Chapter 16Acid-Base Equilibria

AP CHEMISTRY

Acids ◦ Sour, can corrode metals, cause certain dyes to

change colors Bases

◦ Bitter taste, feel slippery, usually used in cleaning products

Arrhenius Acid◦ substances when dissolved in water, increase the

concentration of H+ ions Arrhenius Base

◦ Substances, when dissolved in water, increase the concentration of OH- ions

16.1 Acids and Bases

Arrhenius definition is restricted to aqueous solutions

Brønsted-Lowery Acid◦ Substance (molecule or ion) that can donate a

proton to another substance Brønsted-Lowery Base

◦ Substance that can accept a proton

16.2 Brønsted-Lowry Acids and Bases

In any acid-base equilibrium, both the forward and reverse reactions involve protong transfers.

An acid and base in an equation that differ only in the presence or absence of a proton are called conjugate acid-base pairs.◦ Every acid has a conjugate base and every base

has a conjugate acid

Conjugate Acid-Base Pairs

What is the conjugate base of each of these acids?◦ HClO4, PH4

+, HNO2

◦ ClO4-, PH3, NO2

-

What is the conjugate acid of each of these bases?◦ CN-, H2O, HCO3

-

◦ HCN, H3O+, H2CO3

Identify the conjugate base pairs.NH3(aq) + H2O(l) ⇆ NH4

+(aq) + OH-

(aq)

Practice

Some acids/bases are better proton donors/acceptors than others.

The stronger the acid, the weaker its conjugate base (and the stronger the base, the weaker the conjugate base.

Strengths of acids and bases◦ Strong Acids completely transfer their protons to water. (Their

conjugate base have negligible tendency to be protonated.)◦ Weak Acids only partly dissociate and therefore exist as a

mixture of acid molecules and their ions (Conjugate base is a weak base as well.)

◦ Negligible Acidity contain hydrogen but do not demonstrate any acidic behavior in water. (Conjugate bases are strong bases, reacting completely in water)

Relative Strengths of Acids and Bases

Relative Strengths of Conjugate Acid-Base pairs

Think of proton-transfer reactions as being governed by the relative abilities of two bases to abstract protons

HC2H3O2(aq) + H2O(l) ⇆ H3O+(aq) + C2H3O2

-(aq)

C2H3O2- is a stronger base than H2O and therefore

abstracts the proton from H3O+.

In every acid-base reaction the position of the equilibrium favors transfer of the proton to the stronger base.

Acid-Base Equilibrium

One of the most important chemical properties of water is its ability to act as either a Brønsted-Lowery acid or base.

This is called the autoionization of water◦ Reaction is rapid and no individual molecules

remains ionized for long (only 2 in every 109 molecules is ionized at any moment)

16.3 Autoionization of Water

Because water ionizes, we can write its equilibrium constant. Keq = [H3O

+] [OH-]◦ Kw = 1.0 x 10-14 (at 25⁰C)◦ Ion-product constant◦ Commit this to memory!!

Important because this equilibrium applies to any dilute aqueous solution and can be used to calculate the [H+] or [OH-]◦ The product of [H+] and [OH-] = 1.0 x 10-14

◦ [H+] = [OH-] is said to be neutral.◦ [H+] exceeds [OH-] is acidic◦ [OH-] exceeds [H+] is basic

[H3O+] [OH-] = 1.0 x 10-14

Ion Product of Water

Calculate the [H+] in a solution in which [OH-] is 0.010 M. Is this solution acidic or basic?◦ [H3O

+] [OH-] = 1.0 x 10-14

◦ [H+] = = = 1.0 x 10-12 M◦ Because [H+] <[OH-], this is a basic solution.

Practice

The molar [H+] is very small in aqueous solutions so we express [H+] in terms of pH, which is the negative logarithm in base 10 of [H+].

pH = -log[H+]

16.4 The pH Scale

What is the pH for a solution that has [H+] = 5.6 x 10-6 M?

pH = -log(5.6 x 10-6) = 5.25

A sample of apple juice has a pH of 3.76. Calculate the [H+].

pH = -log[H+] = 3.76 [H+] = 10-3.76 = 1.7 x 10-4 M

Practice

pOH = -log[OH-] pH + pOH = 14 (at 25˚C)

Other “p” Scales

7 Common Strong Acids◦ 6 Monoprotic: HCl, HBr, HI, HNO3, HClO3, and HClO4

◦ 1 Diprotic: H2SO4

Strong acids completely dissociate into their ions◦ HNO3(aq) + H2O(l) H3O+

(aq) + NO3-(aq)

[H+] = the original concentration of acid.◦ If 0.20 M HNO3 was used, [H+] = 0.20 M◦ pH = -log(0.20) = 0.7◦ (diprotic acid is a little more complex)

16.5 Strong Acids

Common Soluble Strong Bases◦ Alkali metal hydroxides and heavier alkaline earth metal

hydroxides LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2

Strong Bases also completely dissociate so calculating pH is also straightforward.

◦ What is pH of 0.028 M solution of NaOH? pOH = -log(0.028) = 1.55 pH = 14.00 – 1.55 = 12.45

OR [H+] = 1.0 x 10-14 = 2.35 x 10-13

0.028 pH = -log(3.57 x 10-13) = 12.45

Strong Bases

Weak acids only partially ionize in aqueous solutions

We can use the equilibrium constant for the ionization reaction to express the extent to which an acid ionizes.

Ionization reaction can be written in two ways:◦ HA(aq) + H2O(l) ⇌ H30+

(aq) + A-(aq)

◦ HA(aq) ⇌ H+(aq) + A-

(aq)

16.6 Weak Acids

HA(aq) + H2O(l) ⇌ H30+(aq) + A-

(aq)

HA(aq) ⇌ H+(aq) + A-

(aq)

Ka = [H30+][A-] or [H+][A-]

[HA] [HA]

Ka is acid-dissociation constant The larger the value of Ka, the stronger the

acid

Acid-dissociation Constant

A 0.10 M solution of formic acid (HCHO2) has a pH of 2.38 at 25⁰C. Calculate the Ka for formic acid at this temperature.◦ HCHO2(aq) ⇌ H+

(aq) + CHO2-(aq)

◦ Ka = [H+][CHO2-]

[HCHO2]◦ pH = -log[H+] = 2.38◦ [H+] = 10-2.38 = 4.2 x 10-3 M

◦ 0.10 -4.2 x 10-3 M ≈ 0.10

◦ Ka = [H+][CHO2-] = (4.2 x 10-3 M)(4.2 x 10-3 M) = 1.8 x 10-4

[HCHO2] 0.10

Calculating Ka from pH

HCHO2 H+ CHO2-

I 0.10 M 0 0

C -4.2 x 10-3 M +4.2 x 10-3 M +4.2 x 10-3 M

E 0.10 -4.2 x 10-3 M 4.2 x 10-3 M 4.2 x 10-3 M

Know the value of Ka and the initial concentration of the weak acid, we can calculate the [H+] in a solution of a weak acid.

Calculate the pH of a 0.30 M solution of acetic acid.1. Write the ionization for acetic acid:

HC2H3O2 ⇆ H+(aq) + C2H3O2

-(aq)

2. Write the acid-dissociation constant expression and calculate the value: Ka = [H+][C2H3O2-] = 1.8 x 10-5 (from table 16.2 page 628)

[HC2H3O2]

3. Use an ICE box to express equilibrium concentrations

Using Ka to Calculate pH

4. Substitute equilibrium concentrations into acid dissociation constant expression (0.30 – x ≈ 0.30)

Ka = [H+][C2H3O2-] = (x)(x) = 1.8 x 10-5

[HC2H3O2] 0.30

5. Solve for x x2 = (0.30)(1.8 x 10-5) = 5.4 x 10-6

x = 2.3 x 10-3

6. Solve for [H+] [H+] = x = 2.3 x 10-3 M

7. Solve for pH pH = -log(2.3 x 10-3) = 2.64

HC2H3O2 H+ C2H3O2-

I 0.30 M 0 0

C -x M +x M +x M

E (0.30 –x) M x M x M

Percent Ionization = [H+]equilibrium x 100

[Acid]initial

Example Percent Ionization of HC2H3O2

= 0.0023 M x 100 = 0.77% 0.30 M

Percent Ionization

In polyprotic acids, H atoms ionize in successive steps:◦ H2SO4(aq) ⇌ H+ + HSO3

-(aq) Ka = 1.71 x 10-2

◦ HSO3-(aq) ⇌ H+

(aq) + SO32-

(aq) Ka = 6.4 x 10-8

It is always easier to remove the first proton from a polyprotic acid than the second.

As long as successive Ka values differ by a factor of 103 or more, it is possible to obtain a satisfactory estimate of the pH of polyprotic acid solutions by considering Ka1

Polyprotic Acids

The most commonly encountered weak base is NH3

NH3(aq) + H20 ⇌ NH4+

(aq) + OH-(aq)

Base-dissociation constant The constant Kb always refers to the

equilibrium in which a base reacts with H2O to form the corresponding conjugate acid and OH-

Table 16.4 on page 636 provides formulas and Kb values for several weak bases in water.

16.7 Weak Bases

Weak Bases fall into 2 general categories1. Neutral substances that have an atom with a

nonbonding pair of electrons that can serve as a proton acceptor.

2. Anions of weak acids ClO-

(aq) + H2O(l) ⇌ HClO(aq) + OH- Kb = 3.33 x 10-7

Types of Weak Bases

The product of the acid-dissociation constant for and acid and the base-dissociation constant for its conjugate base is the ion-product for water◦ (Ka)(Kb) = Kw

16.8 Relationship Between Ka and Kb

Acid Ka Base Kb

HNO3 (strong acid) NO3- Negligible

HF 6.8 x 10-4 F- 1.5 x 10-11

HC2H3O2 1.8 x 10-5 C2H3O2- 5.6 x 10-10

H2CO3 4.3 x 10-7 HCO3- 2.3 x 10-8

NH4+ 5.6 x 10-10 NH3 1.8 x 10-5

HCO3- 5.6 x 10-11 CO3

2- 1.8 x 10-4

OH- Negligible O2- (Strong base)

Sometimes acid- and base- dissociation constants are expressed as pKa or pKb

◦ pKa = -log Ka

◦ pKb = -log Kb

◦ pKa + pKb = pKw = 14.00 at 25˚C

Salts completely dissociate in water Many ions are able to react with water to

generate H+ or OH- ions. ◦ This type of reaction is called hydrolysis◦ The pH of an aqueous salt solution can be

predicted by considering the ions of which the salt is composed.

16.9 Acid-Base Properties of Salt Solutions

In general, an anion, X-, in solution can be considered the conjugate base of an acid.◦ Whether or not an anion will react with water to produce

OH- depends on the strength of the acid the anion would create.

◦ To identify the acid, add a proton to the anion◦ X- + a proton = HX

◦ If the acid is a strong acid, the anion will have a negligible tendency to abstract a proton from water.

◦ If the acid is a weak acid, X- will react to a small extent. He pH would be higher (more basic)

◦ If the ion has an ionizable proton like HSO3-, it is amphoteric.

The behavior is determined by the magnitude of Ka and Kb

An Anion’s Ability to React with Water

Most metal ions can react with water to decrease the pH of an aqueous solution◦ Ions of alkali metals and of the heavier alkaline

earth metals do not react with water and therefore do not affect pH. (same cations that make strong bases)

A Cation’s Ability to React with Water

1. An anion that is conjugate base of a strong acid will not affect the pH of solution

2. An anion that is conjugate base of a weak acid will cause an increase in pH

3. A cation that is the conjugate acid of a weak base will cause a decrease in pH

4. With the exception of group 1A and heavy 2B, metal ions will decrease pH

5. When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the ion with the largest ionization constant will have the greatest affect on pH

Combined Effect of Cation and Anion in Solution

Predict whether the salt Na2HPO4 will form an acidic or basic solution on dissolving in water.◦ Na2HPO4 ⇆ 2Na+ + HPO4

2-

◦ HPO42- can act like an acid or base

1. HPO42- ⇆ H+ + PO4

3- (acid) 2. HPO4

2- + H2O ⇆ H2PO4- + OH- (base)

The reaction with the larger ionization constant will determine whether it is acidic or basic

Ka1 = 4.2 x 10-13 (Table 16.3) (Ka2)(Kb2) = Kw

(6.2 x 10-8)(Kb2) = 1.0 x 10-14

Kb = 1.6 x 10-7

Since Kb > Ka, reaction is basic

Example

Factors that Affect Acid Strength1. A molecule containing H will transfer a proton

only if the H-X bond is polarized in the following way

In ionic hydrides (NaH) H has a negative charge and acts as a proton acceptor

2. Strong bonds (Table 8.4) are more difficult to dissociate than weak bonds

3. The greater the stability of the acid’s conjugate base, the stronger the base

16.10 Acid-Base Behavior and Chemical Structure

H-X strength is the most important factor in determining acid strength in binary acids

The H-X strength decreases as element X increases in size

Binary Acids

Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom are called oxyacids.

Consider -Y-O-H

◦ When Y is a metal, they are sources of OH- ions and behave as bases

◦ When Y is a nonmetal, the OH bond is more polar so it will donate a H+ and behave as an acid

Oxyacids

For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom.◦ Example: Strength of H-O-Y, increases as

electronegativity of Y increases For oxyacids that have the same central atom

Y, acid strength increases as the number of oxygen atoms attached to Y increases.◦ Example: the strength of the oxyacids of chlorine

increases from HClO to HClO2 to HClO3 to HClO4

Oxyacids

Another group of acids are those with a carboxyl group, often written COOH ◦ Example: acetic acid HC2H3O2 can also be written

CH3COOH

Other examples: Formic Acid Benzoic Acid

Carboxylic Acids

The acid strength of carboxylic acids increase as the number of electronegative atoms in the acid increase.◦ Example: trifluoroacetic acid (CF3COOH) is

stronger than acetic acid (CH3COOH)

Carboxylic Acids

For a substance to be a proton acceptor (base), it must have an unshared pair of electrons for binding a proton.

G.N. Lewis noticed this and proposed definitions for acid and base◦ A Lewis acid is an electron-pair acceptor◦ A Lewis base is an electron-pair donor

◦ **Not using this definition in AP Chem.**

16.11 Lewis Acids and Bases