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4
General, Organic, and
Biochemistry, 7e
Bettelheim,
Brown, and March
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4-1
4 Chapter 4
Chemical Reactions
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4-2
4 Chemical Reactions
• In a chemical reaction, one or more reactants is
converted to one or more products
Reactant(s)
Product(s)
• In this chapter we discuss three aspects of
chemical reactions
(a) mass relationships (stoichiometry)
(b) types of reactions
(c) heat gain and loss accompanying reactions
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4-3
4 Formula Weight
• Formula weight: the sum of the atomic weights in
atomic mass units (amu) of all atoms in a
compound’s formula
Ionic Comp ou nds
Sod ium chlorid e (N aCl)
23.0 amu N a + 35.5 amu Cl = 58.5 amu
Nickel(II) ch loride h yd rate 58.7 amu N i + 2(35.5 amu Cl) +
12(1.0) amu H) + 6(12.0 amu O) = 237.7 amu
(N iCl 2• 6H 2O)
Molecu lar Comp ou nds
Water (H 2O)
Aspirin (C9H 8 O 4)
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2(1.0 amu H) + 16.0 amu O = 18.0 amu
9(12.0 amu C) + 8(1.0 amu H) +
4(16.0 amu O) = 180.0 = amu
4-4
4 Formula Weight
• formula weight can be used for both ionic and
molecular compounds; it tells nothing about
whether a compound is ionic or molecular
• molecular weight should be used only for
molecular compounds
• in this text, we use formula weight for ionic
compounds and molecular weight for molecular
compounds
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4-5
4 The Mole
• Mole (mol)
• a mole of the amount of substance that contains as
many atoms, molecules, or ions as are in exactly 12 g
of carbon-12
• a mole, whether it is a mole of iron atoms, a mole of
methane molecules, or a mole of sodium ions, always
contains the same number of formula units
• the number of formula units in a mole is known as
Avogadro’s number
• Avogadro’s number has been measured experimentally
• its value is 6.02214199 x 1023 formula units per mole
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4-6
4 Molar Mass
• Molar mass: the formula weight of a substance
expressed in grams
• Glucose, C6H12O6
• molecular weight: 180 amu
• molar mass: 180 g/mol
• one mole of glucose has a mass of 180 g
• Urea, (NH2)2CO
• molecular weight 60.0 amu
• molar mass: 60.0 g/mol
• one mole of urea has a mass of 60.0 g
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4-7
4 Molar Mass
• We can use molar mass to convert from grams to
moles, and from moles to grams
You are given one of these
and asked to find the other
Moles of A
Grams of A
Use molar mass (g/mol)
as the conversion factor
• calculate the number of moles of water in 36.0 g water
36.0 g H 2O x
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1 mol H2 O
18.0 g H 2O
= 2.00 mol H2 O
4-8
4 Grams to Moles
• Calculate the number of moles of sodium ions,
Na+, in 5.63 g of sodium sulfate, Na2SO4
• first we find the how many moles of sodium sulfate
• the formula weight of Na2SO4 is
2(23.0) + 32.1 + 4(16.0) = 142.1 amu
• therefore, 1 mol of Na2SO4 = 142.1 g Na2SO4
5.63 g Na2 SO4 x
1 mol Na2SO4
= 0.0396 mol Na2SO4
142.1 g Na2 SO4
• the formula Na2SO4 tells us there are two moles of Na+
ions per mole of Na2SO4
+
0.0396 mol Na2 SO 4
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2 mol Na
x
1 mol Na2SO 4
= 0.0792 mol Na+
4-9
4 Grams to Molecules
• A tablet of aspirin, C9H8O4, contains 0.360 g of
aspirin. How many aspirin molecules is this?
• first we find how many mol of aspirin are in 0.360 g
0.360 g aspirin x
1 mol aspirin
180.0 g aspirin
= 0.00200 mol aspirin
• each mole of aspirin contains 6.02 x 1023 molecules
• the number of molecules of aspirin in the tablet is
0.00200 mole x 6.02 x 1023 molecules = 1.20 x 1021 molecules
mole
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4-10
4 Chemical Equations
• The following chemical equation tells us that
propane gas and oxygen gas react to form
carbon dioxide gas and water vapor
C3 H8 ( g) + O2 (g)
CO2 (g) + H2 O( g)
Propane
Carbon
dioxid e
Oxygen
Water
• But while it tells us what the reactants and
products are and the physical state of each, it is
incomplete because it is not balanced
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4-11
4 Balancing Equations
• To balance a chemical equation
• begin with atoms that appear only in one compound on
the left and one on the right; in this case, begin with
carbon (C) which occurs in C3H8 and CO2
C3 H8 (g) + O2 (g)
3CO2( g) + H2 O(g)
• now balance hydrogens, which occur in C3H8 and H2O
C3 H8 (g) + O2 (g)
3CO2( g) + 4H2 O(g)
• if an atom occurs as a free element, as for example Mg
or O2, balance this element last; in this case O2
C3 H8 (g) + 5O2 ( g)
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3CO2 (g) + 4H2 O( g)
4-12
4 Balancing Equations
• Practice problems: balance these equations
Ca( OH) 2 ( s) + HCl( g)
Calcium
hydroxide
CO2 ( g) + H2 O(l)
photosynthesis
C4 H1 0 ( g) + O2 (g)
Butane
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CaCl2 (s) + H2 O( l)
Calcium
chlorid e
C6 H1 2 O6 (aq) + O2 (g)
Glucose
CO2 (g) + H2 O(g)
4-13
4 Balancing Equations
• Solutions to practice problems
Ca( OH) 2 ( s) + 2 HCl(g)
Calcium
hydroxide
6CO2 (g) + 6H2 O(l)
CaCl2 (s) + H2 O( l)
Calcium
chlorid e
photosynthesis
C4 H1 0 ( l) + 13 O2 (g)
2
Bu tane
C6 H1 2 O6 (aq) + 6O2 (g)
Glucose
4 CO2 (g) + 5 H2 O( g)
• it is common practice to use only whole numbers;
therefore, multiply all coefficients by 2, which gives
2C4 H1 0 ( l) + 13O2 (g)
Butane
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8CO2 ( g) + 10H2 O(g)
4-14
4 Stoichiometry
• Stoichiometry: the study of mass relationships in
chemical reactions
• following is an overview of the the types of calculations
we study
You are given one of these
Grams of A
Moles of A
From grams to moles,
use molar mass (g/mol)
as a conversion factor
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And asked to find one of these
Moles of B
Grams of B
From moles to moles, From moles to grams,
use the coefficients in use molar mass (g/mol)
the balanced equation as a conversion factor
as a conversion factor
4-15
4 Stoichiometry
• Problem: how many grams of nitrogen, N2, are
required to produce 7.50 g of ammonia, NH3
N2 (g) + 3H2 (g)
2NH3 ( g)
• first find how many moles of NH3 are in 7.50 g of NH3
7.50 g NH 3 x
1 mol NH 3
17.0 g NH 3
= mol NH 3
• next find how many moles of N2 are required to
produce this many moles of NH3
7.50 g NH 3 x
1 mol NH 3
17.0 g NH 3
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x
1 mol N2
2 mol NH 3
= mol N 2
4-16
4 Stoichiometry
• Practice problem (cont’d)
• finally convert moles of N2 to grams of N2 and now do
the math
7.50 g NH 3 x
1 mol NH 3
17.0 g NH 3
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x
1 mol N2
2 mol NH 3
x
28.0 g N 2
1 mol N2
= 6.18 g N 2
4-17
4 Stoichiometry
• Practice problems:
• what mass of aluminum oxide is required to prepare 27
g of aluminum?
Al2 O3 ( s) electrolysis
Al( s) + O2 ( g)
• how many grams each of CO2 and NH3 are produced
from 0.83 mol of urea?
( NH2 ) 2 CO(aq) + H2 O
Urea
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urease
2NH3 (aq) + CO2 (g)
4-18
4 Limiting Reagent
• Limiting reagent: the reagent that is used up first
in a chemical reaction
• consider this reaction of N2 and O2
N2 (g) + O2 (g)
before reaction (moles) 5.0
after reaction (moles) 4.0
1.0
0
2 NO( g)
0
2.0
• in this experiment, there is only enough O2 to react with
1.0 mole of N2
• O2 is used up first; it the limiting reagent
• 4.0 moles of N2 remain unreacted
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4-19
4 Limiting Reagent
• Practice Problem
• suppose 12 g of carbon is mixed with 64 g of oxygen
and the following reaction takes place
C(s) + O2 ( g)
CO2 ( g)
• complete the following table. Which is the limiting
reagent?
C
before reaction (g)
12 g
+
O2
CO2
64 g
0
before reaction (mol)
after reaction (mol)
after reaction (g)
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4-20
4 Percent Yield
• Actual yield: the mass of product formed in a
chemical reaction
• Theoretical yield: the mass of product that should
be formed according to the stoichiometry of the
balanced chemical equation
• Percent yield: actual yield divided by theoretical
yield times 100
Actual yield
Percent yield =
x 100
Theoretical yield
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4-21
4 Percent Yield
• Practice problem:
• suppose we react 32.0 g of methanol with excess
carbon monoxide and get 58.7 g of acetic acid
• complete this table
CH3 OH + CO
before reaction (g) 32.0
before reaction (mol)
th eoretical yield (mol)
th eoretical yield (g)
actual yield (g)
percent yield (%)
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excess
CH3 COOH
0
58.7
4-22
4 Reactions Between Ions
• Ionic compounds, also called salts, consist of
both positive and negative ions
• When an ionic compound dissolves in water, it
dissociates to aqueous ions
NaCl(s)
H2 O
Na+ (aq) + Cl- (aq)
• What happens when we mix aqueous solutions of
two different ionic compounds
• if two of the ions combine to form a water-insoluble
compound, a precipitate will form
• otherwise no physical change will be observed
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4-23
4 Reactions Between Ions
• Example:
• suppose we prepare these two aqueous solutions
H2 O
Solution 1 AgNO3 (s)
Ag+ ( aq) + NO3 - (aq)
H2 O
+
Solution 2 NaCl(s)
Na ( aq) + Cl ( aq)
• if we then mix the two solutions, we have four ions
present; of these, Ag+ and Cl- react to form AgCl(s)
which precipitates
+
+
Ag (aq) + NO3 ( aq) + Na (aq) + Cl ( aq)
+
AgCl(s) + Na (aq) + NO3 ( aq)
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4-24
4 Reactions Between Ions
• we can simplify the equation for the formation of AgCl
by omitting all ions that do not participate in the
reaction
Net ionic equation:
Ag+ ( aq) + Cl-(aq)
AgCl( s)
• the simplified equation is called a net ionic equation; it
shows only the ions that react
• ions that do not participate in a reaction are called
spectator ions
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4-25
4 Reactions Between Ions
• In general, ions in solution react with each other
when one of the following can happen
• two of them form a compound that is insoluble in water
• two of them react to form a gas that escapes from the
reaction mixture as bubbles, as for example when we
mix aqueous solutions of sodium bicarbonate and
hydrochloric acid
+
HCO3 ( aq) + H3 O (aq)
Bicarbonate ion
CO2 ( g) + 2 H2 O( l)
Carbon d ioxide
• an acid neutralizes a base (Chapter 8)
• one of the ions can oxidize another (Section 4.7)
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4-26
4 Reactions Between Ions
• Following are some generalizations about which
ionic solids are soluble in water and which are
insoluble
• all compounds containing Na+, K+, and NH4+ are soluble
in water
• all nitrates (NO3-) and acetates (CH3COO-) are soluble in
water
• most chlorides (Cl-) and sulfates (SO42-) are soluble;
exceptions are AgCl, BaSO4, and PbSO4
• most carbonates (CO32-), phosphates (PO43-), sulfides
(S2-), and hydroxides (OH-) are insoluble in water;
exceptions are LiOH, NaOH, KOH, and NH4OH which
are soluble in water
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4-27
4 Oxidation-Reduction
• Oxidation: the loss of electrons
• Reduction: the gain of electrons
• Oxidation-reduction (redox) reaction: any
reaction in which electrons are transferred from
one species to another
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4-28
4 Oxidation-Reduction
• Example: if we put a piece of zinc metal in a
beaker containing a solution of copper(II) sulfate
• some of the zinc metal dissolves
• some of the copper ions deposit on the zinc metal
• the blue color of Cu2+ ions gradually disappears
• In this oxidation-reduction reaction
• zinc metal loses electrons to copper ions
Zn(s)
2+
Zn (aq) + 2 e
Zn is oxidized
• copper ions gain electrons from the zinc
2+
Cu ( aq) + 2 e
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-
Cu( s)
Cu
2+
is red uced
4-29
4 Oxidation-Reduction
• we summarize these oxidation-reduction relationships
in this way
electrons flow
from Zn to Cu2 +
2+
Zn(s)
+
Cu (aq)
loses electrons ; gains electrons ;
is red uced
is oxidized
gives electrons tak es electrons
to Cu 2+ ; is th e from Zn; is th e
oxidizin g agent
red ucing agent
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2+
Zn ( aq) + Cu( s)
4-30
4 Oxidation-Reduction
• Although the definitions of oxidation (loss of
electrons) and reduction (gain of electrons) are
easy to apply to many redox reactions, they are
not easy to apply to others
• for example, the combustion of methane
CH4 (g) + O2 ( g)
Methane
CO2 (g) + H2 O( g)
• An alternative definition of oxidation-reduction is
• oxidation: the gain of oxygen or loss of hydrogen
• reduction: the loss of oxygen or gain of hydrogen
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4-31
4 Oxidation-Reduction
• using these alternative definitions for the combustion
of methane
electrons are
trans ferred from
carb on to oxygen
CH4 (g)
gain s O and los es
H; is oxidized
+
O2 (g)
CO2 (g) + H2 O(g)
gain s H;
is reduced
is the reducin g is th e oxid izing
agent
agent
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4-32
4 Oxidation-Reduction
• Five important types of redox reactions
• combustion: burning in air. The products of complete
combustion of carbon compounds are CO2 and H2O.
• respiration: the process by which living organisms use
O2 to oxidize carbon-containing compounds to produce
CO2 and H2O. The importance of these reaction is not
the CO2 produced, but the energy released.
• rusting: the oxidation of iron to a mixture of iron oxides
• bleaching: the oxidation of colored compounds to
2Fe2 O3 ( s)
+ 3O
2 ( g)
products4Fe(s)
which are
colorless
• batteries: in most cases, the reaction taking place in a
battery is a redox-reaction
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4-33
4 Heat of Reaction
• In almost all chemical reactions, heat is either
given off or absorbed
• example: the combustion (oxidation) of carbon
liberates 94.0 kcal per mole of carbon oxidized
C( s) + O2 (g)
CO2 (g) + 94.0 kcal/mole C
• Heat of reaction: the heat given off or absorbed in
a chemical reaction
• exothermic reaction: one that gives off heat
• endothermic reaction: one that absorbs heat
• heat of combustion: the heat given off in a combustion
reaction; all combustion reactions are exothermic
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4-34
4 Chemical Reactions
End
Chapter 4
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4-35