Redox Equilibria (Edexcel International A Level (IAL) Chemistry): Flashcards

The increased surface area of platinum black increases the rate of equilibrium between hydrogen gas and hydrogen ions at the electrode surface.

1. Ion concentration of 1.00 mol dm^-3

2. Temperature of 298 K

3. Pressure of 100 kPa

The standard electrode potential is measured using a high-resistance voltmeter so that no current flows and the maximum potential difference is achieved.

Standard electrode potential is the potential difference produced when a standard half-cell is connected to a standard hydrogen electrode under standard conditions.

False.

The standard hydrogen electrode is assigned an E^ꝋ value of 0.00 V by convention, serving as the reference point for all other electrode potentials.

A more positive E^ꝋ value indicates the species is more easily reduced relative to the standard hydrogen electrode.

2H⁺ (aq) + 2e^– ⇌ H₂ (g)

This equilibrium is established on the platinum electrode surface at standard conditions.

True.

A more negative E^ꝋ means the species on the right of its half-equation is more easily oxidised compared to species with higher E^ꝋ values.

1. Metal / metal ion half-cell

2. Non-metal / non-metal ion half-cell

3. Ion / ion half-cell (ions in different oxidation states)

A salt bridge is an ionic conductor (typically saturated KCl or KNO₃ in a gel) that connects two half-cells, allowing ions to flow and balance charge without transferring electrons.

A metal wire allows electrons to flow, but only an ionic conductor can complete the circuit by moving ions between half-cells to balance the build-up of charge.

Chloride and nitrate ions are highly soluble, unreactive with most ions in the half-cells, and unlikely to form precipitates, ensuring smooth ion flow.

In a non-metal / non-metal ion half-cell, a platinum electrode is used to make electrical contact with the solution.

E^ꝋ_cell = E^ꝋ_right – E^ꝋ_left = (+0.80) – (0.00) = +0.80 V

The Ag⁺/Ag half-cell is the positive pole.

Reduction occurs at the positive electrode (cathode), while oxidation occurs at the negative electrode (anode).

True.

When both species are in solution, an inert platinum electrode is required to provide a surface for electron transfer without interfering with the redox reaction.

The standard cell potential (E^ꝋ_cell) is calculated by subtracting the less positive E^ꝋ from the more positive E^ꝋ value.

A cell diagram (cell notation) is a shorthand representation of an electrochemical cell that shows the materials and redox processes without drawing the full cell.

The anode (oxidation, most negative E^ꝋ) is on the left, and the cathode (reduction, most positive E^ꝋ) is on the right.

A single vertical line represents a phase boundary between two different states of matter in the same half-cell (e.g. between a solid electrode and an aqueous solution).

The double vertical line represents the salt bridge separating the two half-cells.

When two species in a half-cell are in the same phase, they are separated by a comma in the cell diagram.

Both Fe³⁺ and Fe²⁺ are in aqueous solution, so there is no solid metal to act as an electrode. An inert platinum electrode is needed to carry electrons in and out of the solution.

Zn (s) | Zn²⁺ (aq) ‖ Cu²⁺ (aq) | Cu (s)

E^ꝋ_cell = (+0.34) − (−0.76) = +1.10 V

False.

Cell diagrams are not quantitative — they show the materials and redox processes involved, but do not indicate stoichiometric ratios of species.

By convention, the half-cell with the more positive E^ꝋ (cathode, reduction) is written on the right, and the half-cell with the less positive E^ꝋ (anode, oxidation) is on the left.

Thermodynamic feasibility is the prediction of whether a reaction is likely to occur spontaneously, based on whether the E^ꝋ_cell is positive.

A reaction is feasible (spontaneous) when E^ꝋ_cell is positive. A negative E^ꝋ_cell indicates the reaction is not feasible in the forward direction.

Cl₂ is reduced, as it has the more positive E^ꝋ value. Cu is oxidised because it has the less positive value.

A larger cell potential means a bigger change in total entropy, which is directly proportional to E^ꝋ.

ΔG^ꝋ = −nFE^ꝋ_cell

ΔG^ꝋ = −RT ln K

From these, E^ꝋ is directly proportional to ln K.

False.

A positive E^ꝋ_cell shows a reaction is thermodynamically feasible, but gives no information about the rate of reaction. High activation energy can make the reaction kinetically hindered.

1. E^ꝋ gives no information about the rate of reaction (kinetics).

2. Non-standard conditions (e.g. different concentrations) shift the equilibrium and alter the electrode potential.

The equilibrium shifts to the right, removing electrons from the system and making the electrode potential less negative (more positive) than the standard value.

True.

Because ΔG^ꝋ = −nFE^ꝋ_cell and ΔG^ꝋ = −RT ln K, it follows that cell potential is directly proportional to ln K.

2H₂ (g) + O₂ (g) → 2H₂O (l)

The cell potential is +1.23 V in both cases.

A fuel cell is an electrochemical cell in which a fuel donates electrons at the anode and oxygen gains electrons at the cathode, generating a continuous electrical current.

H₂ (g) + 2OH^– (aq) → 2H₂O (l) + 2e^–

This is oxidation at the negative electrode.

O₂ (g) + 2H₂O + 4e^– → 4OH^– (aq)

This is reduction at the positive electrode.

In a hydrogen-oxygen fuel cell, the anode is the negative electrode and the cathode is the positive electrode.

1. The only product is water, so no CO₂ or harmful pollutants are emitted.

2. No oxides of nitrogen are produced because the reaction occurs at room temperature without high-temperature combustion.

1. Hydrogen is highly flammable, making production and storage hazardous.

2. Hydrogen is currently a by-product of the crude oil industry, relying on a non-renewable resource.

False.

A fuel cell operates continuously as long as hydrogen and oxygen are supplied — energy is not stored within the cell, unlike in conventional batteries.

In a fuel cell, the reaction occurs at room temperature without combustion, so all the bond energy is converted to electrical energy rather than heat and light.

The ratio is 1 : 5 (one mole of MnO₄^– reacts with five moles of Fe²⁺), as balanced by combining the two half-equations.

Dilute sulfuric acid does not react with manganate(VII) ions and does not act as a reducing agent under these conditions, so it does not interfere with the titration.

The formula for percentage uncertainty in a measurement is:

Percentage uncertainty = uncertainty ÷ measured value × 100

The ratio is 2 : 1 (two moles of S₂O₃^2– react with one mole of I₂), from the equation 2S₂O₃^2– (aq) + I₂ (aq) → 2I^– (aq) + S₄O₆^2– (aq).

Starch forms a blue-black complex with iodine, making it easy to detect the precise end-point when the colour disappears as the last iodine is consumed.

Total absolute uncertainty = 0.1 + 0.1 = 0.2 cm³

Percentage uncertainty = (0.2 / 5.0) × 100 = 4.0%

In redox titrations, the manganate(VII) ion is the oxidising agent and the iron(II) ion is the reducing agent.

Dilute the solution in the burette so that a larger volume is required to reach the end-point, reducing the percentage uncertainty in the titre.

MnO₄^– (aq) + 8H⁺ (aq) + 5Fe²⁺ (aq) → Mn²⁺ (aq) + 4H₂O (l) + 5Fe³⁺ (aq)

Mn is reduced from +7 to +2, and Fe is oxidised from +2 to +3.