Comprehensive study guide for Chemistry A Level, made by an Oxford Biochemistry student with all 9s at GCSE and 3 A*s at A Level! Information arranged by spec point. Notes written using past papers, textbooks and more.
14. REDOX II
1. understand the terms ‘oxidation’ and ‘reduction’ in terms of electron
transfer, applied to s-, p- and d-block elements
Oxidation and reduction in terms of electron transfer:
Oxidation – loss of electrons.
Reduction – gain of electrons.
OIL RIG – oxidation is loss, reduction is gain.
2. understand the terms ‘oxidation’ and ‘reduction’ in terms of changes
in oxidation number, applied to s-, p- and d-block elements
Oxidation and reduction in terms of changes to oxidation number:
Oxidation – an increase in oxidation number.
Reduction – a decrease in oxidation number.
Disproportionation is when an element is simultaneously oxidised and
reduced.
3. know what is meant by the term ‘standard electrode potential’, E ϴ
Standard electrode potential, Eϴ – the electron motive force (emf) of a half cell
relative to the standard hydrogen electrode, measured under standard
conditions.
The emf tells you the potential difference between two electrodes when no
current is flowing.
Eϴ tells us about the position of equilibrium and strength of the
reducing/oxidising agent.
o A large positive Eϴ shows that the equilibrium lies far to the right, and
that the reactant is a strong oxidising agent and is easily reduced.
We also combine the Eϴ of two different half equations to give an Eϴcell value.
o This tells us about the direction of electron flow, the position of
equilibrium, species’ relative reducing and oxidising abilities and the
feasibility of the reaction.
Electrochemical cells:
Using an electrochemical cell, we can harness the electron flow of a redox
reaction to convert chemical energy into electrical energy.
o We separate the oxidation and reduction reactions of a redox reaction,
allowing each to occur in isolation in a half cell.
o An electrochemical cell is made by joining these two half cells and
forming a complete circuit with a wire and salt bridge.
The cell reaction is the sum of the reduction and oxidation half-reactions
taking place in the cell; i.e. it is the overall chemical reaction.
,
Anode – the (negatively charged) electrode of the oxidation half cell.
Cathode – the (positively charged) electrode of the reduction half cell.
High-resistance voltmeter – measures the potential difference between the
electrodes.
o Since V = IR, a high resistance means that there is minimal current,
allowing us to measure the emf and thus the E ϴ of the cell.
Salt bridge – completes the circuit by allowing ions to flow between the
solutions, without the solutions mixing. Normally filter paper soaked in KNO 3
(aq) is used.
o KNO3 is inert and will not react with either of the solutions in the half
cells.
4. know that the standard electrode potential, Eϴ, refers to conditions
of:
i) 298 K temperature
ii) 100 kPa pressure of gases
iii) 1.00 mol dm-3 concentration of ions
The standard electrode potential is taken under these standard conditions:
A temperature of 298K.
A pressure of 100 kPa (1 atm) for all gases.
A concentration of 1 mol dm-3 for all aqueous ions (1 mol dm-3 of the reacting
species).
o E.g. the standard hydrogen electrode needs 1 mol dm -3 H+ so we use
0.5 mol dm-3 H2SO4.
5. know the features of the standard hydrogen electrode and
understand why a reference electrode is necessary
Electrode potential and need for a reference electrode:
A potential difference is set up when a metal electrode is dipped into a
solution of its ions. This can occur in two ways:
o The metal atoms form cations, leaving the metal with a surplus of
electrons that gives the electrode a net negative charge.
o The aqueous ions gain electrons from the metal. The electrode gets a
net positive charge.
We can only measure the difference between two electrode potentials, not
the potential difference between an electrode and a solution (the absolute
electrode potential).
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