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, O2O2- (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
donates H+
Acid -----------> conjugate base
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)
accepts H+
base -------------> conjugate acid
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 BrønstedLowry definition
H+ is an electron pair acceptor; a Lewis acid
Soluble metal oxides are strong bases
2-
O
H
O H
H
O
-
+
O H-
NH3 + H2O  NH4+ + OHbase acid
H
H N
H
H
O H
H +
H N H
H
+
O H-
Reactions between electron deficient and electron-rich
molecules
BF3(g)
Lewis acid
+
NH3 (g)  F3B - NH3 (s)
Lewis base
H
H N
H
H
+ B H
H
H
H N
H
H
B H
H
B-N bond is called a coordinate covalent bond; formed by the
coordination of an electron-pair donor to an electron pair
acceptor
Amphoterism
H2O: 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) + CO32- (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) ]
DGro = DGfo(H3O+(aq)) + DGfo(OH-(aq)) - 2 DGfo(H2O(l))
= + 79.89 kJ/mol
DGro = - 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