Download 1 • Introduction The Scientific Method (1 of 20) 1

This is an attempt to state how scientists do science. It is
necessarily artificial. Here are MY five steps:
1 • Introduction
The Scientific Method
(1 of 20)
• Make observations
the leaves on my plant are turning yellow
• State a Problem to be solved
how can I get my plants healthy (non-yellow)
• Form a hypothesis
maybe they need more water
• Conduct a controlled experiment
water plants TWICE a week instead of once a week
• Evaluate results
if it works, good... if not, new hypothesis (sunlight?)
Step 1 of the Scientific Method is Make Observations.
These can be of general physical properties (color, smell,
hardness, etc.) which are called qualitative observations.
1 • Introduction
Observations and Measurements
Qualitative, Quantitative, Inferences
(2 of 20)
These can be measurements which are called quantitative
observations.
There are also statements that we commonly make based on
observations. “This beaker contains water” is an example.
You infer (probably correctly) it is water because it is a
clear, colorless liquid that came from the tap. The
observations are that it is clear, it is colorless, it is a liquid,
and it came from the tap.
Consider: 16.82394 cm
In a measurement or a calculation, it is important to know
which digits of the reported number are significant.
1 • Introduction
Significant Digits I
What do they mean?
(3 of 20)
That means… if the same measurement were repeated again
and again, some of the numbers would be consistent and
some would simply be artifacts.
All of the digits that you are absolutely certain of plus one
more that is a judgment are significant.
If all the digits are significant above, everyone who
measures the object will determine that it is 16.8239 cm, but
some will say …94 cm while others might say …95 cm.
a
1 • Introduction
Significant Digits II
Some examples with rulers.
(4 of 20)
b
1
c
2
(A composite ruler)
a- No one should argue that the measurement is between 0.3
and 0.4. Is it exactly halfway between (.35 cm)… or a little
to the left (.34 cm)? The last digit is the judgment of the
person making the measurement. The measurement has 2
significant digits.
b- The same ruler… so the measurement still goes to the
hundredths place… 1.00 cm (3 significant digits).
c- A ruler with fewer marks reads 1.6 cm (2 sig digits).
1 • Introduction
Significant Digits III
Rules for Recognizing Sig. Digits
(5 of 20)
In a number written with the correct number of sig. digits...
• All non-zero digits are significant.
523 grams (3)
• 0’s in the MIDDLE of a number are ALWAYS significant.
5082 meters (4)
0.002008 L (4)
• 0’s in the FRONT of a number are NEVER significant.
0.0032 kg (2)
0.00000751 m (3)
• 0’s at the END of a number are SOMETIMES significant.
• Decimal point is PRESENT, 0’s ARE significant
2.000 Liters (4) 0.000500 grams (3)
• Decimal point is ABSENT, 0’s are NOT significant
2000 Liters (1) 550 m (2)
NOTE: textbook values are assumed to have all sig. digits
Scientific notation uses a number between 1 and 9.99 x 10 to
some power. It’s use stems from the use of slide rules.
1 • Introduction
Scientific Notation
Useful for showing Significant Digits
(6 of 20)
Know how to put numbers into scientific notation:
5392 = 5.392 x 103
0.000328 = 3.28 x 10 –4
1.03 = 1.03
550 = 5.5 x 102
Some 0’s in numbers are placeholders and are not a
significant part of the measurement so they disappear when
written in sci. notation. Ex: 0.000328 above. In scientific
notation, only the three sig. digits (3.28) are written.
Scientific Notation can be used to show more sig. digits.
Values like 550 ( 2 sig. digits) can be written 5.50 x 102 (3)
When you perform a calculation using measurements, often
the calculator gives you an incorrect number of significant
digits. Here are the rules to follow to report your answers:
1 • Introduction
Significant Digits IV
Significant Digits in Calculations
(7 of 20)
x and ÷: The answer has the same # of sig. digits as the
number in the problem with the least number of sig. digits.
example: 3.7 cm x 8.1 cm = 29.97 ≈ 30. cm2 (2 sig. digits)
+ and –: The last sig. digit in the answer is the largest
uncertain digit in the values used in the problem.
example: 3.7 cm + 8.1 cm = 11.8 cm (3 sig. digits)
Know how to ilustrate why these rules work.
1 • Introduction
Accuracy vs. Precision
(8 of 20)
Accuracy refers to how close a measurement is to some
accepted or true value (a standard).
Ex: an experimental value of the density of Al° is 2.69
g/mL. The accepted value is 2.70 g/mL. Your value is
accurate to within 0.37%
% error is used to express accuracy.
Precision refers to the reliability, repeatability, or
consistency of a measurement.
Ex: A value of 2.69 g/mL means that if you repeat the
measurement, you will get values that agree to the
tenths place (2.68, 2.70, 2.71, etc.)
± and sig. digits are used to express precision.
We generally use three types of measurements:
volume
Liters (mL)
length
meters (km, cm and mm)
mass
grams (kg and mg)
1 • Introduction
Metric System
(9 of 20)
We commonly use the prefixes:
1/
centi100 th
1
milli/1000th
kilo1000
Occasionally you will encounter micro(µ), nano, pico, mega,
and giga. You should know where to find these in chapter 1.
Know that 2.54 cm = 1 inch and 2.20 lb = 1 kg
1 • Introduction
% and ppm
(10 of 20)
Percentage is a mathematical tool to help compare values.
Two fractions, 3/17 and 5/31 are difficult to compare:
If we set up ratios so we can have a common denominator:
3
x
17.65
5
x
16.13
=
=
=
=
17 100
100
31 100
100
3
5
so… we can see that
> .
17 31
There are 17.65 parts per 100 (Latin: parts per centum) or
17.65 percent (17.65 %)… the % is a “1 0 0”
ppm (parts per million) is the same idea, (use 1,000,000
3
x
instead of 100)
=
= 176,470 ppm
17 1 000 000
Consider the metric/English math fact: 2.54 cm = 1 inch
This can be used as the “conversion factor”:
2.54 cm
1 inch
or
1 inch
2.54 cm
1 • Introduction
Unit Analysis
Converting between English and Metric Units
(11 of 20)
You can convert 25.5 inches to cm in the following way:
Given: 25.5 in
2.54 cm
Desired: ? cm
25.5 in x
= 64.77 cm ≈ 64.8 cm
1 in
This is the required way to show your work. You have two
jobs in this class, to be able to perform the conversions and
to be able to prove that you know why the answer is correct.
The important idea is that temperature is really a measure
of something, the average motion (kinetic energy,
KE) of the molecules.
1 • Introduction
Temperature Scales
(12 of 20)
Does 0°C really mean 0 KE? nope... it simply means the
freezing point of water, a convenient standard.
We have to cool things down to –273.15°C before we reach
0 KE. This is called 0 Kelvin (0 K, note: NO ° symbol.)
For phenomena that are proportional to the KE of the
particles (pressure of a gas, etc.) you must use
temperatures in K. K = °C + 273 °C = K – 273
mass is the amount of something...
weight is how much gravity is pulling on the mass.
(Weight will be proportional to the mass at a given spot.)
1 • Introduction
Mass vs. Weight
Theory, Measuring, Conversions
(13 of 20)
Mass is what we REALLY want to use... measured in grams.
You use a balance to measure mass... you compare your
object with objects of known mass.
Weight is measured with a scale (like your bathroom scale
or the scale at the grocery store). If there is no gravity, it
doesn’t work. Note: electronic balances are really scales!
You convert mass / weight using:
1 kg
2.205 lbs
or
2.205 lbs
1 kg
1
You can calculate the KE of an object: KE = mv2
2
m = mass, v = velocity [Note units: 1 J = 1 kg·m 2 ·s–2]
Temperature is a measure of the average kinetic energy.
1 • Introduction
Potential Energy (PE) and
Kinetic Energy (KE)
(14 of 20)
PE = the potential to do work which is due to an object’s
position in a field. For example, if I hold a book 0.5 m
above a student’s head it can do some damage... 1.0 m above
her/his head, more work can be done.
Important ideas:
Objects tend to change from high PE to low PE (downhill).
High PE is less stable than low PE.
Extensive properties depend on the amount of substance.
We measure these properties frequently... (mass &
volume... mostly).
1 • Introduction
Mass, Volume, and Density
Intensive vs. Extensive Properties
(15 of 20)
Intensive properties are independent of the size of the
sample. These are useful for identifying substances...
(melting point, boiling point, density, etc.)
mass
volume
is the ratio of two extensive properties... the size of the
sample sort of “cancels out.” Be able to do density problems
(3 variables) and know the usefulness of specific gravity.
It is interesting that an intensive property, density =
Heat is the total KE while temperature is the average KE.
1 • Introduction
Calorimetry
(16 of 20)
A way to measure heat is to measure the temperature change
of a substance... often water. It takes 1 calorie of heat
energy (or 4.184 J) to heat 1 gram of H 2 O by 1 °C.
cal
J
The specific heat of water = 1
= 4.184
g·°C
g·°C
heat = specific heat x mass H2 O x ∆T
You can heat other substances as well, you just need to
know their specific heats. Notice that this is simply
heating or cooling a substance, not changing its phase.
Equations to symbolize changes: reactants → products
1 • Introduction
Physical and Chemical Properties
Physical and Chemical Changes
(17 of 20)
Physical Properties can be measured from a sample of the
substance alone... (density, MP, BP, color, etc.)
Chemical Properties are measured when a sample is mixed
with another chemical (reaction with acid, how does it
burn in O2 )
Physical Changes imply that no new substances are being
formed (melting, boiling, dissolving, etc.)
Chemical Changes imply the substance is forming new
substances. This change is accompanied by heat, light,
gas formation, color changes, etc.
Matter
Pure Substances
Elements
1 • Introduction
Pure Substances, Elements, & Compounds
Homogeneous & Heterogeneous Mixtures
(18 of 20)
Compounds
Energy
Mixtures
Homogeneous
Heterogeneous
This chart should help you sort out these similar terms.
Be able to use chemical symbols to represent elements and
compounds. For example...
CuSO4 •5H2 O, a hydrate, contains 21 atoms & 4 elements.
Memorize the 7 elements that exist in diatomic molecules:
HONClBrIF or BrINClHOF or “H and the 6 that make a 7
starting with element #7”
1 • Introduction
Separating Mixtures by Filtration,
Distillation, and Chromatography
(19 of 20)
Mixtures are substances the are NOT chemically combined...
so if you want to separate them, you need to exploit
differences in their PHYSICAL properties.
Filtration:
some components of the mixture dissolve and some do
not. The filtrate is what passes through the filter.
Distillation:
some components vaporize at different temperatures or
one component may not vaporize at all (e.g.: salt+water)
complete separation may not be possible.
Chromatography:
differences in solubility vs. adhesion to the substrate.
Substaratemay be filter paper (paper chromatography),
or other substances, GLC, TLC, HPLC, column, etc.
Definite Composition:
samples of the same substance from various sources
(e.g. water) can be broken down to give the same %’s of
elements.
Calculation: percent composition
1 • Introduction
Early Laws: the Law of Definite Composition
& the Law of Simple Multiple Proportions
(20 of 20)
Multiple Proportions:
samples of 2 substances made of the same 2 elements...
(e.g. CO2 & CO or H2 O and H2 O2 or CH4 and C3 H8 )
if you break down each to give equal masses of one
element, the masses of the other element will be in a
simple, whole-number ratio.
Calculation: proportions to get equal amounts of one
element and then simple ratios.
Superscripts
used to show the charges on ions
Mg2+ the 2 means a 2+ charge (lost 2 electrons)
2•Stoichiometry: Chemical Arithmetic
Formula Conventions
(1 of 24)
Subscripts
used to show numbers of atoms in a formula unit
H2 SO4 two H’s, one S, and 4 O’s
Coefficients
used to show the number of formula units
2Br–
the 2 means two individual bromide ions
Hydrates
CuSO4 • 5 H2 O
some compounds have water molecules included
stoichiometry
atomic mass
2•Stoichiometry: Chemical Arithmetic
Stoichiometry Terms
(2 of 24)
formula mass
molecular mass
study of the quantitative relationships
in chemical formulas and equations.
weighted average mass of an atom,
found on the periodic table
sum of the atomic masses of the
atoms in a formula
sum of the atomic masses of the
atoms in a molecular formula
gram molecular mass molecular mass written in grams
molar mass
same as gram molecular mass
empirical formula
formula reduced to lowest terms
Formula or molecular mass is found by simply summing
the atomic masses (on the periodic table) of each atom in a
formula.
2•Stoichiometry: Chemical Arithmetic
Calculating Formula Mass
(3 of 24)
H2 SO4
1.01 + 1.01 + 32.06 + 16.0 + 16.0 + 16.0 + 16.0 = 98.08 u
2(1.01) + 32.06 + 4(16.0) = 98.06 u or 98.06 g/mole
Generally, round off your answers to the hundredths or
tenths place. Don’t round off too much (98.06 g/mol or
98.1 g/mol is OK, but don’t round off to 98 g/mol)
Units
Use u or amu if you are referring to one atom or molecule
A mole (abbreviated mol) is a certain number of things. It
is sometimes called the chemist’s dozen.
A dozen is 12 things, a mole is 6.02 x 1023 things.
2•Stoichiometry: Chemical Arithmetic
Mole Facts
(4 of 24)
Avogadro’s Number
1 mole of any substance contains 6.02 x 1023 molecules
Molar Volume (measured at P = 760 mmHg and T = 0 °C)
1 mole of any gas has a volume of 22.4 Liters
Molar Mass (see gram formula mass)
1 mole
6.02 x 10 23 molecules
1 mole
22.4 L
1 mole
molar mass
A Line Equation is the preferred way to show conversions
between quantities (amount, mass, volume, and number) by
canceling units (moles, grams, liters, and molecules)
2•Stoichiometry: Chemical Arithmetic
Line Equations
(5 of 24)
The line equation consists of the Given Value, the Desired
Unit, and the line equation itself.
Example: What is the mass of 135 Liters of CH4 (at STP)?
Given: 135 L CH4
Desired: ? g CH 4
1 mol CH4
16.0 g CH4
135 L CH4 x
x
= 96.43 g CH 4
22.4 L CH4
1 mol CH4
The “Mole Map” shows the structure of mole problems
Mass
2•Stoichiometry: Chemical Arithmetic
Mole Relationships
(6 of 24)
Mass
➀
Volume
at STP
➀
➂
1 mol
molar mass
2)
Volume
at STP
➂
number of
atoms or
molecules
1)
➁
moles
➁
number of
atoms or
molecules
1 mol
22.4 L
3)
1 mol
6.02 x 10 23 molecules
Percentage Composition quantifies what portion (by mass)
of a substance is made up of each element.
2•Stoichiometry: Chemical Arithmetic
Percentage Composition (by mass)
(7 of 24)
mass of element
mass of molecule
mass of element
Change to percentage: 100 x
mass of molecule
Set up a fraction:
Generally, round off your answers to the tenth’s place.
The percentage compositions of each element should add up
to 100% (or very close, like 99.9% or 100.1%)
2•Stoichiometry: Chemical Arithmetic
Formula from % Composition
(8 of 24)
Given the Percentage Composition of a formula, you can
calculate the empirical formula of the substance.
Step 1
assume you have 100 g of substance so
the percentages become grams
Step 2
change grams of each element to moles
of atoms of that element
Step 3
set up a formula with the moles
example: C2.4 H4.8
Step 4
simplify the formula by dividing moles
2.4
4.8
by the smallest value C 2.4 H2.4 = CH2
Step 5
If ratio becomes… 1:1.5 multiply by 2
1:1.33 or 1:1.66 multiply by 3
equation
condensed statement of facts about a
chemical reaction.
substances that exist before a
chemical rxn. Written left of arrow.
substances that come into existence as
a result of the reaction. Written to the
right of the arrow.
an equation describing a chemical
change using the names of the
reactants and products.
a number preceding atoms, ions, or
molecules in balanced chemical
equationns that showing relative #’s.
reactants →
2•Stoichiometry: Chemical Arithmetic
Equation Terms
(9 of 24)
→ products
word equation
coefficients
The gas density is often converted to molar mass:
2•Stoichiometry: Chemical Arithmetic
Other Mole Problems and Conversions
(10 of 24)
Example :
The gas density of a gas is 3.165 g/Liter (at STP). What is
the molar mass of the gas?
Knowing that 22.4 L is 1 mole, you can set up the ratio:
3.165 g molar mass
=
1 Liter
22.4 L
Other metric conversions you should know:
1000 mL
1 kg
1 Liter
1000 grams
Example: Write the formula equation of...
sodium metal + water → sodium hydroxide + hydrogen gas
Na° + H 2 O → NaOH + H2
2•Stoichiometry: Chemical Arithmetic
Writing Formula Equations
Things To Remember
(11 of 24)
• metals often are written with the ° symbol to emphasize
that the metal is in the neutral elemental state, not an ion.
• some compounds have common names that you should
just know... water, H 2 O; ammonia, NH3 ; methane, CH4
• remember the seven diatomic elements so they can be
written as diatomic molecules when they appear in their
elemental form. Other elemental substances are written
as single atoms (e.g. sodium metal or helium gas, He)
Since every gas takes up the same amount of room (22.4 L
for a mole of a gas at STP), the coefficients in an equation
tell you about the volumes of gas involved.
2•Stoichiometry: Chemical Arithmetic
Coefficients and Relative Volumes of Gases
(12 of 24)
Example:
N2 (g) + 3 H2 (g) → 2 NH 3 (g)
+
The “heart of the problem” conversion factor relates the
Given and the Desired compounds using the coefficients
from the balanced equation.
2•Stoichiometry: Chemical Arithmetic
Heart of the Problem
(13 of 24)
Example: N2 + 3 H2 → 2 NH3
3 moles H
♥ could be 2 moles NH2
3
…which means that every time 2 moles of NH 3 is formed, 3
moles of H2 must react.
The format is always,
moles of Desired
moles of Given
Mass-Mass problems are probably the most common type
of problem. The Given and Desired are both masses (grams
or kg).
2•Stoichiometry: Chemical Arithmetic
Mass-Mass Problems
Mass-Volume Problems
(14 of 24)
2•Stoichiometry: Chemical Arithmetic
Mass-Volume-Particle Problems
(15 of 24)
The pattern is:
molar mass
Given x of Given
x
mass
♥ x molar
of Desired
In Mass-Volume problems, one of the molar masses is
22.4 L
replaced with
depending on whether the Given or
1 mole
the Desired is Liters.
If the Given or Desired is molecules, then the Avogadro’s
6.02 x 10 23 molecules
Number conversion factor,
is used
1 mole
and the problem is a Mass-Particle or Volume-Particle
problem.
The units of the Given and Desired will guide you as to
which conversion factor to use:
Mass
Volume
Particles
grams or kg
Liters or mL
molecules or atoms
In a problem with two Given values, one of the Given’s will
limit how much product you can make. This is called the
limiting reactant. The other reactant is said to be in excess.
2•Stoichiometry: Chemical Arithmetic
Limiting Reactant Problems
(16 of 24)
Solve the problem twice using each Given… the reactant
that results in the smaller amount of product is the limiting
reactant and the smaller answer is the true answer.
Example: N2 + 3 H2 → 2 NH3
When 28.0 grams of N2 reacts with 8.00 grams of H 2 , what
mass of NH3 is produced?
(in this case, the N2 is the limiting reactant)
2•Stoichiometry: Chemical Arithmetic
How Much Excess Reactant is Left Over
(17 of 24)
Example: N2 + 3 H2 → 2 NH3
When 28.0 grams of N2 reacts with 8.00 grams of H 2 , what
mass of NH3 is produced?
(in this case, the N2 is the limiting reactant)
To find out how much H2 is left over, do another line
equation:
Given: 28.0 g N2
Desired: ? g H2
subtract the answer of this problem from 8.00 g H2
It is difficult to simply guess which reactant is the limiting
reactant because it depends on two things:
(1) the molar mass of the reactant and
(2) the coefficients in the balanced equation
2•Stoichiometry: Chemical Arithmetic
Limiting Reactants
(18 of 24)
The smaller mass is not always the limiting reactant.
Example: N2 + 3 H2 → 2 NH3
1 mole (28 g N 2 ) will just react with 3 moles (6.06 g H2 )
so, if we react 28.0 g N2 with 8.0 g H2 , only 6.06 g H2
will be used up and 1.94 g of H2 will be left over.
In this case, N2 is the L.R. and H2 is in X.S.
The answer you calculate from a stoichiometry problem can
be called the Theoretical Yield. Theoretically, you
should get this amount of product.
2•Stoichiometry: Chemical Arithmetic
Theoretical Yield and Percentage Yield
(19 of 24)
In reality , you often get less than the theoretical amount due
to products turning back to reactants or side reactions.
The amount you actually get is called the Actual Yield.
Percentage Yield =
2•Stoichiometry: Chemical Arithmetic
Balancing Chemical Equations
(20 of 24)
Actual Yield
x 100
Theoretical Yield
The balanced equation represents what actually occurs
during a chemical reaction. Since atoms are not created
or destroyed during a normal chemical reaction, the
number and kinds of atoms must agree on the left and
right sides of the arrows.
__Na 2 CO3 + __HCl → __ NaCl + __H2 O + __CO2
To balance the equation, you are only allowed to change the
coefficients in front of the substances... not change the
formulas of the substances themselves.
Reduce the coefficients to the lowest terms.
Fractions may be used in front of diatomic elements.
The burning of fuels made of C, H, and O is called
combustion. You need to memorize O2 , CO2 and H 2 O
Example: The combustion of propane, C3 H8 , is written:
2•Stoichiometry: Chemical Arithmetic
Combustion Equations
(21 of 24)
C 3 H8 + 5O2 → 3CO2 + 4H2 O
Be careful when writing equations for alcohols, such as
butanol, C 4 H9 OH
• don’t forget to add the H’s (a total of 10 of them)
• don’t forget to take account of the O atom in the alcohol
C 4 H9 OH + 6O2 → 4CO2 + 5H2 O
2•Stoichiometry: Chemical Arithmetic
Solutions -- Molar Concentration
(22 of 24)
Many reactions are carried out in solution. Solutions are
convenient and speed up many reactions.
Concentration is often expressed as
moles of solute
Molarity (M) =
Liters of solution
You can calculate the molarity of a solution when given
moles (or grams) of a substance and its volume.
You can use the molarity of a solution as a conversion factor
0.150 moles HCl
1 Liter HCl
0.150 M HCl ≈
or
1 Liter HCl
0.150 moles HCl
to convert moles to Liters and vice versa.
Volumetric flasks are used to make solutions.
You can calculate the moles of a solute using the volume
and molarity of the substance. Since diluting a solution
adds water and no solute, the moles of solute before and
after the dilution remains constant. So...
2•Stoichiometry: Chemical Arithmetic
Dilution Problems
(23 of 24)
Vi · Mi = Vf · Mf
where “I” means “initial” and “f” means “final”
The units of volume or concentration do not really matter as
long as they match on the two sides of the equation.
Acids form the H+ ion. Bases form the OH– ion.
Acids + bases mix to form H 2 O (HOH) and a salt.
The moles of H+ = the moles of OH– in a neutralization.
2•Stoichiometry: Chemical Arithmetic
Acid-Base Titrations
(24 of 24)
An acid-base titration is the technique of carefully
neutralizing an acid with a base and measuring the
volumes used. An indicator (we used phenolphthalein)
allows us to observe when the endpoint is reached.
If a monoprotic acid is neutralized with a base that only has
one OH– ion per formula unit, the simple formula:
Va · Ma = Vb · Mb
allows you to determine the molarity of the unknown.
3 • The Periodic Table & Makeup of Atoms
The Subatomic Particles
(1 of 12)
Name Symbol Mass Charge
protons
p+
1u
1+
neutron
n°
1u
0
1
–
electron e
u
1–
1837
Location
part of the nucleus
part of the nucleus
normally at large
distances from the
nucleus
J.J. Thompson is given credit for discovering electrons
using a Crookes tube and testing many different gases.
Cathode rays were found to be beams of electrons.
Chadwick is given credit for the discovery of the neutron.
3 • The Periodic Table & Makeup of Atoms
Terms I-- Atomic Structure
(2 of 12)
atoms
the smallest particle of an element . It
consists of a central nucleus and electron
clouds outside the nucleus.
nucleus
the dense central portion of an atom.
subatomic
smaller than an atom. The proton,
neutron, and electron are subatomic.
net charge
the difference in the positive charge due to
protons and the negative charge due to
electrons in an atom.
nucleons
the particles that make up the nucleus.
atomic number the number of protons in an atom. This #
determines the identity of an element.
3 • The Periodic Table & Makeup of Atoms
Terms II-- Atomic Structure
(3 of 12)
mass number
the number of protons + neutrons
isotopes
atoms with the same number of protons,
but different numbers of neutrons. Atoms
with the same atomic number, but
different mass numbers.
isotopic notation shorthand notation for a nucleus that
shows the mass #, atomic # and the
238
symbol. U-238 would be 92 U
Any real sample of an element contains more than one
naturally occuring isotope. For instance, boron
isotope
3 • The Periodic Table & Makeup of Atoms
Calculating Atomic Mass
(4 of 12)
10
boron-10 5 B
11
boron-11 5 B
abundance
mass #
isotopic mass
19.78%
10
mass = 10.013 u
80.22%
11
mass = 11.009 u
The atomic mass is the weighted average of the isotopes.
(19.78%)(10.013u) + (80.22%)(11.0009u)
or
100
at. mass = (0.1978)(10.013u) + (0.8022)(11.0009u)
= 10.81 u
at. mass =
33
Consider the following symbol: 16S2–
The 16 is the atomic number which is the number of
protons.
3 • The Periodic Table & Makeup of Atoms
Determining Numbers of Protons, Neutrons,
and Electrons from the Isotopic Notation
(5 of 12)
The 33 is the mass number which is the mass of one of the
isotopes. This mass is due to the protons and neutrons.
The number of neutrons is the mass number - the atomic
number. 33 - 16 = 17 neutrons.
Since the charge is 2-, there are 2 more electrons than
protons. In this case, there are 18 electrons.
3 • The Periodic Table & Makeup of Atoms
Important People in the Development of the
Atomic Theory
(6 of 12)
Democritus [atomos]
philosopher who argued that matter was discontinuous
John Dalton [billiard-ball model]
experimented with gases… different substances are
different combinations of atoms
J.J. Thomson [plum-pudding model]
experimented with gas-discharge tubes… atoms have + and
– parts… the negative e– ’s are the same for any atom
Ernest Rutherford [nuclear model/solar system model]
most of the mass of the atom is concentrated in a tiny,
positively-charged nucleus
Niels Bohr [quantized electron energy levels]
the electrons have only certain allowed energy levels
The elements can be classified as metals, nonmetals, and
metalloids. Memorize the elements classified as metalloids
(also called semi-metals or semiconductors).
3 • The Periodic Table & Makeup of Atoms
Metals, Nonmetals, and Metalloids
(7 of 12)
Properties of metals include:
ductility - the ability to pull a substance into a wire
sectility - the ability to cut with a knife
malleability - the ability to pound substance into a sheet
conductivity - the ability to carry an electrical current
Gold is the most malleable of all the metals.
Ernest Rutherford’s classic gold foil experiment led to the
nuclear model of the atom.
3 • The Periodic Table & Makeup of Atoms
Rutherford’s Gold Foil Experiment
(8 of 12)
α
a few bounced back
at a large angle
gold foil
most alpha's
α came straight
through here
• the nucleus is tiny - because most of the alpha’s missed
the nucleus and went straight through the foil
• the nucleus is positively charged - because the (+)
charged alpha was repelled by the (+) charged nucleus
• the nucleus is incredibly dense - because the nucleus was
able to bounce back at a very large angle
3 • The Periodic Table & Makeup of Atoms
The First Periodic Table
(9 of 12)
Meyer and Mendeleev are given credit for developing the
first version of the periodic table. Mendelleev’s true claim
to fame was that he actually predicted the existence of
several element that had not been discovered. He found
gaps in the table when he tried to organize the atoms and left
spaces for those elements (ekasilicon = "like silicon", etc.)
He predicted Ga, Ge, and Sc.
He also arranged elements in order of atomic number rather
than the previous idea of atomic mass. Several of the
elements change order... (like Te and I).
Horizontal rows of the table are called periods.
Vertical columns are called groups or families.
3 • The Periodic Table & Makeup of Atoms
Families of the Periodic Table
(10 of 12)
Memorize the names of some groups:
IA - the alkali metals
IIA - the alkaline earth metals
VIIA - the halogens
0 - the noble gases
Also know the transition metals, the inner transition metals
(composed of the lanthanide series and the actinide series...
the lanthanides are also called the rare earth metals)
Families IA - VIIIA are called the representative elements.
Radioactivity was discovered by Henri Becquerel (but
named by Marie Curie).
3 • The Periodic Table & Makeup of Atoms
Radioactivity Basics
(11 of 12)
“Becquerel rays” were found to consist of 3 types or
radiation:
alpha particles (α) a helium nucleus - 2 protons + 2 neutrons
…easily stopped by paper or skin
beta particles (β) a high energy electron
…stopped by Al foil (several thicknesses of foil)
gamma radiation (γ) a very high energy form of light (EMR)
…the most penetrating and dangerous of the rays.
3 • The Periodic Table & Makeup of Atoms
JJ Thomson & cathode ray tubes
(12 of 12)
Know the design of a cathode rays tubes. Realize that
cathode rays are really beams of electrons. The cathode rays
are the same for any substance, but the canal rays (the
positive ions left after ionizing the gases) are different for
each gas.
Know how the bending of cathode rays can tell you the
charge-to-mass ratio (e/m) (but not the mass or the charge
of the electron).
Millikan’s oil drop experiment gave evidence for the
charge of the electron. Knowing this and the e/m ratio, you
can calculate the mass of the electron.
4 • Electronic Structure & the Per. Table
Wave Ideas You Should Know
(1 of 16)
EMR
electromagnetic radiation …
oscillating electric & magnetic fields at right angles
wavelength (λ)
the distance from crest to crest or
trough to trough.
amplitude
the distance from the equilibrium
point to the crest or trough.
frequency (ν)
the number of waves that pass a
point per second. (Hz, s–1, 1/s)
continuous spectrum a “normal” rainbow that contains all
of the colors (ROYGBV).
line spectrum
a spectrum that only contains certain
bright lines that result from electron
transitions within an atom.
Exchanging wavelength, frequency, & energy of light:
λ·ν=c
c = 3.00 x 10 8 m/s [speed of light]
4 • Electronic Structure & the Per. Table
Wave Calculations
(2 of 16)
4 • Electronic Structure & the Per. Table
The Balmer Series
(3 of 16)
E = hν
h = 6.626 x 10–34 J·s [Planck’s constant]
hc
E=
λ
Energy of Level “n” in the Hydrogen atom:
A
En = – 2
A = 2.18 x 10–18 J [Arrhenius constant?]
n
NOTE: You can learn the Balmer equation or the Rydberg
equation listed on pp. 108 & 109. I suggest you use the
above equations to calculate the wavelengths or frequencies
of light emitted when electrons change energy levels.
When the electron “drops” to energy level n=2, visible light
is emitted. The bight line spectrum observed is called the
Balmer Series of lines. [Memorize this info...]
3→2
red
4→2
blue-green
5→2
blue-violet
6→2
violet
All the lines result from an electron transition to level n=1
are too high energy to be visible… UV [Lyman Series]
All the lines that result from an electron transition to levels
n=3 [Paschen], n=4 [Brackett] and n=5 [Pfund] are too
LOW energy to be visible.
4 • Electronic Structure & the Per. Table
The De Broglie Wavelength of Electrons
(4 of 16)
Two equations can be combined into one to allow you to
calculate the wavelength of a particle:
hc
E = mc2
E=
λ
h
c
mc2 =
λ
h
h
mc = or more generally, mv =
λ
λ
h
λ=
mv
h = Planck’s constant, m = mass of particle (in kg),
v = velocity of particle (in m/s).
Standing waves are something that waves do…
resulting from having repeating waves that always cancel at
the nodes and always add up at the antinodes.
4 • Electronic Structure & the Per. Table
Standing Waves
(5 of 16)
We have seen standing waves
of strings:
of sound:
of drumheads:
+
+ –
+ –
We hypothesized the standing waves
of electrons: (4-dimensional vibrating wavicles)
the orbitals (s, p, d, f… probability waves)
These are variables in some unseen equation:
The Rules:
n = 1, 2, 3, ...
determines the energy of the e–
4 • Electronic Structure & the Per. Table
Quantum Numbers (n, l, m, s)
(6 of 16)
l = 0 → (n–1)
the type of orbital (subshell)
0 ≈ s, 1 ≈ p, 2 ≈ d, 3 ≈ f
m = –l → +l
which orientation of the orbital
(x, y, z… for p orbitals)
1
1
s = + or –
2
2
the “spin” of the electron
NOTE: m is also called ml and s is ms
The Pauli Exclusion Principle states that no two electrons
in an atom may have the same four quantum numbers. This
translates to the idea that an orbital may contain no more
than two electrons.
4 • Electronic Structure & the Per. Table
Three Rules for Filling Orbitals
(7 of 16)
The Aufbau Principle states that electrons occupy the
lowest energy available orbital. You must memorize the
orbital chart on page 120. [Aufbau = “building up”]
Hund’s Rule states that when you have several orbitals of
the same energy (e.g. three p orbitals or five d orbitals) place
one electron in each orbital before doubling them up.
NOTE: Remember the Aufbau hotel analogy...
This is a shorthand notation for the arrangement of electron
in the orbitals.
1s2 means 2 electrons in the 1s orbital
4 • Electronic Structure & the Per. Table
Electron Configurations
(8 of 16)
Consider Arsenic, As
This is the order in which the electrons fill…
As 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
however, we write the orbitals according to the distance
from the nucleus (n=3’s then n=4’s, etc.)
As 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3
long form: (shown above) show each subshell
short form: (show the last filled NRG level...noble gas core)
As [Ar] 3d10 4s2 4p3
4 • Electronic Structure & the Per. Table
The s-block, p-block... of the Periodic Table
Exceptions to the Filling Rules
(9 of 16)
The shape of the Periodic Table comes from the way the
electrons fill the orbitals.
s-block
Families I and II
p-block
Families III, IV, V, VI, VII, and VIII
d-block
The Transition Elements
f-block
The Inner Transition Elements
Six elements do not follow the rules. They are in the dblock of the periodic table… the transition elements.
Cu, Ag, and Au... instead of a full s-orbital and almost filled
d-orbital, they have a filled d and half-filled s-orbital
Cr, Mo, W... instead of a full s-orbital and an almost halffilled d-orbital, they have d5 and s1
Know the general shapes of the
s (spherical) orbitals,
4 • Electronic Structure & the Per. Table
Shapes of the orbitals
(10 of 16)
4 • Electronic Structure & the Per. Table
Predicting the Atomic Size (radius)
Trends in the Periodic Table
(11 of 16)
p (perpendicular) orbitals, and
d (diagonal) orbitals.
What: As you move ACROSS a period, the size of the atom
DECREASES
Why: As you move ACROSS a period, the number of
protons in the nucleus increases... so the protonelectron attraction increases... the size decreases.
NOTE: The increased e– -e– repulsion that one might expect
is not important, because the outer electrons do not
SHIELD the electrons from the nucleus..
What: As you move DOWN a family, the size of the atom
INCREASES
Why: As the value of “n” increases, the average distance of
the electron from the nucleus increases. The size of
the atom IS the electron cloud... so the size increases
The change in size of ions (compared to the neutral element)
depends on the electron-electron repulsion.
4 • Electronic Structure & the Per. Table
Explaining The Sizes of Ions
The Lanthanide Contraction
(12 of 16)
If an atom gains electrons, the increased repulsion increases
the size of the electron cloud. So... negative ions are larger
than the neutral atom.
Positive ions form by losing electrons (less repulsion) and
get smaller.
The Lanthanide contraction makes the third row of the
transition elements about the same size as the second row...
This causes these elements to be much more dense.
Following the decrease in size, the ionization energy
GENERALLY increases as you move across a period. If the
atom is smaller, the electron being removed is closer to the
nucleus and therefore feels a stronger attraction.
4 • Electronic Structure & the Per. Table
Ionization Energy
Trends Across a Period
(13 of 16)
The Jags: (why are some electrons EASIER to remove)
In Family III, the electron being removed comes from the porbital rather than the s-orbital. The p-orbital electron is at a
higher energy and requires LESS energy to ionize.
In Family VI, the electron being removed comes from an
orbital with TWO electrons. The e– - e– repulsion felt by the
electron allows it to be ionized with less added energy.
4 • Electronic Structure & the Per. Table
Ionization Energy & Reactivity
Trends Down a Family
(14 of 16)
4 • Electronic Structure & the Per. Table
Ionization Energy
Trends in Successive Ionizations
(15 of 16)
What:
Why:
Down a family, Ionization energy DECREASES
The trend here goes right with the size of the atom.
Since “n” increases, the distance of an electron
from the nucleus increases and is easier to remove.
What:
Why:
Down family I and II, the reactivity INCREASES.
These families lose electrons to become + ions.
The easier it is (lower ionization energy) the more
reactive they are.
What:
Why:
Down family VII, the reactivity DECREASES.
These elements GAIN electrons to become – ions.
Smaller atoms mean the attracted electrons will be
closer to the nucleus... more effective attractions.
What:
You can tell the family of an element by observing
when the ionization energies get very large.
For example: family III elements
Al + energy = Al+ + e–
(3p electron)
Al+ + energy! = Al2+ + e–
(3s electron)
Al2+ + energy!! = Al3+ + e–
(3s electron)
Al3+ + ENERGY = Al4+ + e– (2p electron!)
Why:
The “easy” electrons to remove are in an orbital
with a higher value of n. When n decreases, the
average distance of the electron from the nucleus
decreases and the attractions between the proton
and electrons increase.
Electron affinity is the energy involved when an atom gains
an electron to become a negative ion.
F + e – → F– + energy
4 • Electronic Structure & the Per. Table
Electron Affinity
(16 of 16)
Elements in the upper right corner of the periodic table have
the greatest electron affinity (greatest attraction for
electrons).
The electron affinity may be + or –. Negative values mean
energy is released and also counts as a greater electron
affinity.
Electron affinity data is not complete, but it gives SOME
evidence for trends in the periodic table.
5 • Chemical Bonding: Gen Concepts
Some Properties of
Ionic and Molecular Compounds
(1 of 12)
Compound
Conducts as Solid
Conducts as Liquid
Conducts in Solution
Conducts as Gas
Molecular
NO
NO
NO
NO
Ionic
NO
YES
YES
YES
Hardness
soft
hard
MP / BP
low
high
Bonding
covalent
ionic
Examples
He, CH 4 , CO2 ,
C 6 H12O6
NaCl, KI,
AgNO 3
Lewis symbols conisist of the atomic symbol surrounded by
valnece electrons. The four sides represent the four valence
orbitals. Atoms are usually shown in their excited states (II,
III, IV)
5 • Chemical Bonding: Gen Concepts
Lewis Symbols of Atoms and Ions
(2 of 12)
Li
Be
B
C
N
O
F
Ne
Ions include brackets. Positive ions show no valence
electrons while negative ions usually have an octet.
[Li]+
[Mg]
2+
[O]
2–
Many ions can be explained because they have gained or lost
electrons and attain a noble gas configuration. For example:
P 3– S2– Cl– Ar K+ Ca 2+
all have the same electron arrangement: 1s2 2s2 2p 6 3s2 3p 6
5 • Chemical Bonding: Gen Concepts
The Ionic Bond
Noble and Pseudonoble Gas Configurations
(3 of 12)
A pseudonoble gas configuration is:
1s2 2s2 2p 6 3s2 3p 6 3d 10
This is found in Cu+ Zn 2+ Ga3+ and Ge4+
Similar configurations are found in the next two periods.
The importance of this configuration is that there is more
than one reason why ions form what they do. Many ions are
not explained.
Know the 5 steps that can be thought to occur when an ionic
bond forms. Note whether each is exo- or endothermic...
whether a larger energy helps or hinders the bond formation.
5 • Chemical Bonding: Gen Concepts
Factors that Influence
the Formation of Ionic Bonds
(4 of 12)
Overall:
1.
2.
3.
4.
5.
Li(s) + 1/2F2 (g) → LiF(s)
heat of vaporization
heat of decomposition
ionization energy
electron affinity
lattice energy
Li(s) + NRG → Li(g)
1/2F2 (g) + NRG → F(g)
Li(g) + NRG → Li+(g) + e–
F(g) + e– → F– (g)
Li+(g) + F– (g) → LiF(s)
Large energy values for 1,2,3 hinder ionic bond formation.
The covalent bond between two atoms depends on the
balanc e of attractions between one atom’s + nucleus and
the other atom’s – electrons and the proton-proton
repulsions as well as electron-electron repulstions.
5 • Chemical Bonding: Gen Concepts
The Covalent Bond
Attractions and Repulsions
(5 of 12)
PE
Distance between nuclei
If two atoms have half-filled orbitals , the interactions
balance at a small enough distanc eso the e – ’s can be close
to both nucle i at the same time... this is a covalen tbond.
Count up your valence electrons.
5 • Chemical Bonding: Gen Concepts
Groves’ Electron Dot System
Multiple & Extended Valence Bonds
(6 of 12)
Give every atom who “wants” and octet an octet.
[the first 5 elements do not need octets... too small]
[Family I, II, and III do not form octets]
If you have drawn too man y electrons...
“Take away a lone pair... take away a lone pair...
make these two atoms share”
If you have drawn too fe w electrons... place the extra
electrons on the central atom (extended valence shell)
5 • Chemical Bonding: Gen Concepts
Bond Order: Bond Length, Strength, &
Vibrational Frequency
(7 of 12)
Bond orde r is the number of pairs of electrons bonding two
atoms together.
single bond
bond order = 1
double bond
bond order = 2
triple bond
bond order = 3
single bonds have the longest bond length
single bonds have the weakest bond strength
single bonds have the lowest vibrational frequency
(think of single bonds as soft, springy springs...
triple bonds are tight springs...sproinnnnng)
Bonds in resonance structure must be averaged... the S-O
bond in SO 2 has a bond order of 1.5. C-O in CO3 2– is 1.33
When you draw a Lewis structure (SO2 , O3 , CO3 2–, etc.) in
which you must make a choic e as to who gets a double
bond, the structure is actually a blend of two or three
structures.
5 • Chemical Bonding: Gen Concepts
Resonance
(8 of 12)
We “say” that the structure “resonate s” or we say that the
structure contains contributions from each of the resonance
structures.
Resonance occurs simply because the electron-dot model
(while very useful) is too limited to show how the electrons
are being shared between the atoms... wait for π bonding.
Coordinate covalent bond: When a covalent bond is
formed by sharing a pair of electrons BUT the electron pair
belonged to only one of the atoms.
5 • Chemical Bonding: Gen Concepts
Coordinate Covalent Bonds
(Preview: Lewis Acids)
(9 of 12)
Classic Example: NH3 + BF3 → NH3 BF3
The bond between the N and the B is coordinate covalent.
The lone pair donor is called a Lewis Base.
(this atom has a lone pair of electrons)
The lone pair acceptor is called a Lewis Acid.
(this atom has an empty orbital)
“Have Pair Will Share” --Lewis Base
5 • Chemical Bonding: Gen Concepts
Electronegativity and Polar Bonds
(10 of 12)
5 • Chemical Bonding: Gen Concepts
Naming Ionic Compounds
Traditional and Stock Names
(11 of 12)
You will be given a chart of electronegativity values.
Memorize the most electronegative element s (F = 4.0)
then oxygen (O = 3.5) and chlorine (Cl = 3.0). The noble
gases have no electronegativity values… no bonds.
Trend is large electronegativity in the upper right of the
per. table and small in the lower left portion of the table.
Classify the bond between any two atoms by subtracting
their electronegativity values (∆e)
Non-polar covalent
0 < ∆e < 0.5
Polar covalent
0.5 ≤ ∆e ≤ 1.7
Ionic
∆e > 1.7
The more electronegative atom is more negative.
Polar covalent bonds have partial charges δ+ and δ–
The Stock System of naming compounds is used…
• when a positive ion has more than one possible charge
Traditional: mercurous, Hg2 2+
mercuric, Hg2+
Stock:
mercury(I)
mercury(II)
Traditional: cuprous, Cu+
cupric, Cu 2+
Stock:
copper(I)
copper(II)
• for molecular compounds where the elements have many
different oxidation states (i.e. N in NO 2 , NO, N 2 O, etc.)
Stock Name:
Traditional Name:
NO2
nitrogen(IV) oxide
nitrogen dioxide
NO
nitrogen (II) oxide
nitrogen monoxide
N2 O
nitrogen(I) oxide
dinitrogen monoxide
Acids are ionic formulas in which the positive ion is H+.
Use as many H+ ions as the charge on the negative ion.
5 • Chemical Bonding: Gen Concepts
Naming Acids
(12 of 12)
Three rules for naming:
if the anion ends with:
–ite
–ate
–ide
the acid is named:
********ous acid
********ic acid
hydro********ic acid
• Acids from sulfide, sulfite, and sulfate include a “ur”
H2 S is hydrosulfuric acid, not hydrosulfic acid
• Acids from phosphate and phosphite include a “or”
H3 PO4 is phophoric acid, not phosphic acid
7 • Chemical Reactions & the Per. Table
Solutions and Solubility
(1 of 16)
We learned the terms solute, solvent, and solution.
Solubility (how MUCH solute will dissolve) is measured in
g/100 mL of H2 O or sometimes in M.
This information may be given numerically or graphically.
A saturated solution is one in which any additional solute
added will simply settle on the bottom of the container.
An unsaturated solution is any amount less than saturated.
Supersaturated implies that the solution was saturated at
some higher temperature and then carefully cooled. This
unstable situation can be changed with a “seed” crystal.
Recall the supersaturated solution of NaC2 H3 O2 demo.
The terms concentrated and dilute refer to the amount of
solute and do not necessarily coincide with saturation.
7 • Chemical Reactions & the Per. Table
Electrolytes: Weak and Strong
(2 of 16)
Solutions of acids, bases, and salts contain mobile ions and
conduct electricity. These solutions are called electrolytes.
Salts are ionic compounds that dissociate in water.
Acids are actually molecular compounds (covalently
bonded) the become ions when dissolved in water.
Only 8 acids are strong electrolytes and completely
dissociate when dissolved. All others dissolve completely,
but only partially dissociate into ions.
Only 8 bases are strong electrolytes because they dissolve
completely. All others have low solubility and remain
solids rather than dissolve. One common exception is the
weak base NH4 OH . It dissolves, but partially dissociates.
7 • Chemical Reactions & the Per. Table
Ionic Reactions
(3 of 16)
A common class of chemical reactions occurs when two
ionic solutions are mixed. The double replacement or
metathesis reaction involves the formation of two new
combinations of ions.
AgNO 3 + NaCl → AgCl + NaNO3 (molecular equation)
The new combinations may be more stable than the original
due to low solubility (precipitate forms), weak electrolyte,
gas formation, or complex ion formation.
The reaction is written above as though the substances exist
as molecules. This is the easiest time to balance.
The ionic equation shows strong electrolytes as separate
ions. The net equation eliminates “spectator ions”.
7 • Chemical Reactions & the Per. Table
Arrhenius Acids and Bases
(4 of 16)
Acid: a substance that increases the [H+] in solution.
Base: a substance that increases the [OH–] in solution.
Diprotic acids have more than one removable H. (H 2 SO4 )
However, only the first H is ever easily dissociated.
Oxides of nonmetals (SO2 ) are acid anhydrides.
Oxides of metals (Na2 O)are basic anhydrides.
Just add water... to get the acid or base.
SO2 + H2 O → H2 SO3
Na 2 O + H2 O → 2 NaOH
Acids and bases neutralize each other because the H+ and
OH– ions form the very weak electrolyte... H2 O (and a salt).
Acid salts are the partially neutralized polyprotic acids.
NaH2 PO4 or Na2 HPO4 or NaHSO4 , etc... solid acids.
7 • Chemical Reactions & the Per. Table
Brønsted-Lowry Acids and Bases
(5 of 16)
Acid: a proton (H+) donor. Base: a proton acceptor.
This is a more general definition of acids and bases because
it does not require the substance to be dissolved in water.
Consider the following equations:
The species that accepted the proton (base) can be
considered a donor (conjugate acid) in the reverse reaction.
HF + H 2 O ⇔ H3 O+ + F–
(acid) (base)
(acid) (base)
NH3 + H2 O ⇔ NH4 + + OH–
(base) (acid)
(acid) (base)
Strong base≈weak conjugate acid, (good acceptor≈lousy
donor). Conjugates differ by only a H+ (e.g. HF and F – )
7 • Chemical Reactions & the Per. Table
Ions in Water
Some Metal Ions Make Acidic Solutions
(6 of 16)
Since water molecules are polar, they surround ions in
solution (called hydration). When we write Na+(aq) and
Cl– (aq) we are implying this more complex situation.
Some highly charged ions (Al3+ , Cr3+ , Fe3+ ) tend to tightly
bind the water molecules. We can write them as complex
ions: Al(H2 O) 6 3+
The electron clouds are drawn toward the central ion and
away from the oxygen and therefore the O-H bond. This
extra-polar O-H bond results in the H atom more readily
joining with a passing water molecule... making the solution
acidic. [Note: this is a conjugate acid-base situation.]
Al(H 2 O) 6 3+ + H2 O ⇔ H3 O+ + Al(H2 O) 5 OH2+
7 • Chemical Reactions & the Per. Table
Trend in Strengths of Acids and Bases
Three cases to explain
(7 of 16)
7 • Chemical Reactions & the Per. Table
Lewis Acids and Bases
(8 of 16)
Case 1: The more oxygens on an oxoacid, the stronger the
acid. H2 SO4 > H2 SO3 .. HClO4 > HClO 3 > HClO 2 > HClO
Why?: The electronegative oxygens draw electron density
away from the central atom and therefore from the H-O
bond... making it more polar. H leaves more easily.
Case 2: The more electronegative the central atom, the
stronger the oxoacid. H 3 PO4 < H2 SO4 < HClO 4
Why? Same as above... the more electronegative atom in
the center makes the O-H bond more polar.
Case 3: Binary acid strength depends on the SIZE of the
atom... HF < HCl < HBr < HI... not the electronegativity.
Why?: The greater distance means a weaker attraction.
Consider the formation of a coordinate covalent bond:
Acid: electron pair acceptor [Note: proton donor]
Base: electron pair donor [Note: proton acceptor]
This definition is more general than the other two because
this covers cases that don’t even involve hydrogen (protons).
Classic case: HN3 + BF3 → NH3 BF3
Other important cases:
oxide of metal + oxide of nonmetal → salt
(base anhydride + acid anhydride)
CaO + SO2 → CaSO3 (s)
and the oft-confusing reactivity of CO2
OH– + CO 2 → HCO3 –
7 • Chemical Reactions & the Per. Table
Oxidation Numbers
(9 of 16)
7 • Chemical Reactions & the Per. Table
Balancing Redox Equations
Oxidation Number Change Method
Step-By-Step
(10 of 16)
7 • Chemical Reactions & the Per. Table
Balancing Redox Equations
Ion-Electron Method
(Half-reaction Method)
Step-By-Step
(11 of 16)
Definition:
Oxidation numbers are the apparent charges atoms have if
shared e– ’s are assigned to the more electronegative atom.
You can assign ox. #’s by studying a Lewis diagram or...
The Rules:
…ox. # of neutral atoms is 0
In compounds…
…simple ions have ox. #’s equal to their charge.
…F (-1), Family I (+1), II (+2), Al (+3)
…O is usually (-2) except peroxide and OF 2
…H is usually (+1) except hydrides of Fam. I, II, Al
…the sum of the ox. #’s of individual atoms equals the
charge on the entire species.
1. Identify the ox. #’s of elements that change PER ATOM.
One element will change UP (oxidation / lose e– ’s) one
element changes DOWN (reduction / gain e– ’s).
2. Adjust coefficients and e – changes for situations where
more than one atom MUST change:
ex: 2HCl + K2 Cr 2 O7 → KCl + 2CrCl3 + Cl 2 + H2 O
Cr: 2 x (3 e – per Cr) Cl: 2 x (1 e– per Cl)
3. Balance these changes in e– ’s (e– gained = e– lost).
4. Balance all elements except H and O.
5. Balance O’s (add H2 O’s as needed).
6. Balance H’s (add H+’s as needed).
7. If solution is basic, see card #12.
1. Identify the substances involved in the oxidation and
reduction changes. Include substances so that all
elements are represented (except H and O).
For each half-reaction…
2. Balance all elements except H and O.
3. Balance O’s (add H2 O’s as needed).
4. Balance H’s (add H+’s as needed).
5. Balance charges (add e– ’s to the more positive side).
6. Balance e– ’s in the two half-reactions.
7. Combine the two half-reactions. Cancel substances that
show up on both sides of the equation (e – ’s must cancel).
8. If solution is basic, see card #12.
If the reaction occurs in a basic solution (usually stated
clearly in the problem) then instead of H+’s and H2 O’s, you
utilize OH– ’s and H2 O’s.
7 • Chemical Reactions & the Per. Table
Balancing Redox Equations
Reactions in Basic Solution
(12 of 16)
An easy method is:
- balance as though the reaction were in an acid solution.
- add OH– ’s to each side of the equation until all H+’s are
turned into H2 O’s.
- cancel H2 O’s that show on both sides of the equation.
In each case, the metal changes to the + ion: M → M+ + e–
Since the metal is oxidized, it is a reducing agent.
Metals that most easily lose e– ’s (those with low ionization
energy and low electronegativity) make the best reducers.
7 • Chemical Reactions & the Per. Table
Metals as Reducing Agents
(13 of 16)
M E M O R I Z E T H I S:
Some metals react with H2 O [to H 2 (g) and OH – ]
Some react with non-oxidizing acids such as HCl, and cold,
dilute H2 SO4 (the H+ is the reacting species) [to H 2 (g)]
Some react only with oxidizing acids:
• dilute HNO3 [to colorless NO(g) + H2 O],
• conc. HNO 3 [to red-brown NO 2 (g) and H2 O], and
• hot, conc. H2 SO4 [to SO2 (g) and H2 O]
FOUR GROUPS OF METALS
7 • Chemical Reactions & the Per. Table
The Activity Series of Metals
(14 of 16)
1 - Most active - Families I and II - great reducing agents
reduction half-reaction: 2H2 O + 2e– → H2 + 2OH–
2 - Most metals... Zn, Fe, Al, etc.
reduction half-reaction: 2H+ + 2e– → H 2
3 - Ag, Cu, Hg
ex. half-reaction: 2H+ + NO3 – + e – → NO2 + H2 O
4 - Noble Metals - Au, Pt, Ir
only changed by “aqua regia” [HNO3 + HCl forms Cl 2 ]
A more active metal can reduce or displace the ion of a less
active metal. Ex. Zn° is more active than Cu°, so…
Zn° + Cu2+ → Zn2+ + Cu° Cu° + Zn2+ → no reaction
Elemental nonmetals (O, Cl, F, S, etc.) form negative ions
by gaining electrons (reduction) and are oxidizing agents.
7 • Chemical Reactions & the Per. Table
Nonmetals as Oxidizing Agents
Oxygen as an Oxidizing Agent
Combustion as a Redox Reaction
(15 of 16)
7 • Chemical Reactions & the Per. Table
Amphiprotic/Amphoteric & Leveling Effect
(16 of 16)
The strongest oxidizing agents are those to the right of each
period (excluding noble gases) and those at the top of each
family. So we can predict that F>Cl>Br>I and O > S.
Example: Cl will displace Br – , but not F–
Cl2 (g) + 2NaBr → Br2 (l) + 2NaCl • Cl2 + NaF → N.R.
Oxygen, O2 , is a common and powerful oxidizing agent.
Corrosion of metals, formation of oxides, and combustion
are all examples of redox with O2 as the oxidizer.
Anything with a lone pair can act as a proton acceptor
(base) Anything with a H atom can act as a proton donor
(acid). See water on card 5 act as either an acid or a base.
See acetic acid on page 241 act as both acid and base. We
say these are amphiprotic or amphoteric substances.
Leveling effect...
You cannot tell which strong acid is strongest in water,
because donating a proton to water is not a good enough
challenge. Strong acids completely dissociate in water
because water is a pretty good acceptor of protons. We say
that water has a leveling effect. Acetic acid is amphiprotic,
but a poor proton acceptor and therefore is a great test for
which strong acid is the strongest...
8 • Ionic Reactions in Solution
Driving Forces for Metathesis Reactions
(1 of 12)
8 • Ionic Reactions in Solution
Precipitates as a Driving Force
The Solubility Rules
(2 of 12)
8 • Ionic Reactions in Solution
Weak Electrolyte Formation as a Driving Force
Weak Acids and other Weak Electrolytes
Neutralization Reactions
(3 of 12)
8 • Ionic Reactions in Solution
Gas Formation as a Driving Force
Gases That Commonly Form
(4 of 12)
During a double replacement or metathesis reaction, two
new combinations of ions are produced. We identify four
reasons why these NEW combinations are more stable
than the original combos.
• a precipitate forms
memorize your solubility rules
• a gas forms which leaves the system
memorize the list of gases that form
• a weak electrolyte forms
memorize the strong acid list so you will recognize weak
acids, also H2 O and NH4 OH
• a complex ion forms
learn the structure of complex ions and common ligands
Always Always Soluble compounds with alkali metal ions
(Li+, Na +, K+, Cs +, Rb +), NH4 +, NO3 – , C 2 H3 O2 – , ClO 3 –
& ClO4 –
Usually Soluble
Cl– , Br– , I– [except “AP/H”... Ag+, Pb 2+ , Hg2 2+ ]
SO4 2– [except “CBS”: Ca 2+ , Ba2+ , Sr2+ & “PBS”: Pb2+ ]
Usually NOT Soluble
O2–, OH– [except alkali and “CBS” Ca2+ , Ba2+ , Sr2+ ]
Never Soluble
CO3 2–, SO3 2–, S 2–, PO4 3– [except NH4 + & alkali]
NOTE: some of these insoluble compounds WILL dissolve
in acid solutions because of gas formation... useful idea!
Weak electrolytes are substances that break up into ions only
a LITTLE in solution... therefore, the two ions are MOSTLY
in a combined state... not likely to re-form the reactants.
H2 O, weak acids, NH4 OH
Memorize the 8 strong acids so you can recognize a weak
acid when you see one...
HCl, HBr, HI, HNO3 , H2 SO4 , HClO3 , HClO4 , HIO4
Acids (forming H+ ions) and bases (forming OH – ions)
combine to form a salt (an ionic compound) and H2 O... the
very weak electrolyte. Neutralization of the acid and base
occur because the H+ & OH– ions are “tied up” as H2 O.
If you see the following substances formed during
metathesis, realize that they will breakup into gases and
leave the system (preventing re-formation of the reactants).
Watch For…
It Turns Into…
H2 CO3
→ CO2 (g) + H 2 O
H2 SO3
→ SO2 (g) + H 2 O
H2 S
→ H 2 S(g) [rotten egg smell]
NH4 OH
→ NH3 (g) + H 2 O
2 HNO2
→ NO(g) + NO 2 (g) + H 2 O
NOTE: these compounds are formed from acids with
carbonates, sulfites, sulfides, nitrites, and bases with
ammonium compounds.
If you want to make the ionic compound, XY, you mix
AY + XB to make XY + AB
Either XY or AB need to drive the reaction (ppt, gas, etc.)
8 • Ionic Reactions in Solution
Preparation of Salts
(5 of 12)
8 • Ionic Reactions in Solution
Comparing Driving Forces
(6 of 12)
8 • Ionic Reactions in Solution
More Concentration Units
Weight Percent, ppm, and ppb
(7 of 12)
You may need to do this in two steps... make a carbonate (or
sulfite, sulfide, hydroxide or oxide) of the cation (+ ion) you
need and then react it with an acid that has the proper anion.
The Practical Side:
Keep in mind how you could recover the product you want...
could you filter the product mixture? Do you want what is
in the filter paper or what is in the filtrate? If you need the
filtrate, you need to be careful not to have excess ions in it.
Gas formation is a very strong driving force... even
compounds that exist as insoluble solids will react (slowly)
to form gases because gases leave the system and CANNOT
re-form reactants.
CaCO3 (s) + 2HCl → CO 2 (g) + H 2 O + Ca2+ + 2 Cl –
The tendency to form H2 O is very strong. Insoluble oxides
will react with acids.
ZnO(s) + 2HCl → Zn2+ + 2Cl– + H2 O
Sometimes, one insoluble solid can change into another even
MORE insoluble solid... but you need more than the
solubility rules to predict this (you need Ksp’s).
AgCl(s) + Br– → AgBr(s) + Cl–
Percent means “parts per 100”
96
96% means
, that is, 96 out of every 100.
100
96g
96g
Weight Percent (w/w) means
… (w/v) means
100g
100mL
ppm means “parts per million”
96
96 ppm means
, 96 out of every 1 million.
1 000 000
ppb means “parts per billion”
96
96 ppb means
, 96 out of every BILLION!
1 000 000 000
8 • Ionic Reactions in Solution
Chemical Analyses
Precipitations, Combustions,
and Titrations
(8 of 12)
Real chemistry often deals with testing what is in a
particular reaction mixture, environmental sample, etc.
Stoichiometry is used to analyze the compositions.
• You can precipitate an ion you are interested in, filter the
precipitate and then determine from its mass the amount
of compound in the original sample.
• You can burn a sample and collect the combustion
products (CO2 & H2 O) to determine the amount of C and
H in the original sample.
• You can carefully measure the volumes of solutions used
during a titration. The endpoint must have some sort of
indicator to allow you to recognize when the correct
amounts of reactants have been added.
8 • Ionic Reactions in Solution
Titration Terminology
Acid-Base and Redox Titrations
(9 of 12)
A titration is a volumetric analysis because you carefully
measure the volume of titrant, dispensing it from a buret.
When you have added just enough titrant to completely react
with the sample, you have reached the endpoint. This is
usually apparent because of the color change of some
indicator molecule (such as phenolphthalein). The endpoint
can also be tracked because of changes in pH or changes in
voltage due to the amount of some ion.
Acid-Base & Redox titrations follow the formula: V·N=V·N
where the N indicates [H+] or [OH– ] in acid-base titrations
and [Oxidizer]· e– gained or [Reducer] · e– lost in redox
titrations.
8 • Ionic Reactions in Solution
Three Most Common Oxidizing Agents
and what they turn into
(10 of 12)
purple permanganate ion
acid solution MnO4 – + 8H+ 5e– → Mn2+ + 4H2 O
Mn2+ ion is colorless
neutral/basic MnO4 – + 2H2 O + 3e– → MnO2 (s) + 4OH –
MnO2 (s) is a black solid
yellow chromate / orange dichromate ion depends on [H+]
→ Cr2 O7 2– + H2 O
2CrO4 2– + 2H+ ←
2–
acid solution Cr2 O7 + 14H+ + 6e– → 2Cr3+ + 7H2 O
slightly basic CrO4 2– + 4 H2 O + 3e– → Cr(OH)3 + 5OH–
Cr(OH)3 is a solid
very basic
CrO4 2– + 2H2 O + 3e– → CrO2 – + 4OH–
8 • Ionic Reactions in Solution
Common Reducing Agents
and what they turn into
(11 of 12)
Tin(II) (a gentle reducing agent)
Sn 2+ → Sn4+ + 2 e –
Sulfites and Bisulfites
acidic solution: HSO3– + H2 O → SO4 2– + 3H+ + 2e–
basic solution: SO3 2– + 2OH– → SO 4 2– + H2 O + 2e–
Thiosulfate ion (also called “hypo” in photography)
strong oxidizer: S2 O3 2– + 5H2 O → 2SO4 2– + 10H+ + 8e–
half-reaction w/I2 :
2S 2 O3 2– → S4 O6 2–+ 2e –
complete:
I2 + 2S2 O3 2– → 2I– + S4 O6 2–
–
excess I :
I2 + I– → I3 –
→ starch•I2 complex (blue-black)
I2 + starch ←
8 • Ionic Reactions in Solution
Equivalents, Equivalent Weights,
and Normality
(an old-fashioned, but useful idea)
(12 of 12)
equivalents = H+, OH– , or electrons gained or lost
n = # of equivalents in the balanced chemical equation
example: I2 + 2S2 O3 2– → 2I– + S4 O6 2– n = 2
equivalent weight = molar mass ÷ n
…mass of a chemical that provides 1 mole of equivalents.
moles equivalents
Normality, N, = n · M, =
Liter solution
This idea is useful in acid-base and redox titrations because
it takes into account the differences of acids and bases or
oxidizers and reducers. This concept allows the use of the
simple formula: V·N = V·N
9 • Properties of Gases
Boyle’s Law (P and V)
(1 of 12)
General:
Formula:
When P↑, V↓ (inversely proportional)
P·V = constant or P1 V1 = P2 V2
Restrictions:
P 1 and P2 must be in the same units
V1 and V 2 must be in the same units
Convert pressures using conversion factors using the fact
that 1 atm = 760 mmHg = 760 torr = 101.3 kPa = 14.7 psi
lb
psi = 2
in
101.3 kPa
Example:
730 mmHg x
= 97.3 kPa
760 mmHg
Graphically:
V
1/V
9 • Properties of Gases
Boyle’s Law Lab
(2 of 12)
P
P
In our lab, we had to add the atmospheric pressure to our
measurements because tire gauges only measure the
pressure ABOVE atmospheric pressure.
1
Consistent (“ good”) data form a straight line (P vs. ).
V
K = °C + 273
°C = K – 273
Examples: 0 °C + 273 = 237 K
25 °C + 273 = 298 K
100 °C + 273 = 373 K
9 • Properties of Gases
Kelvin Temperature Scale
(3 of 12)
300 K – 273 = –27 °C
The Kelvin scale is used in gas law problems because the
pressure and volume of a gas depend on the kinetic energy
or motion of the particles.
The Kelvin scale is proportional to the KE of the
particles… that is, 0 K (absolute zero) means 0 kinetic
energy. 0 °C is simply the freezing point of water.
Charles’ Law
General:
9 • Properties of Gases
Charles’ Law (V and T)
Gay-Lussac’s Law (P and T)
(4 of 12)
When T↑, V↑ (directly proportional)
V
V
V
Formula:
= constant or 1 = 2
T
T1 T2
Restrictions:
T must be in Kelvins
V1 and V 2 must be in the same units
Gay-Lussac’s Law
General:
When T↑, P↑ (directly proportional)
P
P
P
Formula:
= constant or 1 = 2
T
T1 T2
Restrictions:
T must be in Kelvins
P 1 and P2 must be in the same units
Formula:
Restrictions:
9 • Properties of Gases
The Combined Gas Law
(5 of 12)
P ·V P ·V
P·V
= constant or 1 1 = 2 2
T
T1
T2
T must be in Kelvins
V1 and V 2 must be in the same units
P 1 and P2 must be in the same units
STP (“standard temperature and pressure”) is often used as
one of the two conditions
T = 0 °C = 273 K P = 1 atm = 760 mmHg = 101.3 kPa
Each of the three gas laws is really a special case of this
law.
Example: If T1 = T2 , the law becomes P1 V1 = P2 V2
Formula:
where
9 • Properties of Gases
The Ideal Gas Law
(6 of 12)
9 • Properties of Gases
Dalton’s Law of Partial Pressure
(7 of 12)
P·V = n·R·T or PV = nRT
P = pressure
V = volume
n = number of moles
R = the ideal gas constant
T = temperature (in Kelvins)
The value of R depends on the P and V units used.
PV
R=
so you can use the molar volume info to calculate R
nT
(101.3 kPa)(22.4 L)
L·kPa
R=
= 8.31
(1 mole)(273 K)
mol·K
L·mmHg
L·atm
R = 62.4
= 0.0821
mol·K
mol·K
When you have a mixture of gases, you can determine the
pressure exerted by each gas separately. This is called
the partial pressure of each gas.
Since each gas has the same power to cause pressure (see
card #8) the partial pressure of a gas depends on how
much of the mixture is composed of each gas (in moles)
Example: Consider air, a mixture of mostly O2 and N 2
P O2
P N2
moles O 2
moles N 2
=
=
moles total
P total
moles total P total
Also: P total = P O2 + P N2
This idea is used when a gas is collected over water
P atm = P gas + P H2 O P H2 O is found on a chart
The gas laws work (to 3 significant digits) for all gases…
that is, all gases have the same power to cause pressure.
9 • Properties of Gases
Why Do All Gases Cause the Same Pressure?
(8 of 12)
At the same temperature, the KE of each gas is the same.
KE = 1 /2 mass·velocity2 … if two particles have different
masses, their velocities are also different. So…
SMALL
particles move FAST
LARGE particles move SLOWLY
v2
m
m
v2
We can use this idea with numbers as well: (Graham’s Law)
KEA = KEB
mAvA2 = mBvB2
[another version of this formula is on the next card]
mAvA2 = mBvB2 can also be used as the equation…
rate of effusion of A
rate of effusion of B =
9 • Properties of Gases
Graham’s Law of Effusion
(9 of 12)
MB
MA
Notice that the A is in the numerator in the ratio of the
rates and in the denominator in the radical.
“Effusion” is similar to diffusion. It means to escape
through a small opening.
The ratio of the rates (or velocities) of CH4 (mass=16 u) to
SO2 (mass=64 u) is
64
=
16
4 = 2
Ideal gases have no volume & no attractions for each
other. Luckily, real gases act pretty much like ideal gases at
room temperature and pressure. The most ideal of real gases
is He.
9 • Properties of Gases
The Real Gas Law
(10 of 12)
9 • Properties of Gases
Kinetic Molecular Theory
(11 of 12)
The REAL GAS Law is:
n2a
(P + 2 )(V - nb) = nRT
V
where:
a corresponds to the attractions between real gas particles
b corresponds to the size of the real gas particle
Explaining the behavior of gases involves the kinetic
molecular theory. Here are the main ideas:
• all particles are in constant, random motion
• temperature is a measure of the average kinetic energy
• pressure is due to collisions of gas particles with the walls
of the container
• increased temperature causes more collisons as well as
harder collisions
• some particles are moving fast, some are moving slowly
F
P=A
Pressure is proportional to the force pushing and inversely
proportional to the area over which that force pushes.
9 • Properties of Gases
Pressure = Force ÷ Area
(12 of 12)
P=F
A
P=F
A
Properties of matter depend on the model that gas particles
are spread out; liquids are close together, but random; and
solids are close together and arranged in a crystal lattice.
10 • States of Matter & IMFs
Comparing Gases, Liquids, and Solids
(1 of 16)
gas
liquid
solid
Volume and shape, compressibility, and the ability of
substances to diffuse depend on these models.
Gases have no set shape or volume.
Liquids have a constant volume, but no set shape.
Solids have constant volume and shape.
Surface tension- a measure of the amount of energy needed
to expand the surface area of a liquid.
10 • States of Matter & IMFs
Surface Tension
(2 of 16)
An interior molecule is surrounded by molecules to which it
is attracted… no net attraction. A surface molecule feels a
net attraction toward the interior. To move a molecule to the
surface (i.e. increase the surface area), energy must be used,
work must be done. The potential energy of the liquid is
increased.
Substances tend toward the lowest potential energy so
liquids tend toward the minimum surface area. A sphere is
the smallest surface area for given volume.
10 • States of Matter & IMFs
KE Distributions and Evaporation
(3 of 16)
In any sample of liquid, the
distribution of KE varies.
Particles to the right of the line
(the “threshold energy” have
enough KE to escape the IMF’s
holding them in the liquid.
Increasing the temperature (average KE) of the liquid moves
the curve to the right. The line depends on the IMF of the
liquid.
Only particles at the surface of the liquid may escape
(evaporate.)
10 • States of Matter & IMFs
Molecular Crystals and IMFs
(4 of 16)
Substances that exist as molecules (as opposed to ionic,
metallic, or covalent network crystals) are in three groups:
nonpolar
polar molecules
polar with H-O-,
molecules
Dipole-dipole
H-N-, or H-F
London Forces
attractions
H-bonding
Weak IMF. Due
+ end of one
Strong dipole
to polarizable emolecule
because of high
clouds & temp.
attracting - end
electroneg. /small
attraction
of other molecule size of O, N, & F
In London Forces--larger atoms and larger molecules have
stronger London forces due to more sites or more
polarizable electron clouds .
Heat of Vaporization, ∆ Hvap, can be thought of as the energy
needed to vaporize a mole of a liquid.
It can be used as a conversion factor in a calculation of heat
during a phase change. [Calorimetry is used for temperature
changes between the phase changes.]
10 • States of Matter & IMFs
Hvap and IMFs
(5 of 16)
Ex: 111g H 2 O ×
1 mol H 2 O
40.6 kJ
×
= 250. kJ
18.0 g H 2 O 1 mol H 2 O
It can also serve as a indicator of the strength of the IMF
(intermolecular forces of attraction) in the liquid.
Ex: CH4 (9.20 kJ/mol) vs. C3H8 (18.1 kJ/mol)
Larger molecule… greater IMF… greater Hvap
In a closed container, the number of particles changing from
liquid → vapor will eventually equal the number of particles
changing from vapor → liquid.
10 • States of Matter & IMFs
Vapor Pressures of Liquids
(6 of 16)
The amount of vapor when this balance is reached depends
on the IMF and the KE of the liquid (& not on the volume of
the container). The pressure exerted by this vapor is called
the equilibrium vapor pressure (VP) of the liquid.
As temperature increases, the VP increases. (This is
important for why/when a liquid boils.)
VP is another indicator of the strength of the IMF. The
stronger the IMF, the smaller the VP.
Boiling occurs when the vapor pressure (VP) of a substance
= the air pressure above the liquid.
You can boil a liquid by increasing the VP of the liquid
(heating) or by lowering the pressure above the liquid.
10 • States of Matter & IMFs
Boiling Point and IMFs
(7 of 16)
The temperature at which a liquid reaches 760 mmHg is called
the “normal boiling point” of the liquid.
---Again, BP is an indicator of IMF.--↑ boiling points (BP) ~ ↑ IMF’s ~ ↓ vapor pressures.
Altitude (low air pressure) lowers the boiling temperature of
water in an open container (increases cooking time).
Pressure cookers ↑ BP by ↑ the pressure above the liquid.
Freezing Point and Melting Point are the same
temperature, just opposite directions. That is, a substance
will freeze and melt at the same temperature.
10 • States of Matter & IMFs
Freezing Point, Melting Point, ∆ Hfusion
(8 of 16)
When you fuse two metals, you MELT them… thus the term
fusion means melting. ∆ Hfusion is the energy needed to melt
a mole of solid into a mole of liquid.
As with ∆Hvap, ∆Hfus can be used as a conversion factor as
well as an indicator of IMF strength.
Note: Freezing is more complicated than vaporization
because the process of forming a crystal causes some subtle
considerations which we will not deal with in this course.
10 • States of Matter & IMFs
Common Crystal Structures and Unit Cells
(9 of 16)
Three common crystal structures are Simple Cubic; BodyCentered Cubic (which means there is an atom, “a body,” in
the center of the cubic structure; and Face-Centered Cubic
(with an atom in the center of each side or “face”).
A unit cell is the theoretical
arrangement of atoms that, if
repeated, will recreate the crystal.
This topic, although interesting,
is no longer on the AP curriculum
and will not be dealt with here.
One resource is a page by OSU
From Dr. John Gelder’s
Solid State Chemistry Page chemist, Dr. John Gelder. (See Card 16)
Crystal
lattice
points:
IMF’s
10 • States of Matter & IMFs
Four Types of Crystals -- Summary
(10 of 16)
Props.
Ex
Molecular
molecules
or atoms
London,
dipole,
H-bonding
soft, low
MP, nonconduct
Ionic
+ and - ions
Covalent
atoms
Metallic
positive ions
attraction
between +
& - ions
hard, brittle,
high MP, (l)
(aq) conduct
covalent
bonds
I2 H 2O HI
NaBr
attr. between
+ ions & “sea
of e-‘s”
high luster,
conductor,
variable MP,
soft/hard
Na° Fe° Cu°
v. hard, high
MP,
nonconduct
(graphite)
C SiC WC
Students are often confused between molecular crystals in
which covalent bonds hold the molecules together (but the
IMF = London forces, dipole-dipole attractions or hydrogen
bonding) and covalent crystals in which covalent bonds hold
the crystal together (the IMF = covalent bonds).
10 • States of Matter & IMFs
Crystal Types -- Further Notes
(11 of 16)
Substances can conduct electricity for two reasons:
freely moving ions or delocalized electrons.
Ionic compounds have freely ions in the liquid state and
when dissolved in water.
Metals have delocalized electrons -- the “sea of electrons.”
Graphite has a chicken-wire shaped π-bond above & below
each sheet of sp 2-hybridized C atoms, allowing it to conduct.
(g)
(l)
10 • States of Matter & IMFs
Heating and Cooling Curves
(12 of 16)
(s) & (l)
(s)
PE
KE
(l) & (g)
PE
KE
KE
KE = kinetic
energy changes
which are times
when the heat
energy speeds up
the molecules.
Time (min)
PE = potential energy changes which are times when
the heat energy separates the molecules from solid to
liquid or liquid to gas.
The phase diagram shows the
phases of a substance at all
temperatures and pressures.
10 • States of Matter & IMFs
Phase Diagrams -I
(13 of 16)
Moving across the diagram
gives you the MP and then
BP of a substance.
There is a point above which it is no longer possible to
liquefy a substance (the critical point, C).
Moving vertically you can see the effect of pressure on the
phase of the substance.
This is a diagram for a substance like CO2, in which the liquid
can be compressed into the solid. (Unlike H2O.)
Water’s phase diagram is unique
because the liquid phase is less
dense than the solid phase. To
maximize hydrogen bonding, the
solid must expand.
10 • States of Matter & IMFs
Phase Diagrams -II
(14 of 16)
10 • States of Matter & IMFs
Name of the Phase Changes
(15 of 16)
The B-D boundary of the phase
diagram has a negative slope.
The “triple point” is the temperature and pressure in which
the solid, liquid, and vapor phases of a substance can coexist. I visualize this as a boiling glass of ice water.
By increasing the pressure, dry ice can melt. By decreasing
the pressure, solid water can sublime.
NRG is REQUIRED
solid → liquid
melting or fusion
NRG is RELEASED
liquid → solid
freezing
liquid → gas
vaporization
evaporation or boiling
gas → liquid
condensation
solid → gas
sublimation
gas → solid
solidification
The energy involved in the phase change is calculated using
heat of fusion (solid → liquid or liquid → solid)
heat of vaporization (liquid → gas or gas → liquid)
Searching the Internet, I found an interesting set of topic
reviews. These are from Purdue University (Indiana)
I was looking at the topic, LIQUIDS, but there are many
topics to choose from.
10 • States of Matter & IMFs
More Internet Resources
(16 of 16)
chemed.chem.purdue.edu/genchem/topicreview/
The unit cell is a frame from an online movie file on Dr. John
Gelder’s Solid State Chemistry page. (OK State Univ.)
www.okstate.edu/jgelder/solstate.html
thermodynamics
system
12 • Chemical Thermodynamics
Commonly Used Terms
(1 of 12)
surroundings
adiabatic
isothermal
state functions
Study of NRG changes & flow of NRG.
Tells whether a reaction is possible.
That portion of the universe on which
we focus.
Everything outside the system.
A change without heat transfer between
the system and its surroundings.
A change that occurs at constant
temperature.
Properties that depend on the initial and
final state, not on how the change was
made. ex: ∆H, temp, but not work
Energy changes can be measured using calorimetry. Often
this involves heating water under controlled conditions (a
bomb calorimeter).
12 • Chemical Thermodynamics
Heat Capacity and Specific Heat
(2 of 12)
Three closely related terms are:
heat capacity is the amount of heat needed to change a
system by 1°C.
molar heat capacity is the amount of heat needed to change
a mole of a substance by 1°C.
specific heat is the amount of heat needed to change 1 gram
of a substance by 1°C. (Water: 1 cal/g°C = 4.184 J/g°C)
Note that heat capacity is an extensive property whereas the
other two are intensive.
The first law: “if a system undergoes some series of changes
that ultimately brings it back to its original state, the net
energy change is zero.”
∆E = Efinal - Einitial
12 • Chemical Thermodynamics
First Law of Thermodynamics
(3 of 12)
∆E = 0 when Efinal = Einitial
The usefulness of this idea is that the internal energy, ∆E,
depends only on the initial and final state, not on how you
get there. Every path you take from Einitial to Efinal takes the
same amount of energy. (no perpetual motion)
There are two ways for a system to exchange energy with the
surroundings, heat and work.
∆E = q - w
[q = heat absorbed by system, w = work done by system]
From physics, work = force × distance = F × d.
In chemistry there is electrical work (e-‘s through a wire)
and P∆
∆ V work as a gas expands.
12 • Chemical Thermodynamics
Work and PV Work
(4 of 12)
Use pressure = force/area and area × distance = volume to
derive work = P∆
∆ V from work = F × d.
Note that units of work are units of energy. Energy and work
are two forms of energy. Doing work on a system increases
the potential energy of the system.
P∆V work is in L⋅⋅ atm
1 L⋅atm = 24.2 cal = 101.3 J
PV = nRT… work can be calculated as work = ∆ nRT for
chemical reactions where the # of moles of gas change, ∆n.
12 • Chemical Thermodynamics
Work is NOT a State Function
(It depends on HOW you do the work.)
(5 of 12)
Q: Gas in a piston at a pressure of 10 atm is allowed to
expand from 1 L to 10 L. How much work done?
A: It depends on the resisting force.
If it expands against 0 atm pressure,
P∆V = 0 atm × 9 L = 0 L⋅⋅ atm
Expanding against 1 atm pressure, Vfinal is 10 L
P∆V = 1 atm × 9 L = 9 L⋅⋅ atm
Expanding in two steps, first, against 2 atm, Vfinal = 5 L
P∆V = 2 × 4 L = 8 L⋅atm
then in a second step against 1 atm, Vfinal = 10 L
P∆V = 1 × 5 L = 5 L⋅atm… Total = 13 L⋅⋅ atm
From the P∆V work example, you can see that the more steps
(and smaller increments) you use to get the work from the
system, the more work you can get.
12 • Chemical Thermodynamics
Reversible Processes
(6 of 12)
There is an upper limit to how much work can be derived
from any system. That maximum work is called the Gibb’s
Free Energy, ∆G.
The theoretical maximum work can be achieved if the steps
are so small that they can go either way, if they are
reversible.
If a chemical reaction occurs in a closed container ∆ V = 0
and so P∆
∆ V work = 0.
∆ E = q - w becomes ∆ E = q.
12 • Chemical Thermodynamics
∆ H = ∆ E + P∆
∆V
(7 of 12)
In a system at constant pressure (a common situation) the
energy change involves both heat (q) & work (P∆V).
∆ E = q - P∆
∆ V [P∆V = work done BY the system] or
q = ∆ E + P∆
∆ V q at constant pressure is called ∆ H.
∆ H = ∆ E + P∆
∆ V or ∆ E = ∆ H - P∆
∆V
The work (P∆V) is generally insignificant unless the # of
moles of gas (∆n = n final - n initial) is changing. P∆
∆ V = ∆ nRT
Recall that if ∆n is negative (# moles decreasing) work is
negative (work is being done ON the system).
If several reactions add up to give an overall reaction, the
∆ H’s of the reactions will add up to the overall ∆ H.
12 • Chemical Thermodynamics
Hess’s Law of Heat Summation
(8 of 12)
Standard Heats of Formation, ∆ Hf° , are useful for this
purpose. This is the energy involved in making a mole of a
substance from its elements at 25°C and 1 atm pressure.
This law is often written as:
∆ H°° reaction = Σ ∆ Hf° products - Σ ∆ Hf° reactants
Note: ∆Hf° for elements is 0.
If you are NOT using heats of formation, you need to write
out the equations to see how they combine.
Bond energy is the amount of energy needed to BREAK a
certain bond.
12 • Chemical Thermodynamics
Bond Energies
(9 of 12)
You can determine the approximate energy change in a
chemical reaction by summing the bond energies of the
reactants and subtracting the bond energies of the products.
Note: this is opposite to the Hess’s Law (products reactants) because bond energy involves breaking bonds
whereas ∆H°f involve forming bonds.
Things in the world tend toward lowest energy (-∆H) and
also tend toward greatest disorder (+∆S).
12 • Chemical Thermodynamics
Two Big Driving Forces of the Universe
Enthalpy (∆
∆ H) and Entropy (∆
∆ S)
(10 of 12)
More disorder can be recognized as:
• greater # of moles of gas formed
• gas > liquid > solid and (aq) > (s)
• greater volume formed
• mixed molecules (HI) formed from diatomic molecules
example: H2 (g) + I2 (g) → 2HI(g)
+∆S
Note: ∆S° can be calculated using Hess’s Law, but S° of
elements is NOT 0. Also, S° is often reported in J/mol⋅K
and cal/mol⋅K, not kJ and kcal… watch your units!
The second law defines a value called Gibb’s Free Energy
symbolized as ∆G. This represents the theoretical
maximum work that can be done by a system.
12 • Chemical Thermodynamics
The Second Law of Thermodynamics
∆ G = ∆ H - T∆
∆S
(11 of 12)
A reaction is spontaneous when ∆G is negative (∆G<0). The
reverse reaction is spontaneous when ∆G is positive.
When ∆G = 0, the reaction is at equilibrium.
∆H


+
+
∆S
+

+

Spontaneous…
… at all temperatures
… at LOW temperatures
… at HIGH temperatures
the reverse reaction is spontaneous
We have been calculating ∆G° values, that is, ∆G under
standard conditions. Once a reaction begins, the
concentrations change and the value of ∆G is different.
12 • Chemical Thermodynamics
∆ G (thermodynamics) and Keq (equilibrium)
(12 of 12)
If ∆ G°° < 0, then the reaction will proceed in a forward
direction. As it does, however, the value of ∆G will increase
until the forward and reverse reactions are balanced and ∆ G
= 0. This is called equilibrium.
If ∆ G°° > 0, then the reverse reaction will proceed and the
value of ∆G will decrease until equilibrium is reached.
∆G° gives information about where we are in relation to
equilibrium. There is a formula that combines ∆ G and Keq.
22 • Nuclear Chemistry
The People
(1 of 16)
22 • Nuclear Chemistry
Terms I-- Radioactivity
(2 of 16)
22 • Nuclear Chemistry
Terms II--Radioactivity
(3 of 16)
• Wilhelm Roentgen (1845-1923) discovered X-rays, a high
energy form of light. (1895)
• Henri Becquerel (1852-1909) found that uranium ores emit
radiation that can pass through objects (like x-rays) and
affect photographic plates. (1896)
• Marie Sklodowska Curie (1867-1934) Marie and Pierre
worked with Becquerel to understand radioactivity. The
three shared a Nobel Prize in Physics in 1903. Marie won a
second Nobel Prize in Chemistry in 1911 for her work with
radium and its properties.
• E. O. Lawrence invented the cyclotron which was used at
UC Berkeley to make many of the transuranium elements.
radioactivity
the spontaneous breakdown of atomic
nuclei, accompanied by the release of
some form of radiation (also called
radioactive decay)
half-life
time required for half of a radioactive
sample to decay
transmutation
one element being converted into another
by a nuclear change
nuclides
isotopes of elements that are identified by
the number of their protons and neutrons
decay series
the sequence of nuclides that an element
changes into until it forms a stable nucleus
radioactive
dating
using half-life information to determine
the age of objects. C-14/C-12 is common
for organic artifacts. Uranium is common
for rocks.
nuclear fission
large nucleus breaking down into pieces of
about the same mass
nuclear fusion
two or more light nuclei blend to form one
or more larger nuclei
Alpha particles are the same as a helium nucleus,
4
2
He, with
a mass of 4 amu. It travels about 1/10th the speed of light
and is the most easily stopped of the three particles (a sheet
of paper will stop them). It is the least dangerous.
22 • Nuclear Chemistry
Types of Radiation
(4 of 16)
0
Beta particles are high speed electrons, −1 e, with a mass of
0.00055 amu and travel at nearly the speed of light. They can
be stopped by a sheet of aluminum. It is more penetrating
and therefore more dangerous than alpha.
Gamma rays are extremely high energy light, γ, with no mass,
and are the most penetrating (several cm’s of lead are needed
to stop them). They can cause severe damage.
In each half-life problem there are basically four variables:
• total time
• half-life
• starting amount
• ending amount
64g
22 • Nuclear Chemistry
Half-Life Problems
(5 of 16)
32g 16g
8g
4g
2g
1g
0.5g 0.25g
Question:
If you have 0.25 g of a radioactive substance with a half life
of 3 days, how long ago did you have 64 grams?
Answer: Draw the chart to determine the number of halflives to get from the ending amount to the starting amount…
each half-life is worth 3 days…24 days.
Half-Life
The time it takes for half of a radioactive
substance to decay.
The decay graph has a characteristic shape:
#
22 • Nuclear Chemistry
Half-Life
(6 of 16)
time
The time it takes for the amount of substance or the activity
of the substance to drop to half is the same WHEREVER you
start on the graph. This is a first-order reaction.
Half-lives can range from microseconds to thousands of
years and is characteristic of each substance.
Memorize the symbols for the important particles
alpha
beta
positron
neutron
4
2
22 • Nuclear Chemistry
Nuclear Equations
(7 of 16)
He
0
−1
e
0
+1
1
0n
e
Decay means the particle is on the right side of the equation:
example: alpha decay of U-238
238
4
92 U → 2
He +
234
90 Th
The 234 and 90 are calculated… the Th is found on the
periodic table (find the element with atomic # = 90).
Several neutrons can be shown together and written as…
1
3
3( 0 n) and would be counted as 0 n in the equation.
22 • Nuclear Che mistry
How Each Type of Decay Can Stabilize an
Unstable Nucleus
(8 of 16)
Certain values of p +’s and n°’s in the nucleus are stable. A
nucleus can be unstable (radioactive) for 3 reasons:
• the nucleus has too many protons compared to neutrons
solution: positron decay
(change a proton into a neutron and a positive electron…
…a positron)
• the nucleus has too many neutrons compared to protons
solution: beta decay
(change a neutron into a proton and a negative beta particle)
• the nucleus is too big (too many protons and neutrons)
solution: alpha decay (lose 2 p + and 2 n°)
22 • Nuclear Chemistry
Uses of Radioactivity
(9 of 16)
Radioactive Dating: In every living thing there is a constant
ratio of normal C-12 and radioactive C-14. You can calculate
the time needed to change from what is expected to what is
actually found.
Radioisotopes: Many substances can be radioactive and
then followed as they move through the body.
Fission Reactors: Current nuclear reactors use fission
reactions to produce heat which is used to turn water into
steam and drive turbine engines that produce electricity.
The Sun and Stars are powered by nuclear fusion… this is
related to the fact that the most abundant element in the
universe is hydrogen… followed by helium.
U-235 is “fissionable” which means it can be split when
bombarded by neutrons.
235
1
92 U + 0 n
22 • Nuclear Chemistry
Fission and Fusion Reactions
(10 of 16)
→
141
56 Ba
92
36
+
1
Kr + 3 0 n + energy
The fact that each splitting nucleus can emit neutrons that
can split other nuclei is the basis for the “chain reaction.”
“Breeder reactors” use different isotopes.
Fusion in the Sun involves several steps that can be
1
4
0
summed up as: 4( 1 H) → 2 He + 2 1 e + energy
Thermonuclear devices use isotopes of hydrogen
(deuterium and tritium):
2
1
3
H+ 1 H→
4
2
1
He + 0 n + energy
Einstein’s famous equation, E = mc2 , is the basis for
explaining where the energy associated with nuclear changes
comes from.
22 • Nuclear Chemistry
Energy–Mass Conversions
(11 of 16)
When a nuclear change occurs, the mass of the products is
slightly less than the mass of the reactants. This loss in
mass is called the mass defect.
E = the energy
m = the mass defect
c = the speed of light, 3.00 x 108 m/s
1 kg of mass converted into energy would be equivalent to
burning 3 billion kg of coal!
During beta decay,
1 neutron changes into 1 proton + 1 negative beta particle
(The atomic # increases by one due to the new proton. The
mass # is unchanged… a neutron is gone. To maintain
electrical neutrality, a negative beta particle is also formed.)
22 • Nuclear Chemistry
What Happens During
Beta and Positron Decay
(12 of 16)
Example:
235
0
92 U → −1
e+
235
93 Np
During positron decay,
1 proton changes into 1 neutron + 1 positron particle
(The atomic # decreases by one due to the loss of a proton.
Since it changed into a neutron, the mass # is unchanged.)
Example:
235
0
92 U → +1 e
+
235
91 Pa
When a problem involves whole numbers of half-lives,
divide by 2 to determine the amounts involved.
For other situations, the following equations are useful:
ln
22 • Nuclear Chemistry
Calculating Half-Lives
(13 of 16)
[A] 0
= kt and the special case for half-life, t ½, where by
[A] t
definition, [A]t = ½[A]0 ln 2 = 0.693= kt ½
[A] is the concentration (or activity) of the radioactive
substance, t = time, k = the rate constant (the same that is in
Rate Laws). Note: if you know the half-life, you can calculate
the rate constant and vice-versa.
Once a nucleus decays, the daughter isotope is often
unstable as well. Many decays may occur before a stable
nucleus is formed.
22 • Nuclear Chemistry
Radioactive Decay Series
(14 of 16)
A classic example is U-238 that decays through 14 steps into
stable Pb-206. Each step has a characteristic decay particle
and half-life.
This characteristic decay series is the method used to
verify the identity of newly formed atoms.
The fact that daughter products can be even more
radioactive than the parent isotope adds to the problem of
nuclear waste and its storage/disposal.
An useful characteristic of decay particles are that they
ionize the air they pass through by striking atoms and
knocking off electrons.
22 • Nuclear Chemistry
Geiger-Muller Tubes, Smoke Detectors, and
Brushes for Cleaning Negatives
(15 of 16)
Geiger counters use this idea. As radioactive particles pass
through a chamber with two electrodes, ionized particles
migrate to the + and - electrodes and complete the circuit.
Smoke Detectors use a tiny piece of radioactive Am to keep
a circuit flowing due to ionized particles. Smoke particles
attract ionized particles, break the circuit, & set off the alarm.
Brushes are kept ionized by tiny bits of radioactive material
to more easily attract tiny bits of dust.
Uranium, Z=92, is the largest naturally-occurring element.
Larger atoms were manufactured. Elements 93 and 94 were
formed in atomic bomb tests and identified by Seaborg.
Glenn Seaborg and Al Ghiorso at UC Berkeley were able to
use E. O. Lawrence’s cyclotron to make larger atoms (95-103).
22 • Nuclear Chemistry
Extending the Periodic Table
(16 of 16)
Some of these new elements have uses in the medical field as
well as helping to further the understanding of the nucleus.
For many of the larger elements, however, only a few atoms
or even one atom formed. They were identified by their
characteristic decay series.
As of July 2000, 118 is the largest element.
Chemicals from living things were thought to contain a “vital
force” that could not be duplicated in the lab. This changed
with Friedrich Wöhler who mixed cyanic acid (HCNO) with
ammonium hydroxide making ammonium cyanate (NH4CNO).
23 • Organic Chemistry
Historical Ideas
(1 of 16)
O
C
NH2
H2 N
urea
He usually allowed the salt solution to
evaporate overnight, but tried heating it to
hurry the process. The result was a crystal
that he recognized as urea (H2NCONH2).
The modern view of organic chemistry is the chemistry of
carbon compounds. C is the key element. It can form four
bonds and that are very strong bonds due to its small size.
The alkanes (paraffins) follow the formula: CnH2n+2:
These molecules contain ONLY single bonds. They are said
to be “saturated” with hydrogens.
23 • Organic Chemistry
Alkane Series -- Saturated Hydrocarbons
(2 of 16)
Memorize these prefixes also used with alkenes & alkynes.
CH4
methane
C6H14 hexane
C2H6
ethane
C7H16 heptane
C3H8
propane
C8H18 octane
C4H10 butane
C9H20 nonane
C5H12 pentane
C10H22 decane
Given a formula, you can tell that it contains only single
bonds because it fits the alkane formula.
As the molecules increase in size, they tend to be liquids and
23 • Organic Chemistry
Structural Formulas Can Be Misleading
(3 of 16)
CH4, can be drawn using a structural formula.
This can be misleading. The molecule is not flat
with bond angles of 90°. You must be aware of
the 3-D structure and the 109.5 ° bond angles.
Cl
H
C
Cl
H
H
C
Cl
Cl
H
H
H
C
H
H
For example, there is only one isomer
of dichloromethane, but you can
draw it at least two ways.
Building models of the molecules is an important part of
strengthening this skill.
H
Alkenes contain 1 double bond. The formula
is CnH2n. They are said to be “unsaturated”
(like unsaturated fats). The double bond can
be broken and more hydrogens added.
23 • Organic Chemistry
Alkenes and cis- /trans- Isomerism
(4 of 16)
H
C
C
H
ethene
Since double bonds cannot easily rotate (due to the pi
bonding) cis- and trans- isomers can be formed.
Example: 1,2-dichloroethene can be built two ways.
Cl
Cl
C
H
Cl
C
H
C
H
cis -1,2-dichloroethene
(a polar molecule)
H
C
Cl
trans-1,2-dichloroethene
(a nonpolar molecule)
H
23 • Organic Chemistry
Alkynes, Alkadienes, and Cyclic Hydrocarbons
(5 of 16)
H C C H
Alkynes contain 1 triple bond
(unsaturated). Formula: CnH2n-2.
ethyne (acetylene)
The triple bond is linear, so no cis/trans isomerism occurs.
Alkadienes are molecules with two double bonds. They
have the same formula as the alkynes, CnH2n-2.
H
Example: C4H6 is named 1,3-butadiene
C
CH2
because the double bonds start on
H2C
C
carbons #1 and #3.
H
H2
C
H2C
Cyclic compounds contain rings having the
same formula as the alkenes, CnH2n.
Example: cyclopropane, C3H6.
CH2
The basic idea is to name the molecule after the longest
continuous chain of carbon atoms. Side groups are listed
with #’s to indicate the C atom to which they are attached.
23 • Organic Chemistry
Naming Organic Compounds
(Organic Nomenclature Using IUPAC Rules)
(6 of 16)
23 • Organic Chemistry
Common Errors in Drawing/Naming Structures
1-methylsomething & 2-ethylsomething
(7 of 16)
Side Groups :
-CH3 methyl
-Cl chloro
-C2H5 ethyl
di- = 2 groups
tri- = 3 groups
2,2,3-tribromobutane (not 2,3,3-)
Note that we # the carbons from
whichever end results in the
smallest numbers.
While drawing the isomers of
pentane, C5H12, students draw this
structure, naming it 2-ethylpropane.
(a chain of 3 C’s with an ethyl group)
The longest chain is four C’s, and
should be named 2-methylbutane.
tetra- = 4 groups
H
H
Br
Br H
C
C
C
H
H
Br H
H
H
H
C
C
C
H
CH2 H
H
C
H
H
CH3
A similar error is to draw and name “1-methylsomething”.
H
H
H
C
C
H
H
O
5 bonds to C
H2C
H2
C H CH2
H3 C
C
23 • Organic Chemistry
Optical Isomers
Chiral Compounds
(8 of 16)
-Br bromo
-I iodo
-C3H7 propyl, etc.
CH3
Two more tips … double check that
each C has four and only four bonds.
Also, remember that N and O atoms
have lone pairs of e-‘s although they
are seldom drawn. (Impt. for steric #!)
CH3
Some molecules have the ability to
rotate polarized light.
These molecules can be recognized
by a C atom (the chiral carbon)
bonded to four different groups .
3-methylhexane
This carbon is bonded to H, methyl, ethyl, & propyl groups.
You can build two versions of this molecule that are
“nonsuperimposable mirror images of each other.” One will
rotate light clockwise, one counterclockwise.
In biology, these are called dextro- and levo- (D and L) forms.
ethene is also called ethylene
propene is also called propylene
H3C
23 • Organic Chemistry
Common Names You Should Know About
(9 of 16)
CH3
CH3
H3 C
2-methylbutane is also called isopentane.
“Iso-“ means the same… the same two
methyl groups come branch from C #2.
2-methylpentane is isohexane, etc.
H2
C
HC
H3 C
C
CH3
CH3
2,2-dimethylpropane is called neopentane.
These common names show up
occasionally in names… such as in
isopropyl alcohol.
H
C
H
C
HC
HC
23 • Organic Chemistry
Aromatic Compounds
Benzene and its Derivatives
(10 of 16)
C
H
CH
HC
CH
HC
Benzene, C6H6, is unique. It
can be drawn as shown, but
the actual structure involves
a circular pi bond (sp 2
orbitals & delocalized e-‘s).
CH
C
H
CH
two resonance structures
Benzene is also shown with a circle as the pi bond.
1
The carbon #’s can be used for substituted
6
2
benzene. Example: dichlorobenzene
1,2- is known as the ortho- position
5
3
1,3- is known as the meta- position
4
1,4- is known as the para- position
Paradichlorobenzene: the main ingredient in some moth balls.
General formula: R-O-H [R ≈ Rest of molecule]
C atoms are classified as primary (1),
C
secondary (2), or tertiary (3) by the
number of C atoms it is bonded to.
1
C
C
3
1C
C
A primary alcohol has the -OH group
2
bonded to a primary carbon, etc.
This is not a base because the -OH is covalent, not ionic.
Naming: group + “alcohol” (e.g. ethyl alcohol or ethanol)
Ethers
General formula: R-O-R’ [R’ can = R, but not H]
Naming: two groups + “ether”
H2
H2
C
C
diethyl ether was the 1st effective
H3 C
O
CH3
surgical and dental anesthetic.
Alcohols
1
23 • Organic Chemistry
Functional Groups I
Alcohols and Ethers
(11 of 16)
Aldehydes
Ketones
O
O
General formula:
C
C
R
23 • Organic Chemistry
Functional Groups II
Aldehydes and Ketones
(12 of 16)
Naming:
H
R
R'
names end in “al”
names end in
or “aldehyde”
“one”
methanaldehyde
propanone
(formaldehyde)
(acetone)
Aldehydes and ketones both have a C=O group (carbonyl
group). Aldehydes have it on an end carbon. Ketones have
it on a middle carbon. Reactions: Primary alcohols can be
oxidized into aldehydes. Secondary alcohols into ketones.
Carboxylic Acids
Esters
O
General formula:
O
C
R
23 • Organic Chemistry
Functional Groups III
Carboxylic Acids and Esters
(13 of 16)
C
OH
R
OR
Naming:
names end - “oic names end - “ate”
acid”
ethyl acetate
ethanoic acid
(acetic acid +
(acetic acid)
ethyl alcohol)
Reactions: Acids can be made by oxidizing aldehydes.
Esters are formed (“esterification”) from a carboxylic acid &
an alcohol. Water is removed (a “condensation” reaction).
Esters often have pleasant, agreeable odors (e.g. banana.)
Amines
H
General formula:
Amides
H
O
N
C
R
23 • Organic Chemistry
Functional Groups IV
Amines & Amides
(14 of 16)
23 • Organic Chemistry
Polymers I
Monomers & Addition Polymerization
(15 of 16)
R
NH2
Naming:
names contain
names end in
“amino” or end in
“amide”
“amine”
acetamide
aminomethane
(methylamine)
The N may have 1 or 2 or all 3 H atoms replaced with groups.
The lone pair on the N atom makes these molecules basic.
Your body needs certain amines “vital amines” ≈ “vitamins.”
Monomer = one part
Polymer = many parts
H
H
One kind of polymer is made up of
C C
monomers that contain a double bond.
The double bond can break and we can
H
H
“ethylene” ADD to it… “Addition polymerization.”
*
H
H
C
C
H
H
H
+
*
H
C
H
→
C
*
H
H
H H H
C
C
H
H H H
C
C
*
Different monomers form different polymers. This polymer
would be called polyethylene. Replace on H on the monomer
with Cl and you can make polyvinyl chloride, “PVC.”
Another polymerization involves condensation reactions.
O
O
C
23 • Organic Chemistry
Polymers II
Copolymers & Condensation Polymerization
(16 of 16)
HO
R
H H
HO
C
C
OH
H
OH
C
R
H
a di-alcohol (a glycol)
a di-acid
Esters form from an acid and an alcohol. Using a di-acid and
a di-alcohol, you can make a continuous chain by removing
water molecules. The resulting polymer is called a polyester.
Soda bottles are made from a polyester, polyethylene
terephthalate ester (PETE).
Nylon (a polyamide) can be made from a di-amine & a di-acid.