Standard Electrode Potentials E⦵, Standard Cell Potentials E⦵cell & the Nernst Equation (Cambridge (CIE) A Level Chemistry): Flashcards

A salt bridge is a strip of filter paper soaked in a saturated solution of KNO₃ or KCl that connects two half-cells, completing the circuit by allowing mobile ions to flow between the half-cells.

False.

A platinum electrode is used instead. Platinum is inert and does not participate in the reaction. The redox equilibrium is established on the platinum surface.

Electrons flow through the external wire from the negative pole to the positive pole. The half-cell with the less positive E^ɔ is the negative pole (oxidation occurs there).

A reaction is feasible when the standard cell potential E_cell^ɔ is positive, and not feasible when it is negative.

False.

Oxidation occurs at the negative pole (anode). Reduction occurs at the positive pole (cathode).

Nitrates and chlorides are generally soluble, so using these salts ensures that no precipitates form in the half-cells. Precipitates would disturb the equilibrium position of the half-cells and give inaccurate results.

A redox equation is a balanced chemical equation for a reaction involving both oxidation and reduction, constructed by combining and balancing the relevant half-equations so that the number of electrons cancels.

In a MnO₄^- / Mn²⁺ ion-ion half-cell connected to a standard hydrogen electrode, MnO₄^- / Mn²⁺ acts as the positive pole because its E^ɔ value of +1.52 V is more positive.

E_cell^ɔ = E_right^ɔ – E_left^ɔ = (+0.34) – (−0.76) = +1.10 V

Cu²⁺/Cu is the positive pole as it has the more positive E^ɔ value.

An electrochemical series is a list of redox equilibria arranged in order of decreasing standard electrode potential (E^ɔ), used to compare the relative strengths of oxidising and reducing agents.

A less positive E^ɔ means the species is more easily oxidised (loses electrons more readily), making it a stronger reducing agent.

In a conventional cell diagram, the half-cell with the more positive E^ɔ value is placed on the right and is the positive pole.

E_cell^ɔ = (+1.36) – (+0.34) = +1.02 V

The cell potential is positive, so the forward reaction is feasible (spontaneous). A negative E_cell^ɔ would indicate the reaction is not feasible.

False.

A species at the top of the series has the most positive E^ɔ and is the strongest oxidising agent (most easily reduced). The strongest reducing agent is at the bottom of the series.

Electrons flow from the negative pole to the positive pole through the external wire. The negative pole has the less positive E^ɔ value, where oxidation occurs.

A feasibility in electrochemistry is a measure of whether a reaction is likely to occur spontaneously. A reaction is feasible when the standard cell potential E_cell^ɔ is positive.

The standard cell potential formula is: E_cell^ɔ = E_reduction^ɔ – E_oxidation^ɔ

A non-standard electrode potential (E) is the electrode potential measured under conditions that differ from standard conditions (1.00 mol dm^-3, 298 K, 1 atm), symbolised as E rather than E^ɔ.

Increasing the concentration of the species on the left shifts the equilibrium to the right, making that species more easily reduced. The electrode potential E becomes more positive (less negative).

False.

Increasing the concentration of the species on the right shifts equilibrium to the left, making reduction less favourable. The electrode potential becomes less positive (more negative).

For the half-cell Zn²⁺ (aq) + 2e^- ⇌ Zn (s), if [Zn²⁺] is increased, the equilibrium shifts to the right and the E value becomes less negative.

Temperature alters the position of the redox equilibrium in the half-cell. Changing temperature changes the extent of the forward and reverse reactions, so the ratio of oxidised to reduced species changes and the electrode potential E shifts from E^ɔ.

False.

Under non-standard conditions, the symbol E (without the degree sign) is used to distinguish the measured electrode potential from the standard electrode potential E^ɔ.

Equilibrium shifts to the left (more Fe³⁺ forms). The species on the left is less easily reduced, so the electrode potential E becomes less positive than +0.77 V.

A redox equilibrium is a reversible reaction in which the forward and reverse processes involve transfer of electrons, establishing a dynamic equilibrium between oxidised and reduced species.

Increasing the concentration of products (right-hand species) in a half-cell shifts equilibrium to the left, making the species on the left less easily reduced and the E value more negative.

A Nernst equation is an equation that relates the electrode potential under non-standard conditions (E) to the standard electrode potential (E^θ) by accounting for the ratio of oxidised to reduced species concentrations, temperature and the number of electrons transferred.

E = E^θ + (0.059 / z) × log₁₀ ([oxidised species] / [reduced species])

where z is the number of electrons transferred.

True.

Solids do not appear in equilibrium expressions. Their concentrations are fixed at 1.0 mol dm^-3, so they do not affect the Nernst equation calculation.

In the Nernst equation at 298 K, the term 0.059 / z arises because the constants R, T and F are combined and ln is converted to log₁₀ using ln x = 2.303 log₁₀ x.

E = 0.77 + (0.059 / 1) × log₁₀(0.034 / 0.64)

E = 0.77 + (−0.075)

E = +0.69 V

False.

The simplified form applies only at 298 K (25 °C). At other temperatures, the full Nernst equation using RT/zF must be used.

Oxidised species: Cu²⁺ (higher oxidation state, +2)

Reduced species: Cu (lower oxidation state, 0)

Cu is a solid so its concentration is set to 1.0 in the Nernst equation.

A natural logarithm (ln) in the Nernst equation is a mathematical function that can be converted to log₁₀ using ln x = 2.303 log₁₀ x, giving the simplified form used at 298 K.

For Cu²⁺ (aq) + 2e^- ⇌ Cu (s) at 298 K with [Cu²⁺] = 0.001 mol dm^-3 and E^θ = +0.34 V:

E = 0.34 + (0.059 / 2 ) × log₁₀(0.001 / 1.0)

A standard Gibbs free energy change (ΔG^ɔ) is the maximum energy available to do useful work when a reaction occurs under standard conditions (298 K, 100 kPa, unit activity). It combines enthalpy and entropy: ΔG^ɔ = ΔH^ɔ − TΔS^ɔ. A negative value indicates a spontaneous reaction.

ΔG^ɔ = −n × E_cell^ɔ × F

• ΔG^ɔ = standard Gibbs free energy (J mol^-1)

• n = moles of electrons transferred

• E_cell^ɔ = standard cell potential (V)

• F = Faraday constant (96 500 C mol^-1)

True.

From ΔG^ɔ = −n × E_cell^ɔ × F, when E_cell^ɔ is positive and n and F are always positive, ΔG^ɔ is negative, indicating a spontaneous reaction.

For the cell 2Fe³⁺ (aq) + Cu (s) ⇌ 2Fe²⁺ (aq) + Cu²⁺ (aq) with E_cell^ɔ = +0.43 V and n = 2:

ΔG^ɔ = − 2 × 0.43 × 96 500

ΔG^ɔ = −n × E_cell^ɔ × F

ΔG^ɔ = −2 × 0.43 × 96 500

ΔG^ɔ = −82 990 J mol^-1

ΔG^ɔ = −83 kJ mol^-1

False.

The equation gives ΔG^ɔ in J mol^-1. To convert to kJ mol^-1, divide by 1000.

ΔG^ɔ accounts for both the cell potential and the number of electrons transferred, giving the total free energy released. A reaction with a large n and small E_cell^ɔ may release more energy than one with a small n and larger E_cell^ɔ.

To determine n in ΔG^ɔ = −n × E_cell^ɔ × F, identify the half-cell with the less positive E^ɔ value, which undergoes oxidation, and count the electrons it transfers.

A Faraday's constant (F = 96 500 C mol^-1) is the charge carried per mole of electrons. In the equation ΔG^ɔ = −n × E_cell^ɔ × F, it converts electrical potential into energy.