Unit 5 Electrochemistry

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Electrochemistry and Redox Equilibria - Answers
1. In a redox reaction, one substance is oxidised and the other is reduced:
Oxidation is the gain of oxygen
Reduction is loss of oxygen
Neither oxidation nor reduction can take place on its own and they happen simultaneously, e.g. when methane is
burnt in oxygen, methane is oxidised and oxygen is reduced.
2. The definition of oxidation and reduction was expanded to include the removal of hydrogen as an example of
oxidation and the gain of hydrogen as an example of reduction. For example, when chlorine gas is bubbled into
aqueous hydrogen sulphide, the hydrogen sulphide is oxidised to sulfur by the removal of hydrogen. The chlorine
is reduced to HCl by the addition of hydrogen:
H2S + Cl2 S + 2HCl
3. In terms of electron transfer:
Oxidation is loss of electrons
Reduction is gain of electrons ­ remember it as OIL RIG.
E.g. when a piece of zinc is placed in a solution of copper (II) sulphate, it is oxidised to zinc ions and the copper
ions are reduced to copper metal:
Zn(s) Zn2+(aq) + 2e-
Cu2+(aq) + 2e- Cu(s)
4. An oxidising agent is a species that oxidises another species by removing one or more electrons. When an
oxidising agent reacts, it is itself reduced.
A reducing agent reduces another species by giving it one or more electrons. When a reducing reagent reacts, it
is itself oxidised.
5. Oxidation reactions can be written as half-equations, which show the loss of electrons from a single species and
the oxidation product. For example, the oxidation half-equation for the oxidation of zinc atoms to zinc ions is:
Zn (s) Zn2+ + 2e-
6. There are four points to note about half-equations:
In an oxidation reaction the electrons are on the RHS of the equation and in a reduction reaction the
electrons are on the LHS of the equation
If the reaction takes place in acid solution, add H+ ions to the LHS and water molecules to the RHS
The equation must balance for numbers of atoms
Change the number of electrons so that the equation balances for charges
7. Half equations can be combined to give the overall equation for a redox reaction:
Multiply one or both half-equations by integers so that the number of electrons becomes the same in both
Add the two half-equations together and cancel the electrons to obtain the overall equation
Check that both reactants are on the LHS of the overall equation
8. The oxidation number of an element in a compound or ion is the charge that the element would have if the
compound were fully ionic. The oxidation is calculated on the basis that the bonding electrons are assigned to the
more electronegative atom in a covalent bond.
9. Oxidation numbers can be deduced using a series of rules:
1) The oxidation number of an uncombined element is zero
2) The oxidation number of the element in a simple ion is the charge on the ion
3) When combined, group 1 elements are +1, group 2 elements +2 and group 3 elements +3
4) The sum of the oxidation umber in a mole is zero
5) The sum of all the oxidation umbers in a polyatomic ions is the charge on the ion
6) Fluorine in compounds always has the oxidation -1
7) Hydrogen when combined is always +1 (except metal hydrides)
8) Oxygen, when combined, is always -2, except in peroxides (-1) and when combined with fluorine (+2)
9) Chlorine, when combined, is always -1, except when combined with oxygen or fluorine when it is positive
10. An increase in the oxidation number of an element means that it has been oxidized, e.g. Fe2+ Fe3+ + e-
A decrease in the oxidation number of an element means that it has been reduced, e.g.
MnO4- + 8H+ + 5e- Mn2+ + 4H2O. The manganese has gone from the +7 state to +2
11. The number of electrons in a half-equation is equal to the total change in oxidation number of the element, e.g.
In acid solution, FeO42- can be reduced to Fe2+ ions. The oxidation number of iron changes from +6 to +2. This is a
decrease of four, so there are four electrons on the LHS of the reduction half-equation.

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Disproportionation is a redox reaction in which an element in a single species is simultaneously oxidized and
reduced. E.g. when a chlorine gas is bubbled into aqueous NaOH, a disproportionation reaction takes place. One
atom in the Cl2 molecules is oxidized to NaOCl and on is reduced to NaCl. The equation with the oxidation number
of chlorine is: Cl2 + 2NaOH NaOCl + NaCl + H2O
0 +1 -1
13.…read more

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Comparing values of E°:
The more negative the E° value, the more the position of equilibrium lies to the left - the more readily
the metal loses electrons. The more negative the value, the stronger reducing agent the metal is.
The more positive the E° value, the more the position of equilibrium lies to the right - the less readily
the metal loses electrons, and the more readily its ions pick them up again.…read more

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Iron (II) ions are the reducing agent:
a) Reduction of vanadium (V) to vanadium (IV). The reactants are VO2+ and Fe2+, so the first equation must be
reversed:
VO2+ + 2H+ + e- VO2+ + H2O E= +1.00V
Fe2+ (aq) Fe3+ (aq) + e- E= -0.77V
The overall equation is therefore:
VO2+ + 2H+ + Fe2+ VO2+ + H2O + Fe3+ (aq) Ecell = +1.00V + (-0.77V) = +0.23V +ve so feasible.
b) Subsequent reduction to vanadium (III).…read more

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Absorption within a solid.
Adsorption & absorption allow a lot of hydrogen to be stored at room temperature, but current
absorbers and absorbers have limited lifetimes.
Hydrogen has to be manufactured; e.g. by electrolysis of water but this requires energy.
Fuel cells have limited lifetimes and need to be replaced regularly whilst having high disposal costs.
Toxic chemicals are used in the manufacture of fuel cells.
32.…read more

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