5.1.3 - acids, bases, buffers

• A bronsted-Lowry acid is a proton donor
• a bronsted-Lowry base is a proton acceptor
• Monobasic acids release one proton when they dissociate in water

  eg, HCl
• Dibasic acids release two protons when they dissociate in water

  eg, H2SO4
• tribasic acids release three protons when they dissociate in water

  eg, H3PO4
• CH3COOH + H2O <——> CH3COO- + H3O+

  acid 1 base 2 base 1 acid 2
• A strong acid completely dissociates
• examples of strong acids are:

  hydrochloric acid

  sulfuric acid

  nitric acid
• a weak acid partially dissociates
• examples of weak acids are:

  methanoic acid

  any organic acid
• acid dissociation constant is used to measure the extent of acid dissociation
• symbol of acid dissociation constant is Ka
• for acid HA, HA <——> H+ + A-

  Ka = [H+][A-] / [HA]
• Ka = products / reactants
• the larger the Ka value the greater the extent of dissociation - the stronger the acid
• to convert Ka into pKa:

  pKa = -log10Ka
• to convert pKa into Ka:

  Ka = 10^-pKa
• the smaller the pKa value, the stronger the acid
• A buffer solution is a mixture that minimises changes in pH when small amounts of acid or base are added
• buffer solutions can be made by either:

  • weak acid and its conjugate base
  • weak acid and a strong alkali
• when an acid is added to a buffer solution, equilibrium shifts to the left because the [H+] increases and the conjugate base reacts with the H+ to remove most of the H+
• when an alkali is added to a buffer solution, equilibrium shifts to the right because [OH-] increases and the small concentrations of H+ react with OH-. to restore the H+ ions HA dissociates shifting the equilibrium.
• To calculate the [H+] of buffer solution:

  [H+] = Ka x [HA] / [A-]
• Blood is maintained at pH of 7.4 by hydrogencarbonate buffer
• when an acid/alkali is added to the blood:

  H+ + HCO3 <——> CO2 + H2O

  adding OH- reacts with H+ to form H2O then shifts equilibrium to the left to restore H+ lost.

  adding H+ equilibrium shifts to the right removing excess H+
• A Lewis acid is an electron pair acceptor
• a Lewis base is an electron pair donor
• a H+ ion causes a solution to become more acidic
• an OH- ion causes a solution to become more alkaline
• Equation for ionisation of water:

  H2O <——> H+ + OH-
• concentrated means many mol per dm3 whereas strong refers to amount of dissociation
• Ionisation of water equation:

  H2O <——> H+ + OH-
• The expression for ionic product of water:

  Kw = [H+][OH-]
• the units for Kw are mol2dm-6
• the value of Kw at 298 K is 1.0 x 10-14
• temperature affects Kw.
• in Kw if temperature is increased, equilibrium moves to the right so Kw increases and the pH of pure water decreases
• Indices of [H+] and [OH-] always add up to -14
• A strong base 100% dissociates in water
• examples of strong bases are:

  NaOH

  KOH

  Ca(OH)2