Acids and Bases

• calculating the pH of acids:

  A) -log

  B) H+

  
• Ka - acid dissociation constant

  • finding the pH of weak acids

  
• calculating pH of a weak acid without a buffer:

  A) H+

  B) 2

  C) HA

  
• calculating pKa:

  A) -log

  B) Ka

  
• inverse pKa equation:

  A) 10

  B) -pKa

  
• bronsted-lowry acids are proton donors
• bronsted-lowry bases are proton acceptors
• when acids and bases react with water, the reaction is reversible
• strong acids dissociate completely, whereas weak acids only partially dissociate
• strong acids dissociate completely because:

  • forwards reaction is favoured strongly
  • lots of H+ ions produced
• weak acids partially dissociate because:

  • backwards reaction is favoured
  • not many H+ ions produced
• • water acts as a base when acid is added to it
  • it is amphoteric - can act as a base or an acid
• Kw:

  • the ionic product of water
  • at room temperature, the value of Kw is 1 x 10-14 mol2 dm-6
• the value of Kw changes with temperature
• • water: [H+] = [OH-]
  • pure water: Kw = [H+]^2
• pH is a measure of the concentration of H+ ions in a solution
• monoprotic acids dissociate to produce one H+ ion per acid molecule
• • diprotic acids dissociate to produce 2 H+ ions per acid molecule
• calculating the pH of bases:

  1. [H+] = Kw  ÷ [OH-]
  2. pH = -log[H+]
• • pKa measures the strength of an acid, like pH
  • the lower the value of pKa, the stronger the acid
• • acidic buffers resist the change in pH to keep the solution below pH 7
  • they are made from a weak acid and its salt
• Kw expression:

  A) H+

  B) OH-

  
• inverse pH equation:

  A) 10

  B) -pH

  
• why is [H2O] not shown in the Kw expression:

  • it is constant
• Explain why the value of Kw increases as the temperature increases:

  • forwards reaction is endothermic
  • equilibrium position shifts to the right hand side, in the endothermic direction, to oppose the change and lower the temperature
• why is the pH probe washed with distilled water between each of the calibration measurements:

  • to avoid contamination between solution
  • so that residual solution doesn't interfere with the pH reading
• how to determine Ka value from a titration curve:

  • find volume at half-equivalence point
  • read off pH value
  • Ka = 10^-pH
• titration steps:

  • Place a fixed volume of alkali in a flask or beaker
  • Add acid in small portions from a burette
  • Stir with magnetic stirrer
  • use a pH meter to record the pH after each addition of acid
• uses of buffers:

  • shampoo
  • washing powder
  • blood
• explain why a burette is more suitable than a pipette

  it can release variable volumes
• use information from the curve in figure 1 to explain why the end point would be difficult to judge accurately using an indicator

  the change in pH is gradual at the end point

  indicator wouldn't change colour rapidly

  
• calculate the pH of the solution when half of the acid has reacted

  • pH at the half equivalence point is EQUAL TO pKa
  • pKa = -log (Ka)
• calculate the concentration of HX in the original solution

  • use the graph - at the end point [OH-] = [HX]