Measuring Standard Electrode Potential
- There are three different types of half-cells that can be connected to a standard hydrogen electrode to measure standard electrode potential
- 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
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) Eꝋ = + 0.80 V
2H+ (aq) + 2e- ⇌ H2 (g) Eꝋ = 0.00 V
- Since the Ag+/ Ag half-cell has a more positive Eꝋ value, this is the positive pole and the H+/H2 half-cell is the negative pole
- The standard cell potential (Ecellꝋ) is Ecellꝋ = (+ 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 Eꝋ value
- Reduction occurs at the positive electrode
- Oxidation occurs at the negative electrode
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
- Br2 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 (aq) + 2e- ⇌ 2Br- (aq) Eꝋ = +1.09 V
2H+ (aq) + 2e- ⇌ H2 (g) Eꝋ = 0.00 V
- The Br2 / Br- half-cell is the positive pole and the H+ / H2 is the negative pole
- The Ecellꝋ is: Ecellꝋ = (+ 1.09) - (0.00) = + 1.09 V
- The Br2 molecules are more likely to get reduced than H+ as they have a greater Eꝋ value
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) Eꝋ = +1.52 V
2H+ (aq) + 2e- ⇌ H2 (g) Eꝋ = 0.00 V
- The H+ ions are also present in the half-cell as they are required to convert MnO4- into Mn2+ ions
- The MnO4- / Mn2+ half-cell is the positive pole and the H+ / H2 is the negative pole
- The Ecellꝋ is Ecellꝋ = (+ 1.52) - (0.00) = + 1.52 V
Ions in solution half cell
The Salt Bridge
- A salt bridge has mobile ions that complete the circuit
- Ions must be able to flow between the half-cells or solutions
- This should be made on metal wire, even if the metal is inert
- Metal wire allows the flow of electrons but not the flow of ions
- Potassium chloride and potassium nitrate are commonly used to make the salt bridge as chlorides and nitrates are usually soluble
- This should ensure that no precipitates form which can affect the equilibrium position of the half cells