Acid-Base Equilibria
Bronsted-Lowry Acid→ proton donor
Bronsted-Lowry→ proton acceptor
pH = -log[H+] ⇒ [H+] = 10-pH
Always give pH values to 2 d.p.
Ionic Product of Water
Kc = [H+][OH-]/[H2O] ⇒ Kc * [H2O] = [H+][OH-] ⇒ Kw = [H+][OH-]
Because [H2O] is larger than the concentration of the ions it is a constant so Kc * [H2O]
is a new constant Kw.
@ 25*C Kw = 10-14
Pure water/neutral solutions are neutral because [H+]=[OH-] ∴ Kw = [H+]2
H2O ⇋ H+ + OH- is endothermic so increasing temperature pushes equilibrium to favour
forward reaction so [H+] increases and pH is more acidic/lower
Dissociation Constant
Weak acids partially dissociate ∴ Ka = [H+][A-] / [HA]
pKa = -log Ka ⇒ Ka = 10-pKa
The larger Ka the stronger the acid
Pure weak acid solution→ [H+]=[A-] ∴ Ka = [H+]2 / [HA]
Acid and Base Neutralisation
Bronsted-Lowry Acid→ proton donor
Bronsted-Lowry→ proton acceptor
pH = -log[H+] ⇒ [H+] = 10-pH
Always give pH values to 2 d.p.
Ionic Product of Water
Kc = [H+][OH-]/[H2O] ⇒ Kc * [H2O] = [H+][OH-] ⇒ Kw = [H+][OH-]
Because [H2O] is larger than the concentration of the ions it is a constant so Kc * [H2O]
is a new constant Kw.
@ 25*C Kw = 10-14
Pure water/neutral solutions are neutral because [H+]=[OH-] ∴ Kw = [H+]2
H2O ⇋ H+ + OH- is endothermic so increasing temperature pushes equilibrium to favour
forward reaction so [H+] increases and pH is more acidic/lower
Dissociation Constant
Weak acids partially dissociate ∴ Ka = [H+][A-] / [HA]
pKa = -log Ka ⇒ Ka = 10-pKa
The larger Ka the stronger the acid
Pure weak acid solution→ [H+]=[A-] ∴ Ka = [H+]2 / [HA]
Acid and Base Neutralisation
