A2: Redox and electrode potentials

A2: Redox and electrode potentials

1. Redox reactions

   • Manganate (VII) – MnO4- titration (under acidic conditions – self-indicating). It can be used to analyse Fe2+ ions and ethanedioic acid (COOH)2. 1. KMnO4- is added to the burette. 2. Add a measured volume of the solution being analysed and H2SO4 to the conical flask (provide H+ for reduction). 3. The end point is the first permanent pink colour – excess of MnO4-. • Repeat to get concordant results. • The meniscus is read from the top due to the deep purple colour.
   • Iodine/thiosulfate titrations. It can be used to analyse CIO- content in household bleach, Cu2+ content in copper(II) compounds and Cu content in copper alloys. • S2O32- are oxidised and I2 is reduced. • Overall: 2S2O32-(aq) + I2(aq) → 2I-(aq) + S4O62-(aq). 1. Add Na2S2O3 to the burette. 2. Add a measured solution of the oxidising agent to the conical flask then excess KI. This reacts to form iodine, yellow-brown colour. 3. Titrate this solution against the Na2S2O3. The iodine is reduced back to I- ions. When the end point is approaching, the solution will appear a pale straw colour. 4. Starch indicator is then added to form a deep blue-black colour. As more sodium thiosulfate is added, the colour fades so the end point will be when all of the blue-black colour disappears.

2. Electrode potentials

   • Electrochemical cells • Voltaic cell: converts chemical energy into electrical energy. • Half-cell: contains the species present in a redox half-equation. The two half-cells must be kept separate otherwise if mixed, the electrons would flow in an uncontrollable way and heat energy would be released rather than electrical energy.
   • Metal/metal ion half-cell: • A metal rod dipped into a solution of its aqueous metal ion. • An equilibrium will be set up at the phase boundary where the metal is in contact with its ions. Forward reaction = reduction and reverse reaction = oxidation. • In an isolated half-cell, there is not net transfer of electrons into or out of the metal. • When two half-cells are connected, the direction of electron flow depends upon the relative tendency of each electrode to release electrons. • The more reactive metal releases electrons more readily and is oxidised.
   • • The electrode with the less reactive metal gains electrons = reduced – positive electrode. • Standard electrode potential: the e.m.f of a half-cell compared with a standard hydrogen half-cell under standard conditions, 298K, solution concentrations of 1 mol dm-3 and gas pressure of 100 kPa. • Two electrodes are connected by a wire to allow a controlled flow of electrons. • The two solutions are connected with a salt bridge which allows ions to flow e.g. KNO3(aq). • The more negative value: lose electrons and undergo oxidation. • The more positive value: gain electrons and undergo reduction. • Metals = negative values and lose electrons whereas non-metals = positive values and gain electrons. • More negative value = greater the reactivity of a metal in losing electrons. • More positive value = greater the reactivity of a non-metal in gaining electrons. • E cell = E(positive electrode) – E(negative electrode).
   • Ion/ion half-cell: • Contains ions of the same element in different oxidation states. • There is no metal to transport electrons into or out of the cell, so an inert metal electrode made out of platinum is used.
   • Predicting redox reactions: • Most negative system has the greatest tendency to be oxidised and lose electrons. • Most positive system has the greatest tendency to be reduced and gain electrons. • The strongest reducing agent is at the top on the right. • The strongest oxidising agent is at the bottom on the left. • The redox system with the most positive value will react from left to right and gain electrons. • The redox system with the more negative value will react from right to left, and lose electrons.
   • Limitations: • Reactions have a large activation energy resulting in a very slow rate. The electrode potentials may indicate the thermodynamic feasibility of a reaction, but they give no indication of the rate of a reaction. • Concentrations may not be at 1 mol dm-3. • Actual conditions used may be different to the standard conditions used to calculate E values.

3. Storage and fuel cells

   • Primary cells: • Non rechargeable and are designed to be used once only. • Electrical energy is produced by oxidation and reduction at the electrodes in an irreversible reaction. • The chemicals will be used up, voltage will fall, battery will go flat and the cell will be discarded or recycled. • Used in low-current, long-storage devices such as wall clocks and smoke detectors. • Most are based on Zn/MnO2 and a potassium hydroxide alkaline electrolyte.
   • Secondary cells: • Rechargeable – the cell reaction producing electrical energy can be reversed during recharging. • Examples include lead-acid in car batteries, lithium-ion (and polymer) in laptops, phones etc.
   • Fuel cells: • Use the energy from the reaction of a fuel with oxygen to create a voltage. • Fuel and oxygen flow into the fuel cell and the products flow out. The electrolyte remains in the cell. • Can operate continuously as long as the fuel and oxygen are supplied into the cell. • Don’t have to be recharged. • Hydrogen fuel cells are the most common, as they produce no carbon dioxide during combustion, with water being the only combustion product. • Can have an acid or alkali electrolyte. The cell voltage still stays the same at 1.23V.