1.12 Acids and Bases Flashcards
Brønsted-Lowry acid
Proton (H+) donor
Brønsted-Lowry base
Proton (H+) acceptor
Brønsted-Lowry acid-base reaction
Reaction involving the transfer of a proton
Monoprotic acid
Acid that releases one H+ ion per molecule e.g. HCl
Diprotic acid
Acid that releases two H+ ions per molecule e.g. H2SO4
Definition of pH, and useful rearrangement
pH = – log [H+]
[H+] = 10^-pH
Ionic Product of Water (Kw)
(Kw) H2O ⇌ H+ + OH-
ΔH = endothermic
Kc = [H+] [OH-]/[H2O]
So Kc [H2O] = [H+] [OH-]
so Kc [H2O] = a constant = Kw
Kw = [H+][OH-] where Kw = 1x10^-14
The effect of temperature on the pH of water and the neutrality of water
As the temperature increases, the equilibrium moves right to oppose the increase in temperature
Therefore [H+] and [OH-] increase
So Kw increases and therefore pH decreases
However, the water is still neutral as [H+] = [OH-]
Calculating the pH of water
In pure water, [H+] = [OH-]
So Kw = [H+]^2
Therefore [H+] = √Kw
pH of a Strong Base
Find [OH-] and then substitute into: [H+] = Kw/[OH-]
Mixtures of Strong Acids and Strong Bases
1) Calculate moles H+
2) Calculate moles OH-
3) Calculate moles XS H+ or OH-
4) Calculate XS [H+] or XS [OH-]
5) Calculate pH
Weak Acids
HA ⇌ H+ + A-
The acid dissociation constant, Ka
Ka = [H+] [A-] / [HA]
pKa = -log Ka
Ka = 10^-pKa
These expressions hold for weak acids at all times
The acid dissociation constant, Ka, notes
• Ka – has units mol dm-3
• Ka – the bigger the value, the stronger the acid
• pKa – the smaller the value, the stronger the acid
In a solution of a weak acid in water, with nothing else added:
a) [H+] = [A-]
b) [HA] ~ [HA]initial
Ka = [H+]^2 / [HA]
This expression ONLY holds for weak acids in aqueous solution with nothing else added