1. For each of the following reactions, identify the Bronsted acid, the Bronsted base, the conjugate base, and the conjugate
acid. Also, for each reaction predict whether the equilibrium lies predominantly on the reactants side or the products
side. Refer to Figure 15.12 in the textbook.
Remember a reaction will proceed from the stronger acid and the stronger base to the weaker acid and the weaker base.
(a) HF (aq) + H2O (l) W H3O+ (aq) + F– (aq)
Bronsted acid Þ HF
Bronsted base Þ H2O
Conjugate base Þ F –
Conjugate acid Þ H3O +
H3O+ is a stronger acid than HF; F – is a stronger base than H2O
Equilibrium lies predominantly on the reactants side.
(b) NH3 (aq) + H3PO4 (aq) W H2PO4– (aq) + NH4+ (aq)
Bronsted acid Þ H3PO4
Bronsted base Þ NH3
Conjugate base Þ H2PO4 –
Conjugate acid Þ NH4 +
H3PO4 is a stronger acid than NH4+; NH3 is a stronger base than H2PO4 –
Equilibrium lies predominantly on the products side.
2. Identify the Lewis acid and the Lewis base in each of the following reactions.
Lewis acid Þ electron-pair acceptor Lewis base Þ electron-pair donor
(a) Ag + (aq) + 2 NH3 (aq) W Ag(NH3)2 + (aq)
Lewis acid Þ Ag+
Lewis base Þ NH3
(b) FeBr3 (s) + Br – (aq) W FeBr4 – (aq)
Lewis acid Þ FeBr3
Lewis base Þ Br –
3. Write an equation and the corresponding equilibrium-constant expression to describe the proton transfer reaction that
occurs when each of these acids is added to water.
(a) HCN
HCN(aq) + H2O(l) W H3O+(aq) + CN – (aq)
Ka = ([H3O+ ][CN – ])/[HCN]
(b) CH3NH3+
CH3NH3+ (aq) + H2O(l) W H3O+(aq) + CH3NH2 (aq)
Ka = ([H3O+ ][CH3NH2 ])/[CH3NH3+ ]
4. (a) Calculate the pH of a solution with [H3O+] = 4.5 x 10-5 M.
pH = – log10[H3O+ ]
pH = – log10(4.5 x 10-5)
pH = 4.35
(b) Calculate [H3O+] for a solution with pH = 2.52.
[H3O+ ] = 10 – pH
[H3O+ ] = 10 – 2.52
[H3O+ ] = 3.0 x 10 -3 M
(c) Calculate the pH of a solution with [OH–] = 3.2 x 10-3 M
There are two ways to solve this problem.
(1) pOH = – log10[OH – ]
pOH = – log10(3.2 x 10 -3)
pOH = 2.49
pH = 14.00 – pOH = 14.00 – 2.49
pH = 11.51
(2) Kw = [H3O+ ][OH – ] = 1.0 x 10 -14
[H3O+ ] = (1.0 x 10 -14)/[OH – ]
[H3O+ ] = (1.0 x 10 -14)/(3.2 x 10 -3) = 3.1 x 10 -12 M
pH = – log10[H3O+ ] = – log10(3.1 x 10 -12)
pH = 11.51
5. Identify the major and minor (if present) species present in the following aqueous solutions.
(a) HCOOH(aq)
Weak acid
major species Þ HCOOH
minor species Þ H3O+ and HCOO –
(b) HNO3(aq)
Strong acid
major species Þ H3O+ and NO3 –
no minor species present
6. What mass of HBr should be present in 0.500 L of solution to obtain a solution with each of the following pH values?
Since HBr is a strong acid we can use the following flow diagram for this problem.
pH ® [H3O+] ® [HBr] ® moles HBr ® mass HBr
Remember that moles = MV. The molar mass of HBr is 80.91 g/mol.
(a) pH = 2.75
[H3O+ ] = 10 – pH = 10 – 2.75 = 0.0018M = [HBr]
(0.0018 mol/L)(0.500 L)(80.91 g/mol) = 0.073 g
(b) pH = 1.60
[H3O+ ] = 10 – pH = 10 – 1.60 = 0.025M = [HBr]
(0.025 mol/L)(0.500 L)(80.91 g/mol) = 1.0 g
(c) pH = 0.75
[H3O+ ] = 10 – pH = 10 – 0.75 = 0.18M = [HBr]
(0.18 mol/L)(0.500 L)(80.91 g/mol) = 7.3 g
