In response to the part of your question regarding half-cells:
Yes, you can continue to use these solutions, although we suggest using iron (II) sulfate, rather than iron (III) nitrate, in an Fe/Fe2+ half-cell (or using a mixture of iron (II) sulfate and iron (III) sulfate, if you are preparing an Fe2+/Fe3+ half-cell).
Concentrations of 0.5 M achieve a better voltmeter reading than 0.1 M solutions. Aqueous solutions of nitrates are usually used in half-cells as the nitrates ions will not interfere with the electrochemical reaction by being oxidised or reduced and will not form any precipitates.
A half-cell consists of an electrode partially immersed in an aqueous solution. A potential difference is developed between the positively charged solution and the negatively charged electrode when connected to a standard hydrogen electrode. A number of different half-cells can be used.
- Copper metal in an aqueous solution of copper (II) nitrate
- Zinc metal in an aqueous solution of zinc (II) nitrate
- Magnesium metal in an aqueous solution of magnesium nitrate
- Aluminium metal in an aqueous solution of aluminium nitrate
Note: Aqueous solutions of the metal sulfates can also be used for the copper, zinc, magnesium and aluminium half-cell.
- Iron metal in an aqueous solution of iron (II) sulfate
- Lead metal in an aqueous solution of lead (II) nitrate
Note: Before using lead or its compounds, a risk assessment should be prepared taking into consideration its toxicity, level of student compliance, quantities used and clean up procedures.
Science ASSIST recommendations
- Consult SDSs before using any chemicals.
- Safety glasses and lab coats should be worn throughout the laboratory activity.
- General laboratory hygiene, such as washing hands before leaving the laboratory, should be observed.
- All used solutions should be collected and disposed of in a heavy metal wastes container.
- Metal strips can be cleaned and reused.
An electrochemical cell comprises of two half-cells joined by a salt bridge. In one half-cell, oxidation of a metal electrode occurs, while in the other half-cell there is the reduction of metal ions in solution. The half-cell with the most negative electrode potential forms the negative terminal (anode). Oxidation occurs at the anode. Reduction occurs at the positive terminal (cathode). The salt bridge usually contains a saturated solution of sodium or potassium nitrate or chloride. The salt bridge allows the flow of ions from one half-cell to another without the mixing of the two solutions. The salt bridge also maintains electrical neutrality of the solutions in the two half–cells and completes the electrical circuit.
Consider the electrochemical cell consisting of the 2 half-cells expressed in the following half equations:
Cu2+(aq) + 2e- Cu(s) E0 = +0.34 V (strip of copper partially immersed in aqueous solution of copper (II) nitrate)
Zn2+(aq) + 2e- Zn(s) E0= -0.76 V (strip of zinc partially immersed in aqueous solution of zinc (II) nitrate)
Copper (II) ions have a higher standard electrode potential value and are a stronger oxidant than zinc ions. At the zinc anode, Zn2+ ions are released into the solution of the half-cell while at the copper cathode, Cu2+ ions are reduced to metallic copper.
The overall equation is:
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
C Lewis and P Lewis, Chemistry for WA2 Stage 3: Units 3A and 3B (Pearson, 2013) ch9.