Acid-Base Equilibria

Acids and bases are some of the more commonly encountered chemicals

Acids and Bases control composition of blood and cell fluids, affect flavors, involved in digestion

Bases used in house hold cleaners (NH3-based cleansers)

Acid rain is an environmental problem

Acids and bases are involved in reactions that produce polymers, synthetic fibers, dyes.

Arrhenius Acid & Base

Acid: produces H+ in aqueous solution

Base: produces OH- in aqueous solution

HCl(aq) H+ (aq) + Cl- (aq)

NaOH(aq) Na+ (aq) + OH- (aq)

Acid + base neutralization: H+(aq) + OH-(aq) H2O(l)

However, an H+ cannot exist by itself in water

Brønsted-Lowry Acids and Bases

Acid: proton donor

Base: proton acceptor

H+: PROTON since the H+ consists of 1 proton and 0 electron

HCN(aq) + NH3(aq) NH4+(aq) + CN-(aq)

acid base

The H+ is transferred from HCN to NH3

HCN is said to have an acidic H, a hydrogen that can be donated as a H+

HCl has an “acidic” H+, but by itself cannot act as an acid

However HCl(aq):

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

H3O+: hydronium ion

HCN - hydrogen cyanide

HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)

Only a fraction of HCN donate their H+ to H2O

HCN is a weak acid

At equilibrium there is both CN- and un-dissociated HCN

In the Brønsted-Lowry theory:

a strong acid is fully deprotonated in solution

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

a weak acid is only partially deprotonated in solution

HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)

Typically the solvent is water, but not necessarily.

An acid that is strong in water, may be weak in another solvent

Brønsted-Lowry Base

A proton acceptor. In most cases the molecule possesses a lone pair of electrons to which a H+ can bond to.

Example: Oxide, O2-

O2- (aq) + H2O(l) 2 OH- (aq)

Strong base since all O2- (aq)

forms OH- (aq)

NH3: a Brønsted base. The lone pairs on N in NH3 can bond with a H+.

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

NH3(aq) is a weak base; at equilibrium both undissociated NH3 (aq) and NH4

+ (aq) exist.

A strong base is completely protonated in solution

O2- (aq) + H2O(l) 2 OH- (aq)

A weak base is partially protonated in solution

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

Strength depends on solvent

Solvent Leveling

Since all strong acids are completely de-protonated in water (behave as though they were solutions of H3O+) strong acids are “leveled” in water

To compare acidity of acids that are strong acids in water, need to use a solvent in which the “acidity” of the acids differ

Strong bases are leveled in water in the same way as strong bases.

Arrhenius definition restricted to water as a solvent

However Brønsted-Lowry theory includes non-aqueous solvents

CH3COOH (l) + NH3 (l) CH3COO- (am) + NH4+(am)

am - denotes a species dissolved in ammonia

Brønsted-Lowry includes acid/base in the absence of solvent

Protons can be transferred in the gas phase:

HCl(g) + NH3(g) NH4Cl(s)

Acid-base reaction does not have to involve the solvent

HCN(aq) + NH3(aq) NH4+(aq) + CN-(aq)

Conjugate Acids & Bases

HCN(aq) + H2O(l) H3O+(aq) + CN-(aq)

acid conjugate base

CN- (aq) is the conjugate base of HCN

Brønsted-Lowry acids form conjugate bases

Acid -----------> conjugate basedonates H+

Brønsted-Lowry bases form conjugate acids

NH3 (aq) is the base; NH4+ (aq) is the conjugate acid

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

base -------------> conjugate acidaccepts H+

An acid is a proton donor and a base is a proton acceptor.

The conjugate base of an acid is the base formed when the acid has donated a proton.

The conjugate acid of a base is the acid that forms when the base has accepted a proton.

Lewis Acids & Bases

A Lewis base donates a lone pair of electrons

A Lewis acid accepts a lone pair of electrons

Lewis acids/bases are a broader definition than the Brønsted-Lowry definition

H+ is an electron pair acceptor; a Lewis acid

Soluble metal oxides are strong bases

O2-

O H

H O H

O H-

-

+

NH3 + H2O NH4+ + OH-

base acid

Reactions between electron deficient and electron-rich molecules

BF3(g) + NH3(g) F3B - NH3 (s)

Lewis acid Lewis base

O H

H

O H-+

NH

HH

NH

HH H

+

B-N bond is called a coordinate covalent bond; formed by the coordination of an electron-pair donor to an electron pair acceptor

NH

HH B

H

HH+ N

H

HH B

H

HH

AmphoterismH2O: acts as both an acid and a base - amphoteric

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

OH- conjugate base of H2O

H3O+ conjugate acid of H2O

HCO3- is amphoteric

HCO3- (aq) + H2O(l) H3O+ (aq) + CO3

2- (aq)

HCO3- (aq) + H2O(l) H2CO3

(aq) + OH- (aq)

Water is amphiprotic - both an acid and a base

When one molecule transfers a proton to another molecule of the same kind - autoprotolysis or autoionization

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

An O-H bond is strong; the fraction of protons transferred is very small.

Calculate the equilibrium constant for the autoionization of H2O(l)

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

Kw = [H3O+(aq) ] [OH- (aq) ]

Gro = Gf

o(H3O+(aq)) + Gfo(OH-(aq)) - 2 Gf

o(H2O(l))

= + 79.89 kJ/mol

Gro = - R T ln Kw

Kw = 1.0 x 10-14 at 298 K

Kw = 1.0 x 10-14 at 298 K

Kw = [H3O+(aq) ] [OH- (aq) ]

[H3O+(aq) ] [OH- (aq) ] = 1.0 x 10-14

Kw is an equilibrium constant; the product of the concentrations of H3O+ and OH- is always equal to Kw.

In pure water [H3O+(aq) ] = [OH- (aq) ] = 1.0 x 10-7 M at 298 K

If the concentration of [OH-(aq)] in increased, then [H3O+(aq) ] decreases to maintain Kw.

What are the molarities of H3O+ and OH- in 0.0030 M Ba(OH)2 at 25oC?

Ba(OH)2 (aq) Ba2+ (aq) + 2 OH- (aq)

Molarity of [OH- (aq)] = 0.0060 M

[H3O+ (aq)] = Kw/[OH- (aq)] = 1.7 x 10-12 M

pH Scale

The concentration of H3O+ can vary over many orders of magnitude

A log scale allows a compact description of the H3O+ concentration.

-

-

-

-

pH = - log [H3O+]

[H3O+] = 10- pH mol/L

For pure water at 25oC

pH = - log (1.0 x 10-7) = 7.00

For a change in pH by 1, H3O+ concentration changes by 10

Higher pH, lower H3O+ concentration

pH of pure water is 7

pH of an acidic solution is less than 7

pH of a basic solution is greater than 7