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Chapter 9
Charge-Transfer Reactions:Acids and Bases and
Oxidation-Reduction
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
9.1 Acids and Bases• Acids: Taste sour, dissolve some metals, cause
plant dye to change color• Bases: Taste bitter, are slippery, are corrosive.• Two theories that help us to understand the
chemistry of acids and bases.
1. Arrhenius Theory
2. Brønsted-Lowry Theory
Arrhenius Theory of Acids and Bases• Acid - a substance, when dissolved in water, dis-
sociates to produce hydrogen ions– Hydrogen ion: H+ also called “protons”
HCl is an acid:
HCl(aq) H+(aq) + Cl-(aq)• Base - a substance, when dissolved in water,
dissociates to produce hydroxide ions.
NaOH is a base
NaOH(aq) Na+(aq) + OH-(aq)
• Where does NH3 fit? When it dissolves in water it is basic but it does not have OH- ions in it.
Brønsted-Lowry Theory of Acids and Bases• Acid - proton donor• Base - proton acceptor
Notice that it is not defined using water.• When writing the reactions, both accepting and
donation are evident.
HCl(aq) + H2O(l) Cl-(aq) + H3O+(aq)
• Now let us look at NH3 and see why it is a base.
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
baseacid
base acid
Conjugate Acids and Bases• The acid base reaction can be written in the general
form:
HA + B A- + HB+
• Notice the reversible arrows.• The products are also an acid and base called the
conjugate acid and base.• Conjugate Acid - what the base becomes after it
accepts a proton. Conjugate Base - what the acid becomes after it donates its proton.
• Conjugate Acid-Base Pair - The acid and base on the opposite sides of the equation.
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acid base base acid
Strong and Weak Acids• The reversible arrow isn’t always written. Some
acids or bases essentially dissociate 100% and a one way arrow is used.
• Example: HCl + H2O Cl- + H3O+
• HCl is called a strong acid-an acid that dissociates 100%
• Weak acid - one which does not dissociate 100%.
9.2 Solutions of Acids and Bases
Strength of Acids and Bases• Acid and base strength - degree of dissociation
– Not a measure of concentration, different thing
• Strong acids and bases - reaction with water is virtually 100% (Strong electrolytes)
• Strong Acids:– HCl, HBr, HI Hydrochloric Acid, etc.
– HNO3 Nitric Acid
– H2SO4 Sulfuric Acid
• Strong Bases:– NaOH, KOH, Ba(OH)2 (all are metal hydroxides)
• Weak acids and bases - only a small percent dissociates. (Weak electrolytes)
• Weak acid example:– Acetic acid:
• Weak base example:– Ammonia:
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
The Dissociation of Water• Pure water is virtually 100% molecular.• Very small number of molecules dissociate
– Dissociation of acids and bases is often called ionization.
• Called autoionization. Very weak electrolyte.• H+ is called the hydrogen ion. In pure water at room
temperature:
[H+] = 1 x 10-7 M
[OH-] = 1 x 10-7 M
H2O(l) H+(aq) + OH-(aq)
• Therefore the equilibrium expression for:
• Remember, liquids are not included.• This constant is called the ion product for water and
has the symbol Kw
• Since [H+] = [OH-] = 1.0 x 10-7 M, what is the value for Kw? 1.0 x 10-14.
– Remember, it is without units.
]OH][[HK -eq
H2O(l) H+(aq) + OH-(aq)
The pH Scale• pH scale - a scale that indicates the acidity or
alkalinity of a solution.– Ranges from 0 (very acidic) to 14 (very basic)
• As we do the problems, keep in mind that since 1 x 10-14 = [H+][OH-], – if we know one concentration, can calculate the other,
– if add an acid, [H+] and [OH-] – if add a base, [OH-] and [H+]
• The pH of a solution is defined as:
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pH = -log[H+]
9.3 Reactions between Acids and Bases• Neutralization reaction - the reaction of an acid
with a base to produce a salt and water.
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
• An analytical technique to determine the
concentration of an acid or base is the titration.• Titration involves the addition of measured amount
of a standard solution (solution of known concentration) to neutralize the second, unknown solution.
• The equivalence point is when the moles H+ and
OH- are equal.
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Polyprotic Substances
• The previous examples have the acid and base at a 1:1 combining ratio.
– Not all acid-bases do this.
• Polyprotic substance - donates or accepts more than one proton per formula unit.
H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l)
• Other polyprotics: Nitric Acid, Sulfuric Acid, and Phosphoric Acid.
9.4 Acid-Base Buffers• Buffer solution - solution which resists large
changes in pH when either acids or bases are added.
• The Buffer Process• Buffers consist of either
– a weak acid and its salt or– a weak base and its salt
• Examples:– Acetic acid (CH3COOH) with sodium acetate
(CH3COONa).– An equilibrium is established in solution between the
acid and the salt anion.
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CH3COOH(aq) + H2O(l) CH3COO-(aq) + H+(aq)
Addition of Base (OH-) to a Buffer Solution.
• The OH- will react with the H+, removing it from the above equilibrium.
• Which way will the equilibrium shift? To the right.
Addition of Acid (H+) to a Buffer solution.
• The acid increases the concentration of H+.
• Which way will the equilibrium shift? To the left.
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H+(aq)
• Buffer Capacity - a measure of the ability of a solution to resist large changes in pH when a strong acid or strong base is added.
Preparation of a Buffer Solution
COOH]CH[
]COOCH][OH[K
3
-33
a
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H+(aq)
9.5 Oxidation-Reduction Reactions• Oxidation - defined by one of the following
– loss of electrons
– loss of hydrogen atoms
– gain of oxygen atoms
• Example: NaNa+ + e-
– Oxidation half of the reaction
• Reduction - defined by one of the following:– gain of electrons
– gain of hydrogen
– loss of oxygen
• Example: Cl + e- Cl-
– Reduction half of the reaction
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Na Na+ + e-
Cl + e- Cl-
Na + Cl Na+ + Cl-
Oxidizing Agent• Is reduced• Gains electrons• Causes
oxidation
Reducing Agent• Is oxidized• Loses electrons• Causes
reduction
9Applications of Oxidation and Reduction
• Corrosion - the deterioration of metals caused by an oxidation-reduction process.
– Example: rust (oxidation of iron)
4Fe(s) + 3O2(g) 2Fe2O3(s)
• Combustion of Fossil Fuels
– Example: natural gas (methane) furnaces.
CH4(g) + 2O2(g) CO2(g) + 2H2O(g)
• Bleaching - Most bleaching agents are oxidizing agents. The oxidation of the stains produces compounds that do not have color.