10. Simple Cells and batteries
(A) Introduction
How a simple cell can be
used as a battery is explained, using the different reactivities of two
metal strips.
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
cell.
These revision notes on how simple cells and batteries
work
should prove useful for the new AQA chemistry, Edexcel chemistry & OCR
chemistry GCSE (9–1, 9-5 and 5-1) science courses.
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In electrolysis,
electrical energy is taken in (endothermic) to enforce the oxidation and
reduction to produce the products at the electrodes.
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The chemistry of simple
voltaic
cells or batteries is in principle the opposite of electrolysis.
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Inside an electrochemical cell or battery are
chemicals
that react together to produce electricity.
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The reactants constitute a
supply of chemical potential energy to be converted into electrical energy.
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A cell will produce a voltage until one
of the reactants is all used up.
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An oxidation-reduction (redox) reactions occurs
at electrodes to produce products and energy is given out because it is an exothermic
reaction,
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A
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
electrolyte.
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The electrolyte is a solution of
charged particles - ions, that can carry an electric current - can be a
salt solution of dilute acid.
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The external wire and voltmeter
completes the circuit - as in physics!
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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.
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The arrangement is shown in the simple
diagram of simple cell (right)
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If you connect several cells together in
series, the voltage is increased.
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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.
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You need is a
solution of charged positive and negative particles called ions e.g.
sodium Na+,
chloride Cl–, hydrogen H+, sulphate SO42–
in the electrolyte solution
etc.
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The greater the difference in
reactivity of the two metals, the bigger the cell voltage produced.
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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.
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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!
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You can predict the potential difference (p.d. in volts) by
subtracting one metal potential from another to give the theoretical cell
voltage.
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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
TOP OF PAGE
and sub-index
(B) A simple cell experiment to investigate
the effects of using different pairs of metals
TOP OF PAGE
and sub-index
(C) An early practical battery cell
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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.
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e.g. connecting strips of zinc and copper (plus
voltmeter) and placing in the electrolyte of zinc sulfate, .
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This simple cell system 'sort of' works
for a few minutes and then the voltage drops away..
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So this set-up is of not practical
enough for any use!
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It won't work effectively as a battery
with just one electrolyte.
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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.
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The diagram below explains the chemistry
behind one of the first practical battery systems.
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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.
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This 'voltaic 'or
galvanic' electrochemical cell uses a half–cell
of copper metal dipped in copper(II) sulphate,
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and in electrical
contact with another half–cell of zinc metal dipped in zinc sulphate solution.
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The zinc is the more
reactive, and is the negative electrode, releasing electrons because
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The less reactive
metal copper, is the positive electrode, and acts as an inert electrode.
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Instead copper(II) ions gain electrons
from the negative electrode through the external wire connection and are
reduced to copper metal ..
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the copper
ions are reduced to copper atoms:
Cu2+(aq) + 2e–
===> Cu(s)
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Overall the reactions is:
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Zn(s) + CuSO4(aq)
===> ZnSO4(aq) + Cu(s)
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or
ionically: Zn(s) + Cu2+(aq)
===> Zn2+(aq) + Cu(s)
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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.
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The theoretical p.d. created is 0.34
- (-0.76) = 1.10 V.
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The electrode potential of each
metal is a measure of its chemical reactivity - the more negative or
less positive, the more reactive the metal.
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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.
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In a working Daniel cell two
salt solutions are separated by a porous barrier that ions can diffuse through to
complete the electrical circuit.
TOP OF PAGE
and sub-index
(D)
More on investigation experiments and how to predict the
cell voltage
(E)
Practical batteries for commercial and domestic use - rechargeable and
non-rechargeable
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The simple cells described above do not
make a
satisfactory 'battery' for producing even a small continuous d.c. current.
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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.
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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.
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In non-rechargeable cells and batteries
the chemical reactions must stop when one of the reactants has been used up.
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You can't produce electricity if one of
the reactants is no longer present!
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Its all changed to the 'product' and
there is no longer any chemical potential energy to be transferred as useful
work - electrical energy.
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The common zinc-carbon and acid paste
battery comes into this category, so don't try and recharge it!
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AND most alkaline batteries are
non-rechargeable too.
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See also 11.
Fuel Cells e.g. the hydrogen - oxygen fuel cell
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Electrolysis and cell-battery theory-examples
for Advanced Level Chemistry Students
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