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TSR Wiki > Study Help > Subjects and Revision > Revision Notes > Chemistry > Oxidation and Reduction

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20.1 Redox equations

20.1.1

Balancing: Copy down both complete half equations, reversing one so the electrons are on opposite sides and so the reactants for both are on the left, and the products on the right. Multiply the equations by the appropriate factors so the number of electrons in each equation is the same. Vertically add the equations, and cancel out any molecules which appear on both sides.

20.2 Standard electrode potentials

20.2.1

Standard electrode potential --> The potential difference between a given half cell (at 1 mol dm^-3 conc) and the standard hydrogen electrode

20.2.2

The standard hydrogen electrode consists of a solution of H₃O⁺ ions at 1 mol dm^-3 in a beaker. Placed into this is a platinum electrode surrounded by a gas tube submerged in the solution, with hydrogen at 1 atm inside. The circuit to the other half cell is then attached to the platinum electrode, and a salt bridge saturated in potassium chloride. The entire process should take place at 298K and 1 atm pressure.

20.2.3

The potential difference between half cells is a relative value, dependent on both half cells, and so a standard is required - the standard hydrogen electrode (why they didn't use a nice simple metal one escapes me...perhaps this is more accurate of something :)

20.2.4

Cell potential --> the potential difference between two half cells (if one half equation is reversed and the two equations are added, the cell potential will be given - it should be positive if you reversed the right one, if it's negative the reaction occurs in the opposite direction to the one you're writing).

Super-secret-bonus stuff they just slipped in without a sub-topic number...

Be able to draw a labeled diag of a half cell and how to link them together...which is basically described above. Cells should also be written in the form:

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which describes both half cells involved. also, the direction of current flow should also be deduced based of the standard cell potentials, as should the actual reaction occurring. The anode will lose electrons, and so the electrons must flow towards the cathode...that should allow everything else to be worked out.

20.2.5

The cell potential will be the potential difference between two half cells (and will be positive, unless the reaction occurs backwards). The magnitude is defined by the difference between the E-zero values of each half cell. One of the half equations will have to be reversed (the one which makes the total positive) and adding these two half equations will give you the overall reaction occurring.

20.2.6

Most reactions with positive E-zero values will occur, however it is possible that at non-standard conditions reactions may not occur, or that some reaction may have very high activation energy, and so will no occur at any reasonable rate.

20.3 Electrolysis

20.3.1

Electrolysis is where the above cells are forced to run in reverse by attaching an electricity source to overcome the potential difference. In aqueous solutions, however, water is also present, and will sometimes be oxidized/reduced in preference to the dissolved salts (or whatever). It is possible to use the standard electrode potentials to predict this, in that species above water (when it is on the left) will not be oxidized, and species below water (on the right) will not be reduced in an aqueous solution. If necessary, this can be checked by working out the cell potential for all possible combinations (involving the, presumably, two elements and water)...the reaction with the smallest negative potential difference will be the one which occurs. Highly concentrated solutions may overcome this to some degree however...ie it is possible for Cl₂ to be oxidized in a concentrated solution.

20.3.2

The Faraday constant is the charge (in magnitude because it should really be negative) of 1 mole of electrons.

20.3.3

Faraday's law states that the mass of product produced will be proportional to the charge passed. (Nb...the equation charge = current x time , or may be necessary). Faraday's law may also be restated as...the number of Faraday's required to discharge 1 mol of an ion at an electrode equals the number of charges on that ion.

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