10. Simple Cells and batteries
How a simple cell can be
used as a battery is explained, using the different reactivities of two
How can you make a simple battery, how can you use a simple
cell to investigate the reactivity series of metals.
What is the
difference between rechargeable and non-rechargeable batteries, how to
make a simple copper-zinc cell, how to make a simple copper-magnesium
These revision notes on how simple cells and batteries
should prove useful for the new AQA chemistry, Edexcel chemistry & OCR
chemistry GCSE (9–1, 9-5 and 5-1) science courses.
electrical energy is taken in (endothermic) to enforce the oxidation and
reduction to produce the products at the electrodes.
The chemistry of simple
cells or batteries is in principle the opposite of electrolysis.
Inside an electrochemical cell or battery are
that react together to produce electricity.
The reactants constitute a
supply of chemical potential energy to be converted into electrical energy.
A cell will produce a voltage until one
of the reactants is all used up.
An oxidation-reduction (redox) reactions occurs
at electrodes to produce products and energy is given out because it is an exothermic
simple electrochemical cell can be
made by dipping two different pieces of metal (must
be of different reactivity - different potential), connected by a wire, into
a solution of ions e.g. a salt or dilute acid which will act as an
The electrolyte is a solution of
charged particles - ions, that can carry an electric current - can be a
salt solution of dilute acid.
The external wire and voltmeter
completes the circuit - as in physics!
The two pieces of metal can be held in
crocodile clips and acts as electrodes - the electrical contacts with the
electrolyte solution - at least one must react, one may be inert, but they
both usually react as part of the electrochemical cell chemistry.
The arrangement is shown in the simple
diagram of simple cell (right)
If you connect several cells together in
series, the voltage is increased.
If the metals have different
reactivities, then an electrical current is generated as long as the circuit
is complete as illustrated above on the right.
You need is a
solution of charged positive and negative particles called ions e.g.
chloride Cl–, hydrogen H+, sulphate SO42–
in the electrolyte solution
The greater the difference in
reactivity of the two metals, the bigger the cell voltage produced.
If you use the same metal for
both strips, their chemical potentials 'cancel' each other out, so no potential difference (voltage
= 0 V)
so no current of electrical energy.
If you connect several cells together,
identical or different, you can add up the individual cell voltages
to give the total p.d. in volts AND
you might light up a bulb! having made a crude battery!
You can predict the potential difference (p.d. in volts) by
subtracting one metal potential from another to give the theoretical cell
Examples of how
to think, i.e. predict the voltage of a very simple cell.
Ignore the highly reactive metals like, potassium, sodium
and calcium - impractical because they rapidly react with water, but lots of
other pairs can be used in simple school experiments e.g.
Pairing magnesium with copper will give a potential difference
of 2.69 V (+0.34 - - 2.35 V theoretically!), one of the biggest voltages
possible from the list on the right.
Pairing iron and tin will give a theoretical p.d. of 0.30 V, one
of the lowest possible from the list (-0.15 - -0.45 V).
See other pages for the
full chemistry of the reactivity series of metals
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(B) A simple cell experiment to investigate
the effects of using different pairs of metals
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(C) An early practical battery cell
As we have seen above, the simplest cell to generate electricity can be made
by dipping two externally connected pieces of different metals into an
electrolyte solution of a non-reactive salt.
e.g. connecting strips of zinc and copper (plus
voltmeter) and placing in the electrolyte of zinc sulfate, .
This simple cell system 'sort of' works
for a few minutes and then the voltage drops away..
So this set-up is of not practical
enough for any use!
It won't work effectively as a battery
with just one electrolyte.
SO, you do need something a bit more
sophisticated than this simple cell and the actual set-up in
principle for the Daniel Cell is shown below, one of the first simple, but effective,
batteries used in laboratories as an electrical supply for other experiments.
The diagram below explains the chemistry
behind one of the first practical battery systems.
The 'fuel' is effectively zinc metal and
copper(II) sulfate solution which get consumed when the battery is working
to generate a constant stream of d.c. electrical current.
This 'voltaic 'or
galvanic' electrochemical cell uses a half–cell
of copper metal dipped in copper(II) sulphate,
and in electrical
contact with another half–cell of zinc metal dipped in zinc sulphate solution.
The zinc is the more
reactive, and is the negative electrode, releasing electrons because
The less reactive
metal copper, is the positive electrode, and acts as an inert electrode.
Instead copper(II) ions gain electrons
from the negative electrode through the external wire connection and are
reduced to copper metal ..
ions are reduced to copper atoms:
Cu2+(aq) + 2e–
Overall the reactions is:
Zn(s) + CuSO4(aq)
===> ZnSO4(aq) + Cu(s)
ionically: Zn(s) + Cu2+(aq)
===> Zn2+(aq) + Cu(s)
It is an exothermic reaction,
BUT here, there is no temperature rise, because the energy is released
as electrical energy carried by the flow of electrons.
The theoretical p.d. created is 0.34
- (-0.76) = 1.10 V.
The electrode potential of each
metal is a measure of its chemical reactivity - the more negative or
less positive, the more reactive the metal.
The overall reaction
is therefore the same as displacement reaction, and it is a redox
reaction involving electron transfer and the movement of the electrons
through the external wire to the bulb or voltmeter etc. forms the
working electric current.
In a working Daniel cell two
salt solutions are separated by a porous barrier that ions can diffuse through to
complete the electrical circuit.
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More on investigation experiments and how to predict the
Practical batteries for commercial and domestic use - rechargeable and
The simple cells described above do not
satisfactory 'battery' for producing even a small continuous d.c. current.
Cells or batteries
are useful and convenient portable sources of energy for torches,
radios, shavers and other gadgets BUT they are
expensive compared to what you pay for 'mains' electricity.
With rechargeable cells and batteries, it is
possible to input electrical energy (via a charger) and reverse the
chemistry that produced the electricity in the first place.
In non-rechargeable cells and batteries
the chemical reactions must stop when one of the reactants has been used up.
You can't produce electricity if one of
the reactants is no longer present!
Its all changed to the 'product' and
there is no longer any chemical potential energy to be transferred as useful
work - electrical energy.
The common zinc-carbon and acid paste
battery comes into this category, so don't try and recharge it!
AND most alkaline batteries are
See also 11.
Fuel Cells e.g. the hydrogen - oxygen fuel cell
Electrolysis and cell-battery theory-examples
for Advanced Level Chemistry Students