Brown's Chemistry Advanced A Level Notes - Theoretical–Physical
Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 7
Part 7.1 Half cell
equilibria and half–cell reactions and electrode potential
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The equilibrium between a metal and its
cation solution is first considered to introduce the idea of a half–cell
potential and examples of half–cell reactions given.
Examination of some redox equilibrium situations
(i) A metal
atom and metal ion
When a piece of
metal (M) is dipped into an aqueous solution of its ions (Mn+(aq))
an equilibrium exists between the two chemical states of the metal,
and importantly, in different
oxidation states (0 for the element and +n for the cation).
n is the
numerical value of
the positive charge on the cation and the number of electrons involved in
the oxidation state change.
This is an example
of a half–equation which represents the particular
oxidation (electron loss) or reduction (electron gain) of an atom/ion
related by electron transfer connecting two different oxidation states
of the element.
idea of an equilibrium between two species of the same element in
two different oxidation states expressed as a half–cell equation also applies to many other
A non–metal atom/molecule and associated anion e.g.
(iii) Two different cations of the
same metal e.g.
(iv) An anionic and
cationic species of the same metal e.g.
+ 8H+(aq) + 5e–
manganate(VII) and manganese(II) ions in acid solution, +7 and +2
Note that half–cell equations are
often, but not always, presented as
a reduction (electron gain). For
examples (ii) to (iv), a chemically inert platinum electrode
is dipped into the solution when used in electrochemical cell
construction (see later).
For an example of
situation (i) and its consequences consider the diagram
comparing the relative positions of the metal–metal ion equilibrium
for copper and zinc.
In the case of the
copper atom/copper(II) ion
equilibrium, you get a positive charge on a piece of copper when it is
dipped into an aqueous copper(II) ion solution.
This is due to an
electron deficiency on the copper because the right–hand side of the
equilibrium, the solid copper, is favoured, and copper(II) ions remove
negative electrons to form copper atoms on the surface (reduction
process). The +'s represent the electron deficiency in the copper metal,
which in cells would make it the positive pole, and the
represent surplus anion charge in the solution.
However, in the
case of the zinc atom/zinc(II) ion equilibrium, you get a negative charge
on a piece of zinc when it is dipped into an aqueous zinc(II) ion
This is due to an electron surplus on the zinc metal because the
left–hand side of the equilibrium, the aqueous zinc(II) ion, is
favoured, and zinc atoms lose negative electrons to form zinc(II) ions
in the solution (oxidation process). For Zn/Zn2+:
The +'s represent the extra zinc ion charge in solution, the –'s
represent the surplus electrons in the zinc metal which in cells would
be the negative pole.
thermodynamically, zinc atoms have a much greater tendency or
potential to lose electrons than copper atoms and copper(II) ions
have a much greater tendency or potential to accept electrons
than zinc ions.
Therefore the two half–cell reactions would be
+ 2e– and
+ 2e– ==> Cu(s)
and the overall
reaction is Zn(s)
+ Cu2+(aq) ==> Zn2+(aq) + Cu(s)
reaction can be accomplished in two very different ways.
(a) by adding zinc
metal to blue copper(II) sulphate solution when the blue colour fades
and a red–brown deposit of copper forms. The reaction is exothermic
which can be observed as a significant temperature rise if zinc powder
is added to 1M copper(II) sulphate solution.
(b) By setting up
a simple chemical cell ('battery') in which the two half–reactions are kept physically separated. To complete the electrical
circuit, the zinc and copper(II) ion solutions are connected by a salt
bridge (a supported ion solution). External wires connect strips of zinc and
copper metal to a voltmeter or other device. The energy release from
the exothermic reaction is now in the form of electrical energy. This, so–called Daniel cell is described in detail in the next section
7.2 Simple cells notation and construction.
Advanced Equilibrium Chemistry Notes Part 1. Equilibrium,
Le Chatelier's Principle–rules
* Part 2. Kc and Kp equilibrium expressions and
calculations * Part 3.
Equilibria and industrial processes * Part 4
Partition between two
phases, solubility product Ksp, common ion effect,
ion–exchange systems *
Part 5. pH, weak–strong acid–base theory and
calculations * Part 6. Salt hydrolysis,
acid–base titrations–indicators, pH curves and buffers * Part 7.
Redox equilibria, half–cell electrode potentials,
electrolysis and electrochemical series
pressure, boiling point and intermolecular forces watch out for sub-indexes
to multiple sections or pages
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Chemical Equilibrium Notes Part 7 Index