3.6: ACIDS AND BASES… Equilibrium Constants…Ka and Kb

HOW DO WE MEASURE PH? For less accurate

measurements, one can useLitmus 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.

PRACTICE Calculate the pH of a 0.0750M solution of HI

and a 0.0750M solution of Ca(OH)2(aq).HI solution:

pH = 1.125Ca(OH)2(aq):

pH = 13.18

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Facts about acid and base concentrations:

For a strong monoprotic acid, [H3O+] in solution is equal to the original acid concentration. Likewise, for a strong monoprotic base, [OH-] will be equal to the original base concentration

For a weak acid, [H3O+] will be much less than the original acid concentration. [H3O+] will be smaller than it would have been if it were a

strong acid For a series of weak monoprotic acids (of the type

HA) of the same concentration, [H3O+] will increase (and the pH will decline) as the acids become stronger. Likewise, for a series of weak bases, [OH-] will increase (and the pH will increase) as the bases become stronger

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Relative strength of an acid or base can be

expressed quantitatively with an equilibrium constant Ionization Constant

Still a constant… still the same thing… we’re just talking about the ionization of acids and bases… so we give it a specific term… why? To make your life more difficult, that’s why

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Write the equilibrium constant for the

following reaction, where HA represents a generic acid:

But this is for an acid… so we call the equilibrium constant is called the acid-dissociation constant, Ka

[H3O+] [A-][HA]

Kc =

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

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: The greater the value of Ka, the stronger is

the acid.

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Calculating the Ka from the pH:

The pH of a 0.10 M solution of formic acid, HCOOH, at 25C is 2.38. Calculate Ka for formic acid at this temperature.

Use the pH to calculate the hydronium ion concentration

[H3O+] [COO-][HCOOH]

Ka =

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES:

pH = -log [H3O+] 2.38 = -log [H3O+]-2.38 = log [H3O+]

10-2.38 = 10log [H3O+] = [H3O+] 4.2 10-3 = [H3O+] = [HCOO-]

[4.2 10-3] [4.2 10-3][0.10]

Ka =

= 1.8 10-4

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Calculate the pH of a 0.100M solution of HF.

Ka = 6.8 x 10-4

Use an ICE table to calculate the [H+]. Then use the [H+] to calculate the pH. [H+] = 0.00791pH = 2.1

HF (aq) + H2O (l) F- (aq) + H3O+ (aq)

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Similarly, we can write equilibrium constants

for weak bases:

But this is for an base… so we call the equilibrium constant is called the base-dissociation constant, Kb

[HB] [OH-][B-]Kb =

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: Kb can be used to find [OH-] and, through it,

pH.

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES: What is the pH of a 0.15 M solution of NH3?

ICE TABLE TIME!!!

NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)

[NH4+] [OH-]

[NH3]Kb = = 1.8 10-5

EQUILIBRIUM CONSTANTS FOR ACIDS AND BASES:

(1.8 10-5) (0.15) = x2

2.7 10-6 = x2

1.6 10-3 = x2

[OH-] = 1.6 10-3 MpOH = -log (1.6 10-3)pOH = 2.80pH = 14.00 - 2.80pH = 11.20

(x)2

(0.15)1.8 10-5 =

POLYPROTIC ACIDS:…have more than one acidic proton

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.

AQUEOUS SOLUTIONS OF SALTS Anions that are conjugate bases of strong acids (things like Cl-) will have

effect on the solution pH. There are numerous basic anions (CO3

-2). All are conjugate bases of weak acids.

The acid-base behavior of anions of polyprotic acids depends on the extent of deprotonation. For example, a fully deprotonated anion (like CO3

-2) will be basic. A partially deprotonated anion (such as HCO3

-) is amphiprotic – its behavior will depend on the other species in the reaction.

Alkali metal and alkaline earth cations have no measurable effect on solution pH.

Basic cations are conjugate bases of acids cations [Al(H2O)6] +3. Acidic cations are limited to metal cations with +2 or +3 charges and to

ammonium ions (and their organic derivatives) All metal cations are hydrated in water. They form ions such as

[M(H2O)6] +3. However, only when M is a +2 or +3 ion, particularly a transition metal ion, does the ion act as an acid

© 2009, Prentice-Hall, Inc.

EFFECT OF CATIONS AND ANIONS

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

2. An anion that is the conjugate base of a weak acid will increase the pH.

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

© 2009, Prentice-Hall, Inc.

EFFECT OF CATIONS AND ANIONS

4. Cations of the strong Arrhenius bases will not affect the pH.

5. Other metal ions will cause a decrease in pH.

6. When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the Ka and Kb values.

EFFECT OF CATIONS AND ANIONS Cliff’s Summary:

Salts made of the anion of a strong acid and the cation of a strong base will be neutral salts, pH = 7. Example – Sodium Chloride

Salts made of the anion of a strong acid and the cation of a weak base will be acidic salts. Example – Ammonium Chloride

Salts made of the anion of a weak acid and the cation of a strong base will be basic salts. Example – Sodium Bicarbonate

KA AND KB

Ka and Kb are related in this way:Ka Kb = Kw

PKA

How are the values of Ka used to compare acid strength? pKa! pKa = -logKa

Like pH, as the value of pKa becomes smaller the acid strength increases

KA AND KB

Ka for lactic acid is 1.4 x 10-4-. What is the Kb for the conjugate base of this acid?

Kb =7.1 x 10-11