Acids and Bases

Bronsted-Lowry acid-base theory

• Acids are H⁺ donors.

• Bases are H⁺ acceptors

• Alkalis are soluble bases

Conjugate acid-base pairs

• Conjugate acid-base pairs are related by the gain or loss of a proton.

• e.g. NH₄⁺ + NO₃^- → NH₃+ HNO₃

• NH₄⁺ (acid) and NH₃ (base) are a conjugate acid-base pair.

• NO₃^- (base) and HNO₃ (acid) are also a conjugate acid-base pair.

Strong and weak acids and bases

• Strong acids fully dissociate in solution.

• General equation: HA → H⁺ + A^-

• E.g. HCl → H⁺ + Cl^-

• Strong bases also fully dissociate in water.

• Weak acids partially dissociate in solution.

• General equation: HA ⇋ H⁺ + A^-

• E.g. CH₃COOH ⇋ H⁺ + CH₃COO^-

• Weak bases also only partially dissociate in solution

pH

pH = -log[H⁺]

• Negative logarithmic scale so that the lowest pH corresponds to the highest [H⁺].

• The equation can be rearranged to find [H⁺]

[H⁺] = 10^-pH

pH calculations for strong acids

• Strong acids dissociate fully so for monoprotic acids [HA] = [H⁺], for diprotic acids 2[HA] = [H⁺]…

• The concentration can be directly inserted into the pH equation to find pH.

pH of pure water

• Pure water dissociates to a small extent. This can be represent by an equilibrium.

H₂O ⇋ H⁺ + OH^-

• An K_c expression can be written for this equilibrium.

• K_c = [H⁺][OH^-]

  • [H₂O] is constant so is not included.

• This K_c is referred to as the equilibrium constant of water (K_w)

• K_w = [H⁺][OH^-]

• K_w at 298K is 1.00 × 10^-14 mol²dm^-6 (given on the data sheet).

pH of strong bases

• Strong bases full dissociate, so for monobasic bases [base] = [OH^-], for dibasic bases 2[base] = [OH^-]…

• To calculate the pH of a strong base, K_w can be used to find the [H⁺]

• Rearrange K_w expression: [H⁺] = K_w / [OH^-]

• Once the [H⁺] has been calculated it can be inserted into the pH equation to find pH.

pH of weak acids

• Weak acids do not fully dissociate, so [HA] ≠ [H⁺]

• Because the dissociation of weak acids is an equilibrium, we can define an equilibrium constant.

• HA ⇋ H⁺ + A^-

• K_c = [H⁺][A^-] / [HA]

• This equilibrium constant is referred to as the acid dissociation constant (K_a)

• K_a = [H⁺][A^-] / [HA]

• Because K_a values are usually so small, they are often converted to pK_a values

• pK_a = -logK_a

• K_a = 10^-pK_a

• 2 assumptions are made when calculating the pH of a weak acid:

  1. At equilibrium, the concentration of HA is effectively unchanged as so little dissociates, so [HA]_initial = [HA]_equilibrium

  2. [H⁺] = [A^-]

• K_a formula can be rearranged so: [H⁺] = square root of (K_a x [HA])

• Then pH can be calculated by inserting the [H⁺] into the pH equation

pH of mixing acids and bases

• Identify which reactant is in excess

  • If acid, is it strong or weak

    • For strong acid, use concentration of excess H⁺

    • For weak acid, use K_a formula however [H⁺] ≠ [A^-] in this case as HA will react with OH^- to form some A^-

  • If base use K_w formula