X–(aq) + HOX(aq)
(X = Cl, Br, )
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Reactions and Compounds of the
Halogens
• The most stable oxoacids are the perhalic acids, which
contain the halogens in their highest oxidation state
(+7)
– Acid strengths of the oxoacids of the halogens increase with
increasing oxidation state, but their stability and acid strength
decrease going down the group
– All the oxoacids are strong oxidants, but some react rather
slowly at low temperatures
– Mixtures of the halogen oxoacids or oxoanions with organic
compounds are explosive if heated or agitated
– Oxoacids and oxoanions of the halogens should never be
allowed to come into contact with organic compounds
because of the danger of explosions
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Reactions and Compounds of the
Halogens
• The halogens react with one another to produce
interhalogen compounds
– In all cases, the heavier halogen, with the lower
electronegativity, is the central atom
– The maximum oxidation state and the number of terminal
halogens increase smoothly as the ionization energy of the
central halogen decreases and the electronegativity of the
terminal halogen increases
– Interhalogen compounds are very powerful Lewis acids with
a strong tendency to react with halide ions to give complexes
with higher coordination numbers
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Reactions and Compounds of the
Halogens
• The halogens react with one another to produce
interhalogen compounds
– All the Group-17 elements except fluorine form compounds in
odd oxidation states (–1, +1, +3, +5, +7)
– Interhalogen compounds are potent oxidants and strong
fluorinating agents; contact with water or organic materials
can result in an explosion
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Chemistry: Principles, Patterns,
and Applications, 1e
22.6 The Elements of Group
18 (the Noble Gases)
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22.6 The Elements of Group 18
(the Noble Gases)
• Noble gases were all isolated for the first time within a
period of only 5 years at the end of the nineteenth
century
• Their existence was not suspected until the eighteenth
century, when work on the composition of air suggested
that it contained small amounts of gases in addition to
oxygen, nitrogen, carbon dioxide, and water vapor
• Elements of Group 18 are helium, neon, argon, krypton,
xenon, and radon
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Preparation and General Properties of
Group-18 Elements
• Fractional distillation of liquid air is the only source of all
the noble gases except helium
• Helium is the second most abundant element in the
universe; natural gas contains high concentrations of
helium, and it is the only practical terrestrial source
• Elements of Group 18 all have closed-shell valenceelectron configurations, ns2 np6, except for He, which is
1s2
• These elements have high ionization energies that
decrease smoothly going down the column, as seen in
the following table
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Preparation and General Properties of
Group-18 Elements
Copyright © 2007 Pearson Education, Inc., publishing as Pearson Benjamin Cummings
Preparation and General Properties of
Group-18 Elements
• From their electron affinities, the noble gases are
unlikely to form compounds in negative oxidation states
– A potent oxidant is needed to oxidize noble gases and form
compounds in positive oxidation states
• Xenon and krypton should form covalent compounds
with F, O, and Cl, in which they have even formal
oxidation states (+2, +4, +6, and possibly +8)
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Reactions and Compounds of the Noble
Gases
• Noble gases form stable clathrates, solid compounds in
which a gas—the guest—occupies holes in a lattice
formed by a less volatile, chemically dissimilar
substance—the host
• Ionization energies of helium, neon, and argon are high,
so no stable compounds of these elements are known
• Ionization energies of krypton and xenon are lower but
still high, and only highly electronegative elements (F, O,
and Cl) can form stable compounds with xenon and
krypton without being oxidized themselves
• Xenon reacts directly with only two elements, F2 and Cl2
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Reactions and Compounds of the Noble
Gases
• Halides of the noble gases are powerful oxidants and
fluorinating agents, so they decompose rapidly upon
contact with trace amounts of water and react violently
with organic compounds or other reductants
• Xenon fluorides are Lewis acids and react with fluoride
ion—the only Lewis base that is not oxidized
immediately upon contact—to form anionic complexes
• Xenon has a high affinity for oxygen because of 
bonding between O and Xe, so xenon forms an
extensive series of oxides and oxoanion salts
• Xenon forms stable compounds with fluorine and oxygen
that contain xenon in even oxidation states up to +8
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23
The d-Block
Elements
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CHAPTER OBJECTIVES
• To know periodic trends in the d-block elements
• To be able to use periodic trends to understand the chemistry of the
transition metals
• To understand how metals are extracted from their ores
• To know the most common structures observed for metal complexes
• To be able to predict the relative stabilities of metal complexes
• To understand how crystal field theory can explain the electronic
structures and colors of metal complexes
• To become familiar with some of the roles of transition-metal
complexes in biological systems
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The d-Block Elements
• The d-block elements are also called transition metals
• All d-block elements are metals, so these elements
exhibit significant horizontal and vertical similarities in
chemistry, and all have a common set of characteristic
properties due to partially filled d subshells
• Alloys and compounds of d-block elements are important
components of the materials the modern world depends
on for its continuing technological development; most of
the first-row transition metals are essential for life
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Chemistry: Principles, Patterns,
and Applications, 1e
23.1 General Trends among
the Transition Metals
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Electronic Structure and Reactivity of the
Transition Metals
•
Valence-electron configurations of the first-row
transition metals are given in the following table
–
–
Going across the row from left to right, electrons are added
to the 3d subshell to neutralize the increase in the positive
charge of the nucleus as the atomic number increases
The 3d subshell is filled based on the aufbau principle and
Hund’s rule with two important exceptions:
1. Chromium has a 4s13d5 electron configuration rather than
a 4s23d4 configuration
2. Copper is 4s13d10 rather than 4s23d 9
–
Anomalies due to the extra stability associated with halffilled subshells
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Electronic Structure and Reactivity of the
Transition Metals
• In the second-row transition metals, electron-electron repulsions
within the 4d subshell cause additional irregularities in electron
configurations that are not easily predicted
• Further complications occur among the third-row transition metals, in
which the 4f, 5d, and 6s orbitals are close in energy
• From this point to element 71, added electrons enter the 4f subshell,
giving rise to the 14 elements known as the lanthanides
• After the 4f subshell is filled, the 5d subshell is populated, producing
the third row of the transition metals; the seventh period comes next,
where the actinides have three subshells (7s, 6d, and 5f) that are so
similar in energy that their electron configurations are more
unpredictable
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Electronic Structure and Reactivity of the
Transition Metals
• Reactivity
– The size of neutral atoms of the d-block elements gradually
decreases going from left to right across a row, due to an
increase in the effective nuclear charge (Zeff) with increasing
atomic number
– The atomic radius increases going down a column
– Because of the lanthanide contraction, the increase in size
between the 3d and 4d metals is much greater than between
the 4d and 5d metals
– Effective nuclear charge experienced by valence electrons in
the d-elements does not change as the nuclear charge
increases across a row
– Ionization energies of these elements increase very slowly
going across a given row
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Electronic Structure and Reactivity of the
Transition Metals
–
As seen in the following
table, going from the top
left to the bottom right
corner of the d block,
1. electronegativities
increase
2. densities and electrical
and thermal
conductivities increase
3. enthalpies of hydration
of the metal cations
decrease
–
Transition metals become
steadily less reactive and
more “noble” in character
going from left to right
across a row
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Trends in Transition Metal Oxidation States
• Similarity in ionization energies and the small increase in successive
ionization energies lead to the formation of metal ions with the same
charge for many of the transition metals, which results in extensive
horizontal similarities in chemistry that are most noticeable for the
first-row transition metals and for the lanthanides and actinides
– All first-row transition metals except Sc form stable compounds that
contain the 2+ ion
– Due to the small difference between the second and third ionization
energies for these elements, all except Zn also form stable
compounds that contain the 3+ ion
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Trends in Transition Metal Oxidation States
• Oxidation states
– The small increase in successive ionization energies causes the
transition metals to exhibit multiple oxidation states, separated
by a single electron
– Because of the slow increase in ionization potentials going
across a row, high oxidation states become progressively less
stable for the elements on the right side of the d block
– Occurrence of multiple oxidation states separated by a single
electron causes the compounds of the transition metals to be
paramagnetic, with one to five unpaired electrons
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Trends in Transition Metal Oxidation States
• The electronegativities of the first-row transition metals
increase smoothly from Sc to Cu
• The standard reduction potential Eº for the reaction
M2+(aq) + 2e–  Mº(s) becomes less negative
from Ti to Cu
• Exceptions to overall trends attributable to the stability
associated with filled and half-filled subshells
• The transition metals form cations by the initial loss of
the ns electrons of the metal, even though the ns orbital
is lower in energy than the (n–1)d subshell in the neutral
atom; therefore, all transition-metal cations possess dn
valence-electron configurations for the 2+ ions of the
first-row transition metals, as seen in the following table
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Trends in Transition Metal Oxidation States
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Trends in Transition Metal Oxidation States
• The most common oxidation states of the first-row transition metals
are shown in the following table; the second- and third-row transition
metals behave similarly but with three important differences:
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Trends in Transition Metal Oxidation States
1. The maximum oxidation states observed for the secondand third-row transition metals in Groups 3–8 increase from
+3 for Y and La to +8 for Ru and Os, corresponding to the
formal loss of all ns and (n–1)d valence electrons; going
farther to the right, the maximum oxidation state decreases,
reaching +2 for the elements of Group 12, which
corresponds to a filled (n–1)d subshell
2. Within a group, higher oxidation states become more
stable going down a group
3. Cations of the second- and third-row transition metals in
lower oxidation states (+2 and +3) are more easily oxidized
than the corresponding ions of the first-row transition
metals
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Trends in Transition Metal Oxidation States
• Binary transition-metal complexes such as the oxides
and sulfides are written with idealized stoichiometries,
but these compounds are cation deficient and never
contain a 1:1 cation:anion ratio
• Acid-base character of transition-metal oxides depend
strongly on the oxidation state of the metal and its ionic
radius
– Oxides of metals in lower oxidation states have ionic
character and tend to be basic
– Oxides of metals in higher oxidation states are more covalent
and tend to be acidic, dissolving in strong base to form
oxoanions
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Chemistry: Principles, Patterns,
and Applications, 1e
23.2 A Brief Survey of
Transition-Metal
Chemistry
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23.2 A Brief Survey of Transition-Metal
Chemistry
• Beginning with Group 3 and continuing to Group 12, the
chemistry of the transition metals will be studied
• It will be seen that the two heaviest members of each
group exhibit substantial similarities in chemical behavior
and are quite different from the lightest member
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Groups 3, 4, and 5
• Group 3 (Sc, Y, La, Ac)
– Observed trends in the properties of the Group-3 elements
similar to those of Groups 1 and 2 (see table)
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Groups 3, 4, and 5
– Due to their ns2 (n–1)d1 valence-electron configurations, the
chemistry of all four elements is dominated by the +3 oxidation
state formed by the loss of all three valence electrons
– Elements are highly electropositive metals and powerful
reductants, with La (and Ac) being the most reactive
– React with water to produce the metal hydroxide and hydrogen
gas
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Groups 3, 4, and 5
– All dissolve readily in acidic solutions to produce hydrogen gas
and a solution of the hydrated metal ion, M3+(aq)
– React with nonmetals to form compounds that are ionic in
character
– All the Group-3 elements react with air to form an oxide coating,
and all burn in oxygen to form the sesquioxides, M2O3, which
react with water or CO2 to form the corresponding hydroxides or
carbonates, respectively
– Commercial uses of the Group-3 metals are limited
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Groups 3, 4, and 5
• Group 4 (Ti, Zr, Hf)
– Have a high affinity for oxygen so all three metals
occur naturally as oxide ores that contain the metal in
the +4 oxidation state resulting from the loss of all
four ns2(n–1)d2 valence electrons; isolated by initial
conversion to the tetrachlorides followed by reduction
of the tetrachlorides with an active metal
– Group-4 elements have important applications
– Become denser, higher melting, and more
electropositive going down the column
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Groups 3, 4, and 5
– +4 oxidation state is the most important for all three metals; only
titanium has a chemistry in the +2 and +3 oxidation states
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Groups 3, 4, and 5
– Reaction of the Group-4 metals with excess halogen forms the
corresponding tetrahalides, MX4
– Titanium, the lightest element in the group also forms dihalides
and trihalides; covalent character of the titanium halides
increases as the oxidation state of the metal increases because
of increasing polarization of the anions by the cation as its
charge-to-radius ratio increases
– All three metals react with excess oxygen or the heavier
chalcogens (Y) to form the corresponding dioxides, MO2, and
dichalcogenides, MY2
– React with hydrogen, nitrogen, carbon, and boron to form
hydrides, nitrides, carbides, and borides
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Groups 3, 4, and 5
• Group 5 (V, Nb, Ta)
– Found in nature as oxide ores that contain the metals in their
highest oxidation state (+5)
– Trends in properties of the Group-5 metals are similar to
those of Group 4
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Groups 3, 4, and 5
– Only vanadium, the lightest element, has any tendency to
form compounds in oxidation states lower than +5
– All three metals react with excess oxygen to produce the
corresponding oxides in the +5 oxidation state, M2O5, in
which polarization of the oxide ions by the high-oxidationstate metal is so extensive that the compounds are covalent
in character
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Groups 3, 4, and 5
– React with the heavier chalcogens to form a complex set of
binary chalcogenides; the most important are the
dichalcogenides, MY2
– Form binary nitrides, carbides, borides, and hydrides whose
stoichiometries and properties are similar to those of
corresponding Group-4 compounds
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Groups 6 and 7
• Group 6 (Cr, Mo, W)
– Encounter a metal (Mo) that occurs naturally as a sulfide ore
rather than as an oxide
– Molybdenum is the only second- or third-row transition
element that is essential for humans
– Group 6 metals are less electropositive than those of the
three preceding groups, and the two heaviest metals are the
same size because of the lanthanide contraction
– All three elements have a total of six valence electrons,
resulting in a maximum oxidation state of +6; compounds in
the +6 oxidation state are highly covalent
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Groups 6 and 7
– The lightest element, Cr, exhibits variable oxidation states; for
Mo and W, the highest oxidation state (+6) is the most important
– Group-6 halides become more covalent as the oxidation state of
the metal increases, their volatility increases, and their melting
points decrease
– As the oxidizing strength of the halogen decreases, the
maximum oxidation state of the metal decreases
– All three metals form hexafluorides
– React with oxygen to form the covalent trioxides, which are
acidic, dissolving in base to form the corresponding oxoanions
MO42–
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Groups 6 and 7
• Group 6 (Cr, Mo, W)
– The sesquioxide of the lightest element in the group is
amphoteric
– At low pH, molybdate and tungstate form a series of
polymeric anions called isopolymetallates
– Reaction of molybdenum or tungsten with heavier
chalcogens gives binary chalcogenide phases, most of
which are nonstoichiometric and electrically conducting
– Elements of Group 6 form binary nitrides, carbides, and
borides whose stoichiometries and properties are similar to
those of the preceding groups
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Groups 6 and 7
• Group 7 (Mn, Tc, Re)
– All three Group-7 elements have seven valence electrons
and can form compounds in the +7 oxidation state; the
lightest element (Mn) exhibits multiple oxidation states, has
a low electronegativity, and is unreactive
– Reaction with less oxidizing halogens produces metals in
lower oxidation states
– Tc and Re have similar size and electronegativity; they form
high-valent oxides, called heptoxides, and form disulfides
and diselenides with layered structures
– Group-7 metals form binary nitrides, carbides, and borides
that are stable at high temperatures and exhibit metallic
properties
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Groups 6 and 7
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Groups 8, 9, and 10
• Elements of these three groups exhibit many horizontal similarities in
their chemistry, in addition to the similarities within each column
• Horizontal similarities are due to the fact that the ionization potentials
of the elements have become so large that the oxidation state
corresponding to the formal loss of all valence electrons is
encountered only rarely (Group 8) or not at all (Groups 9 and 10)
• Chemistry of all three groups is dominated by intermediate oxidation
states, especially +2 and +3 for the first-row metals (Fe, Co, and Ni)
• Heavier elements of these three groups are called precious metals
because they are rare in nature and are chemically inert
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Groups 8, 9, and 10
• Group 8 (Fe, Ru, Os)
– Chemistry of Group 8 is dominated by iron, whose high
abundance in Earth’s crust is due to the extremely high
stability of its nucleus
– Ruthenium and osmium are extremely rare elements
• Group 9 (Co, Rh, Ir)
– Cobalt is one of the least abundant of the first-row transition
metals
– Heavier elements of Group 9 are rare and are found in
combination with the heavier elements of Groups 8 and 10 in
Ni-Cu-S ores
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Groups 8, 9, and 10
• Group 10 (Ni, Pd, Pt)
– Nickel silicates are easily processed
– Palladium and platinum are rare but are more abundant than
the heavier elements of Groups 8 and 9
• Trends in Groups 8, 9, and 10
– Properties of the elements of Groups 8–10 are summarized in
the following table
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Groups 8, 9, and 10
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Groups 8, 9, and 10
– Similarities in size and electronegativity between the two
heaviest members of each group result in similarities in
chemistry
– No longer a clear correlation between the valence-electron
configuration and the preferred oxidation state
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Groups 8, 9, and 10
– Most important oxidation states are +2 and +3 for Group 8, +3
and +1 for Group 9, and +2 and +4 for Group 10
– Higher oxidation states become less stable going across the
d-block elements and more stable going down a group
1. Fe and Co form trifluorides, but Ni forms only the difluoride
2. Ru and Os form a series of fluorides
3. Hexafluorides of Rh and Ir are powerful oxidants
4. Pt is the only element in Group 10 that forms a hexafluoride
– Similar trends are observed among the oxides
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Groups 8, 9, and 10
– The tendency of the metals to form the higher oxides
decreases rapidly going farther across the d block
– Reaction of metals in Groups 8, 9, and 10 with the heavier
chalcogens is complex
1. Oxidation state of Fe, Ru, Os, Co, and Ni in their disulfides is +2
because of the disulfide ion, S22–
2. Disulfides of Rh, Ir, Pd, and Pt contain the metal in the +4
oxidation state together with isolated sulfide ions, S2–
3. Combination of highly charged cations and easily polarized
anions results in substances that are not simply ionic but have
significant covalent character
– Groups 8–10 metals form a range of binary nitrides, carbides,
and borides
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Groups 11 and 12
• Group 11 (Cu, Ag, Au)
– Coinage metals—copper, silver, and gold—occur naturally and
were probably the first metals used by ancient humans;
deposits of gold and copper are widespread and numerous,
while deposits of silver are less common
– Properties of the coinage metals are listed in the following
table
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Groups 11 and 12
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Groups 11 and 12
1. Electronegativity of gold is close to that of the nonmetals sulfur
and iodine, so the chemistry of gold is unusual for a metal
2. Have the highest electrical and thermal conductivities of all the
metals and also the most ductile and malleable
3. Have a ns1(n–1)d10 valence-electron configuration
4. Chemistry dominated by the +1 oxidation state due to loss of
the single ns electron
5. Higher oxidation states are also known because of the low
values of the second and third ionization energies
6. All three elements have significant electron affinities due to the
half-filled ns orbital in the neutral atoms
– All the Group-11 elements are unreactive, and their reactivity
decreases from Cu to Au
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Groups 11 and 12
– Copper reacts with O2 at high temperatures to produce Cu2O
and reacts with sulfur to form Cu2S
– Neither silver nor gold reacts directly with oxygen
– Silver reacts with sulfur compounds to form Ag2S
– Gold is the only metal that does not react with sulfur, nitrogen,
carbon, or boron
– All the coinage metals react with oxidizing acids, and all three
metals dissolve in basic cyanide solutions in the presence of
oxygen to form stable [M(CN)2]– ions
– Known: all the monohalides except CuF and AuF, all the
copper () halides except the iodide, the dihalide of silver, and
all the gold trihalides except the triiodide
– No binary nitrides, borides, or carbides are known
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Groups 11 and 12
• Group 12 (Zn, Cd, Hg)
– Are similar in abundance to those of Group 11 and are always
found in combination with sulfur
– Zinc and cadmium are chemically similar, so zinc ores contain
cadmium
– All three metals are commercially important
– None of the elements has a partially filled (n–1)d subshell, so
they are not strictly “transition metals”
– Much of their chemistry is similar to that of the elements that
immediately precede them in the d block
– Properties of Group-12 metals are shown in the following table
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Groups 11 and 12
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Groups 11 and 12
– More electropositive than the elements of Group 11 and have
less “noble” character
– Have lower melting and boiling points than the preceding
transition metals
– Zn and Cd are similar to each other but are very different from
the heaviest element (Hg)
– Zn and Cd are active metals and mercury is not—mercury is
the only metal that is liquid at room temperature and can
dissolve metals by forming amalgams
– All three elements have ns2(n–1)d10 valence-electron
configurations and the +2 oxidation state, corresponding to the
loss of the two ns electrons
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Groups 11 and 12
– Mercury forms a series of compounds in the +1 oxidation state
that contains the diatomic mercurous ion, Hg22+
– All the possible dihalides, MX2, are known and range from ionic
to highly covalent
– Zinc and cadmium react with oxygen to form amphoteric MO,
and mercury forms HgO within a narrow temperature range
– Zinc and cadmium dissolve in mineral acids, and mercury
dissolves only in oxidizing acids
– All three metals react with sulfur and the other chalcogens to
form the binary chalcogenides, and mercury has a high affinity
for sulfur ligands
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Chemistry: Principles, Patterns,
and Applications, 1e
23.3 Metallurgy
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23.3 Metallurgy
• Very few of the transition metals are found in nature as
the free metals; all metallic elements must be isolated
from metal oxide or metal sulfide ores
• Metallurgy is the set of processes by which metals are
extracted from their ores and converted to more useful
forms
• Metallurgy consists of three general steps:
1. Mining the ore
2. Separating and concentrating the metal or the metalcontaining compound
3. Reducing to the metal
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23.3 Metallurgy
• After an ore is mined, the first step in processing is to
crush it, because the rate of chemical reactions
increases with an increase in surface area
• Three general strategies are used to separate and
concentrate the compound(s) of interest:
1. Settling and flotation—based on differences in density
between the desired compound and impurities
2. Pyrometallurgy—uses chemical reductions at high
temperatures
3. Hydrometallurgy—employs chemical or electrochemical
reduction of an aqueous solution of the metal
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Pyrometallurgy
• In pyrometallurgy, an ore is heated with a reductant in
order to obtain the metal
• A reductant must be used that forms stable compounds
with the metal of interest
• Pyrometallurgy is used in the iron and steel industries
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Hydrometallurgy
• The most selective methods for separating metals
from their ores are based on the formation of metal
complexes
• In hydrometallurgy, metals are separated via the
formation of metal complexes by using chemical or
electrochemical reduction of an aqueous solution
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Chemistry: Principles, Patterns,
and Applications, 1e
23.4 Coordination
Compounds
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23.4 Coordination Compounds
• One of the most important properties of metallic
elements is their ability to act as Lewis acids that form
complexes with a variety of Lewis bases
• Metal complex—consists of a central metal atom or ion
that is bonded to one or more ligands; metal complexes
can be neutral, positively charged, or negatively charged
• Ligands—are ions or molecules that contain one or
more pairs of electrons that can be shared with the metal
• Electrically charged metal complexes are called
complex ions; a coordination compound contains one
or more metal complexes
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23.4 Coordination Compounds
• Coordination compounds are important for three
reasons:
1. Most of the elements in the periodic table are metals, and
almost all metals form complexes
2. Many industrial catalysts are metal complexes, and these
catalysts are important as a way to control reactivity
3. Transition-metal complexes are essential in biochemistry
a. Hemoglobin, an iron complex that transports oxygen in blood
b. Cytochromes, iron complexes that transfer electrons in cells
c. Complexes that are components of enzymes, the catalysts for
all biological reactions
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History of Coordination Compounds
• Coordination compounds have been known and used
since ancient times, but their chemical nature was
unclear
• Modern theory of coordination chemistry is based on the
work of Werner, who studied the properties of several
series of metal halide complexes with ammonia
• Data led Werner to postulate that metal ions have two
different kinds of valence:
1. A primary valence (oxidation state) that corresponds to the
positive charge on the metal ion
2. A secondary valence (coordination number) that is the total
number of ligands bound to the metal ion
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Structures of Metal Complexes
• Coordination numbers of metal ions in metal
complexes can range from 2 to 9
• The differences in energy between different
arrangements of ligands are greatest for
complexes with low coordination numbers and
decrease as the coordination number increases
• Only one or two structures are possible for
complexes with low coordination numbers, and
several different energetically equivalent
structures are possible for complexes with high
coordination numbers (n > 6)
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Structures of Metal Complexes
• Coordination number 2
– Rare for most metals, but is common for the d10 metal ions,
especially Cu+, Ag+, Au+, and Hg2+
– Based on VSEPR, these complexes have the linear L–M–L
structure
• Coordination number 3
– Encountered with d10 metal ions such as Cu+ and Hg2+
– Based on VSEPR, 3-coordinate complexes have the
trigonal planar structure
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Structures of Metal Complexes
• Coordination number 4
– Two common structures observed for 4-coordinate metal
complexes: tetrahedral and square planar
1. Tetrahedral structure is observed for all 4-coordinate
complexes of nontransition metals and d10 ions; also found
for 4-coordinate complexes of the first-row transition metals,
especially those with halide ligands
2. Square planar structures are observed for 4-coordinate
complexes of second- and third-row transition metals with d8
electron configurations, such as Rh+ and Pd2+, and are also
encountered in some complexes of Ni2+ and Cu2+
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Structures of Metal Complexes
• Coordination number 5
– This coordination number is less common than 4 and 6, but is
found in two different structures: trigonal bipyramidal and
square pyramidal
– Many 5-coordinate complexes have distorted structures that lie
somewhere between the two extremes
• Coordination number 6
– The most common
– The six ligands are at the vertices of an octahedron or a
distorted octahedron
– The only other 6-coordinate structure is the trigonal prism,
which is uncommon in simple metal complexes
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Structures of Metal Complexes
• Coordination number 7
– Uncommon and is generally encountered for only
large metals (such as the second- and third-row
transition metals, lanthanides, and actinides)
– Three different structures are known, two of which
are derived from an octahedron or a trigonal prism
by adding a ligand to one face of the polyhedron to
give a “capped” octahedron or trigonal prism; most
common is the pentagonal bipyramid
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Structures of Metal Complexes
• Coordination number 8
– Common for larger metal ions
– The simplest structure is the cube, which is rare because it
does not minimize interligand repulsive interactions
– Common structures are the square antiprism and the
dodecahedron, both of which can be generated from the cube
• Coordination number 9
– Found for larger metal ions
– Most common structure is the tricapped trigonal prism
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Stability of Metal Complexes
• The thermodynamic stability of a metal complex depends
on the properties of the ligand and the metal ion, and on
the type of bonding
• The metal-ligand interaction is an example of a Lewis
acid-base interaction
• Lewis bases can be divided into two categories:
1. Hard bases—contain small, relatively nonpolarizable donor
atoms (such as N, O, and F)
2. Soft bases—contain larger, relatively polarizable donor atoms
(such as P, S, and Cl)
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Stability of Metal Complexes
• Metal ions with the highest affinities for hard bases are
hard acids, and metal ions with the highest affinity for
soft bases are soft acids
– Hard acids are cations of electropositive metals, are relatively
nonpolarizable, and have higher charge-to-radius ratios
– Soft acids are cations of less electropositive metals, have
lower charge-to-radius ratios, and are more polarizable
• Can predict the relative stabilities of complexes formed
by the d-block metals with a remarkable degree of
accuracy by using a simple rule: Hard acids prefer to
bind to hard bases, and soft acids prefer to bind to soft
bases
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Stability of Metal Complexes
• The interaction between hard acids and hard bases is
electrostatic in nature, so the stability of complexes
involving hard acids and hard bases increases as the
positive charge on the metal ion increases and as its
radius decreases
• The stability of complexes of divalent first-row transition
metals with a given ligand varies inversely with the
radius of the metal ion
• Because a hard metal interacts with a base in the same
way as a proton—by binding to a lone pair of electrons
on the base—the stability of complexes of hard acids
with hard bases increases as the ligand becomes more
basic
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Stability of Metal Complexes
• The interaction between “soft” metals (such as the
second- and third-row transition metals and Cu+) and
“soft” bases is covalent in nature
– Most soft-metal ions have a filled or nearly filled d subshell,
which suggests that metal-to-ligand  bonding is important
– Complexes of soft metals with soft bases are much more
stable than would be predicted based on electrostatic
arguments
• The hard acid-hard base/soft acid-soft base concept
helps to explain why metals are found in nature in
different kinds of ore—due to an increase in the “soft”
character of the metals going across the first row of the
transition metals from left to right
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Stability of Metal Complexes
• Ligands
1. Monodentate—they are attached to the metal via only a
single atom
2. Bidentate—they are attached to the metal at two sites
3. Tridentate—they are attached to the metal at three sites
4. Polydentate—they are attached to the metal at several sites
• When a bidentate ligand binds to a metal, a fivemembered ring called a chelate ring is formed
• A polydentate ligand is a chelating agent, and complexes
that contain polydentate ligands are called chelate
complexes
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Stability of Metal Complexes
• Metal complexes of polydentate ligands are more stable than the
corresponding complexes of chemically similar monodentate ligands;
observation is called the chelate effect
• The stability of a chelate complex depends on the size of the chelate
rings
– For ligands with a flexible organic backbone, complexes that contain
five-membered chelate rings and that have no strain are more stable
than complexes with six-membered chelate rings, which are more stable
than complexes with four- or seven-membered rings
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Isomers of Metal Complexes
• The existence of coordination compounds with the same
formula but different arrangements of the ligands is
important in the development of coordination chemistry
– Two or more compounds that have the same formula but
different arrangements of the atom are called isomers
– Isomers have different physical and chemical properties, and it
is important to know which isomer you are dealing with if more
than one isomer is possible
– Coordination compounds exhibit the same types of isomers as
organic compounds, as well as several kinds of unique
isomers
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Isomers of Metal Complexes
• Structural isomers
– Isomers that contain the same number of atoms of each kind
but differ in which atoms are bonded to one another are called
structural isomers
– A trivial kind of isomerism consists of two compounds that
have the same empirical formula but differ in the number of
formula units present in the molecular formula
• Geometrical isomers
– Differ only in the arrangement of ligands around the metal ion
– Metal complexes that differ only in which ligands are adjacent
to one another (cis) or directly across from one another (trans)
in the coordination sphere of the metal are called geometrical
isomers
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Isomers of Metal Complexes
• Geometrical isomers are most important for
square planar and octahedral complexes
1. Square planar complexes
– Because all vertices of a square are equivalent, it does not matter
which vertex is occupied by the ligand B in a square planar MA3B
complex; only a single geometrical isomer is possible
– For an MA2B2 complex, there are two possible isomers: either the A
ligands can be adjacent to one another (cis), in which case the B
ligands must also be cis, or the A ligands can be across from one
another (trans), in which case the B ligands must also be trans; cis
and trans structures are different arrangements in space
– Square-planar complexes that contain symmetrical bidentate
ligands have only one possible structure
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Isomers of Metal Complexes
2. Octahedral complexes
– Only one structure is possible for octahedral complexes in
which only one ligand is different from the other five (MA5B)
since all six vertices of an octahedron are equivalent
– If two ligands in an octahedral complex are different from
the other four (MA4B2), two isomers are possible; the two B
ligands can be cis or trans
– Replacing another A ligand by B gives an MA3B3 complex
for which there are two isomers: fac and mer
1. Fac—the three ligands of each kind occupy opposite
triangular faces of the octahedron
2. Mer—the three ligands of each kind lie on the meridian
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Chemistry: Principles, Patterns,
and Applications, 1e
23.5 Crystal Field Theory
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23.5 Crystal Field Theory
• Crystal field theory (CFT) is a bonding model that
explains many important properties of transition-metal
complexes, including their colors, magnetism, structures,
stability, and reactivity that cannot be explained using
valence bond theory
• Central assumption of CFT is that metal-ligand
interactions are purely electrostatic in nature
• In CFT, complex formation is assumed to be due to
electrostatic interactions between a central metal ion and
a set of negatively charged ligands or ligand dipoles
arranged around the metal ion
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d-Orbital Splittings
• Crystal field theory focuses on the interaction of the five
(n–1)d orbitals with ligands arranged in a regular array
around a transition-metal ion
• According to CFT, an octahedral metal complex forms
because of the electrostatic interaction of a positively
charged metal ion with six negatively charged ligands, or
with the negative ends of dipoles associated with the six
ligands and the ligands interact with one electrostatically
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d-Orbital Splittings
• The energies of the d orbitals of a transition-metal ion
are affected by an octahedral arrangement of six
negative charges
– The five d orbitals are initially degenerate (have the same energy)
– When the six negative charges are distributed uniformly over the
surface of a sphere, the d orbitals remain degenerate, but their
energy will be higher due to repulsive electrostatic interactions
between the spherical shell of negative charge and electrons in the
d orbitals
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d-Orbital Splittings
– If the six negative charges are placed at the vertices of an
octahedron, the average energy of the d orbitals does not
change, but it does remove their degeneracy and the five d
orbitals split into two groups whose energies depend on their
orientations
– The dx2 – y2 and dz2 orbitals (the eg orbitals) point directly at the
six negative charges, which increase their energy compared
with a spherical distribution of negative charge; the dxy, dxz, and
dyz (t2g orbitals) are all oriented at a 45º angle to the coordinate
axes and point between the six negative charges, which
decreases their energy compared with a spherical distribution
of charge
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d-Orbital Splittings
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d-Orbital Splittings
• The difference in energy between the two sets of d
orbitals is called the crystal field splitting energy
– Given the symbol o, where the subscript “o” stands for
“octahedral”
• The magnitude of the splitting depends on the charge on
the metal ion, the position of the metal in the periodic
table, and the nature of the ligands
• The splitting of the d orbitals in a crystal field does not
change the total energy of the five d orbitals
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Electronic Structures of Metal Complexes
• Using the d-orbital energy-level diagram, the electronic
structures and some of the properties of transition-metal
complexes can be predicted
– Start with the Ti3+ ion, which contains a single d electron, and
proceed across the first row of the transition metals by adding a
single electron at a time
– Additional electrons are placed in the lowest-energy orbital
available while keeping their spins parallel
– For d1-d3 systems, the electrons successively occupy the three
degenerate t2g orbitals with their spins parallel, giving one, two,
and three unpaired electrons, respectively
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Electronic Structures of Metal Complexes
• Using the d-orbital energy-level diagram, the electronic
structures and some of the properties of transition-metal
complexes can be predicted
– Reaching the d4 configuration, there are two possible choices
for the fourth electron: either in one of the empty eg orbitals or
in one of the singly occupied t2g orbitals
– Placing an electron in an occupied orbital results in
electrostatic repulsions that increase the energy of the system;
this is called the spin-pairing energy (P)
– If o is less than the spin-pairing energy, then the lowestenergy arrangement has the fourth electron in one of the
empty eg orbitals; this results in four unpaired electrons and is
called a high-spin configuration—a complex with this
configuration is a high-spin complex
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Electronic Structures of Metal Complexes
– If o is greater than the spin-pairing energy, the lowest-energy
arrangement has the fourth electron in one of the occupied t2g
orbitals, which results in two unpaired electrons and is called a
low-spin configuration; a complex with this electron
configuration is called a low-spin complex
– Metal ions with the d5, d6, or d7 electron configurations can be
either high spin or low spin, depending on the magnitude of o
– Only one arrangement of d electrons is possible for metal ions
with d8–d10 electron configurations
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Electronic Structures of Metal Complexes
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Factors That Affect the Magnitude of o
• The magnitude of o dictates whether a complex
with four, five, six, or seven d electrons is high
spin or low spin, which affects its magnetic
properties, structure, and reactivity
– Large values of o yield a low-spin complex; small
values of o produce a high-spin complex
• Magnitude of o depends on three factors:
1. The charge on the metal ion
2. The principal quantum number of the metal
3. The nature of the ligand
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Factors That Affect the Magnitude of o
• Charge on the metal ion
– Increasing the charge on a metal ion has two effects:
1. The radius of the metal ion decreases
2. Negatively charged ligands are more strongly attracted to it
– Both these factors decrease the metal-ligand distance, which
causes the negatively charged ligands to interact more
strongly with the d orbitals
– The magnitude of o increases as the charge on the metal ion
increases
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Factors That Affect the Magnitude of o
• Principal quantum number of the metal
– For a series of complexes of metals from the same group in
the periodic table with the same charge and the same ligands,
the magnitude of o increases with increasing quantum
number:
o (3d) << o (4d) < o (5d)
– Increase in o with increasing principal quantum number is due
to the larger radius of valence orbitals going down a column
– Repulsive ligand-ligand interactions are important for smaller
metal ions, which results in shorter M–L distances and stronger
d-orbital-ligand interactions
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Factors That Affect the Magnitude of o
• The nature of the ligands
– For a series of chemically similar ligands, the magnitude of o
decreases as the size of the donor atom increases because
smaller, more localized charges interact more strongly with the
d orbitals of the metal ion
– A small neutral ligand with a highly localized lone pair results in
larger o values
– The experimentally observed order of the crystal field splitting
energies produced by different ligands is called the
spectrochemical series
1. Strong-field ligands interact strongly with the d orbitals of the
metal ions and give a large o
2. Weak-field ligands interact more weakly and give a smaller o
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Colors of Transition-Metal Complexes
• The striking colors exhibited by transition-metal complexes are
caused by the excitation of an electron from a lower-lying d orbital to
a higher-energy d orbital, which is called a d-d transition
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Colors of Transition-Metal Complexes
• For a photon to affect the transition, its energy must be equal to the
difference in energy between the two d orbitals, which depends on
the magnitude of o
• Color that is observed is due to transmitted or reflected light that is
complementary in color to the light that is absorbed
• The energy of a photon of light is inversely proportional to its
wavelength, so the color of a complex depends on the magnitude of
o, which depends on the structure of the complex
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Crystal Field Stabilization Energies
• If the lower-energy set of d orbitals (the t2g orbitals) is
selectively populated by electrons, then the stability of
the complex increases
• The additional stabilization of a metal complex by
selective population of the lower-energy d orbitals is
called its crystal field stabilization energy (CFSE)
• CFSE of a complex can be calculated by multiplying the
number of electrons in t2g orbitals by the energy of those
orbitals, multiplying the number of electrons in eg orbitals
by the energy of those orbitals, and summing the two
• CFSE is highest for low-spin d6 complexes
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Crystal Field Stabilization Energies
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Crystal Field Stabilization Energies
• CFSEs are important for two reasons:
1. The existence of CFSE accounts for the difference between
experimentally measured values for bond energies in metal
complexes and values calculated based solely on electrostatic
interactions
2. CFSEs represent large amounts of energy, which has
important chemical consequences
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Tetragonal and Square-Planar Complexes
• If two trans ligands in an octahedral complex are either
chemically different from the other four or a different
distance from the metal than the other four, the result is
a tetragonally distorted octahedral complex
• Moving the two axial ligands away from the metal ion
along the z axis gives an elongated octahedral complex
and eventually produces a square-planar complex
• Axial elongation causes the dz2, dxz, and dyz orbitals to
decrease in energy and the dx2–y2 and dxy orbitals to
increase in energy; change in energy is not the same for
all five d orbitals
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Tetrahedral Complexes
• In a tetrahedral complex, none of the five d orbitals
points directly at or between the ligands
• The dxy, dxz, and dyz orbitals interact more strongly with
ligands than do the dx2 – y2 and dz2 orbitals, so the order of
orbital energies in a tetrahedral complex is the opposite
of the order in an octahedral complex
• The splitting of the energies of the orbitals in a
tetrahedral complex, o, is smaller than that in an
octahedral complex for two reasons:
1. The d orbitals interact less strongly with the ligands in a
tetrahedral arrangement
2. There are only four negative charges rather than six, which
decreases the electrostatic interactions
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Consequences of d-Orbital Splitting
• The splitting of the d orbitals because of their
interactions with the ligands in a complex has important
consequences for the chemistry of transition-metal
complexes, these can be divided into structural effects
and thermodynamic effects
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Consequences of d-Orbital Splitting
• Structural effects—two major kinds: effects on the ionic
radius of metal ions with regular octahedral or
tetrahedral geometries, and structural distortions that are
observed for specific electron configurations
1. Ionic radii
– A plot of the ionic radii of the divalent first-row transition metal
ions versus atomic number shows that only dº , high-spin d5,
and d10 fall on the smooth curve calculated based on the
effective nuclear charge; this assumes that the distribution
of d electrons is spherically symmetrical, which is effective at
screening the ligands from the nuclear charge, giving a larger
ionic radius, and making the metal-ligand distances longer
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Consequences of d-Orbital Splitting
– All the other divalent ions fall below the curve because they have
asymmetrical distributions of d electrons; this makes a metal ion
smaller because of poor shielding of the ligands from the nuclear
charge and the metal-ligand distance is short
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Consequences of d-Orbital Splitting
2. The John-Teller effect
– Refers to the distortions observed for d4 and d9 complexes;
occurs in systems that have an odd number of electrons in the eg
orbitals
– Electrons can occupy either one of two degenerate eg orbitals
and have a degenerate ground state
– States that such systems are not stable and that they will
undergo a distortion that makes the complex less symmetrical
and splits the degenerate states, which decreases the energy of
the system
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Consequences of d-Orbital Splitting
• Thermodynamic effects—CFSEs are important factors
in determining the magnitude of hydration energies,
lattice energies, and other thermodynamic properties of
the transition metals
1. Hydration energies
– Hydration energy of a metal ion is defined as the change in
enthalpy for the reaction
M2+(g) + H2O(l)  M2+ (aq)
– Cannot be measured directly but can be calculated from
experimentally measured quantities using thermochemical
cycles
– CFSEs are responsible for the differences between the
measured and calculated values of hydration energies
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Consequences of d-Orbital Splitting
2. Lattice energies
– Defined as the negative of the enthalpy change for the reaction
M2+ (g) + 2Cl– (g)  MCl2 (s)
– Determined indirectly by using a thermochemical cycle
– Transition-metal dichlorides are more stable due to CFSE
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Chemistry: Principles, Patterns,
and Applications, 1e
23.6 Transition Metals in
Biology
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Uptake and Storage of Transition Metals
• There are three possible dietary levels for any
essential element: deficient, optimal, and toxic
– If the concentration of an essential element in the diet is
too low, an organism must be able to extract the element from
the environment and concentrate it
– If the concentration of an essential element in the diet is too
high, an organism must be able to limit its intake to avoid
toxic effects
– Organisms must be able to store essential elements for future
use
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Uptake and Storage of Transition Metals
• There are three distinct steps involved in transition-metal
uptake:
1. Mobilization—the metal must be “mobilized” from the environment
and brought into contact with the cell in a form that can be absorbed
2. Transport—the metal must be transported across the cell membrane
into the cell
3. Transfer—the element must be transported to its point of utilization
within the cell or to other cells within the organism
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Uptake and Storage of Transition Metals
• Process of iron uptake
– To overcome the insolubility of Fe(OH)3, bacteria use organic
ligands called siderophores (cyclic compounds that use
bidentate ligands) that have high affinity for Fe() and are
secreted into the surrounding medium to increase the total
concentration of dissolved iron
– The iron-siderophore complex is absorbed by the cell, and the
iron is released by reduction to Fe()
– Mammals use the low pH of the stomach to increase the
concentration of dissolved iron
– Iron is absorbed in the intestine, where it forms an Fe()
complex with a protein called transferrin that is transferred to
other cells for immediate use or storage in the form of ferritin
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Metalloproteins and Metalloenzymes
• Metalloprotein is a protein that contains one or more
metal ions tightly bound to amino acid side chains
• Metalloenzyme is a metalloprotein that catalyzes a
chemical reaction
• All metalloenzymes are metalloproteins, but the
converse is not true
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Electron-Transfer Proteins
• Proteins whose function is simply to transfer electrons
from one place to another are called electron-transfer
proteins
– They do not catalyze a chemical reaction, so they
are not enzymes
– They are biochemical reductants or oxidants that are
consumed in an enzymatic reaction
– General reaction for an electron-transfer protein is
Mn+ + e–  M(n–1)+
– Many transition metals can exist in more than one
oxidation state, so electron-transfer proteins usually
contain one or more metal ions that can undergo a
redox reaction
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Electron-Transfer Proteins
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Electron-Transfer Proteins
• Incorporating a metal ion into a protein has three
important biological consequences:
1. The protein environment can adjust the redox potential, Eº´, of
the metal ion over a large potential range, whereas the redox
potential of the simple hydrated metal ion, Mn+(aq), is ssentially
fixed
2. The protein can adjust the structure of the metal complex to
ensure that electron transfer is rapid
3. The protein environment provides specificity, ensuring that the
electron is transferred to only the desired site
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Electron-Transfer Proteins
• Three important classes of metalloproteins transfer
electrons: blue copper proteins, cytochromes, and ironsulfur proteins, which generally transfer electrons at high,
intermediate, and low potentials, respectively
• Although these electron-transfer proteins contain
different metals with different structures, they are all
designed to ensure rapid electron transfer to and from
the metal
• When the protein collides with its physiological oxidant
or reductant, electron transfer can occur before the two
proteins diffuse apart; the metal sites in the oxidized and
reduced forms of the protein must have similar structures
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Reactions of Small Molecules
• Small molecules such as O2, N2, and H2 do not react with
organic compounds under ambient conditions, but they
do react with many transition-metal complexes
• All organisms use metalloproteins to bind, transport, and
catalyze the reactions of these molecules
• Hemoglobin, hemerythrin, and hemocyanin, which
contain heme iron, nonheme iron, and copper,
respectively, are used by different kinds of organisms to
bind and transfer O2
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Enzymes Involved in Oxygen Activation
• Many of the enzymes involved in the biological reactions
of oxygen contain metal centers with structures that are
similar to those used for O2 transport
• Many of these enzymes also contain metal centers that
are used for electron transfer, which have structures
similar to those of the electron-transfer proteins
• Two important enzymes that insert oxygen into an
organic molecule are dioxygenases and methane
monooxygenase
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Metal Ions as Lewis Acids
• Reactions catalyzed by metal ions that do not change
their oxidation states during the reaction are group
transfer reactions, in which a group is transferred
• Metalloenzymes use transition-metal ions as Lewis acids
to catalyze group transfer reactions
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Enzymes That Use Metals to Generate
Organic Radicals
• An organic radical is an organic species that contains
one or more unpaired electrons
• Organic radicals are essential components of several
important enzymes, almost all of which use a metal ion
to generate the organic radical within the enzyme
• These enzymes are involved in the synthesis of
hemoglobin and DNA, and they are the targets of
pharmaceuticals for the treatment of diseases
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Enzymes That Use Metals to Generate
Organic Radicals
• Some metalloenzymes use homolytic cleavage of the
cobalt-carbon bond in certain derivatives of vitamin B12
to generate an organic radical that can abstract a
hydrogen atom and thus cause molecular
rearrangements to occur
• The metal is not involved in the actual catalytic reaction;
it only provides the enzyme with a convenient
mechanism for generating an organic radical, which
does the actual work
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24
Organic
Compounds
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CHAPTER OBJECTIVES
• To learn how the three-dimensional structure of organic
compounds with the same molecular formula can vary
• To understand how variations in structure can lead to
differences in the reactivity of related organic
compounds
• To become familiar with the common classes of organic
reactions
• To understand the general properties of functional
groups and their differences in reactivity
• To be able to identify the common structural units found
in important biologically relevant molecules
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Chemistry: Principles, Patterns,
and Applications, 1e
24.1 Functional Groups
and Classes of Organic
Compounds
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24.1 Functional Groups and Classes of
Organic Compounds
• Organic compounds are covalent compounds composed
primarily of carbon and hydrogen
• Carbon is unique among the elements in its ability to
catenate, forming long chains and cyclic structures in a
wide variety of compounds
• Functional groups are structural units that determine
the chemical reactivity of a molecule under a given set of
conditions
– Can consist of a single atom or a group of atoms
– Organic compounds are classified into several major
categories based on the functional groups they contain
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24.1 Functional Groups and Classes of
Organic Compounds
• The following table summarizes five families, gives
examples of compounds that contain each functional
group, and lists the suffix or prefix used in the systematic
nomenclature of compounds that contain each functional
group
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24.1 Functional Groups and Classes of
Organic Compounds
1. First family is the hydrocarbons, which include alkanes, with
the general molecular formula CnH2n+2 where n is an integer;
alkenes represented by CnH2n; alkynes represented by CnH2n–2;
and arenes (CnHn)
2. Second family is the halogen-substituted alkanes, alkenes,
and arenes, which include the alkyl halides and aryl halides
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24.1 Functional Groups and Classes of
Organic Compounds
3. Third family is the oxygen-containing organic compounds,
which are divided into two main types:
a. Those that contain at least one C–O single bond, which
include alcohols, phenols, and ethers
b. Those that contain a carbonyl group (> CO), which include
aldehydes, ketones, and carboxylic acids
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24.1 Functional Groups and Classes of
Organic Compounds
4. Fourth family is the carboxylic acid derivatives; these are
compounds in which the H atom on the –CO2H functional
group is replaced either by an alkyl group, producing an ester,
or by an amine, forming an amide
5. Fifth family is the nitrogen-containing organic compounds;
these include amines, nitriles (which have a CN triple bond)
and nitro compounds (which contain the NO2 group)
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24.1 Functional Groups and Classes of
Organic Compounds
• Systematic and common nomenclature of
organic compounds
– In the systematic nomenclature of organic compounds, the
positions of substituents are indicated using the lowest
numbers possible to identify their locations in the carbon chain
of the parent compound
– Many organic compounds are referred to by common names
rather than by systematic names
1. Common nomenclature uses the prefix form- for a compound that
contains no carbons other than those in the functional group
2. Uses acet- for those that have one carbon atom in addition
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24.1 Functional Groups and Classes of
Organic Compounds
– In the systematic nomenclature of aromatic compounds, the
positions of groups attached to the aromatic ring are
indicated by numbers, starting with 1 and proceeding around
the ring in the direction that produces the lowest possible
numbers
– In common nomenclature, the prefixes ortho-, meta-, and
para- are used to describe the relative positions of groups
attached to an aromatic ring
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Chemistry: Principles, Patterns,
and Applications, 1e
24.2 Isomeric Variations
in Structure
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24.2 Isomeric Variations in Structure
• Isomers are different compounds that have the same
molecular formula
• Three main types of isomers:
1. Conformational
2. Constitutional (structural)
3. Stereoisomers
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Conformational Isomers
• The C–C single bonds in alkanes are formed by the
overlap of an sp3 hybrid orbital on one carbon atom with
an sp3 hybrid orbital on another carbon atom, forming a 
bond
• Differences in three-dimensional structure resulting from
rotation about a  bond are called differences in
conformation, and each different arrangement is called a
conformational isomer
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Conformational Isomers
• Differences between the conformations are
depicted in drawings called Newman projections
– A Newman projection represents the view along a C–C bond
axis, with the carbon that is in front shown as a point and the
carbon that is bonded to it shown as a circle; the C–H bonds to
each carbon positioned at 120º from each other; the
hydrogen atoms nearest the viewer are shown bonded to the
front carbon, and the hydrogen atoms farthest from the viewer
are shown bonded to the circle
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Conformational Isomers
– In one extreme, the eclipsed conformation, the C–H bonds on
adjacent carbon atoms are parallel and lie in the same plane
– In the other extreme, the staggered conformation, the
hydrogen atoms are positioned as far from one another as
possible
– Rotation about the C–C bond produces an infinite number of
conformations between these two extremes with the staggered
conformation being the most stable because electrostatic
repulsion between the hydrogen atoms on adjacent carbons is
minimized
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Conformational Isomers
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Conformational Isomers
• Newman projections are useful for predicting the stability
of conformational isomers
– The eclipsed conformation is higher in energy than the
staggered conformation because of electrostatic repulsions
between hydrogen atoms
– The staggered conformation is the most stable because
electrostatic repulsion between the hydrogen atoms on
adjacent carbons is minimized
• Longer-chain alkanes can also be represented by
Newman projections and rotation can occur about each
C–C bond in the molecule; Newman projections are
useful for revealing steric barriers to rotation at a
particular C–C bond due to the presence of bulky
substituents
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Constitutional (Structural) Isomers
• Constitutional (structural) isomers differ in
the connectivity of the atoms
– The two alcohols, 1–propanol and 2–propanol, have the same
molecular formula (C3H8O), but the position of the –OH group
differs, which causes differences in their physical and
chemical properties
• In the conversion of one constitutional isomer to
another, at least one bond must be broken and
reformed at a different position in the molecule
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Stereoisomers
• Stereoisomers are molecules that have the same connectivity but
whose component atoms have different orientations in space
• Two types of stereoisomers:
1. Geometric isomers differ in the relative placement of substituents in a
rigid molecule; members of an isomeric pair are either cis or trans,
with interconversion between the two forms requiring breaking and
reforming one or more bonds; their structural differences causes
them to have different physical and chemical properties and to exist as
two distinct chemical compounds
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Stereoisomers
2. Optical isomers are molecules that are mirror
images but cannot be superimposed on one
another in any orientation
a. Optical isomers have identical physical
properties, although their chemical properties
may differ
b. Molecules that are nonsuperimposable mirror
images of each other are said to be chiral; an
achiral object is one that can be superimposed
on its mirror image
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Stereoisomers
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Stereoisomers
• Most organic molecules that are chiral have at least one
carbon atom that is bonded to four different groups
– This carbon is designated by an asterisk in structural drawings
and is called a chiral center, chiral carbon atom, asymmetric
carbon atom, stereogenic center, or stereocenter
• A molecule and its nonsuperimposable mirror image are
called enantiomers
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Stereoisomers
• Optical activity of enantiomers
– Enantiomers have identical density, melting and boiling points,
color, and solubility in most solvents
– Differ in their interaction with plane-polarized light, which
consists of electromagnetic waves oscillating in a single plane
– If plane-polarized light is passed through a solution, the
electromagnetic radiation interacts with the solute and solvent
molecules
– If the solution contains an achiral compound, the plane-polarized
light enters and leaves the solution unchanged because achiral
molecules cause it to rotate in random directions and the solute
is optically inactive
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Stereoisomers
– If the solution contains a single enantiomer of a chiral compound,
the plane-polarized light is rotated in only one direction, and the
solute is said to be optically active; clockwise rotation is called
dextrorotatory and is indicated in the compound’s name by (+); a
counterclockwise rotation is called levorotatory and is designated
by (–)
– Magnitude of the rotation of plane-polarized light is directly
proportional to the number of chiral molecules in the solution and
depends on their detailed molecular structure, the temperature,
and the wavelength of the light
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Stereoisomers
– Every chiral compound has a specific rotation
defined as the amount (in degrees) by which the
plane of polarized light is rotated when the light is
passed through a solution containing 1 g of solute
per 1 mL of solvent in a tube 10 cm long
– A chiral solution that contains equal concentrations
of a pair of enantiomers is called a racemic mixture,
where the rotations cancel one another and the
solution is optically inactive
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Stereoisomers
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Stereoisomers
• Interactions of enantiomers with other chiral
molecules
– In living organisms, every molecule with a stereocenter is found
as a single enantiomer, not a racemic mixture
– At the molecular level, our bodies are chiral and interact
differently with the individual enantiomers of a particular
compound
– Only one enantiomer of a chiral substance interacts with a
particular receptor, initiating a response; the other enantiomer
may not bind at all, or it may bind to another receptor, producing
a different response
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Stereoisomers
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Chemistry: Principles, Patterns,
and Applications, 1e
24.3 Reactivity of
Organic Molecules
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24.3 Reactivity of Organic Molecules
• The reactivity of a molecule is affected by the degree of
substitution of the carbon bonded to a functional group;
the carbon is designated as primary, secondary, or
tertiary
– Primary carbon is bonded to only one other carbon and a
functional group
– A secondary carbon is bonded to two other carbons and a
functional group
– A tertiary carbon is bonded to three other carbons and a
functional group
• Identifying the transient species formed in a chemical
reaction, some of which are charged, enables chemists
to predict the mechanism and products of the reaction
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Reactive Intermediates
• When cleaving a C–H bond, the most common species
formed is C+, called a carbocation, which has only six
valence electrons and is electron deficient
– A carbocation is an electrophile, a species that needs
electrons to complete its octet
– A tertiary carbocation is more stable than one that is primary
because it increases electron density at the carbocation
• Adding one electron to a carbocation produces a neutral
species called a free radical, which is electron deficient
and is an electrophile; free radicals can be stabilized by
carbon substituents that can donate electron density, so
a tertiary free radical is more stable than a primary one
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Reactive Intermediates
• Adding an electron to a free radical produces a
carbanion, a negatively charged carbon with eight
valence electrons
– A carbanion is a nucleophile, an electron-rich species that has
a pair of electrons available to be shared with another atom
– Carbanions are destabilized by groups that donate electrons, so
their reactivity is the opposite of carbocations and free radicals;
therefore, a tertiary carbanion is less stable than a primary one
• Electrophiles seek to gain electrons and have a strong
tendency to react with nucleophiles, which are negatively
charged species or substances with lone pairs of
electrons
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Reactive Intermediates
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Chemistry: Principles, Patterns,
and Applications, 1e
24.4 Common Classes of
Organic Reactions
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24.4 Common Classes of Organic
Reactions
• Five common types of organic reactions:
1. Substitution
2. Elimination
3. Addition
4. Free-radical reactions
5. Oxidation-reduction reactions
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Substitution
• In a substitution reaction, one atom or group of atoms
in a substance is replaced by another atom or group of
atoms from another substance
• A typical substitution reaction is the reaction of hydroxide
ion with methyl chloride
CH3Cl + OH–  CH3OH + Cl–
– Methyl chloride has a polar C–Cl bond, so the carbon atom has
a partial positive charge
– Electronegative Cl atom is replaced by another electronegative
species that is a stronger nucleophile, OH–
– Reactions of this type are called nucleophilic substitution
reactions; for this reaction to occur, the nucleophilic
reactant must possess a pair of electrons and have a greater
affinity for the electropositive carbon atom than the original
substituent does
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Elimination
• Reactions in which adjacent atoms are removed, or
“eliminated,” from a molecule with the formation of a
multiple bond and a small molecule are called
elimination reactions
• General form:
A
B
CH2–CH2  CH2CH2 + A–B
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Addition
• A reaction in which the components of a species A–B are
added to adjacent atoms across a carbon-carbon
multiple bond is called an addition reaction
• An example is the reaction of HCl with ethylene to give
chloroethane:
HCl + CH2CH2  CH3CH2Cl
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Addition
• Addition reaction
– Although a multiple bond is stronger than a single bond, the 
bonds of the multiple bond are weaker than the  bond
– The high electron density located between multiple-bonded
carbon atoms causes alkenes and alkynes to behave like
nucleophiles, where nucleophilic attack occurs from the more
weakly bound  electrons; therefore alkenes and alkynes are
regarded as functional groups
– Nucleophilic attack occurs on the H+ atom of the polar HCl
bond, producing a carbon with only three bonds, a carbocation
– In a second nucleophilic attack, Cl–, the electrophile, attacks the
carbocation
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Addition
• Alcohols are produced by addition reactions
– Initial attack by the  bond of an alkene on a H+ of H3O+
produces a carbocation; which undergoes nucleophilic attack by
a lone pair of electrons from H2O, forming the alcohol
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Free-Radical Reactions
• Many important organic reactions involve free radicals,
and the best known is the reaction of a saturated
hydrocarbon with a halogen:
CH3CH3 + Br2  CH3CH2Br + HBr
• Free radical reactions occur in three stages: initiation,
propagation, and termination
– At high temperature or in the presence of light, the weak Br–Br
bond is broken in an initiation step that produces a number of
Br atoms
– During propagation, a bromine atom attacks ethane, producing
a free radical, which then reacts with another bromine molecule
to produce ethyl bromide; the sum of the propagation steps
corresponds to the overall balanced equation for the reaction
– Three possible termination steps: the combination of two
bromine atoms, of two ethyl radicals, or of an ethyl and a
bromine radical
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Oxidation-Reduction Reactions
• Oxidation-reduction reactions are common in organic
chemistry and can be identified by changes in the
number of oxygens at a particular position in the
hydrocarbon skeleton or in the number of bonds
between carbon and oxygen at that position
– An increase in either is an oxidation, whereas a decrease is a
reduction
– An increase in the number of hydrogens in a hydrocarbon is an
indication of a reduction
– In compounds with a carbon-nitrogen bond, the number of
bonds between the C and N atoms increases as the oxidation
state of the carbon increases
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Chemistry: Principles, Patterns,
and Applications, 1e
24.5 General Properties
and Reactivity of
Functional Groups
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24.5 General Properties and Reactivity of
Functional Groups
• The functional groups characteristic of each class of
organic compounds determine the general properties
and reactivity of that class
• There are strong connections among the structure,
physical properties, and reactivity for the compounds
that contain the major functional groups
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Alkanes, Alkenes, and Alkynes
• Alkanes
– Boiling points of alkanes increase smoothly with increasing
molecular mass and are similar to those of the corresponding
alkenes and alkynes because of similarities in molecular mass
between analogous structures
– Melting points of alkanes, alkenes, and alkynes with similar
molecular masses show a wider variation because melting
points depend on how the molecules stack in the solid state
and are therefore sensitive to small differences in structure
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Alkanes, Alkenes, and Alkynes
– Consist of C–C and C–H bonds that are strong, not polar, and
not easily attacked by nucleophiles or electrophiles, so
reactivity is limited
– Undergo catalytic cracking, which converts straight-chain
alkanes to highly branched alkanes
– Catalytic cracking is an example of a pyrolysis reaction, in
which the weakest bond is cleaved at high temperature,
producing a mixture of free radicals
– Free radicals are also produced during the combustion of
alkanes, with CO2 and H2O as the final products
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Alkanes, Alkenes, and Alkynes
• Alkenes
– The multiple bond of an alkene produces geometric isomers (cis
and trans)
– Cis and trans isomers of alkenes behave as distinct compounds
with different chemical and physical properties
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Alkanes, Alkenes, and Alkynes
• Alkynes
– Terminal alkynes are alkynes in which the triple bond is located at one
end of a carbon chain and contain a hydrogen atom attached directly
to a triply bonded carbon (R–CC–H)
– The hydrogen atom can be removed as H+, forming an acetylide ion
(R–CC–)
– Acetylide ions are potent nucleophiles used for making longer carbon
chains by a nucleophilic substitution reaction
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Alkanes, Alkenes, and Alkynes
• Alkenes and alkynes
– Rotation about the carbon-carbon multiple bonds of alkenes
and alkynes cannot occur without breaking a  bond, which
constitutes a large energy barrier to rotation
– Alkenes and alkynes are prepared by elimination reactions
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Arenes
• Most arenes with a single six-membered ring are volatile
liquids
• Some arenes with substituents on the ring are solids at
room temperature
• Undergo substitution rather than elimination because of
enhanced stability from delocalization of their  electron
density; the –H on the arene is replaced by a group –E,
such as –NO2, –SO3H, a halogen, or an alkyl group
• Are poor nucleophiles
• Many substituted arenes have potent biological activity
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Alcohols and Ethers
• Alcohols
– Have “bent” structures and are able to hydrogen-bond
– Have higher boiling points than alkanes or alkenes of
comparable molecular mass
– Good solvents for organic compounds
– Prepared by the addition of water to the carbons of a double
bond or by substitution of an alkyl halide by hydroxide, a
potent nucleophile
– Also prepared by the reduction of compounds containing the
carbonyl functional group (> CO)
– Are classified as primary, secondary, or tertiary, depending on
whether the –OH group is bonded to a primary, secondary, or
tertiary carbon
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Alcohols and Ethers
– Undergo two major types of reactions: those involving
cleavage of the O–H bond, which produces an acid, and
those involving cleavage of the C–O bond occurring under
acidic conditions where the –OH is first protonated followed
by a nucleophilic substitution
– Phenols are more acidic than alcohols because of
interactions between the oxygen atom and the ring
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Alcohols and Ethers
• Ethers
– Have “bent” structures
– Good solvents for organic compounds
– Prepared by a substitution reaction in which the highly
nucleophilic alkoxide ion (RO–) attacks the carbon of the
polarized C–X bond of an alkyl halide (R´ X)
– Unreactive because they lack the –OH unit
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Aldehydes and Ketones
• Aldehydes and ketones
– Contain the carbonyl functional group
– Have higher boiling points than alkanes or alkenes of
comparable molecular mass because of strong intermolecular
interactions
– Prepared by the oxidation of alcohols
– Characterized by nucleophilic attack at the carbon atom of the
carbonyl functional group and electrophilic attack at the
oxygen atom
– React with organometallic compounds that contain stabilized
carbanions such as the Grignard reagents (RMgX, where X =
Cl, Br, ), which convert the carbonyl functional group to an
alcohol and lengthen the carbon chain
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Aldehydes and Ketones
– Can be prepared by reducing a carboxyl group (–CO2H) to a
carbonyl group, which requires a good reducing agent
– Aromatic aldehydes have intense and characteristic flavors and
aromas, and many ketones also have intense aromas
– Ketones are found in hormones responsible for sex
differentiation in humans
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Carboxylic Acids
• Carboxylic acids
– Pungent odors
– High boiling points due to strong hydrogen-bonding
interactions between molecules
– Four lightest carboxylic acids are miscible with water, but as
the alkyl chain lengthens, their solubility in water decreases
– Compounds that contain the carboxyl functional group are
weakly acidic because of delocalization of the  electrons,
which causes them to lose a proton and form the carboxylate
anion
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Carboxylic Acids
– Prepared from the oxidation of alcohols and aldehydes or
through the reaction of a Grignard reagent with CO2, followed
by acidification
– Reactions of carboxylic acids are dominated by their polar
carboxyl group and their acidity
– Reactions with strong bases produce carboxylate salts
– Less susceptible to nucleophilic attack due to delocalization of
 bonding over three atoms (O–C–O)
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Carboxylic Acid Derivatives
• Substitution of the –OH of a carboxylic acid produces
derivative compounds with different tendencies to
participate in resonance with the CO functional group
• Resonance structures have significant effects on the
reactivity of carboxylic acid derivatives, but their
influence varies, being least important for halides and
most important for the nitrogen of amides
• Two important carboxylic acid derivatives are esters and
amides
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Carboxylic Acid Derivatives
• Esters
– Have the general formula RCO2R´, where R and R´ can be
any alkyl or aryl group
– Prepared by the reaction of an alcohol (R´OH) with a
carboxylic acid (RCO2H) in the presence of a catalytic amount
of strong acid (an electrophile); this protonates the doubly
bonded oxygen atom of the carboxylic acid (a nucleophile) to
give a species that is more electrophilic than the parent
carboxylic acid
– The nucleophilic oxygen atom of the alcohol attacks the
electrophilic carbon atom of the carboxylic acid and a new C–
O bond is formed
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Carboxylic Acid Derivatives
– General overall reaction
OH+
R–C
OH
O
+
R´OH  R–C
+ H 2O
OR´
– If an ester is heated with water in the presence of a
strong acid or base, the reverse reaction will occur,
producing the parent alcohol, R´OH, and either the
carboxylic acid, RCO2H (under acidic conditions), or
the carboxylate anion, RCO2– (under basic conditions)
– Have sweet aromas
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Carboxylic Acid Derivatives
•
Amides
– The two substituents on the amide nitrogen can be
hydrogen atoms, alkyl groups, aryl groups, or any
combination of two of those species
O
R1–C–N–R2
R3
– Prepared by the nucleophilic reaction of amines with
other, more electrophilic carboxylic acid derivatives,
such as esters
– Unreactive because of  bonding interactions between
the lone pair of electrons on nitrogen and the carbonyl
group, which inhibits free rotation about the C–N bond
– Stability of amide bond is important in biology because
they form the backbones of peptides and proteins
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Amines
• Amines
– Derivatives of ammonia in which one or more hydrogen
atoms have been replaced by alkyl or aryl groups
– Analogous to alcohols and ethers
– Classified as primary, secondary, or tertiary, depending on
the number of alkyl bonded to nitrogen
1. Primary amines—the nitrogen is bonded to two hydrogen
atoms and one alkyl group
2. Secondary amines—the nitrogen is bonded to hydrogen and
two alkyl groups
3. Tertiary amines—the nitrogen is bonded to three alkyl
groups
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Amines
– Ammonia and simple amines have lower boiling points than
does water
– Primary amines have boiling points intermediate between
those of the corresponding alcohol and alkane
– Secondary and tertiary amines have lower boiling points than
primary amines of comparable molecular mass
– Tertiary amines form cations in which all four H atoms are
replaced by alkyl groups and are called quaternary
ammonium salts, which can be chiral if all four substituents
are different
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Amines
– Alkylamines can be prepared by nucleophilic substitution
reactions of polar alkyl halides with ammonia or other amines
– Reactions of amines are dominated by two properties: the
ability of amines to act as weak bases and their tendency to
act as nucleophiles, both resulting from the lone pair of
electrons on the nitrogen atom
1. Amines behave as bases by accepting a proton from
an acid to form an ammonium salt
2. Amines can react with any electrophile
– Aryl amines are weaker bases than alkylamines because the
lone pair of electrons on nitrogen interacts with the  bonds of
the aromatic ring
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Chemistry: Principles, Patterns,
and Applications, 1e
24.6 The Molecules of
Life
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24.6 The Molecules of Life
• All the functional groups described are found in the
organic molecules that constitute and maintain every
living organism on Earth
• The most abundant substances found in living systems
belong to four major classes:
1. Proteins
2. Carbohydrates
3. Lipids
4. Nucleic acids
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Proteins
• Proteins are biologically active polymers formed from
amino acids linked together by amide bonds; in addition
to an amine group and a carboxylic acid group, each
amino acid contains a characteristic R group
– The nature of the R group determines the particular chemical
properties of each amino acid
• All the amino acids found in proteins are chiral
compounds except glycine, which suggests that their
interactions with other chiral compounds are selective
• Some proteins are enzymes that catalyze biological
reactions
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Carbohydrates
• Carbohydrates are the most abundant of the organic
compounds found in nature; they constitute a substantial
portion of food that is consumed to provide energy
needed to support life
• Carbohydrates are polyhydroxy aldehydes or
polyhydroxy ketones
– The simplest carbohydrates consist of unbranched chains of
three to eight carbon atoms; one carbon is a carbonyl carbon
and the others are bonded to hydroxyl groups
– The structure of a carbohydrate can be drawn either as a
hydrocarbon chain, known as a Fisher projection, or as a ring,
or cyclic form, called a Haworth projection
– The two cyclic forms in a Haworth projection are called
anomers
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Carbohydrates
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Carbohydrates
•
Carbohydrates are classified according to the
number of single saccharide units they contain
1. The simplest are monosaccharides
– Contain several chiral carbons and exist in several
isomeric forms
– An example is glucose
2. A disaccharide consists of two linked
monosaccharide units and an example is sucrose
3. A trisaccharide is three linked monosaccharide
units
4. Polysaccharides contain more than 10
monosaccharide units
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Carbohydrates
•
Because carbohydrates have a carbonyl functional
group and several hydroxyl groups, they can undergo a
variety of biochemically important reactions
– The carbonyl group can be oxidized to form a carboxylic acid or
can be reduced to form an alcohol
– The hydroxyl groups can undergo substitution reactions,
resulting in derivatives of the original compound
– Can eliminate hydroxyl groups, producing alkenes
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Carbohydrates
• Two familiar polysaccharides are starch and cellulose,
which both hydrolyze to produce thousands of glucose
units and differ only in the connection between glucose
units and the amount of branching in the molecule
– Starches are branched or unbranched
– Cellulose, the structural material of plants, is unbranched and
cannot be digested by humans
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Lipids
• Lipids
– Characterized by their insolubility in water
– Form a family of compounds that includes fats, waxes, vitamins,
and steroids
– Fatty acids are the simplest lipids and have a long hydrocarbon
chain that ends with a carboxylic acid functional group
1. Saturated fatty acids—the hydrocarbon chains contain only C–C
single bonds that stack in a regular array
2. Unsaturated fatty acids—have a single double bond in the
hydrocarbon chain (monounsaturated) or more than one double
bond (polyunsaturated); double bonds give fatty acid chains
a kinked structure, which prevents the molecules from
packing tightly
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Lipids
• Unsaturated fatty acids
– Melting point lower than that of a saturated fatty acid of
comparable molecular mass
– Double bonds can be hydrogenated in an addition reaction that
produces a saturated fatty acid or oxidized to produce an
aldehyde or carboxylic acid
– Are the starting compounds for the biosynthesis of
prostaglandins, hormonelike substances
• Waxes are esters produced by nucleophilic attack of an
alcohol on the carbonyl carbon of a long-chain carboxylic
acid
• Triacylglycerols are esters that are used by the body to
store fats and oils and are formed from one molecule of
glycerol and three fatty acid molecules
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Lipids
• Steroids are lipids whose structure is made up of three
cyclohexane rings and one cyclopentane ring fused
together; presence of various substituents on the basic
steroid ring structure produces a family of steroid
compounds with different biological activity
– Cholesterol is a steroid found in cellular membranes and is the
starting point for the biosynthesis of steroid hormones,
including testosterone, the primary male sex hormone, and
progesterone, which helps maintain pregnancy
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Nucleic Acids
•
Nucleic acids are the basic structural components of
DNA and RNA, the biochemical substances found in the
nuclei of cells that transmit the information needed to
direct cellular growth and reproduction
• Structures are derived from nitrogen-containing cyclic
compounds called pyrimidines and purines, which can
hydrogen-bond through the lone electron pair on
nitrogen (in pyrimidine and purine) or through the
hydrogen of the amine (in purine)
• When a pyrimidine or purine is linked to a sugar by a
bond called a glycosidic bond, a nucleoside is formed;
addition of a phosphoric acid group to the sugar
produces a nucleotide
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Nucleic Acids
•
Nucleotides link to form a polymeric chain that consists
of alternating sugar and phosphate groups that are the
backbone of DNA and RNA
• The function of DNA is to preserve genetic information,
and RNA translates the genetic information in DNA and
carries that information to cellular sites where proteins
are synthesized
– Antibiotics function by interfering with the synthesis of proteins
in one or more kinds of bacteria
– Mutations in an organism’s DNA lead to the synthesis of
defective proteins
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