Electrode potentials and cells (A-level only) — AQA A-Level Chemistry
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Electrode potentials and cells (A-level only) explained
By IUPAC convention every electrode half-equation is written as a reduction, with the electrons on the left-hand side: Ox + ne⁻ ⇌ Red.
Read the full explanation
The zinc half-cell is therefore written Zn²⁺(aq) + 2e⁻ ⇌ Zn(s), E⦵ = −0.76 V, never as the oxidation. The tabulated E⦵ belongs to that direction; reverse the equation and the sign of E⦵ reverses with it. A half-equation must balance for atoms and for total charge. In acidic conditions oxygen is balanced with H₂O and hydrogen with H⁺, then electrons are added until the charge on each side is equal: MnO₄⁻ + 8H⁺ + 5e⁻ ⇌ Mn²⁺ + 4H₂O, where both sides carry a charge of +2. Writing every half-equation the same way round is what makes the electrochemical series comparable.
The conventional representation of cells.
A cell diagram is a shorthand for a whole cell. A single vertical line marks a boundary between two phases, a double vertical line marks the salt bridge, and a comma separates two species that share the same solution. The half-cell with the more negative electrode potential is written on the left, and within each half-cell the species are ordered so that the most reduced form sits on the outside: Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s). Where neither species is a solid metal that can conduct, an inert platinum electrode is written in, as in Cu(s) | Cu²⁺(aq) ‖ Fe³⁺(aq), Fe²⁺(aq) | Pt. Read this way, EMF = E⦵(right) − E⦵(left), and oxidation is always happening at the left-hand electrode.
Cells are used to measure electrode potentials by reference to the standard hydrogen electrode.
An electrode potential cannot be measured alone, because a measurement needs a complete circuit and therefore a second half-cell. The standard hydrogen electrode is the agreed reference and is defined as exactly 0.00 V. It is platinum foil coated in platinum black, dipping into 1.00 mol dm⁻³ H⁺(aq) with hydrogen at 100 kPa bubbled over it at 298 K. Connect the half-cell under test to it through a salt bridge and the voltmeter reading is that half-cell's electrode potential. The salt bridge, usually saturated KNO₃, completes the circuit by letting ions move between the two solutions and keeps each electrically neutral; it does not carry electrons. A high-resistance voltmeter is used so that almost no current flows and neither equilibrium is disturbed.
The importance of the conditions when measuring the electrode potential, E (Nernst equation not required).
An electrode potential is the position of an equilibrium, so anything that shifts that equilibrium changes E. Standard values are quoted at 298 K with every solution at 1.00 mol dm⁻³ and every gas at 100 kPa. Lowering [Cu²⁺] in Cu²⁺(aq) + 2e⁻ ⇌ Cu(s) pulls the equilibrium to the left, so the electrode accepts electrons less readily and E becomes less positive. Changing the temperature or the pressure of a gaseous reactant moves the equilibrium too. What does not change E is the surface area of the electrode: a larger platinum surface lets the equilibrium be reached faster but does not alter where that equilibrium lies. Rate and position are separate ideas, and only the position sets the potential.
Standard electrode potentials can be listed as an electrochemical series.
List every half-equation as a reduction, order them by E⦵, and you have the electrochemical series. The most positive values sit at one end: the species on the left of those equations, such as F₂ at +2.87 V and MnO₄⁻ in acid at +1.51 V, are the strongest oxidising agents, because they take electrons most readily. The most negative values sit at the other end, and the species on the right of those, such as lithium and potassium metal, are the strongest reducing agents. The rule that follows is simple: an oxidising agent will oxidise the reduced species of any half-equation with a less positive E⦵. That is the basis of every feasibility prediction. The series tells you what is energetically possible, not what is fast.
Students should be able to: Ɵ use E values to predict the direction of simple redox reactions calculate the EMF of a cell write and apply the conventional representation of a cell.
Three skills sit together here. To predict direction, compare E⦵ values: the more positive half-equation runs forwards as a reduction and the more negative one is forced backwards as an oxidation. To calculate EMF, subtract: EMF = E⦵(more positive) − E⦵(more negative). For zinc and copper, EMF = +0.34 − (−0.76) = +1.10 V. To build the overall equation, scale the two half-equations so the electrons cancel and then add them. E⦵ itself is never multiplied by that scaling factor. For conventional cell representation, place the more negative half-cell on the left (oxidation), use a single vertical line for phase boundaries, a double line for the salt bridge, and include a Pt electrode if no solid metal is present: Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s).
Required practical 8 Measuring the EMF of an electrochemical cell.
In Required practical 8 you build two half-cells and measure the potential difference between them. Each metal strip is cleaned with emery paper to remove the oxide layer, rinsed, and dipped into a 1.00 mol dm⁻³ solution of its own ions at 298 K. The two solutions are joined by a salt bridge, usually filter paper soaked in saturated potassium nitrate, which completes the circuit by letting ions move between the half-cells and keeps each solution electrically neutral. It does not let electrons through, and answers that say it does gain no credit. Potassium nitrate is chosen because neither K⁺ nor NO₃⁻ reacts with or precipitates the ions present, whereas potassium chloride would ruin a silver half-cell. A high-resistance voltmeter draws almost no current, so the reading is the EMF.
Your focus
- Write the Fe³⁺/Fe²⁺ half-equation as a reduction and show that the charge is +2 on each side.
- Balance the half-equation for MnO₄⁻ in acid by adding 8H⁺, 4H₂O and 5 electrons.
- Deduce what happens to the sign of E⦵ when a half-equation is written in the reverse direction.
Show all 21 objectives
- Write the cell diagram for a zinc half-cell joined to a copper half-cell.
- Justify the inclusion of a platinum electrode in the cell diagram for an Fe³⁺/Fe²⁺ half-cell.
- Calculate the EMF of a cell directly from its cell diagram using E⦵(right) − E⦵(left).
- Draw and label a standard hydrogen electrode, marking the platinum, the gas pressure and the acid concentration.
- Explain why potassium nitrate rather than potassium chloride is chosen for the salt bridge in a cell containing silver ions.
- Deduce which electrode is negative when a standard zinc half-cell is connected to a standard hydrogen electrode.
- Predict how E for Cu²⁺(aq) + 2e⁻ ⇌ Cu(s) changes when the copper ion concentration is halved.
- Identify, from a list of changes to a hydrogen electrode, the one that leaves the electrode potential unchanged.
- Explain why a potential measured with 0.10 mol dm⁻³ solutions may not be quoted as a standard value.
- Rank zinc, iron and copper as reducing agents from E⦵ values of −0.76 V, −0.44 V and +0.34 V.
- Deduce from a table of E⦵ values whether acidified manganate(VII) will oxidise chloride ions to chlorine.
- Explain why a redox reaction with a positive EMF may still not be observed in the laboratory.
- Calculate the EMF of the cell Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s) from E⦵ values of −0.76 V and +0.34 V.
- Combine two half-equations so that the electrons cancel and write the balanced overall redox equation.
- Write the conventional representation of a cell, including phase boundaries, salt bridge, and Pt electrodes where necessary.
- Describe how to set up a zinc and copper cell and measure its EMF with a high-resistance voltmeter.
- Justify the choice of potassium nitrate over potassium chloride for the salt bridge in a cell containing a silver half-cell.
- Suggest why a measured EMF differs from the value calculated from standard electrode potentials.
Electrode potentials and cells (A-level only) exam tips
Marking Points
- one mark for writing the half-equation as a reduction, with the electrons on the left-hand side
- one mark for balancing oxygen with H₂O and hydrogen with H⁺ in acidic conditions
- one mark for adding electrons until the total charge is equal on both sides
- one mark for using the equilibrium sign and correct state symbols
- one mark for placing the half-cell with the more negative electrode potential on the left
- one mark for a single vertical line at each phase boundary and a double line for the salt bridge
- one mark for including a platinum electrode where neither species is a solid metal conductor
- one mark for ordering the species so that the most reduced form is furthest from the salt bridge
- one mark for stating that the standard hydrogen electrode is assigned a potential of exactly 0.00 V
- one mark for the conditions 1.00 mol dm⁻³ H⁺(aq), 100 kPa H₂ and 298 K
- one mark for stating that the salt bridge allows ions to move so that the circuit is complete
- one mark for naming a suitable salt bridge electrolyte such as KNO₃ that does not react with the ions present
- one mark for using a high-resistance voltmeter so that negligible current flows
- one mark for stating the standard conditions of 298 K, 100 kPa and 1.00 mol dm⁻³ solutions
- one mark for explaining the effect of a concentration change using the position of the half-cell equilibrium
- one mark for stating that the surface area of the electrode has no effect on the electrode potential
- one mark for deducing whether E becomes more or less positive from the direction the equilibrium shifts
- one mark for identifying the species on the left of the most positive half-equation as the strongest oxidising agent
- one mark for identifying the species on the right of the most negative half-equation as the strongest reducing agent
- one mark for stating that the reaction is feasible when the oxidising agent comes from the more positive half-equation
- one mark for using a positive overall EMF as the test of feasibility
- one mark for identifying which half-equation runs in the reverse, oxidation, direction
- one mark for subtracting the less positive E⦵ from the more positive E⦵
- one mark for scaling the half-equations so the electrons cancel before adding them
- one mark for correctly applying conventional cell representation with phase boundaries and salt bridge
- one mark for stating that the salt bridge completes the circuit by allowing ions to move between the half-cells
- one mark for naming a soluble salt such as potassium nitrate that does not react with the ions in either half-cell
- one mark for cleaning the metal electrodes with emery paper before use
- one mark for using a high-resistance voltmeter so that negligible current is drawn
- one mark for using 1.00 mol dm⁻³ solutions at 298 K so the value may be quoted as standard
Examiner Tips
- 💡Put the electrons on the left every time, then count the total charge on each side as a check before moving on.
- 💡Balance in the order: atoms other than H and O, then O with H₂O, then H with H⁺, then charge with electrons.
- 💡Build the diagram outwards from the salt bridge: oxidised species next to the double line on both sides.
- 💡Check your diagram by calculating E⦵(right) − E⦵(left); a positive answer means you have the halves the right way round.
- 💡Learn the three standard conditions as a set; a question asking for the conditions usually wants all three.
- 💡Say 'allows ions to move' or 'completes the circuit' for the salt bridge, never 'allows electrons to pass'.
- 💡Treat every half-equation as an equilibrium and apply Le Chatelier's principle to it; that answers almost every 'effect of' question here.
- 💡Watch for distractors about electrode size, shape or surface area; these change rate, not potential.
- 💡Mark which half-equation runs backwards before you do any arithmetic; that single decision fixes the sign of the EMF.
- 💡If a predicted reaction is not observed, the mark is usually for saying the activation energy is high, so the rate is very slow.
- 💡Write both half-equations as reductions first, then decide which one to reverse; do not reverse anything on the page until you have compared the values.
- 💡When writing cell diagrams, build outwards from the salt bridge: oxidised species next to the double line.
- 💡Learn one complete sentence for the salt bridge: it allows ions to move so the circuit is complete and each solution stays electrically neutral.
- 💡If a measured EMF is lower than the calculated value, non-standard concentration or temperature is the expected explanation.
Common Mistakes
- writing Zn(s) → Zn²⁺(aq) + 2e⁻ and still quoting E⦵ as −0.76 V for that direction; the sign must be reversed
- reversing a half-equation, changing the sign of E⦵, and then subtracting as well, so the sign is changed twice
- balancing the atoms but leaving the charges unequal, for example MnO₄⁻ + 8H⁺ + 4e⁻ → Mn²⁺ + 4H₂O; check total charge on both sides
- adding OH⁻ to balance oxygen when the question states acidic conditions; use H₂O and H⁺ instead
- writing the more positive half-cell on the left and then reporting a negative EMF; the more negative half-cell goes on the left
- using a single vertical line where the salt bridge belongs; use a double vertical line
- omitting Pt from the Fe³⁺/Fe²⁺ or H⁺/H₂ half-cell, so the diagram has no conductor; always include an inert electrode if no solid metal is present
- separating two species in the same solution with a vertical line instead of a comma; use a comma for species in the same phase
- writing that the salt bridge allows electrons to flow, which examiners specifically refuse
- choosing potassium chloride for the salt bridge in a cell containing silver or lead ions, where a chloride precipitates
- giving the hydrogen ion concentration but leaving out the gas pressure or the temperature
- leaving the platinum out of contact with either the hydrogen or the acid in a labelled diagram
- claiming that increasing the surface area of the platinum electrode makes the potential more positive, when it only changes how fast equilibrium is reached
- saying temperature has no effect because E⦵ is a constant
- treating E as a fixed property of the half-cell whatever the concentrations are
- quoting 273 K or 100 °C as the standard temperature instead of 298 K
- ignoring the sign, so −0.76 V is treated as larger than −0.44 V
- reading the most negative value as belonging to the strongest oxidising agent
- assuming a positive EMF guarantees an observable reaction, when the rate may be immeasurably slow
- applying standard values to a mixture that is nowhere near 1.00 mol dm⁻³ and still calling the prediction certain
- adding the two E⦵ values instead of subtracting them to find the EMF
- multiplying E⦵ by the factor used to balance the electrons; E⦵ is independent of stoichiometry
- omitting the Pt electrode in a cell diagram when neither species is a solid metal; a conductor is always required
- writing that the salt bridge lets electrons flow between the half-cells, which examiners do not credit
- choosing potassium chloride when the cell contains a silver half-cell, because silver chloride precipitates
- replacing the salt bridge with a metal wire
- leaving an oxide layer on the metal strip, so the reading drifts and reads low
- describing the salt bridge as 'maintaining electrical neutrality' only, with no mention of ion movement