Redox Reaction

• Reduction and oxidation involve the transfer of electrons
• Oxidation


  Loss of electrons
• Reduction


  Gain of electrons
• The reaction where calcium is completely burnt in oxygen is a redox reaction
• Redox reaction


  1. metal is oxidized by losing electrons

  2. halide is reduced by gaining electrons
• Reducing agent


  A chemical that loses electrons and oxidizes itself
• Oxidizing agent


  A chemical that gains electrons and is itself reduced
• Oxidation number


  A number assigned to each element that depends on a set of rules
• Rules for determining oxidation numbers

  • Uncombined elements are always 0
  • Ions have an oxidation number equal to their charge
  • Group 1 elements are always +1
  • Group 2 elements are always +2
  • Aluminium is always +3
  • Hydrogen is +1 except in hydrides where it is -1
  • Chlorine is -1 except when bonded to F or O
  • Fluorine is always -1
  • Oxygen is -2 except in peroxides where it is -1
• Transition metals have variable oxidation states
• Iron has different oxidation states, very typical of transition metals
• Vanadium oxide, vanadium must be +4 to balance the charge
• Another vanadium compound, vanadium must be +5 to balance the charge
• Vanadium is a transition metal, it has different oxidation states
• Reduction


  Decrease in oxidation number
• Oxidation


  Increase in oxidation number
• Sodium is the reducing agent, chlorine is the oxidising agent
• Steps to balance half equations


  1. Write species before and after

  2. Balance atoms apart from O and H

  3. Balance O with H2O

  4. Balance H with H+

  5. Balance charges with e-
• Half equations must have electrons
• Combining two half equations forms a full ionic equation, electrons must balance
• Full ionic equation shows both reduction and oxidation, no electrons should remain
• Redox reactions involve the transfer of electrons between two species, where the electron donor is oxidised and the electron recipient is reduced
• Ionic equation

  Compounds in aqueous solution are written as dissociated ions
• Spectator ions

  Ions that do not take part in the chemical reaction and have an unchanged oxidation state
• Oxidation

  Increase in oxidation state
• Reduction

  Decrease in oxidation state
• Disproportionation reaction

  An element in a single species is simultaneously oxidised and reduced
• Electrode Potentials

  Measure the tendency of a species to gain or lose electrons
• Representing Electrochemical Cells

  1. Identify the anode and cathode

  2. Write the overall cell reaction

  3. Calculate the cell potential
• Half Cells

  Contain a single electrode immersed in a solution of its own ions
• Acids

  Substances that release hydrogen ions (H+) in aqueous solution
• Bases

  Substances that release hydroxide ions (OH-) in aqueous solution
• Titrations

  Measure the volume of one solution required to react completely with a known volume of another solution
• Buffer Action

  Ability of a solution to resist changes in pH when small amounts of acid or base are added
• pH of Strong Bases

  pH = 14 - pKb
• Ionic Product of Water

  Kw = [H+][OH-] = 1.0 x 10-14 at 25°C
• Calculating Ka

  Use the equilibrium equation and concentrations to find the acid dissociation constant
• Acid Dissociation Constant, Ka

  Measure of the strength of an acid
• pH Scale

  Measure of the acidity or basicity of a solution
• Strong Acids

  Completely dissociate in water to produce a high concentration of H+ ions