AQA A Level Chemistry

Revision Notes

5.4.2 Standard Electrode Potentials

Standard Electrode Potentials

  • There are three different types of half-cells that can be connected to a standard hydrogen electrode
    • A metal / metal ion half-cell
    • A non-metal / non-metal ion half-cell
    • An ion / ion half-cell (the ions are in different oxidation states)

Metal/metal ion half-cell

Principles of Electrochemistry - Example of a Metal_Metal Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Example of a metal / metal ion half-cell connected to a standard hydrogen electrode

  • An example of a metal/metal ion half-cell is the Ag+/ Ag half-cell
    • Ag is the metal
    • Ag+ is the metal ion
  • This half-cell is connected to a standard hydrogen electrode and the two half-equations are:

Ag+ (aq) + e⇌ Ag (s)        E= + 0.80 V

2H+ (aq) + 2e⇌ H2 (g)        E= 0.00 V 

  • Since the Ag+/ Ag half-cell has a more positive Evalue, this is the positive pole and the H+/H2 half-cell is the negative pole
  • The standard cell potential (Ecell) is Ecell = (+ 0.80) – (0.00) = + 0.80 V
  • The Ag+ ions are more likely to get reduced than the H+ ions as it has a greater Evalue
    • Reduction occurs at the positive pole
    • Oxidation occurs at the negative pole

Non-metal/non-metal ion half-cell

  • In a non-metal/non-metal ion half-cell platinum wire or foil is used as an electrode to make electrical contact with the solution
    • Like graphite, platinum is inert and does not take part in the reaction
    • The redox equilibrium is established on the platinum surface
  • An example of a non-metal/non-metal ion is the Br2/Br half-cell
    • Br is the non-metal
    • Br is the non-metal ion
  • The half-cell is connected to a standard hydrogen electrode and the two half-equations are:

Br2 (l) + 2e⇌ 2Br (aq)        E = +1.09 V

2H+ (aq) + 2e⇌ H2 (g)        E = 0.00 V   

  • The Br2/Br half-cell is the positive pole and the H+/H2 is the negative pole
  • The Ecellis: Ecell = (+ 1.09) – (0.00) = + 1.09 V
  • The Br2 molecules are more likely to get reduced than H+ as they have a greater Evalue

Principles of Electrochemistry - Example of a Non-Metal_Non-Metal Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Example of a non-metal / non-metal ion half-cell connected to a standard hydrogen electrode

Ion/Ion half-cell

  • A platinum electrode is again used to form a half-cell of ions that are in different oxidation states
  • An example of such a half-cell is the MnO4/Mn2+ half-cell
    • MnO4 is an ion containing Mn with oxidation state +7
    • The Mn2+ ion contains Mn with oxidation state +2
  • This half-cell is connected to a standard hydrogen electrode and the two half-equations are:

MnO4 (aq) + 8H+ (aq) + 5e⇌ Mn2+ (aq) + 4H2O (l)       E = +1.52 V

2H+ (aq) + 2e⇌ H2 (g)       E= 0.00 V   

  • The H+ ions are also present in the half-cell as they are required to convert MnO4into Mn2+ ions
  • The MnO4/Mn2+ half-cell is the positive pole and the H+/H2 is the negative pole
  • The Ecell is Ecell = (+ 1.09) – (0.00) = + 1.09 V

Principles of Electrochemistry - Example of an Ion_ Ion Half-Cell, downloadable AS & A Level Chemistry revision notes

Ions in solution half cell

Standard Hydrogen Electrode

  • When a metal rod is placed in an aqueous solution, a redox equilibrium is established between the metal ions and atoms
    • For example, the copper atoms get oxidised and enter the solution as copper ions

Cu(s) → Cu2+(aq) + 2e

Principles of Electrochemistry - Oxidation of Copper, downloadable AS & A Level Chemistry revision notes

Oxidation of copper ions

    • The copper ions gain electrons from the metal rod and deposit as metal atoms on the rod

Cu2+(aq) + 2e→ Cu(s)

Principles of Electrochemistry - Reduction of Copper, downloadable AS & A Level Chemistry revision notes

Reduction of copper ions

  • When equilibrium is established, the rate of oxidation and reduction of copper is equal
  • The position of the redox equilibrium is different for different metals
    • Copper is more easily reduced, thus the equilibrium lies further over to the right

Cu2+ (aq) + 2e⇌ Cu (s)

    • Vanadium is more easily oxidised, thus the equilibrium lies further over to the left

V2+ (aq) + 2e⇌ V(s)

  • The metal atoms and ions in solution cause an electric potential (voltage)
  • This potential cannot be measured directly however the potential between the metal/metal ion system and another system can be measured
  • This value is called the electrode potential (E) and is measured in volts
    • The electrode potential is the voltage measured for a half-cell compared to another half-cell
    • Often, the half-cell used for comparison is the standard hydrogen electrode

Standard hydrogen electrode

  • The standard hydrogen electrode is a half-cell used as reference electrodes and consists of:
    • Hydrogen gas in equilibrium with H+ ions of concentration 1.00 mol dm-3 (at 1 atm)

2H+ (aq) + 2e ⇌ H2 (g)

    • An inert platinum electrode that is in contact with the hydrogen gas and H+ ions
  • When the standard hydrogen electrode is connected to another half-cell, the standard electrode potential of that half-cell can be read off a voltmeter

Principles of Electrochemistry - Standard Hydrogen Electrode, downloadable AS & A Level Chemistry revision notes

The standard electrode potential of a half-cell can be determined by connecting it to a standard hydrogen electrode

Calculating EMF

  • Once the standard electrode potentials (E) of the half-cells are determined, the standard cell potential (Ecell) can be calculated by subtracting the less positive Efrom the more positive Evalue
    • The half-cell with the more positive Evalue will be the positive pole
    • The half-cell with the less positive Evalue will be the negative pole

Worked Example

Calculating the standard cell potential

Calculate the standard cell potential for the electrochemical cell below and explain why the Cu2+/Cu half-cell is the positive pole.

The half-equations are as follows:

Cu2+(aq) + 2e⇌ Cu(s)      E= +0.34 V

Zn2+(aq) + 2e⇌ Zn(s)      E= −0.76 V

Electrochemistry Calculations - Electrochemical Cell, downloadable AS & A Level Chemistry revision notes

Answer

  • Step 1: Calculate the standard cell potential

Ecell = (+0.34) – (-0.76)

= +1.10 V

The voltmeter will therefore read off a value of 1.10 V

  • Step 2: Determine the positive and negative poles
    The Cu2+/Cu  half-cell is the positive pole as its Eis more positive than the Evalue of the Zn2+/Zn half-cell
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