Acids and Bases - the Three Definitions

1. Measurement of pH - the pH meter

2. Bronsted-Lowry definition of acids and bases - an acid is a proton donor - a base is a proton acceptor - conjugate acid/conjugate base pairs

- relationship of Ka of a conjugate acid and Kb of a conjugate base

3. Lewis definition of acids and bases- a base is an

electron pair donor - an acid is an electron pair acceptor

- some examples of Lewis acids and Lewis bases

Ionization Constants (1)

Ionization Constants (2)

Exercise on Acid/Bases Strength

For each conjugate acid/base pair,

(1) Write the reactions defining Ka

and Kb.

(2)Find the values of pKa, pKb, and

Kb.

(3)Which species is the strongest conjugate acid, which is the strongest conjugate base?

nitrous acid: HNO2 / NO2

-

oxalic acid (2): HC2O4- /

C2O42-

arsenic acid (2): HAsO42- /

AsO43-

carbonic acid (1):H2CO3 / HCO3

-

Hydrolysis Reactions

Which salts undergo hydrolysis? Is the resulting solution acidic, basic, or neutral? Write the hydrolysis reaction (if any). Calculate the pH of a 0.10 M solution.

1. sodium acetate

(basic (pH=8.88), acetate (pKb=9.25)

hydrolyses to produce OH-)

2. ammonium chloride

(acidic (pH=5.12), ammonium (pKa=9.25)

hydrolyses to produce H3O+)

3. calcium chloride

(neutral, no hydrolysis)

4. sodium monohydrogen phosphate

(basic (pH=10.12), HPO42- (pKb2=6.79)

hydrolyses to produce OH-) (you need to consider two conjugate acid/base pairs..)

pH and % Dissociation of a Monoprotic Weak Acid

CH3COOH = CH3COO- + H+ Ka = 1.75 x 10-5

0.1 - x x x (We let x = [H+])

What is the pH of 0.10 M CH3COOH?

[CH3COO-] [H+]

[CH3COOH]

x2

0.1 - x

1.75 x 10-5

Ka = = =

Approximation Method: Since Ka <<1, assume x<<0.1x2 = 0.1 * 1.75 * 10-5 = 1.75 x 10-6

(x << 0.1)

and x = [H+] = 0.0013 M

Calculating % Dissociation and the pH

CH3COOH = CH3COO- + H+

% dissociation = 100 *

[CH3COO-]=

x

0.1

= 1.3%[CH3COOH]ini

t

[H+] = 1.3 x 10-3 M

pH = - log10[H+] = 2.88

Measurement of pH: the pH Meter

pH varies linearly with output voltage and can be

measured over the range pH 0 to

pH 14

Ka and Acid Strength

The stronger the acid, the larger the Ka and the smaller the pKa:

CH3COOH (aq) = CH3COO- (aq) + H+ (aq) Ka = 1.76 x 10-5

HCN (aq) = CN- (aq) + H+ (aq) Ka = 6.17 x 10-10

pKa = 4.75

HNO2 (aq) = NO2- (aq) + H+ (aq)

Ka = 4.6 x 10-4 pKa = 3.34

pKa = 9.21

stronger

weaker

Weak Acids

=

=

=

=

=

=

=

weaker

stronger

Kb and pKb

Arrhenius bases liberate OH- in solution.

Kb is the equilibrium constant for this reaction.NH4OH (aq) = NH4

+ (aq) + OH- (aq)

Kb =

[NH4+] [OH-]

[NH4OH]

= 1.76 x 10-5

pKb = - log10 Kb (definition)

pKb = - log10 (1.8 x 10-5) = 4.74

Kb and Base Strength

The stronger the base, the larger the Kb and the smaller the pKb:

NH4OH (aq) = NH4+ (aq) + OH- (aq)

Kb = 1.8 x 10-5 pKb = 4.74

stronger

weaker

PO43- (aq) + H2O (l) = HPO4

2- (aq) + OH- (aq) Kb = 4.5 x 10-2 pKb =

1.34

Conclusion: phosphate anion is a stronger base than NH4OH.

Kb's of weak bases Strength (Ranked)

2

1

6

7

3

5

4

Acids and Bases - Three Definitions

Arrhenius Definition:

Acids: increases [H+] in aqueous solution

Bases: increases [OH-] in aqueous solution

Bronsted-Lowry Definition: (based on proton transfer reactions)

Acids: proton (H+) donor

Bases: proton (H+) acceptor

Lewis Definition:

Acids: electron pair acceptor

Bases: electron pair donor

Some Lewis Acids and Bases

Lewis bases are characterized by having an available lone pair. Examples are:

O-H-..

.

.:N-HH

H: O-H

H:..I.

.

.

.: :

.

.

:S :

.

.2-

hydroxide iodide ammonia water sulfide

Lewis acids are electron deficient - i.e., electron pair acceptors

Examples are:

H+ Zn2+ Hg2+ Ag+ BF3metal cations electron

deficient compounds

Lewis Acids/Bases - the Most General Definition

Least general definition Most general definition

Arrhenius Bronsted-Lowry Lewis

H+(aq) + :OH-(aq) = H2O(l)

Electron pair acceptor electron pair donor

The Lewis definition generalizes the acid/base concept:

Every Arrhenius acid/base is also a Lewis acid/base.

Every Bronsted acid/base is also a Lewis acid/base.

Example: A strong acid reacts with a strong base:

The Lewis Acid-Base ReactionLewis definition:

Acid: electron pair acceptor Base: electron pair donor

HNO2 + ClO2

- = HClO2 + NO2-

Bronsted-Lowry: acid1 base2 acid2 base1

The Lewis acid is H+ (the electron deficient species)

There are 2 bases (NO2- and ClO2

-), which compete for the acid

The lone pairs donated by these bases are on oxygen atoms:

O=N- O....

.

.

.

. :.. O-N= O

.

..

...

.

.:..[ ]-O=Cl-

O

.

..

...

.

. :..

H..

O=N- O....

.

.

.

. :H.. O=Cl-

O

.

..

...

.

. : -O-Cl= O

.

...

:[ ]..

.

.

.

.

.

...

.

.+

+

Lewis Acids and Bases

The acid/base concept is further generalized by the Lewis acid/base definition. The driving force is the donation of an electron pair to electron-deficient atom.

Lewis acid - an electron pair acceptor

Lewis base - an electron pair donor

H+ + :O-H- = H-O

Lewis acid Lewis base

.

..

.

.

. :

H

F - B + N - H = F - B : N - H

F

F:

F

H

H

F H

H

Lewis acid Lewis base

Complex Ions in Solution

One example of Lewis acid-base neutralization involves the stepwise complexes formed between Hg(II) and I- ion.

There are four stepwise reactions, each of which is an acid/base neutralization by the Lewis definition:Hg2+(aq) + I-(aq) = HgI+

(aq)HgI+(aq) + I-(aq) = HgI2(s) (red-brown ppt)HgI2(s) + I-(aq) = HgI3

-

(aq)HgI3

-(aq) + I-(aq) = HgI4

2-(aq)Identify the Lewis acid and Lewis base in each reaction.

Complex Ions and Solubility

Stepwise Lewis acid/base complexes form between Al3+(aq) and OH- ion. All charged species are soluble in aqueous solution. Only the uncharged Al(OH)3(s) forms a white precipitate.

There are four stepwise reactions, each of which is an acid/base neutralization by the Lewis definition:

Identify the Lewis acid and Lewis base in each reaction.

Al3+ + :OH- = AlOH2+

AlOH2+ + :OH- = Al(OH)2

+

Al(OH)2+ + :OH- =

Al(OH)3 (s)Al(OH)3 + :OH- = Al(OH)4

-

white precipitate