Electrochemistry

In a redox reaction, oxidation and reduction occur together. Oxidation is electron loss or an increase in oxidation number. Reduction is electron gain or a decrease in oxidation number. The oxidising agent accepts electrons and is itself reduced; the reducing agent supplies electrons and is itself oxidised.

For Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), zinc changes from 0 to +2 and releases two electrons. Copper changes from +2 to 0 and accepts them. Separating the two half-reactions can make those electrons travel through an external wire, delivering electrical energy.

For a half-equation required in alkaline solution, an acid-balanced equation can be converted by adding enough OH^- to both sides to consume every H⁺, then cancelling water. Use the actual stated product: permanganate, for example, does not necessarily form Mn²⁺ outside acidic conditions.

A standard electrode potential, E°, is the potential difference of a half-cell relative to the standard hydrogen electrode, with all species in their standard states and negligible current flowing. Values are quoted for the half-reaction written as a reduction. A more positive value indicates a stronger tendency for that reduction under the stated standard conditions.

For the usual school measurements, aqueous concentrations are 1 mol dm^-3, gases have standard pressure 100 kPa, and pure solids or liquids are used in their standard states. Strictly, dissolved species have unit activity; concentration is the course approximation. Temperature must be stated: the Data Booklet values are at 298 K. Standard state is not another name for 273 K gas conditions.

The standard hydrogen electrode (SHE) uses an inert platinum electrode coated with finely divided platinum in contact with H₂(g) at 100 kPa and aqueous H⁺ at 1 mol dm^-3, at the stated temperature. Its reversible half-equation is 2H⁺(aq) + 2e^- ⇌ H₂(g), and its standard potential is assigned 0.00 V. Platinum provides electrical contact and a catalytic surface; it is not consumed in this equation.

Keep both half-cells at 298 K when using the tabulated values. A suitable salt bridge, such as one containing KNO₃, permits ion movement and prevents charge build-up without rapidly mixing the solutions. The voltmeter draws negligible current so the reading approximates the equilibrium electromotive force. Record which terminal is positive, not only the size of the reading: copper is +0.34 V relative to the SHE, while zinc is -0.76 V.

The standard cell potential E°_cell is the cell electromotive force, or maximum potential difference measured at negligible current, when all components are in standard states at the stated temperature. For a chosen reaction, E°_cell = E°_cathode - E°_anode, with both tabulated values kept as reduction potentials.

At the cathode, reduction consumes electrons. At the anode, oxidation releases them. In a spontaneous galvanic cell, electrons flow through the wire from the negative anode to the positive cathode. Ions carry charge through the electrolyte and salt bridge; electrons do not flow through the bridge.

Electrode potentials depend on the actual composition. Think of each half-equation as a reduction equilibrium: a change favouring reduction generally makes that electrode potential more positive. A change favouring the reverse makes it less positive. These are qualitative predictions at fixed temperature, with other conditions held constant.

In the Zn/Cu cell, diluting Cu²⁺ lowers the copper reduction potential. Increasing Zn²⁺ raises the zinc reduction potential. Either change reduces E_cell = E_Cu - E_Zn. Far enough from standard conditions, an E° prediction can cease to describe the actual direction. Complex formation and precipitation matter because they change the free ion concentrations.

Even a positive actual cell potential says nothing about how quickly a reaction proceeds. Large activation barriers, slow electrode reactions or a protective surface film can make a thermodynamically favoured change difficult to observe. Different temperature, gas pressure, pH or chemical species can also invalidate a direct application of tabulated E° values. Use the exact relevant half-equation and the conditions, not only the names of the elements.

ΔG° = -nFE°_cell. Here n is the number of moles of electrons transferred per mole of the cell reaction as written, and F is the charge per mole of electrons. With F in C mol^-1 and E in V, ΔG° is in J mol^-1 of reaction because 1 V = 1 J C^-1.

When two reduction steps combine to make a new reduction half-equation, their E° values cannot simply be added. Convert each to its corresponding -nFE° energy contribution relative to the same reference, add those contributions with the required signs, then divide by -nF for the combined half-equation.

A cell converts chemical free energy into electrical energy. A rechargeable battery stores reactants internally and uses an external supply to restore them during charging. A fuel cell can operate while fuel and oxidant are supplied continuously and products are removed. Hydrogen-oxygen fuel cells and improved batteries offer possible gains in compactness, lower mass and cell voltage; compare the complete system, not just an attractive electrode reaction.

The ideal standard voltage for this acidic cell is 1.23 - 0.00 = 1.23 V at 298 K; doubling its chemical equation does not give twice the voltage. Real operating voltage is lower because current flow involves losses. Several cells in series add their voltages. The fuel cell produces water at use, but hydrogen production, storage and transport affect the wider environmental and mass comparison.

Electrolysis uses electrical energy to drive a reaction. The supply delivers electrons to the negative cathode and removes them from the positive anode. Cathode still means reduction and anode still means oxidation: those definitions do not change when the electrode signs differ from a galvanic cell.

In aqueous sodium chloride, water reduction is much easier than Na⁺ reduction, so hydrogen forms at the cathode. At an inert anode, water/oxygen and chloride/chlorine compete. Increasing chloride concentration favours chloride oxidation; oxygen formation also has a kinetic barrier at many electrode surfaces. This is why a simple comparison of the standard +1.23 V oxygen and +1.36 V chlorine reduction values cannot, by itself, correctly predict every brine experiment.

Electrolysis can change its own conditions. As Cu²⁺ is depleted near an inert cathode, hydrogen evolution can become important; changes in concentration and pH may alter the products. An electrolyte conducts because its ions move, not because electrons pass through the bulk liquid.

The Faraday constant F is the charge per mole of electrons: F = Le, where L is the Avogadro constant and e is the magnitude of the charge on one electron. The Data Booklet uses L = 6.02 × 10²³ mol^-1, e = 1.60 × 10^-19 C and F = 9.65 × 10⁴ C mol^-1. The small difference in multiplying the printed rounded constants comes from rounding.

For a constant current, Q = It, with current in amperes, time in seconds and charge in coulombs. Then n(e^-) = Q/F. Divide by the number of electrons required per product particle in the half-equation. Finally use mass = amount × molar mass, or gas volume = amount × the stated molar volume. If current varies, Q is the area under the current-time graph.

The net intended process is transfer of copper from the impure anode to the pure cathode. The power supply controls this electrolytic process. Swapping the electrodes would defeat the purpose; simply filtering the original copper cannot remove metallic impurities within it.

In anodising aluminium, the aluminium object is connected as the positive anode in a suitable acidic electrolyte, such as dilute sulfuric acid. Oxidation builds a thicker protective aluminium oxide layer on its surface. An idealised oxide-forming anode equation is 2Al(s) + 3H₂O(l) → Al₂O₃(s) + 6H⁺(aq) + 6e^-.

At the cathode, hydrogen ions can be reduced: 6H⁺(aq) + 6e^- → 3H₂(g). The oxide layer adheres to and protects the object; anodising is not deposition of aluminium metal onto a cathode. Detailed plant design and operating parameters are outside this syllabus.

Always: oxidation at the anode, reduction at the cathode. Galvanic discharge: anode negative, cathode positive. Electrolysis: anode positive, cathode negative. Standard-potential tables predict thermodynamic direction; composition and kinetic barriers explain why a real observation can differ.