An acid is a substance that can donate a proton (H+ ion) and a base is a substance that can accept a proton

Proton Transfer (Reactions of Acids and Bases)

  • Proton transfer is the movement of a proton from an acid to a base
  • Bronsted-Lowry acids are proton donors
  • Bronsted-Lowry bases are proton acceptors
  • g.

Hydrogen Chloride + Ammonia à Ammonium Chloride

HCl(g) + NH3(g) à NH4Cl(s)

  • The HCl is acting as the acid as it is donating a proton to the NH3
  • The NH3is acting as the base as it is accepting a proton from the HCl

Strong and Weak Acids

  • Strong acids dissociate (or ionise) almost completely in water as nearly all of the H+ ions are released

e.g. HCl(g) + H2O(l) à H+(aq) + Cl(aq)

  • Strong bases dissociate (or ionise) almost completely in water too

e.g. NaOH(s) + H2O(l) à Na+(aq) + OH(aq)

  • In both Strong Bases and Acids, the equilibrium lies over to the right
  • Weak acids dissociate minimally in water, as only a few of the H+ ions are released

CH3COOH(aq) ⇋ CH3COO(aq) + H+(aq)

  • Weak bases dissociate minimally in water

NH3(aq) + H2O(l) ⇋ NH4++ OH(aq)

  • In both weak Bases and Acids, the equilibrium lies over to the left

n.b. Strong acids and concentrated acids are different, likewise weak and dilute acids are different. Strong and weak acids determine how much an acid becomes ionises whereas the concentration, whether that be concentrated or dilute, is the number of moles of an acid.

Water as an Acid and a Base

  • HCl is able to donate a proton to a water molecule, therefore the water acts as the base

HCl + H2O à H3O+ + Cl

  • Oxonium Ion are ions with oxygen cation with three bonds in this case it is the H3O+ (This can also be called hydroxonium or hydronium ion)
  • Water can also act as an acid

H2O + NH3 à OH + NH4+

  • Water is donating a proton to the ammonia

The proton in an Aqueous Solution

  • H+ ions have a diameter of 1015m compared to a hydrogen atom which is 1010m, the small size of the H+ ion results in the ion having unusual properties
  • H+ ions are never found isolated
  • In an Aqueous solution, it is always bonded to at least one water molecule to form a H3O+ ion
  • The H+ ion has no electrons of its own and therefore can only form bonds with other species which have a lone pair of electrons

The Ionisation of Water

  • Water is slightly ionised:

H2O(l) ⇋ H+(aq) + OH(aq)


H2O(l) + H2O(l) ⇋ H3O+(aq) + OH(aq)

  • The equilibrium is established by in water and all aqueous solution:

H2O(l) + H2O(l) ⇋ H+(aq) + OH (aq)

  • The equilibrium expression:

Kc = [H+(aq)][OH(aq)]


  • [H2O(l)] is a constant and is incorporated into a modified equilibrium constant:

Kw = Kc x [H2O(l)]

Kw = [H+(aq)]eqm[OH(aq)] eqm

  • This new constant is called the ionic product of water (Kw)
  • The units for Kw is always mol2 dm-6
  • Kw always has the same value for an aqueous solution at a given temperature e.g. at 298K the Kw is always 1.00×10-14 mol2 dm-6
  • In pure water there is awlys one H+ ion for each OH Therefore [H+] = [OH ] so the Kw of pure water is Kw = [H+]2