1B3 Acids and bases

• Acids

  • naturally occurring substances that produce a 'sour' taste
  • found commonly in fruit and vinegar
• Bases

  • react with acids in a process called neutralization
  • alkalis are bases that are soluble in water
  • metal oxides are bases but insoluble in water
• acid + base -> salt + water
• Indicators

  • test to see if a solution is acid or alkali
  • litmus paper: turn red for acids and blue for alkalis
• A concentrated acid or alkali is one which has a lot of the chemical dissolves in a small amount of water
• A dilute acid or alkali is one with only a small amount of the chemical dissolved in a large amount of water
• A Bronsted-Lowry acid is a proton donor
• A Bronsted-Lowry base is a proton acceptor
• Protons in solution

  • regarded as a H+ ion
  • H+ ion is unable to exist in water so combines to a water molecule with a coordinate bond to make a 'true' acid called oxonium
  • H+ + H2O -> H3O+
• A strong acid is the complete ionisation of a substance when it dissolves in water
• A weak acid only partially dissociates in aqueous solution
• Conjugate base

  • A conjugate base is formed when the acid loses a proton on solution
  • It acts as a base by accepting a proton returning to acid form
• Acid-base reactions involve the transfer of protons
• pH = -log10 [H+(aq)]
• pH of a solution is measured using a pH meter or pen which is calibrated using solutions of a known pH called buffer solutions
• pH must always be quoted to 2 decimal places
• [H+] = inverse log10 (-pH)
• Strong acids

  • completely dissociate in aqueous solution
  • monoprotic acids release a single hydrogen
  • diprotic acids release 2 hydrogens
• Ionic product of water, Kw

  • at 25C Kw= [H+][OH-] = 1.00 x 10-14 mol2 dm3
• For neutral solution e.g. pure water

  • Kw= [H+]2
• Strong base equations

  • [H+] = Kw/[OH-]
  • pH = -log[H+]
• Diluting a strong acid or strong base

  1. calculate moles of H+ or OH-
  2. Convert to concentrations by dividing by new volume

  • If H+: pH = -log[H+]
  • If OH-: [H+] = Kw/[OH-] then pH = -log[H+]
• Mixing unequal amount of acid and alkali

  1. calculate moles of H+
  2. calculate moles of OH-
  3. work out what is in excess and convert to concentrations by dividing by the total volume

  • If H+ in excess: pH = -log[H+]
  • If OH- in excess: [H+] = Kw/[OH-] then pH = -log10[H+]
• Water is amphoteric so it can act as an acid or base
• Weak acid equations

  1. Ka = [H+]2/[HA]
  2. pH = -log[H+]
• Ka is the acid dissociation constant and always has units of mol dm-3
• pKa = -log10 Ka
• The weaker the acid the larger the pKa value
• Half neutralisation

  • the concentration of the acid will be the same as the concentration of the salt
  • Ka = [H+]
  • pKa = pH
• A buffer solution is one which maintains a constant pH despite the addition of small amounts of acid or alkali
• Acidic buffer

  • contains a weak acid and its sodium salt
  • pH below 7
  • e.g. ethanoic acid and sodium ethanoate
  • salt dissociates completely but the acid only partially dissociates
• Basic buffer

  • contains a weak base and its salt
  • pH above 7
  • e.g. ammonia and ammonium chloride
  • salt dissociates completely but the base only partially dissociates