Topic 20 - Oxidation and Reduction

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[edit] 20.1 Redox equations

[edit] 20.1.1

1. 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.
2. Multiply the equations by the appropriate factors so the number of electrons in each equation is the same.
3. Vertically add the equations, and cancel out any molecules which appear on both sides.

[edit] 20.2 Standard electrode potentials

[edit] 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.

[edit] 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.

[edit] 20.2.3

The potential difference between half cells is a relative value, dependent on both half cells, and so a standard is required for comparison purposes, for which the standard hydrogen electrode is used. Why they didn't use a nice simple metal one escapes me, though perhaps this one is more accurate or something.

[edit] 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.

...and the super-secret-bonus stuff they just slipped in without a sub-topic number...

You should be able to draw a labelled diagram of a half cell and how to link two cells together which we've already described above. Cells should also be written in the form Cu_(s)/Cu² _(aq)||Zn² _(aq)/Zn_(s), which describes both half cells involved. 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. From that information, everything else should be possible to work out.

[edit] 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.

[edit] 20.2.6

Most reactions with positive E-zero values will occur, however it is possible that under non-standard conditions reactions may not occur, or that some reactions may have very high activation energy, meaning the will not occur at any significant rate.

[edit] 20.2.7

For non-standard conditions, use the Nernst equation. Shown below;

E_cell = E_(cell under standard conditions) - (RT/zF)(ln Q)

Where R is the gas constant R = 8.314472

     T is temperature in Kelvin. To convert from Celcius to Kelvin 
       add 276.15
     z is the number of electrons transferred. Refer to half equation for this.
     F is Faraday's constant which is equal to the product of Avogrado's number         and the charge of an electron. F = 96485
     Q is the equilibirm expression, that is [product] / [reactants]

[edit] 20.3 Electrolysis

[edit] 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 water is also present, and will sometimes be oxidised/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 oxidised, 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 (i.e. it is possible for Cl₂ to be oxidised in a concentrated solution).

[edit] 20.3.2

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

[edit] 20.3.3

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