Doc
Brown's A Level Chemistry Revision Notes
Theoretical–Physical
Advanced Level
Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 7.1
7.1 Half cell
equilibria and half–cell reactions and electrode potential
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.
Chemical Equilibrium Notes Part 7 Index
7.1
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).
-
Mn+(aq)
+ ne–
M(s)
-
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.
-
The same
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
situations e.g.
-
(ii)
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.
-
MnO4–(aq)
+ 8H+(aq) + 5e–
Mn2+(aq)
+ 4H2O(l)
-
Interchange of
manganate(VII) and manganese(II) ions in acid solution, +7 and +2
ox. states.
-
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
Fig.1 below
comparing the relative positions of the metal–metal ion equilibrium
for copper and zinc.
-
Fig.1
-
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
–'s
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
solution. 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.
-
Quite simply,
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
-
Zn(s)
==> Zn2+(aq)
+ 2e– and
Cu2+(aq)
+ 2e– ==> Cu(s)
-
and the overall
reaction is Zn(s)
+ Cu2+(aq) ==> Zn2+(aq) + Cu(s)
-
This overall
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.
Chemical Equilibrium Notes Part 7 Index
TOP OF PAGE
|
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
Website content © Dr
Phil Brown 2000+. All copyrights reserved on revision notes, images,
quizzes, worksheets etc. Copying of website material is NOT
permitted. Exam revision summaries & references to science course specifications
are unofficial. |
Chemical Equilibrium Notes Part 7 Index
|