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Chapter 14Acid/Base Equilibrium

AP Chemistry Unit 10

Some Definitions

• Arrhenius– An acid is a substance that, when dissolved in

water, increases the concentration of hydrogen ions.

– A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions.

Some Definitions

• Brønsted-Lowry– An acid is a proton donor.– A base is a proton acceptor.

A Brønsted-Lowry acid…

…must have a removable (acidic) proton.

A Brønsted-Lowry base…

…must have a pair of nonbonding electrons.

If it can be either……it is amphiprotic.

HCO3-

HSO4-

H2O

What Happens When an Acid Dissolves in Water?

• Water acts as a Brønsted-Lowry base and abstracts a proton (H+) from the acid.

• As a result, the conjugate base of the acid and a hydronium ion are formed.

Conjugate Acids and Bases

• The term conjugate comes from the Latin word “conjugare,” meaning “to join together.”

• Reactions between acids and bases always yield their conjugate bases and acids.

Example 15.1 Identify the Brønsted–Lowry acids and bases and

their conjugates in:

(a) H2S + NH3 NH4+ + HS–

(b) OH– + H2PO4– H2O + HPO4

2–

Acid and Base Strength

• Strong acids are completely dissociated in water.– Their conjugate bases are

quite weak.

• Weak acids only dissociate partially in water.– Their conjugate bases are

weak bases.

Acid and Base Strength

• Substances with negligible acidity do not dissociate in water.– Their conjugate bases are

exceedingly strong.

Acid and Base Strength

• In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.

HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)

• H2O is a much stronger base than Cl-, so the equilibrium lies so far to the right that K is not measured (K>>1).

Acid and Base Strength

• In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.

• Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1).

CH3CO2H (aq) + H2O (l) H3O+ (aq) + CH3CO2- (aq)

Factors Affecting Acid Strength

• The more polar the H-X bond and/or the weaker the H-X bond, the more acidic the compound.

• So acidity increases from left to right across a row and from top to bottom down a group.

Factors Affecting Acid Strength

In oxyacids, in which an -OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

Factors Affecting Acid Strength

For a series of oxyacids, acidity increases with the number of oxygens.

Factors Affecting Acid Strength

Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.

Example 15.2Select the stronger acid in each pair:

(a) nitrous acid, HNO2, and nitric acid, HNO3

(b) Cl3CCOOH and BrCH2COOH

Strengths of Amines as Bases

• Aromatic amines are much weaker bases than aliphatic amines.

• This is due in part to the fact that the electrons in the benzene ring of an aromatic molecule are delocalized and can involve the nitrogen atom’s lone-pair electrons in the resonance hybrid.

• As a result, the lone-pair electrons are much less likely to accept a proton.

• Electron-withdrawing groups on the ring further diminish the basicity of aromatic amines relative to aniline.

Example 15.3Select the weaker base in each pair:

Lewis Acids

• Lewis acids are defined as electron-pair acceptors.

• Atoms with an empty valence orbital can be Lewis acids.

Lewis Bases

• Lewis bases are defined as electron-pair donors.• Anything that could be a Brønsted-Lowry base is

a Lewis base.• Lewis bases can interact with things other than

protons, however.

Autoionization of Water

• As we have seen, water is amphoteric.

• In pure water, a few molecules act as bases and a few act as acids.

• This is referred to as autoionization.

H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)

Ion-Product Constant

• The equilibrium expression for this process is

Kc = [H3O+] [OH-]

• This special equilibrium constant is referred to as the ion-product constant for water, Kw.

• At 25C, Kw = 1.0 10-14

pH

pH is defined as the negative base-10 logarithm of the concentration of hydronium ion.

pH = -log [H3O+]

pH

• In pure water,

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

• Since in pure water [H3O+] = [OH-],

[H3O+] = 1.0 10-14 = 1.0 10-7

pH

• Therefore, in pure water,pH = -log (1.0 10-7) = 7.00

• An acid has a higher [H3O+] than pure water, so its pH is <7.

• A base has a lower [H3O+] than pure water, so its pH is >7.

pH

These are the pH values for several common substances.

Other “p” Scales

• The “p” in pH tells us to take the negative base-10 logarithm of the quantity (in this case, hydronium ions).

• Some similar examples are– pOH: -log [OH-]

– pKw: -log Kw

Watch This!

Because

[H3O+] [OH-] = Kw = 1.0 10-14,

we know that

-log [H3O+] + -log [OH-] = -log Kw = 14.00

or, in other words,

pH + pOH = pKw = 14.00

Example 15.4

By the method suggested in Figure 15.5, a student determines the pH of milk of magnesia, a suspension of solid magnesium hydroxide in its saturated aqueous solution, and obtains a value of 10.52. What is the molarity of Mg(OH)2 in its saturated aqueous solution? The suspended, undissolved Mg(OH)2(s) does not affect the measurement.

Example 15.5

A Conceptual Example

Is the solution 1.0 x 10–8 M HCl acidic, basic, or neutral?

How Do We Measure pH?

• For less accurate measurements, one can use– Litmus paper

• “Red” paper turns blue above ~pH = 8

• “Blue” paper turns red below ~pH = 5

– Or an indicator.

How Do We Measure pH?

For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

Strong Acids

• You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.

• These are, by definition, strong electrolytes and exist totally as ions in aqueous solution.

• For the monoprotic strong acids,

[H3O+] = [acid].

Strong Bases

• Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).

• Again, these substances dissociate completely in aqueous solution.

Dissociation Constants

• For a generalized acid dissociation,

the equilibrium expression would be

• This equilibrium constant is called the acid-dissociation constant, Ka.

[H3O+] [A-][HA]

Kc =

HA (aq) + H2O (l) A- (aq) + H3O+ (aq)

Dissociation ConstantsThe greater the value of Ka, the stronger is the acid.

Weak Bases

Bases react with water to produce hydroxide ion.

Weak Bases

The equilibrium constant expression for this reaction is

[HB] [OH-][B-]

Kb =

where Kb is the base-dissociation constant.

Weak Bases

Kb can be used to find [OH-] and, through it, pH.

Ka and Kb

Ka and Kb are related in this way:

Ka Kb = Kw

Therefore, if you know one of them, you can calculate the other.

Equilibrium in Solutions of Weak Acids and Weak Bases

These calculations are similar to the equilibrium calculations performed in Unit 10.

1. An equation is written for the reversible reaction.

2. Data are organized, often in an ICE format.

3. Changes that occur in establishing equilibrium are assessed.

4. Simplifying assumptions are examined (the “5% rule”).

5. Equilibrium concentrations, equilibrium constant, etc. are calculated.

Example 15.6Ordinary vinegar is approximately 1 M CH3COOH and as shown in Figure 15.6, it has a pH of about 2.4. Calculate the expected pH of 1.00 M CH3COOH(aq), and show that the calculated and measured pH values are in good agreement.

Example 15.7What is the pH of 0.00200 M ClCH2COOH(aq)?

Example 15.8What is the pH of 0.500 M NH3(aq)?

Example 15.9The pH of a 0.164 M aqueous solution of dimethylamine is 11.98. What are the values of Kb and pKb? The ionization equation is

(CH3)2NH + H2O (CH3)2NH2+ + OH– Kb = ?

Dimethylamine Dimethylammonium ion

Example 15.10 A Conceptual ExampleWithout doing detailed calculations, indicate which solution has the greater [H3O+], 0.030 M HCl or 0.050 M CH3COOH.

Calculating Percent Ionization

• Percent Ionization = 100[H3O+]eq

[HA]initial

Polyprotic Acids

• A monoprotic acid has one ionizable H atom per molecule.

• A polyprotic acid has more than one ionizable H atom per molecule.– Sulfuric acid, H2SO4 Diprotic

– Carbonic acid, H2CO3 Diprotic

– Phosphoric acid, H3PO4 Triprotic

• The protons of a polyprotic acid dissociate in steps, each step having a value of Ka.

• Values of Ka decrease successively for a given polyprotic acid. Ka1 > Ka2 > Ka3 , etc.

Polyprotic Acids

If the difference between the Ka for the first dissociation and subsequent Ka values is 103 or more, the pH generally depends only on the first dissociation.

Example 15.11

Calculate the following concentrations in an aqueous solution that is 5.0 M H3PO4:

(a) [H3O+] (b) [H2PO4–] (c) [HPO4

2–] (d) [PO43–]

Example 15.12

What is the approximate pH of 0.71 M H2SO4?

Ions as Acids and Bases

• HCl is a strong acid, therefore Cl– is so weakly basic in water that a solution of chloride ions (such as NaCl) is virtually neutral.

• Acetic acid, CH3COOH, is a weak acid, so acetate ion, CH3COO–, is significantly basic in water.

• A solution of sodium acetate (which dissociates completely into sodium and acetate ions in water) is therefore slightly basic:

CH3COO– + H2O CH3COOH + OH–

Carbonate Ion as a

BaseA carbonate ion accepts a proton

from water, leaving behind an OH– and making the solution basic.

Reactions of Cations with Water

• Cations with acidic protons (like NH4

+) will lower the pH of a solution.

• Most metal cations that are hydrated in solution also lower the pH of the solution.

Reactions of Cations with Water

• Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water.

• This makes the O-H bond more polar and the water more acidic.

• Greater charge and smaller size make a cation more acidic.

Ions as Acids and Bases (cont’d)

• Salts of strong acids and strong bases form neutral solutions: NaCl, KNO3

• Salts of weak acids and strong bases form basic solutions: KNO2, NaClO

• Salts of strong acids and weak bases form acidic solutions: NH4NO3

• Salts of weak acids and weak bases form solutions that may be acidic, neutral, or basic; it depends on the relative strengths of the cations and the anions: NH4NO2, CH3COONH4.

Example 15.13 A Conceptual Example(a) Is NH4I(aq) acidic, basic,or neutral? (b) What conclusion can you draw from Figure 15.8d about the equilibrium constants for the hydrolysis reactions in CH3COONH4(aq)?

Ions as Acids and Bases (cont’d)

• In order to make quantitative predictions of pH of a salt solution, we need an equilibrium constant for the hydrolysis reaction.

• The relationship between Ka and Kb of a conjugate acid–base pair is:

Ka x Kb = Kw

• If instead we have values of pKa or pKb:

pKa + pKb = pKw = 14.00 (at 25 °C)

Example 15.14Calculate the pH of a solution that is 0.25 M CH3COONa(aq).

Example 15.15What is the molarity of an NH4NO3(aq) solution that has

a pH = 4.80?