CCEA A Level Chemistry 5.6 Electrode Potentials Exam Questions

About CCEA A Level Chemistry 5.6 Electrode Potentials

Section 5.6 of Unit A2 2 covers electrochemistry. You need to define standard electrode potential, describe the standard hydrogen electrode, and use electrode potentials to predict the direction and feasibility of redox reactions and calculate cell emf. You must also explain the limitations of these predictions, since non-standard concentrations and slow kinetics can mean a predicted reaction does not occur. CCEA requires conventional cell notation, the classification of cells as non-rechargeable, rechargeable or fuel cells, the electrode reactions in a lithium cell and in an alkaline hydrogen-oxygen fuel cell, and the environmental issues of cells. This builds on redox (1.7) and is used with vanadium chemistry in 5.5.

What you need to know

• Standard electrode potential The emf of a half-cell measured against a standard hydrogen electrode under standard conditions: 298 K, 100 kPa and 1.00 mol dm-3 solutions.
• Hydrogen electrode Hydrogen gas at 100 kPa bubbled over a platinum electrode in 1.00 mol dm-3 H+ at 298 K; its potential is defined as 0.00 V.
• Cell emf and notation Emf = E(right, more positive) - E(left, more negative). Cell notation: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s), with the salt bridge shown as a double line.
• Feasibility The species on the left of the half-equation with the more positive E is the stronger oxidising agent and is reduced. A reaction is predicted to be feasible if the cell emf is positive. It may still not happen if the activation energy is high or concentrations are not standard.
• Cells Non-rechargeable cells (for example zinc-carbon) are used once; rechargeable cells (for example lithium-ion) can be reversed by applying a current; fuel cells use a continuous supply of fuel. Lithium cell negative electrode: Li gives Li+ + e-.
• Fuel cells Alkaline hydrogen-oxygen fuel cell: negative electrode H2 + 2OH- gives 2H2O + 2e-; positive electrode O2 + 2H2O + 4e- gives 4OH-; overall 2H2 + O2 gives 2H2O. Issues: hydrogen production often uses fossil fuels, and storing hydrogen is difficult. Cells contain toxic or valuable metals and should be recycled.

Common mistakes in 5.6 Electrode Potentials

• Multiplying E values when balancing equations. Electrode potentials are not multiplied by the number of moles.
• Calculating emf as left minus right. Use the more positive value minus the less positive value.
• Saying a positive emf means the reaction will definitely happen. It may be too slow, or conditions may not be standard.
• Using a platinum electrode for metal/metal ion half-cells. Platinum is used when neither species is a solid metal, such as Fe3+/Fe2+.
• Writing the salt bridge as a single line. Use a double vertical line for the salt bridge and single lines for phase boundaries.

Try three 5.6 Electrode Potentials exam-style questions