8.4    Theory of Acids and Bases

The Arrhenius theory

• According to Arrhenius theory:
  1. An acid is a substance which ionises in water to give hydrogen ions, H⁺.
  2. A base is a substance which ionises in water to give hydroxide ions, OH⁻.

• Limitations of this theory:
  1. When ammonia gas reacts with hydrogen chloride gas, ammonium chloride is produced. Although this really is an acid-base reaction but it contradicts with the Arrhenius theory because no H⁺ or OH⁻ is produce

The Brønsted-Lowry theory

• According to the Brønsted-Lowry theory:
  1. An acid is a proton donor.
  2. A base is a proton acceptor.
• A proton is a hydrogen ion, H⁺.
• For example, when hydrogen chloride dissolves in water to form hydrochloric acid, the following reaction occurs:

HCl(g) + H2O(l) → H3O⁺(aq) + Cl⁻(aq)…………(2)

HCl is acting as an acid because it has donated a proton. H2O is acting as  a base because it has accepted a proton.

 

• When the acidic solution reacts with a base, what is actually functioning as an acid is the hydroxonium ion, H3O⁺.

H3O⁺ + OH⁻ → 2H2O

H3O⁺ is acting as an acid because it has donated a proton. OH⁻ Is acting as a base because it has accepted a proton.

 

• Brønsted-Lowry acids and bases do not nave to involve aqueous solutions

 

Conjugate pairs

• When ammonia gas dissolves in water, the reaction that occurs is reversible: NH3(aq) + H2O(l) ⇌ NH4⁺(aq) + OH⁻(aq)…………..(3)

In the forward reaction, H2O is acting as an acid because it has donated a proton and NH3 is acting as a base because it has accepted a  proton.

In the backward reaction, OH⁻ is acting as a base because it has accepted a proton and NH4⁺ is acting as an acid because it has donated a  proton.

• Therefore OH⁻ is the conjugate base of the acid H2O while NH4⁺ is the conjugate acid of the base NH3.
• In general:
  1. Every acid has a conjugate base, this is the particle left when the acid has given away its proton
  2. Every base has a conjugate acid, this is the particle left when the base has accepted a proton
• Alternatively, the acid-I, base-II terminology can also be used:
  1. HA is acid-I and A⁻ is base–I, they are one conjugate pair.
  2. H2O is base-II and H3O⁺ is acid-II, they are another conjugate pair
• Substances which can behave as an acid as well as base are described as amphoteric. One example is water:
  1. In reaction (2), water is behaving as a base
  2. In reaction (3), water is behaving as an acid

 

Strength of  acids and bases

• A strong acid is one which dissociates completely in a solution. HCl(g) + H2O(l) → H3O⁺(aq) + Cl⁻(aq)
• This produces high concentration of hydroxonium Therefore the pH of the solution is very low, pH ≈ 1.
• Examples of strong acids are HCl, H2SO4 and  HNO3.

 

• A strong base is one which dissociates completely in a solution. NaOH(s) + aq → NaOH(aq)
• This produces high concentration of hydroxide Therefore the pH of the solution is very high, pH ≈ 14.
• Examples are Group I metal hydroxides

 

• A weak acid is one which dissociates partially in a solution. CH3COOH(l) + H2O(l) ⇌ CH3COO⁻(aq) + H3O⁺(aq)
• This produces very low concentration of hydroxonium ions, the position of equilibrium is far over the Therefore the pH of the solution is higher, pH ≈ 3.
• Examples are organic acids

 

• A weak base is one which dissociates partially in a NH3(aq) + H2O(l) ⇌ NH4⁺(aq) + OH⁻(aq)
• This produces very low concentration of hydroxide ions, the position of equilibrium is far over the Therefore the pH of the solution is lower, pH ≈ 12.
• Examples are ammonia, amines and some hydroxides of transition metals

 

• Note:
  1. Strength of acids and bases is defined in terms of degree of dissociation while concentration is defined as the number of moles per unit volume.
  2. Therefore a weak acid in high concentration is still classified as a weak acid
  3. Also, a strong acid in low concentration is still classified as a strong acid