Chemistry: Electrode Potentials and Transition Metals
Electrode potentials and transition metals is the A-level Chemistry topic covering two linked areas of Year 13 inorganic chemistry. The electrode potentials strand extends standard cell EMF calculations to predicting the feasibility of redox reactions from electrode potential data, explaining the effect of concentration changes on an electrode potential using Le Chatelier's principle, and evaluating commercial cells (rechargeable cells and fuel cells). The transition metals strand covers the properties that define a transition metal (a partially filled d-subshell in at least one of its ions), variable oxidation states, ligand substitution reactions and the resulting colour changes, and the role of transition metal ions and their compounds as catalysts.
Before you start
Make sure you're comfortable with these topics first:
Method
- To predict whether a redox reaction is feasible from standard electrode potentials, identify the two half-equations, calculate E cell = E(more positive electrode, as written for reduction) - E(more negative electrode, as written for reduction); a positive E cell means the reaction as written is energetically feasible, though it may still be slow, or not observed, for kinetic reasons.
- Apply Le Chatelier's principle to predict the effect of a concentration change on an electrode potential: increasing the concentration of the oxidised species in a half-equation (the species on the side with the electrons) shifts the half-equation's equilibrium to produce more electrons, making that electrode potential more positive (or less negative); increasing the concentration of the reduced species has the opposite effect.
- When evaluating a rechargeable cell or a fuel cell, weigh a specific named advantage (such as no toxic combustion products for a hydrogen fuel cell, or no emissions at the point of use) against a specific named disadvantage (such as the difficulty and cost of producing, storing and transporting hydrogen safely, or the energy cost and emissions from generating the hydrogen in the first place).
- Identify a transition metal using the definition precisely: an element that forms at least one stable ion with a partially filled d-subshell; check specific ions against this definition rather than assuming every d-block element or ion qualifies, since Sc3+ (empty 3d subshell) and Cu+ (full 3d10 subshell) are both d-block ions that do not count as transition metal ions by this definition.
- For a ligand substitution reaction, identify the ligands being exchanged and any change in coordination number (for example, an octahedral 6-coordinate aqua complex exchanging for a tetrahedral 4-coordinate chloro complex), and link the colour change observed to the change in ligand and geometry around the central metal ion.
- Explain the colour of a transition metal complex ion in terms of the partially filled d-subshell: the ligands split the d-orbitals into two energy levels, and an electron can absorb a specific frequency of visible light to jump between them; the colour seen is the complementary colour to the light absorbed, since the remaining wavelengths are transmitted or reflected.
- For catalysis questions, name a specific transition metal catalyst and its specific reaction (heterogeneous: iron in the Haber process, vanadium(V) oxide in the Contact process; homogeneous: Fe2+/Fe3+ or Mn2+ catalysing specific redox reactions by an intermediate-oxidation-state or autocatalysis mechanism), rather than answering only in general terms.
- For qualitative tests, add NaOH solution dropwise then in excess to distinguish transition metal ions by the colour of the hydroxide precipitate formed, and note which precipitates redissolve in excess ammonia solution to form a soluble complex ion.
Worked example
The half-equation Fe3+(aq) + e- <=> Fe2+(aq) has a standard electrode potential of E = +0.77 V, measured using standard (1 mol/dm3) concentrations of both Fe3+(aq) and Fe2+(aq). Predict and explain the effect on this electrode potential of increasing the concentration of Fe3+(aq) while keeping the concentration of Fe2+(aq) constant.
- Identify the equilibrium position of the half-equation as written: Fe3+(aq) + e- <=> Fe2+(aq).
- Increasing the concentration of Fe3+(aq), a reactant in the forward (reduction) direction as written, shifts the position of this equilibrium to the right (towards Fe2+), by Le Chatelier's principle, since the system responds to oppose the increase by using up some of the added Fe3+.
- Shifting the equilibrium to the right means the half-cell has a greater tendency to proceed in the reduction (electron-gaining) direction as written.
- A greater tendency towards the reduction reaction corresponds to a more positive electrode potential, since electrode potential is a direct measure of how readily a half-equation proceeds in the reduction direction as written.
- Final answer: the electrode potential becomes more positive than +0.77 V. Increasing [Fe3+] shifts the equilibrium towards Fe2+ (more reduction), which by definition corresponds to a more positive (more oxidising) electrode potential for this half-cell.
Practice questions
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Q1State the standard electrode potential data needed to calculate the standard EMF of a cell made from two half-cells.Show answer
Answer: The standard electrode potential (E value) of each of the two half-cells (half-equations) that make up the cell.
Q2Standard electrode potentials: E(Cu2+/Cu) = +0.34 V; E(Zn2+/Zn) = -0.76 V. Calculate the standard EMF of a cell made from these two half-cells.Show answer
Answer: E cell = E(positive, Cu2+/Cu) - E(negative, Zn2+/Zn) = +0.34 - (-0.76) = +1.10 V.
Q3State which electrode, copper or zinc, is the negative terminal in the cell described above, and explain why.Show answer
Answer: Zinc; it has the more negative electrode potential, so it is oxidised more readily, releasing electrons into the external circuit (it is the electrode at which oxidation occurs).
Q4State the reason a d-block ion such as Cu+ does not count as a transition metal ion.Show answer
Answer: A transition metal ion requires a partially, not completely, filled d-subshell; Cu+ has an electron configuration ending in 3d10, a full subshell, so it is not classed as a transition metal ion.
Q5Write an equation for the ligand substitution reaction that occurs when excess concentrated hydrochloric acid is added to a solution containing the hexaaquacopper(II) ion, [Cu(H2O)6]2+, and state the colour change observed.Show answer
Answer: [Cu(H2O)6]2+ + 4Cl- <=> [CuCl4]2- + 6H2O; the colour changes from pale blue to yellow (or yellow-green).
Q6Name a transition metal used as a heterogeneous catalyst in the Haber process, and state the gases it catalyses the reaction between.Show answer
Answer: Iron; it catalyses the reaction between nitrogen and hydrogen to form ammonia.
Q7A student adds NaOH solution dropwise to a solution suspected to contain Fe2+ ions. Describe the observation that would confirm this, including any colour change on standing in air.Show answer
Answer: A green precipitate (of iron(II) hydroxide) forms; on standing in air, this precipitate gradually turns from green to red-brown, as the Fe2+ is oxidised to Fe3+ by oxygen in the air.
Q8State one advantage and one disadvantage of a hydrogen-oxygen fuel cell compared with a petrol/diesel internal combustion engine, as a means of powering a vehicle.Show answer
Answer: Advantage: a fuel cell produces only water as a product at the point of use, with no carbon dioxide or other combustion pollutants released locally. Disadvantage: hydrogen gas is difficult and costly to produce, store and transport safely (it requires high pressure or very low temperature storage), and producing the hydrogen may itself release carbon dioxide unless a low-carbon method is used.
Exam-style questions
Written in the style of a A Level Science exam paper, with a full mark scheme.
Standard electrode potentials: E(MnO4-/Mn2+, acidic) = +1.51 V; E(Cl2/Cl-) = +1.36 V. (a) Predict, with a calculation, whether acidified potassium manganate(VII) is able to oxidise chloride ions to chlorine under standard conditions. (b) Write the overall ionic equation for this reaction, given the half-equations: MnO4- + 8H+ + 5e- -> Mn2+ + 4H2O, and Cl2 + 2e- -> 2Cl-.
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When aqueous ammonia is added dropwise to a solution of [Cu(H2O)6]2+, a pale blue precipitate initially forms. On addition of excess concentrated ammonia, this precipitate dissolves to give a deep blue solution. (a) Identify the pale blue precipitate formed in the first step, and write an equation for its formation. (b) Write an equation for the formation of the deep blue solution in excess ammonia, naming the complex ion formed and stating its shape.
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The reaction between acidified potassium manganate(VII) and ethanedioate ions (C2O4^2-) is observed to start slowly, then speed up part-way through, before slowing again as the reactants are used up. (a) Explain, in terms of catalysis, why the reaction speeds up part-way through, even though no catalyst was added at the start. (b) State the specific term used to describe this type of catalysis.
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