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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
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