Download Electrochemical Cells – Voltage (Electric potential)

Electrochemical Cells –
Voltage (Electric potential)
The half cells
Standard electrode potentials
Calculating voltages
Examples
Author: J R Reid
Half Cells
Previously we have learnt that every redox reaction has an
oxidation half and a reduction half
Cells also have two halves – (oxidation (left) and reduction
(right)). We call these half cells – strangely enough
Every half cell has a rated ability hold on to its electrons:
Some grab electrons very well e.g.
F2 → FCl2 → Cl-
Some find it very hard to hold their electrons e.g.
Li → Li+
Na → Na+
This ability is called a standard electrode potential and is
measured in volts
Standard Electrode Potential
A half cell doesn’t have any ability to give
or take electrons by itself – it needs to
have something else to give the electrons
to (or to take them from). In this case we
compare every half cell with its ability to
take electrons from a hydrogen half cell:
H 2 → H+
The hydrogen half cell is the oxidation
(left) half cell
This potential to push (or pull) electrons
is measured in volts (as mentioned earlier)
It is called a standard electrode potential
because it is measured under standard
conditions:
Temperature: 25°C
Concentrations: 1.0 molL-1
Atmospheric pressure: 1 atmosphere
(101.3kPa)
Calculating Voltages of Cells
Because all of the half cells potentials are relative to Hydrogen, we
can work out the difference in potential when we combine any two
half cells
Because each half cell has a different potential to give or take
electrons there will be a flow from the best giver to the best taker of
electrons
We use the following formula to calculate the difference in voltage
between two half cells:
E°(cell) = E°(RHE) – E° (LHE)
In other words the total Electrode potential for the cell = the
Electrode potential of the right hand electrode – the Electrode
potential of the left hand electrode
If the E°(cell) value is a positive voltage then the electrons are flowing
in the correct direction therefore we say that the reaction is
spontaneous (it happens without us having to add energy)
Examples of Standard
Electrode Potentials
Note that the potentials of the following reactions
are all for the oxidation reactions – this is because
they were all compared to the reduction half cell of
H+
MnO4-,Mn2+
Au3+/Au
Cl2/ClAg+/Ag
Fe3+,Fe2+
Cu2+/Cu
H+/H2
=
=
=
=
=
=
=
1.51V
1.50V
1.44V
0.80V
0.77V
0.34V
0.00V
H+/H2
Pb2+/Pb
Fe2+/Fe
Zn2+/Zn
Al3+/Al
Mg2+/Mg
Na+/Na
=
=
=
=
=
=
=
0.00V
-0.13V
-0.47V
-0.76V
-1.66V
-2.36V
-2.71V
Exercises: Calculating The
Voltage
MnO4-,Mn2+
Au3+ /Au
Cl2/ClFe3+,Fe2+
Cu2+/Cu
H+/H2
Fe2+/Fe
Zn2+/Zn
Al3+/Al
Mg2+/Mg
Na+/Na
=
=
=
=
=
=
=
=
=
=
=
1.51V
1.50V
1.44V
0.77V
0.34V
0.00V
-0.47V
-0.76V
-1.66V
-2.36V
-2.71V
Examples of Electrochemical
Cells – The Dry Cell
The dry cell consists of three main
parts:
The zinc case – which gets oxidised
The black paste which gets reduced
The carbon electrode
The black (oxidant) paste contains
two main chemicals which undergo
the following reaction:
MnO2 + NH4+ + H2O + e- → Mn(OH)3 +
NH3
The zinc half equation is:
Zn → Zn2+ + 2e-
Oxidant paste
Carbon anode
Zinc cathode
Example II – Lead Acid
Cells
A lead-acid battery supplies
electrical current through the
following reactions:
Lead oxidising
Pb → Pb2+ + 2eLead oxide being reduced
PbO2 + 4H+ + 2e→ Pb2+ + 2H2O
In each half equation the Pb2+ builds
up on the outside of the sheets as
PbSO4
Each cell generates about 2 volts so
6 cells are placed in series to create
a 12V battery
This type of cell is also
‘rechargeable’ - the reactions can be
reversed by running current through
the cell in the opposite direction to
what it normally travels
The diagram below
shows alternating lead
and lead dioxide sheets
in a “pool” of H2SO4
Exam Practice - 2008
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Exam Practice - 2007
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Exam Practice - 2006
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Exam Practice - 2005
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Exam Practice - 2004
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