Oxidation-reduction: electrode potentials and cells

Oxidation numbers, oxidants and reductants

• Oxidation is loss of electrons, seen as a rise in oxidation number. Reduction is gain of electrons, seen as a fall. The oxidant is reduced; the reductant is oxidised.
• Oxidation numbers: elements are 0, simple ions equal their charge, O is −2 (−1 in peroxides), H is +1 with non-metals, and the total equals the charge on the particle.
• In S₂O₃²⁻: 2x + 3(−2) = −2, so each S is +2. In MnO₄⁻, Mn is +7; in Cr₂O₇²⁻, each Cr is +6.
• Write the sign first (+7) for an oxidation number, and last (2+) for an ion's charge.

Try it: electrons moving

Standard reduction potentials

• Each half-cell is written as a reduction, such as Cu²⁺ + 2e⁻ ⇌ Cu, and its standard reduction potential, E°, is measured against the standard hydrogen electrode, set at 0.00 V.
• Standard conditions are 1 mol L⁻¹ solutions, 25 °C and 100 kPa for gases.
• The more positive E° is, the more the left-hand species wants electrons, so the stronger the oxidant. The more negative E° is, the more readily the right-hand species gives electrons away, so the stronger the reductant.
• E° does not depend on how many moles are in the equation, so it is never multiplied when a half-equation is.

How an electrochemical cell works

• An electrochemical (galvanic) cell keeps the oxidant and reductant apart, so the electrons have to travel through an external wire, where they can do work.
• Oxidation happens at the anode, the negative electrode: Zn → Zn²⁺ + 2e⁻. Reduction happens at the cathode, the positive electrode: Cu²⁺ + 2e⁻ → Cu. Electrons flow through the wire from anode to cathode.
• The salt bridge, such as filter paper soaked in KNO₃, lets ions move to keep each beaker electrically neutral: anions drift towards the anode beaker and cations towards the cathode beaker. Without it the current stops at once.
• A cell diagram puts the anode on the left: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). A single line is a phase boundary; the double line is the salt bridge.
• If neither species in a half-cell is a metal, as in Fe³⁺/Fe²⁺, an inert platinum or graphite electrode carries the electrons: Pt(s) | Fe²⁺(aq), Fe³⁺(aq).

Predicting whether a reaction happens

• Write the proposed reaction as a reduction half and an oxidation half, look up both E° values, and calculate E°cell = E°(reduced) − E°(oxidised).
• If E°cell is positive, the reaction is spontaneous under standard conditions. If it is negative, it is not, but the reverse reaction is.
• On a table in this order, a species on the left reacts spontaneously with any species on the right that sits below it.
• A positive E°cell does not mean the reaction is fast; some spontaneous reactions are too slow to see. And under non-standard conditions, such as very dilute solutions, the cell potential differs from E°cell.

Try it: spontaneous or not?

Electrolytic cells

• In an electrolytic cell a power supply pushes electrons to make a non-spontaneous reaction happen, one with a negative E°cell. The supply must provide at least that voltage.
• Oxidation still happens at the anode, but here the anode is the positive electrode; reduction happens at the cathode, which is negative.
• In molten NaCl the only ions are Na⁺ and Cl⁻: sodium metal forms at the cathode and chlorine gas at the anode.
• In water, water itself competes. At the cathode, the species with the most positive E° is reduced: Cu²⁺ (+0.34) is reduced before water (−0.83), but Na⁺ (−2.71) is not, so H₂ forms instead. At the anode, water is oxidised to O₂ unless a more easily oxidised ion is present in high concentration.
• At Tiwai Point in Southland, aluminium is made by electrolysis of aluminium oxide dissolved in molten cryolite. With E° = −1.66 V for Al³⁺/Al, aluminium is too strong a reductant to be freed from its ore by heating with carbon, as iron is.

Galvanic and electrolytic cells compared

• Both kinds of cell have oxidation at the anode and reduction at the cathode. What changes is the sign of each electrode and where the energy comes from.

How this internal standard is assessed

• AS 91393 is assessed by your school, not in the exam. Your teacher sets a task, often a set of practical observations such as metals in solutions, cells you build and measure, or an electrolysis, followed by written questions.
• You explain what you saw using half-equations, overall equations, E° values and E°cell calculations. Calculations are limited to electrode potentials.
• Schools set their own conditions, but the written part is usually done in class under supervision, with an E° table supplied. If your teacher allows it, one resubmission can fix a small error; it is not a second attempt at the whole task.
• The grade depends on how well you link processes, observations, equations and calculations, not on how much you write.

What lifts it to Merit and Excellence

• Achieved describes processes and does the calculations. Merit explains links between what was seen, the equations and the E° values. Excellence compares and justifies, explaining why one process happens and another does not.