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Standard Grade Chemistry
Summary Notes
Topic 10: Making Electricity
General
Learning Outcomes
 In a battery, electricity comes from a chemical reaction.
 Electricity passing along metal wires is a flow of electrons.
 Batteries need to be replaced as chemicals are used up in the reaction.
 Some batteries are rechargeable, e.g. the lead-acid battery.
 Ammonium chloride in a cell is an example of an electrolyte.
 The purpose of an electrolyte is to complete the circuit.
 Electricity can be produced by connecting different metals together (with an
electrolyte) to form a cell.
 The voltage between different pairs of metals varies and that this leads to the
electrochemical series.
 Use an electrochemical series showing metals only to explain why displacement
reactions occur and describe the experimental observations.
 Electricity can be produced in a cell by connecting two different metals in
solutions of their metal ions.
 The purpose of the ‘ion-bridge’ is to complete the circuit.
 Compare batteries and mains electricity sources in relation to ease of transport,
safety, costs and uses of finite resources.
Credit
Learning Outcomes
 Use an electrochemical series to predict whether or not displacement reactions
occur; and predict the experimental observations.
 The reactions with metals and acid can establish the position of hydrogen in an
electrochemical series.
 Electricity can be produced in a cell when at least one of the half-cells does not
involve metal ions.
 The movement of ions in the ion bridge is to provide ions to complete the circuit.
 A metal element reacting to form a compound is an example of oxidation.
 Oxidation is the loss of electrons by any reactant in a reaction.
 A compound reacting to form a metal element is an example of reduction.
 Reduction is a gain of electrons by a reactant in any reaction.
 In a redox reaction, oxidation and reduction reactions go on together,
 You must be able to apply the terms oxidation and reduction to more complex
ion– electron equations e.g. SO42-/SO32-
Making Electricity
General
You probably use several things every day which need batteries e.g. radio, MP3 player,
calculator, watch. There are several different kinds of battery, but they all do the
same thing – they provide us with a convenient and portable source of electricity. This
topic allows us to study how electricity is made from the chemicals in a battery.
Chemical reactions, as you know, often result in changes in appearance or temperature.
Chemical reactions also take place when batteries produce electricity.
Apparatus which produces electricity from chemical reactions is usually called a cell.
A battery is really two or more cells joined together.
Cells
General
The lead-acid ‘battery’ is an example of a wet cell. These are not very suitable for
many purposes.
The batteries we use in most everyday appliances are dry cells, where the liquid or
solution of the wet cell is replaced by a thick paste. This paste is often in the form of
ammonium chloride.
All batteries change chemical energy into electrical energy (electricity). This
electricity can pass along metal wires as a flow of electrons.
All cells, whether wet or dry, become run down. This usually requires the batteries
having to be replaced as the chemicals which produced the electricity have been used
up.
Some kinds of cell can be recharged and used again. These cells are called
rechargeable cells. The lead-acid cell is an example of a rechargeable cell.
When a cell is charged, electrical energy is changed into chemical energy in the cell
because of the chemical reactions which take place.
When the cell is discharged, different chemical reactions take place, and the
chemical energy stored in the cell is changed into electrical energy.
Joining Metals
General
Cells produce electricity because of chemical reactions. So what happens in a cell?
In a zinc/copper cell, the ammeter shows that a current is flowing from the zinc strip
to the copper strip. Zinc atoms lose electrons and change into ions. This reaction can
be represented by this ion-electron equation.
Zn  Zn2+ + 2eThe electrons produced flow through the external wires and the ammeter to the
copper strip. The salt solution acts as an electrolyte and completes the electrical
circuit.
A cell can be made by connecting two different metals together in an electrolyte.
When this is done electricity is produced because one metal reacts and produces
electrons which flow to the other metal as an electric current. This current can be
measured by an ammeter.
However, most batteries are described by their voltage and not their current.
Voltage is the force which the cell produces to push the electrons around the
circuit. The voltage of a cell can be measured by replacing the ammeter with a
voltmeter.
A simple cell can also be made by sticking two different strips of metal into a citrus
fruit such as a lemon and joining the metals with a wire.
Electrochemical Series
General/Credit
Two metals connected in a cell produce a voltage. The size of the voltage depends on
the metals used.
Different voltages are obtained when different metals are joined to copper in a cell.
When the metals are placed in decreasing order of voltage, this leads to the
Electrochemical Series. The Electrochemical Series is found on page 7 of your Data
Booklet.
The higher a metal in the Electrochemical Series the more easily it loses electrons to
form ions.
The Electrochemical Series allows us to understand things about a cell.
 Electrons flow through the external wire from the metal higher in the
Electrochemical Series to the metal lower.
 Metals which are far apart in the Electrochemical Series produce higher
voltages than metals which are close together.
Some metals react with the hydrogen ions of acids to form hydrogen gas. In fact the
metal atoms are displacing the hydrogen ions.
Mg (s)  Mg2+(aq) + 2e2H+(aq) + 2e-  H2(g)
These reactions allow us to work out the position of hydrogen in the Electrochemical
Series.
Metals which react with acids lie above hydrogen in the Electrochemical Series; metals
which do not react with acids lie below hydrogen in the Electrochemical Series.
Displacement Reactions
General/Credit
The Electrochemical Series placed metals in order of their ability to convert their
atoms into ions by losing electrons.
A displacement reaction happens because one metal is better at losing electrons than
the other. The metal higher in the series will displace lower metals from solutions of
their ions.
When powdered zinc is added to a solution containing Cu2+ ions, the zinc atoms ‘push’
the copper ions out of solution by changing them to copper atoms.
Cu2+(aq) + 2e-  Cu(s)
At the same time the zinc atoms change into zinc ions.
Zn (s)  Zn2+(aq) + 2eThis provides the electrons the copper ions need. This is called a displacement
reaction.
The overall reaction is
Zn (s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
The rule for displacement reactions is :
A metal higher in the Electrochemical Series can displace a metal lower in the
series from a solution of its ions.
We can use this rule and the Electrochemical Series to predict if a reaction will occur.
If you have enough information, you could also predict what you will see when a
displacement reaction occurs.
More Cells
Displacement reactions can also be set up as a cell.
General/Credit
Again a displacement reaction occurs. The zinc is better at losing electrons than copper
(it is higher in the Electrochemical Series), so the zinc atoms change into zinc ions.
Zn (s)  Zn2+(aq) + 2eThe electrons formed here flow through the external wire and the voltmeter and are
taken up by the copper ions in the solution to make copper atoms.
Cu2+(aq) + 2e-  Cu(s)
i.e. the electric current flows from the zinc to the copper
The rule is :- the electrons flow from the metal higher in the Electrochemical Series
through the wire to the metal lower in the series.
The movement of ions in the ‘ion bridge’ allows the ion bridge to complete the
electrical circuit.
Cells Without Metals
General/Credit
It is possible to produce electricity by joining non-metal half cells with other half-cells.
The ion-electron equations for the half-cells above are
Zn (s)  Zn2+(aq) + 2eI2(aq) + 2e-  2I-(aq)
Again the reaction which gives out the electrons lies higher in the Electrochemical
Series.
Redox
Credit
In the displacement reaction involving zinc being added to a solution containing Cu2+
ions the ion-electron equations which represent the reactions taking place are
Zn (s)  Zn2+(aq) + 2eCu2+(aq) + 2e-  Cu(s)
The reaction Zn (s)  Zn2+(aq) + 2e- is an example of an oxidation reaction because it
involves the loss of electrons by a reactant.
The reaction Cu2+(aq) + 2e-  Cu(s) is an example of a reduction reaction because it
involves the gain of electrons by a reactant.
The overall reaction Zn (s) + Cu2+(aq)  Zn2+(aq) + Cu(s) is a combination of the
reduction and oxidation steps and is called a redox reaction.
The terms oxidation or reduction can be applied to any ion-electron equation.
In an oxidation reaction the electrons always lie on the right hand side of the ion –
electron equation.
SO32- + 2H+  SO42- + H2O + 2e-
In a reduction reaction the electrons always lie on the left hand side of the ionelectron equation.
Fe3+ + e-  Fe2+
The memory aid
Oxidation
Is
Loss of electrons
Reduction
Is
Gain of electrons
can be useful in helping to remember the meaning of the terms oxidation and reduction.
Oxidation and reduction can also be defined in more general terms as follows :
A metal element reacting to form a compound is an example of oxidation
A compound reacting to form a metal element is an example of reduction.
e.g. the conversion of iron ore to iron metal in a blast furnace is often described as the
reduction of iron ore (also TOPIC 11)