Chem revise electrode potentials

• Electrochemical cells


  • A cell has two half–cells
  • The two half cells have to be connected with a salt bridge
  • Simple half cells will consist of a metal (acts an electrode) and a solution of a compound containing that metal (eg Cu and CuSO4)
  • These two half cells will produce a small voltage if connected into a circuit (i.e. become a Battery or cell)
• Salt bridge


  • The salt bridge is used to connect up the circuit
  • The free moving ions conduct the charge
  • A salt bridge is usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate
  • The salt should be unreactive with the electrodes and electrode solutions
• A wire is not used because the metal wire would set up its own electrode system with the solutions
• Why does a voltage form?

  1. When connected together the zinc half-cell has more of a tendency to oxidise to the Zn2+ ion and release electrons than the copper half-cell

  2. More electrons will therefore build up on the zinc electrode than the copper electrode

  3. A potential difference is created between the two electrodes

  4. The zinc strip is the negative terminal and the copper strip is the positive terminal

  5. This potential difference is measured with a high resistance voltmeter, and is given the symbol E
• Why use a high resistance voltmeter?

  • The voltmeter needs to be of very high resistance to stop the current from flowing in the circuit
  • In this state it is possible to measure the maximum possible potential difference (E)
  • The reactions will not be occurring because the very high resistance voltmeter stops the current from flowing
• What happens if current is allowed to flow?

  1. If the voltmeter is removed and replaced with a bulb or if the circuit is short circuited, a current flows

  2. The reactions will then occur separately at each electrode

  3. The voltage will fall to zero as the reactants are used up

  4. The most positive electrode will always undergo reduction

  5. The most negative electrode will always undergo oxidation
• Cell Diagrams


  • Electrochemical cells can be represented by a cell diagram
  • The solid vertical line represents the boundary between phases e.g. solid (electrode) and solution (electrolyte)
  • The double line represents the salt bridge between the two half cells
  • The voltage produced is indicated
  • The more positive half cell is written on the right if possible (but this is not essential)

  
• Systems that do not include metals

  • If a system does not include a metal that can act as an electrode, then a platinum electrode must be used and included in the cell diagram
  • A platinum electrode is used because it is unreactive and can conduct electricity
• It is not possible to measure the absolute potential of a half electrode on its own. It is only possible to measure the potential difference between two electrodes
• The Standard Hydrogen Electrode


  • The potential of all electrodes are measured by comparing their potential to that of the standard hydrogen electrode
  • The standard hydrogen electrode (SHE) is assigned the potential of 0 volts
  • The hydrogen electrode equilibrium is: H2 (g) 2H+ (aq) + 2e-
  • To make the electrode a standard reference electrode some conditions apply: 1) Hydrogen gas at pressure of 100kPa 2) Solution containing the hydrogen ion at 1.0 mol dm-3 3) Temperature at 298K 4) Platinum electrode
• Secondary standards


  • The standard hydrogen electrode is difficult to use, so often a different standard is used which is easier to use
  • These other standards are themselves calibrated against the SHE
  • The common ones are: silver / silver chloride (E = +0.22 V) and calomel electrode (E = +0.27 V)
• Standard Electrode Potentials


  • The standard conditions are: all ion solutions at 1 mol dm-3, temperature 298 K, gases at 100 kPa pressure, no current flowing
  • When an electrode system is connected to the hydrogen electrode system, and standard conditions apply the potential difference measured is called the standard electrode potential, E
  • Standard electrode potentials are found in data books and are quoted as: Li+(aq) | Li (s) E= -3.03V or Li+ (aq) + e- Li (s) E= -3.03V
• Note: in the electrode system containing two solutions it is necessary to use a platinum electrode and both ion solutions must be of a 1 mol dm-3 concentration, so [Fe2+] = 1 mol dm-3 and [Fe3+] = 1 mol dm-3
• Calculating the EMF of a cell

  1. Ecell = Erhs - Elhs

  2. The more negative half cell will always oxidise (go backwards)

  3. The more positive half cell will always reduce (go forwards)

  4. A spontaneous change will always have a positive Ecell
• Using series of standard electrode potentials

  • As more +ve, increasing tendency for species on left to reduce, and act as oxidising agents
  • As more -ve, increasing tendency for species on right to oxidise, and act as reducing agents
  • The most powerful reducing agents will be found at the most negative end of the series
  • The most powerful oxidising agents will be found at the most positive end of the series
• O2(g) + 4H+(aq) + 4e– → 2H2O(I) Eo+1.23V


  • F2(g) + 2e– → 2F–(aq) Eo +2.87V
• Cl2(aq) + 2e– → 2Cl–(aq) Eo+1.36V

  • 2HOCl(aq) + 2H+(aq) + 2e– → Cl2(aq) + 2H2O(I) Eo+1.64V
• H2O2(aq) + 2H+(aq) + 2e– → 2H2O(I) Eo +1.77V

  • O2(g) + 4H+(aq) + 4e– → 2H2O(I) Eo +1.23V
• Oxidation number


  The charge on an atom in a compound
• Calculating Ecell from standard electrode potentials

  Ecell = Ered - Eox
• O2(g) + 4H+(aq) + 4e– → 2H2O(I) Eo+1.23V
• F2(g) + 2e– → 2F–(aq) Eo +2.87V
• Fluorine reacts with water


  Fluorine has a higher oxidation number than oxygen
• Cl2(aq) + 2e– → 2Cl–(aq) Eo+1.36V
• 2HOCl(aq) + 2H+(aq) + 2e– → Cl2(aq) + 2H2O(I) Eo+1.64V
• H2O2(aq) + 2H+(aq) + 2e– → 2H2O(I) Eo +1.77V
• Chlorine should undergo a redox reaction with water


  Chlorine has a higher oxidation number than oxygen
• O2(g) + 2H+(aq) + 2e– → H2O2(aq) Eo +0.68V
• Zn2+(aq) + 2e- Zn(s) E= - 0.76V
• Fe2+(aq) + 2e- Fe(s) E= -0.44V
• Zn + Fe2+ Fe + Zn2+ E= +0.32
• Electrochemical cells can be used as a commercial source of electrical energy
• Cells can be non-rechargeable (irreversible), rechargeable and fuel cells
• Ecell =+1.51V
• Ecell= +2.04V
• Li + MnO2 → LiMnO2
• Ecell =+2.91V
• Ecell = +1.40V
• Hydrogen can be stored in fuel cells as a liquid under pressure, adsorbed on the surface of a solid material, or absorbed within a solid material
• Limitations of hydrogen fuel cells

  • Expensive
  • Storing and transporting hydrogen is difficult
  • Limited lifetime and high production costs
  • Use of toxic chemicals in production