Acids and Bases

• Why the expression for Kw does not include the concentration of water?


  -very few H+ and OH- ions. Only slightly dissociates

  • H2O is effectively constant
• Why Kw increases as the temperature of water increases?


  • Dissociation is endothermic
  • Equilibrium moves to the RHS to absorb heat
• Burette is a more suitable piece of apparatus to measure out NaOH solution than pipette
• Why is burette better than pipette to measure out solutions?


  It can deliver variable volumes
• Why the end point of a reaction would be difficult to judge accurately using an indicator on a graph?


  change in pH is gradual, not rapid at the end point

  • indicator would change colour over a range of volumes of 'solution'
• suggest a substance in the air that might make the pH of a solution change?

  CO2
• What happens to pH when an basic solution reacts with CO2?

  CO2 + 2OH- -> CO32- +H2O

  pH decreases, reacts with acidic gas
• Why the pH of a acidic buffer solution remains almost constant despite the addition of a small amount of NaOH?


  Added OH- removed by reaction with H+

  equilibrium moves to the right
• Why is pure water not acidic?

  H+ = OH-
• Why small volume of water doesn't affect pH of buffer solution?


  Ratio of (A-):(HA) remain constant

  So, (H+) = Ka(HA)/(A-)

  (H+) remains constant
• Acid is defined as a substance that can donate a proton
• A base is defined as a substance that can accept a proton
• Calc pH of strong acids


  pH = -log(H+)
• Strong acids completely dissociate. The concentration of H+ is will be the same as the concentration of the acid
• Ionic product for water equation

  Kw = (OH-) (H+)
• Monoprotic acid is an acid which can release only one H+ upon dissociation
• Diprotic is an acid which can release 2 H+ ions upon dissociation
• Concentration of hydrogen ions in a monoprotic strong acid will be the same as the concentration of the acid
• Pure water are not acids/bases because (H+)=(OH-)
• Dissociation of water is endothermic so increasing the temperature, equilibrium shifts to the right, increase concentration of H+ and lower pH
• Larger the Ka stronger the acid
• pKa = -logKa
• Ka expression for weak acid 


  Ka = (H+) (A-)/(HA)
• Buffer solution - solutions that resist slight change in pH when small amount of acid, base or water is added
• acidic buffer solution is made from a weak acid and salt of a weak acid
• basic buffer solution is made from a weak base and a salt of that weak base
• weak acids only slightly dissociate so there is a low concentration of (H+) and (OH-) ions
• If small amounts of acid is added to acidic buffer then the equilibrium will shift to the left to remove nearly all the H+ ions added
• In a buffer solution there is a much higher concentration of the salt ions than the pure acid
• If a small amount of alkali is added to acidic buffer, the OH- ions will react with H+ ions to form water. Equilibrium will shift to the right to produce more H+ ions
• If small amount of acid added to basic buffer, H+ will react with OH-, equilibrium will shift to the right to the replace OH- that reacted with H+
• If small amount of alkali added to a basic solution, OH- will react with NH4+, equilibrium shifts to the left
• As there is a large concentration of the salt ion in the buffer, the ratio (HA):(A-) stays almost constant, so the pH stays fairly constant
• Some ethanoic acid molecules are changed to ethanoate ions but as there is a large concentration of salt ion in the buffer, the ration of (HA):(A-) stays almost constant, so pH stays fairly constant
• Use weak acids dissociation (Ka) to calculate the pH of buffer solutions
• Indicators can be considered as weak acids. The acid must have a different colour to its conjugate base.
• An indicator changes colour from Hln to ln- over a narrow range. Different indicators change colour over different ranges.
• Phenolphthalein for strong bases only. Colour changes from colourless to pink
• Methyl orange with strong acids only . Colour changes from red acid to yellow alkali
• HCO3- <-> CO32- + H+

  How can the solution act as a buffer when small amount of acid or alkali are added


  Acid : Concentration of H+ increase, equilibrium shift to the left

  Alkali : OH- reacts with H+, equilibrium shifts to the right

  Concentration of H+ almost remain constant