Oxidation, reduction & redox reactions

• oxidation states/numbers = used to see reduction & oxidation processes more easily (movement/distribution of electrons)

  • reduction = gain of electrons/loss of oxygen
  • oxidation = loss of electrons/gain of oxygen
• Rules to determine an element's oxidation state

  1. Oxi. state of element is 0
  2. overall oxi. state of a compound is 0
  3. overall oxi. state of an ion is equal to it's charge
  4. groups 1, 2 & 3 metals in a compound will have an oxi. state equal to their charge
  5. F is always -1 in a compound
  6. H is normally +1 in a compound
  7. O is normally -2 in a compound
  8. transition metals' oxi. state given by roman numerals
• Redox + writing half equations

  • reduction = increase in oxidation state
  • oxidation = decrease in oxidation state

  Rules for half equations

  1. Balance the element
  2. determine oxidation state change
  3. add in electrons
  4. balance O with H2O
  5. Balance with H+
• Electro-potentials

  • when solid metal added to solution of its own ions -> equilibrium is set up
  • potential = how readily electrons are lost by a metal
  • how good a reducing agent the metal is
  • more negative pd -> better reducing agent
  • difference in potentials caused by different reactivities of separate half-cells
• Standard Hydrogen Electrode

  • Potential of SHE set to 0.00V -> used as reference point for finding potentials of other half cells
  • always positioned on the left
  • Solution contains H+ ions (i.e. an acid) with 1 mol/dm3 concentration & inert platinum solid used to increase SA for contact
• Cell Conventions

  • Reduction on the right (positive electrode) & oxidation on the left (negative electrode)
  • +ve value -> material is reduced (RHS)
  • -ve value -> material is oxidised (LHS)
  • Ecell has to be positive to work
  • Ecells can be created between ions of the same substance but must have different concentrations
• Commercial electrochemical cells

  • electrochemical cell = controls electron transfer to produce electrical energy
  • basis of cells is to control the transfer of electrons
• Non-rechargeable cells

  • provide electrical energy until the chemicals have reacted to such an extent that the voltage falls
  • use of irreversible reactions
  • Pros:
  • cheaper than rechargeable
  • lasts longer per charge
  • contains less toxic metals -> less hazardous
  • cons:
  • replaced regularly & difficult disposal
  • don't supply a lot of power
  • Can leak & pollute water
• Rechargeable cells

  • chemicals in cell react to provide electrical energy
  • use of reversible reactions
  • examples: Lead/Acid, Nickel/Cadium, Lithium ion batteries
  • can use graphite or paste as electrolyte medium to prevent leaks
• Fuel cells

  • Create voltage using energy from reaction between a fuel & O2
  • Greater efficiency in converting chemical energy
  • E.g. (Alkali) hydrogen-oxygen cell -> OH- ions from electrolyte oxidise and O2 reduces to form more OH- ions

  2H2 + 4OH- -> 4H2O + 4e- O2 + 2H2O + 4e- -> 4OH-
• Limitations of fuel cells

  • Storage of H2
  • could be stored as a pressurised liquid or adsorbed onto material surface or absorbed within a material
  • alternatives: using methanol or ethanol for cells instead