Acids, bases and pH

• acids dissociate and release H+ ions in aqueous solution
• alkalis dissociate and release OH- ions in aqueous solution
• Acid-base reactions transfer protons, H+

  • An acid donates a proton
  • A base accepts a proton
• H+ ions are neutralised by OH- ions to form water: H+(aq) + OH- (aq) → H2O(I)
• An alkali is a soluble base
• The Bronsted-Lowry model for acids and bases extends this emphasise the role of proton transfer between species.

  • A Bronsted-Lowry acid is a proton donor
  • A Bronsted-Lowry base is a proton acceptor
• HCl is a strong acid
• HCl (aq) and Cl-(aq) are called a conjugate acid-base pair
• A conjugate acid-base pair contains two species that can be interconverted by transfer of a proton
• In the forward direction, HCl releases a proton to form its conjugate base, Cl-
• In the reverse direction, Cl- accepts a proton to forms its conjugate acid, HCl
• HNO3 is a monobasic acid
• H2SO4 is a dibasic acid
• H3PO4 is a tribasic acid
• Ka = [H+][A-] / [HA]
• pH = -log[H+]
• pKa = -logKa
• pH is a measure of the concentration of H+
• the unit of Ka is mol dm-3
• Kc = products/reactants
• Kp is partial pressure equilibrium
• In the weak acid equation, the forward reaction is usually endothermic (+ve)
• [HA] equilibrium = [HA] start
• Ka is temperature dependent, different for different weak acids
• Ka is only for weak acids because it has a reversible reaction
• Kw = [H+ (aq)][OH- (aq)]
• Increasing the temperature of pure water will decrease its pH
• As temperature decreases, water becomes less acidic. This is because the backwards reaction is exothermic (-ve), so a decrease in temp will favour the backwards reaction. Therefore less H+ will be reduced and so will be less acidic.