3.1.12 Acids and bases

• Arrhenius Acid : All acids are substances which produce hydrogen ions when added to water

  Arrhenius Base : All bases produce hydroxide ions when added to water.
• Bronsted-Lowry Acid : A proton donor

  Bronsted-Lowry Base : A proton acceptor

  conjugate = formed
• A strong acid is fully dissociated in solution
• A strong base in fully dissociated in solution
• Ionic product of water

  kw = [OH-][H+]

  The concentration of water is very high and water is very weakly dissociated so [H2O] is effectively constant and left out of the expression
• The value of kw increases as temp increases. The dissociation of water is endothermic. As temp increases the position of equilibrium shifts to the right to oppose the increase in temp therefore H+ and OH- increases therefore kw increases
• The pH of pure water

  [H+]=[OH-]

  so kw = [H+]$^2$
• A weak acid partially dissociates in solution
• Weak acid

  ka = [H+][A-]/[HA]/ ka = [H+]^2/[HA]
• pKa = -log10(Ka)
• A buffer solution resists changes in pH when small amounts of acid or base are added to it. It is able to maintain its pH when diluted with water
• Acid Buffers

  Method 1: Mix together a weak acid and its salt e.g. Ethanoic acid and sodium ethanoate

  Method 2: Weak acid and half volume of a strong base of the SAME CONCENTRATION
• Explain how an acidic buffer works

  If a small amount of acid (H ions) is added to the buffer, the equilibrium shifts to the left to oppose the increase in H+ ions. The extra H+ ions combine with CH3COO- forming CH3COOH. 

  This maintains the H+ concentration and the pH.

  If a small amount of base (OH ions) is added to the buffer, the equilibrium shifts to the right to oppose the decrease in H+ ions.

  The equilibrium shifts to the right to oppose the decrease in H+ ions. CH3COOH dissociates to release more H+ ions

  This maintains the concentration and the pH
• Basic Buffer

  Method: Mix together a weak base and its salt e.g ammonia and ammonium chloride
• Explain how basic buffer works

  If a small amount of acid is added to the buffer, the amount of H+ increase, equilibrium shifts to the right to oppose the increase in H+ ions. The extra H+ ions combine with NH3 forming NH4+.

  This maintains the H+ concentration and the pH.

  If a small amount of base is added to the buffer, it reacts with H+ to form H2O which decreases H+ and equilibrium shifts to the left to oppose the decrease in H+ ions. The NH4+ dissociates to release more H+ ions. Maintaining the H+ concentration and the pH
• Weak acid and weak base titrations does not have a sharp change in pH so indicators do not work and you need use a pH meter as it is hard to judge
• pKa=pH at half equivalence of a weak acid
• Half-neutralisation we assume [X-]=[HX]

  so Ka =[H+]
• Pka = -log10 ka