Acid, bases and buffers

• Kw

  ionic product of water
• Monobasic acid

  Acid that donates one mole of a proton
• Brønsted-Lowry acid

  Species that donates a proton
• monobasic acids

  HCl

  HCN

  HNO3

  CH3COOH
• dibasic acid

  H2SO4

  H2CO3
• tribasic acids

  H3PO4

  H3BO3
• conjugate acid-base pair

  pair of two species that can be interconverted by a transfer of protons

  • A1 B2 —> A2 B1
  • acids have a positive charge
  • bases have a negative charge
• larger Ka means more dissociation and a stronger acid
• Ka

  acid dissociation constant
• Ka of weak acids

  only paritally dissociate so theres more molecules of weak acid than of H+ and A-
• Assumption of Ka of weak acids

  H+ = A-

  • HA dissociates to produce equal concentrations of the two
  • concentrant on of H+ due to the ionisation of water is negligible compared to the H+ concentration of the acid
• Brønsted-Lowry base

  species that accepts a proton
• Dibasic acid

  one mole of an acid can donate/release 2 protons
• tribasic acid

  one mole of an acid can release/ donate 3 protons
• acids dissociate and release H+ ions in aqueous solutions and alkalis dissociate and release OH- in aqueous solution
• Calculating pH of strong Monobasic acids

  pH = -log[H+]
• calculating the pH of strong bases

  Kw = [OH-][H+]
• calculating the pH of weak acids 


  Ka= [H+]²/[HA]
• Kw increases with increasing temperature as more bonds are broken (endothermic process)
• assumptions for calculating pH of weak acids

  [HA] at the start is the same as [HA] in equilibrium

  • remains constant as dissociation is small

  this doesn’t apply when there’s a stronger weak acid
• assumption for pH of weak acids do not always work

  -[H+] doesnt equal [A-] if the acid is very dilute

  -[HA] is not the same throughout if there’s a stronger weak acid
• range of weak acids lies between pH 3 and 6
• pH of strong monobasic acids

  complete dissociation in solution

  • [H+] = [HA]

  pH of strong acid can be calculated directly from the concentration of an acid
• Kw

  if [H+(aq)] > [OH- (aq)] then the solution is acidic
• Kw

  if [H+(aq)] < [OH- (aq)] then the solution is alkaline
• pH of a buffer solution


  [H+] = Ka x [HA] / [A-]
• what is a buffer solution?


  solution that minimises pH changes on the addition of small amounts of bases or acids
• buffer

  [H+] / Ka = [HA] / [A-]

  • ratio of acid to its conjugate base
• Kw

  if temperature increases, acid dissociates to a greater extent and so the Kw is higher
• acid + base —> salt + water
• acid + metal oxide —>salt + hydrogen (gas)
• acid + alkali —> salt + water
• How are buffers formed?

  weak acids and the salt of a weak acid

  • weak acid partially dissociates in solution and forms a relatively low concentration of A- ion
  • salt fully ionises in solution

  reserve supplies of acid and its conjugate base
• formation of buffers example

  weak acid and salt of weak acid

  • ethanoic acid and sodium ethanoate
  • CH3COOH / CH3COONa
• How are buffers formed?

  partial neutralisation of the weak acid

  • add aq alkali to excess weak acid
  • weak acid is partially neutralised by the alkali forming the conjugate base
  • some of the weak acid is left over unreacted so a mixture of salt and unreacted acid forms
• buffer formation

  partial neutralisation e.g CH3COOH

  CH3COOH + NaOH —> CH3COONa + H2O

  CH3COOH ⇌ CH3COO- + H+
• action of a buffer

  Addition of an acid:

  • equilibrium shifts to the left
  • added H+ react with conjugate base
  • [H+] increases and pH is lowered
• action of a buffer

  Addition of an alkali:

  • equilibrium shifts to the right (to replenish lost H+)
  • H+ (aq) + OH- (aq) —> H2O (l)
  • HA dissociates
  • [OH-] increases and pH increaases
• buffer is most effective when there are equal concentrations of weak acid and its conjugate base
• calculation of the pH of a buffer when its partial neutralisation

  n(acid used)

  n(alkali) = n(conjugate base formed)

  n(acid remaining) = n(acid used) - n(alkali added)