Electrochemistry

What the exam asks

Expect a table of standard reduction potentials and a cell, and expect to be asked which electrode is which, which way the electrons go, what the cell voltage is, or how much metal plates out. The commonest items are: pick the cathode from the table, compute the cell potential, decide whether the reaction is spontaneous, and run a Faraday calculation from current and time to a mass. Two traps do most of the damage. The first is memorising 'cathode is positive', which is true only for galvanic cells and gets electrolysis exactly wrong; go back to oxidation at the anode and reduction at the cathode, then work out the sign from which cell you are in. The second is multiplying a standard potential by the coefficient used to balance electrons. Voltage is per coulomb, so it does not scale, and a stem that offers a neatly doubled value is offering it on purpose.

Electrochemistry is redox with the two halves pulled apart and a wire run between them. Once the halves are separated, electrons have to travel to get from the species being oxidised to the species being reduced, and that traffic is a current you can measure, drive, or use.

Two cell types, and they are mirror images. A galvanic cell runs a spontaneous reaction and delivers energy: the numbers tell you it will go, and the cell voltage is positive. An electrolytic cell has a power supply forcing a reaction that would never happen on its own, so the cell voltage is negative and you are paying for it. What does not change between them is the definition of the electrodes. Oxidation happens at the anode and reduction at the cathode, always, in both cells. What flips is the sign the electrodes carry.

The exam supplies a table of standard reduction potentials and expects you to use it. The skill is reading the table: the half-reaction with the more positive reduction potential is the one that actually gets reduced, the other one runs backwards as an oxidation, and the difference between them is the cell voltage. The rest is Faraday arithmetic, which is a unit chain from current and time to moles of electrons to grams.

What to hold

• Oxidation is loss of electrons and reduction is gain, and each is defined by what happens to the oxidation number, not by whether oxygen is involved.
• The anode is where oxidation happens and the cathode is where reduction happens, in every cell of either type. Electrons leave the anode and arrive at the cathode through the external circuit.
• In a galvanic cell the anode is negative and the cathode positive; in an electrolytic cell the anode is positive and the cathode negative. The signs swap, the definitions do not.
• Standard reduction potentials are tabulated for the reduction direction, so a half-reaction run as an oxidation has its sign reversed.
• Given a table, the half-reaction with the more positive reduction potential is reduced and becomes the cathode; the other is forced to run in reverse as the anode.
• Cell potential is E(cathode) minus E(anode), using both values as they appear in the reduction table, with no sign flipping on the way in.
• A positive cell potential means the reaction is spontaneous as written, which is the definition of a galvanic cell. A negative one means it needs driving, which is electrolysis.
• Cell potential is intensive: doubling a half-reaction to balance the electrons does not double its voltage, because voltage is energy per unit charge.
• Free energy change equals minus nFE, so a positive cell potential and a negative free energy change say the same thing in different units.
• The standard hydrogen electrode is assigned exactly zero volts by convention, which makes every other potential a comparison against it rather than an absolute quantity.
• The salt bridge keeps each compartment electrically neutral as charge builds, and without it the cell stops almost immediately.
• Faraday's constant is about 96500 coulombs per mole of electrons, and the chain from current to mass is: charge equals current times time, charge divided by Faraday gives moles of electrons, then divide by the electrons per ion to get moles of metal.
• A more positive reduction potential means a stronger oxidising agent, because the species is more willing to take electrons.