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REVIEWTOPIC 9/19: OXIDATION AND REDUCTION
A. KEY DEFINITIONS (MEMORIZE!)
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Oxidation is the loss of electrons and reduction is the gain of electrons. (You could remember this using “OIL RIG”
or “LEO GER”.)
An oxidizing agent causes oxidation and in the process is reduced.
A reducing agent causes reduction and in the process is oxidized.
The standard electrode potential, E°, is the potential difference generated when a half-cell is connected to a
standard hydrogen electrode under standard conditions, which are 298 K, 101 kPa and 1 mol dm-3 solutions
B. YOU SHOULD KNOW:
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Every element has an oxidation number, which is 0 if it is an unreacted element, and in a compound it will be + or –
depending on whether the substance it has joined with is more or less electronegative than it
 The maximum positive oxidation number is the same as the group number and the lowest negative oxidation
number occurs when the element gains enough electrons to fill its outer shell (so, for example, Br in group 7 can
have an oxidation number between +7 and -1 (e.g., Br = +7 in BrO4- and Br = -1 in KBr)
 An oxidation number is not the same as a charge (so, for example, Ca2+ has an oxidation number of +2 and not 2+)
 That reactivity depends on how easily electrons are lost or gained
 The most reactive substances are found at either end of the standard electrode potential table (table 14 in data
booklet)
 The most reactive metals are at the top right and the most reactive non-metals are at the bottom left
VOLTAIC CELLS:
 That if two half-cells are connected to a voltmeter and there is a salt bridge between them, they will produce a
voltage
 The bigger the difference in reactivity between the half-cells, the greater the voltage – the electrode potentials are
given in table 14 of the data booklet
 Oxidation occurs at the Anode (“AN OX”); Reduction occurs at the cathode (“RED CAT”)
 The anode is the NEGATIVE electrode(e- flow from the anode); the cathode is the POSITIVE electrode
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The components of the standard hydrogen electrode that all electrode potentials are measured against - that is,
hydrogen gas at 298 K and 101 kPa bubbling onto a platinum electrode immersed in 1 mol dm -3 H+(aq)
ELECTROLYTIC CELLS:
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Electrolysis is the decomposition of an electrolyte by an electric current and makes elements from compounds
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In an electrolytic cell there is a direct current power source (e.g., a battery) and NOT a voltmeter, and there are
two electrodes and one electrolyte(there might also be an ammeter or bulb in the circuit)
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Electrolysis occurs when the ions can move – so the electrolyte (liquid with ions) must be molten or in aqueous
solution
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The amount of substance produced at the electrode will depend on the charge on the ion, the current that is
flowing and the time for which the current flows (i.e. charge, current, time)
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Oxidation still takes place at the anode BUT the anode is now the POSITIVE electrode – NEGATIVE IONS flow to it;
reduction still takes place at the cathode but cathode is now the NEGATIVE electrode – POSITIVE IONS flow to it
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Electrolysis can be used to electroplate metals – the object to be plated becomes the NEGATIVE electrode (i.e. the
cathode); the metal you are plating it with is the POSITIVE electrode (i.e. the anode)
C. YOU SHOULD BE ABLE TO:
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use oxidation numbers to decide if an element has been oxidized, because its oxidation number increases
(becomes more positive), or has been reduced, because its oxidation number decreases (becomes more negative)
name compounds using oxidation numbers using roman numerals, for example iron(II) oxide for FeO
write half-equations for reactions, balancing them with electrons (e-), H+ and H2O
combine half-equations then balance them by making sure electrons lost = electrons gained
work out a reactivity series if given information about how substances react with the same substance such as
oxygen, or with each other
Apply Faraday’s Law
Apply the Winkler Method
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remember that more reactive substances can displace less reactive substances, for example:
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
VOLTAIC CELLS:
 draw a diagram to explain how a voltage is produced and show the direction of electron flow [from more reactive
metal(anode) to less reactive metal(cathode)]
 calculate the cell potential by adding together the reduction potential and the oxidation potential(remember to
reverse the sign for the oxidation)
 predict whether the reaction written will occur spontaneously by looking at the cell potential – if it is positive, then
the reaction will occur spontaneously
ELECTROLYTIC CELLS:
 work out what will be produced at each electrode during the electrolysis of a molten salt
 work out what will be produced ate each electrode during the electrolysis of an aqueous solution
 write half-equations for the reactions at each electrode, remembering that metal ions gain electrons and nonmetal ions lose electrons
 remember that many non-metals form diatomic molecules
 remember that electrons flow in the wires and ions flow through the electrolyte
D. BE PREPARED
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There are many redox half-equations in table 14 of the data booklet. These are all written as reduction reactions,
and as equilibrium reactions. When you use them, remember to change the equilibrium arrows to a normal arrow,
and to write them the other way around if you want an oxidation reaction.
PROBLEM SET #9 (TOPIC 9/19)