School Science Lessons
2018-07-27
Please send comments to: J.Elfick@uq.edu.au

3.86.0 Electrode potential
Table of contents
Experiments
3.86.1 Electrode potential order of metals

3.84.6 Standard electrode potential of metals, reduction potential, E0

15.7.1 Potential difference from combining half cells, zinc and iron

15.7.2 Potential difference from combining half cells, zinc and copper, zinc and lead

15.7.3 Differences in potential on an iron nail

15.7.1 Potential difference from combining half cells, zinc and iron
To measure the potential difference of a zinc half cell connected to an iron half cell.
Use a strip of zinc metal in a zinc chloride solution and an iron nail in iron (II) sulfate solution.
Connect the two half cells with a strip of filter paper soaked in potassium chloride solution to act as a salt bridge.
Complete the circuit by connecting leads from each metal to a voltmeter.
Read the voltmeter.
Electrons flow with potential difference of 0.32 V.
Zn (s) --> Zn2+ (aq) + 2e- (E0= + 0.76 V)
Fe2+ + 2e- --> Fe (aq) (E0 = -0.44 V)
Zn (s) + Fe2+ --> Zn2+ + Fe (s) (E0= + 0.32 V)

15.7.2 Potential difference from combining half cells, Zn and Cu, Zn and Pb
If Zn E0 = -0.76 V set up cells to measure the E0 values of copper (copper in copper (II) sulfate solution) and lead (lead in lead (II)
nitrate solution).

15.7.3 Differences in potential on an iron nail
Soak 1 gm agar in 100 mL water for two hours then boil until dissolved.
Add phenolphthalein indicator and add acid or alkali until pH = 8.
Add drops of freshly prepared potassium ferricyanide solution and pour into a Petri dish.
Add a very clean nail and place the petri dish on an overhead projector.
After some hours, a pink colour forms around the shaft of the nail because of hydroxide ions and blue-green colour forms around the
head of the nail because of Fe2+ ions.
The stressed head shows positive potential and the unstressed shaft shows negative potential.
At the anode: Fe (s) --> 6 Fe2+ (aq) + 2e-
At the cathode: O2 (aq) + 2H2O (l) + 4e-- --> 4OH- (aq)

3.86.1 Electrode potential order of metals
See diagram 3.86: Electrode potential apparatus
A metal tested, B copper sulfate in filter paper, C copper foil
1. Electrode potentials of metals are calculated from comparisons with the hydrogen cell under standardized conditions.
However, you can use a copper and copper (II) sulfate solution as a standard.
Lay filter paper soaked with copper (II) sulfate solution on clean copper foil.
Use a short length of wire and crocodile clips to connect the copper foil to the positive terminal of the 1 to 5 V voltmeter.
Similarly connect the specimen metal to the negative terminal of the voltmeter.
Clean the surface of the specimen metal and press it firmly on absorbent paper.
Record the voltage for this metal.

2. Test the following metals: magnesium, tin, lead, iron, zinc aluminium, and silver.
After testing a metal, clean the copper again with a fine emery cloth and replace the absorbent paper, then test another metal.
The metal surfaces must be clean and the absorbent paper must contain enough copper (II) sulfate solution for a steady reading on the
voltmeter.
If the voltage starts at a high value and then falls as a deposit forms on the metal, record the highest value.

3. Test aluminium after dipping it in concentrated hydrochloric acid then press it on absorbent paper to remove the layer of aluminium
oxide.
The voltage reading will start at a low value then increase as remaining aluminium oxide dissolves.
Record the maximum value.

3.84.6 Standard electrode potential of metals, reduction potential, E0
See diagram 3.84.1.1: Standard electrode potential apparatus
See diagram 3.2.86: Electrode potentials
1. Electrode potentials of metals are calculated from comparisons with the hydrogen cell under standardized conditions.
However, you can use a copper and copper (II) sulfate solution as a standard.
Lay filter paper soaked with copper (II) sulfate solution on clean copper foil.
Use a short length of wire and crocodile clips to connect the copper foil to the positive terminal of the 1 to 5 V voltmeter.
Similarly connect the specimen metal to the negative terminal of the voltmeter.
Clean the surface of the specimen metal and press it firmly on absorbent paper.
Record the voltage for this metal.

2. Test the following metals: magnesium, tin, lead, iron, zinc aluminium, and silver.
After testing a metal, clean the copper again with a fine emery cloth and replace the absorbent paper, then test another metal.
The metal surfaces must be clean and the absorbent paper must contain enough copper (II) sulfate solution for a steady reading on the
voltmeter.
If the voltage starts at a high value and then falls as a deposit forms on the metal, record the highest value.

3. Test aluminium after dipping it in concentrated hydrochloric acid then press it on absorbent paper to remove the layer of aluminium
oxide.
The voltage reading will start at a low value then increase as remaining aluminium oxide dissolves.
Record the maximum value.

4. Values of electrode potentials of metals are derived from comparisons with the hydrogen cell  under standardized conditions of 1 M
solution at 25oC and 1 atmosphere (101.2 kPa) pressure.
The standard hydrogen cell is hydrogen gas from a platinum electrode in 1 M solution of H+.
If E0 value is +ve, then the preferred direction of electron flow is left to right.
The ion or atom with the greater value of E0 will attract electrons more easily.
A positive value for E0 means that particles in the half cell attract electrons more easily than particles in the hydrogen half cell.
If more than one reaction could occur, the reaction that does occur is the reaction that would form the greatest voltage.
Standard electrode potential of half reactions 1.0 molar solutions of some metals at 25oC,  is compared to the hydrogen half reaction
with electrode potential assumed to be zero.
A negative vale for E0 shows a poorer electron attracting ability than in the hydrogen half-cell.
Half reaction and E0 in volts, V
Zn2+ (aq) + 2e- ---> Zn (s) -0.76 V (Zn2+ / Zn)
2H+ (aq) + 2e- ---> H2 0.00 V <-- Hydrogen (H+ / H2 Standard Hydrogen Electrode, SHE)
Cu2+ (aq) + 2e- ---> Cu (s) +0.34 V (Cu2+ / Cu)
The E0 value = higher value - lower value (+0.34) - (-0.76) = 1.1 V.
The saturated calomel electrode, SCE (Hg2Cl2 / Hg) uses saturated KCl electrolyte
Hg2Cl2 (s) + 2e- --> 2Hg (l) + 2Cl-
E = 0.24 V
The following list show the standard electrode potentials of half reactions 1.0 molar solutions of some metals at 25oC, compared to the
hydrogen half reaction with electrode potential assumed to be zero.
A negative vale for E0 shows a poorer electron attracting ability than in the hydrogen half cell.
For a cell with the following half cell reactions:
Ag+ (aq) + e- --> Ag (s) +0.80 V, and Zn2+ (aq) + 2e- --> Zn (s) -0.76 V
the E0 value = (higher value - lower value)
= [(+0.80) - (-0.76) = 1.56 V].

Half reaction and E0 in volts, V
Table 15.7.0
Li+ (aq) + e- --> Li (s)
E0 = -3.04 V
K+ (aq) + e- --> K (s)  E0 = -2.93 V
Ba2+ (aq) + 2e- --> Ba (s) E0 = -2.91 V
Ca2+ (aq) + 2e- --> Ca (s)  E0 = -2.87 V
Na+ (aq) + e- --> Na (s) E0 = -2.71 V
Mg2+ (aq) + 2e- --> Mg (s) E0 = -2.37 V
Al3+ (aq) + 3e- --> Al (s) E0 = -1.66 V
Mn2+ + 2e- --> Mn E0 = -1.19 V
2H2O (l) + 2e- --> H2 (g) + 2OH- (aq)
E0 = -0.83 V
Zn2+ (aq) + 2e- --> Zn (s)  E0 = -0.76 V
Cr3+ (aq) + 3e- --> Cr (s) (aq) E0 = -0.74 V
Fe3+ (aq) + 3e- --> Fe (s) E0 = -0.44 V
Cd2+ (aq) + 2e- --> Cd (s)
E0 = -0.40 V
Ni2+ (aq) + 2e- --> Ni (s) E0 = -0.25 V
Sn2+ (aq) + 2e- --> Sn (s) E0 = -0.13 V
Pb2+ (aq) + 2e- --> Pb (s) E0 = -0.13 V
2H+ (aq) + 2e- --> H2 (g) E0 = 0.00 V
Cu2+ (aq) + e- --> Cu+ (aq) E0 = +0.16 V
Ag+ (aq) + e- --> Ag (s) E0 = +0.80 V
2Hg2+ (aq) + 2e- --> Hg22+ (aq) E0 = +0.80 V
Pt2+ + 2 e- --> Pt (s)
E0 = +1.19 V
Pb4+ + 2 e- --> Pb (s)
E0 = +1.69
Au+ +  e- --> Au (s)
E0 =  +1.83
Ag2+ +  e --> Ag+
E0 = + 1.98