AQA Chem5: Redox Equilibria

• Oxidation is the loss of electrons
• Reduction is the gain of electrons
• In a redox equation something is oxidised while something else is reduced
• An oxidation state is a charge and number assigned to atoms in molecules or ions to show how many electrons the atom has used in bonding with a different element
• Many elements have variable oxidation states in their compounds
• Some elements have fixed oxidation states in their compounds e.g. halides are -1, group I are +1 and group II are +2
• To make a redox equation you have to balance and then combine the half equations 

• convert chemical energy into electrical energy
• involve redox reactions
• a metal in contact with a solution of its ions is known as a half cell
• when two half cells are connected an electrochemical cell is obtained
• transfer of electrons between two half cells produces an electrical current
• electrons flow from the negative half cell to the positive half cell in the external circuit
• ions conduct electricity in the internal circuit
• a salt bridge has to be present to allow the flow of ions and to complete the circuit

• they aren't actually a diagram but a shorthand way of representing an electrochemical cell
• the cell undergoing reduction (the positive one) is assigned to the right hand side electrode
• generally:   M (s) | Mn+ (aq) || M1n+ (aq) | M1 (s)
• the  single solid line represents the boundary between the solid metal and the metal ions
• the double line represents the salt bridge

• the electrode potential of a half cell is a measure of electromotive force (e.m.f) of the cell
• the standard electrode potential has the symbol EѲ
• when two half cells are combined the cell potential or standard e.m.f of the cell is given by:                             EѲ cell = EѲ R.H.S - EѲ L.H.S
• EѲ values are determined using a standard reference electrode, this is usually the standard hydrogen electrode (S.H.E)

The S.H.E consists of:

• a Pt electrode
• H2 gas at 100kPa
• a 1moldm-3 solution of H+ ions (HCl)
• a mercury contact
• a copper wire
• (it is all at 298K)

The half cell diagram is: Pt(s) | H2(g) | 2H+(aq) ||

• the EѲ is 0.00V
• it is difficult to set up so a secondary reference electrode is used
• this is the calomel electrode which is more convenient to use

• the reactions are written as reductions
• positive potentials accept electrons and are good oxidising agents e.g. Au3+(aq) + 3e- <-> Au(s)
• negative potentials donate electrons and are good reducing agents e.g. Li+(aq) + e- <-> Li(s)
• it can be said that those with positive potentials are 'electron grabbing' and those with negative potentials are 'electron giving'

Use EѲcell = EѲ R.H.S - EѲ L.H.S

• the more positive electrode is on the right hand side (R.H.S)
• if the EѲcell value is negative the cells are in reverse
• when calculating EѲcell do not change the signs of the values (i.e. keep them negative or positive)
• the more positive half cell is always reduced
• electrons flow from the less positive to the more positive