Chapter 9

Charge-Transfer Reactions:Acids and Bases and

Oxidation-Reduction

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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

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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.