Cells

Chemical reactions can produce electricity.

A cell (often referred to as a battery) contains chemicals which react to make electricity.

The energy change is:

In dry cells, the chemicals are used up and the cells then have to be replaced.

In a simple dry cell, the chemicals are shown in the diagram below.

The paste containing ammonium chloride is the electrolyte needed to complete the circuit.

 
Rechargeable cells.

In rechargeable cells the chemicals are not used up and can be regenerated by recharging the cell.

 
Energy Changes in Cells

Using the cell (Discharging):

Charging the cell:

 
Lead-sulphuric acid battery (Car battery)

This is charged using electricity and a brown coating appears on the positive lead plate. When the charger is removed and replaced by a bulb, the bulb lights. Gradually the bulb dims as the chemical energy in the brown coating is used up and the cell is discharged. The cell now has to be recharged.

 
The electrochemical series

Electricity is produced when two different metals are dipped in an electrolyte and connected together with a wire.

There is a flow of electrons in the wire from one metal to the other, and ions move through the electrolyte.

The data book (Page 7) gives a list of metals in a table called the electrochemical series.

When two different metals are joined together as shown in the above cell, electrons always flow through the wire from the metal higher in the electrochemical series to the metal lower in the series e.g. from Zinc to Copper.

If a voltmeter is used in the wire between the two metals, a different voltage is obtained when different metals are used.

The voltage values can be used to put the metals into the order shown in the electrochemical series

e.g. a larger voltage is obtained when magnesium is connected to copper than when zinc is connected to copper, and magnesium is placed above zinc in the electrochemical series.

The reactions of metals with acids (which contain the hydrogen ion) can be used to place hydrogen in the reactivity series.

You can try creating your own cells in the Virtual Lab.

 
Electricity from different metals in solutions of their own ions

Electricity can be produced in a cell by connecting two different metals in solutions of their metal ions.

The ammeter will show a flow of electrons through the wire between the two metals always from the metal higher in the electrochemical series to the one lower in the series.

If the ion bridge (or salt bridge) is removed, the current stops flowing through the ammeter.

This happens because the ion bridge is needed to complete the circuit.

By separating the reaction in this way, the cell is divided into two half cells.

In the cell above the two reactions are:

They explain how electrons made at zinc travel through the wire to copper where they join onto copper(II) ions to make copper metal.

This is similar to the displacement reaction (see notes later in this page) that occurs when zinc metal is added to a solution which contains copper(II) ions e.g. copper(II) sulphate.

The ion bridge completes the circuit by allowing ions to move through it.

 
Electricity from cells where at least one of the half cells does not involve a metal

The electrochemical series in the data book has some reactions that involve non-metals e.g.

The diagram shows that electrons flow from zinc metal to the carbon rod and the reactions are shown below

 
Displacement reactions

When grey magnesium metal is added to blue copper(II)sulphate solution, brown copper metal is made and the solution becomes colourless.

In this reaction magnesium has pushed copper out of copper(II) sulphate solution as shown in the following equation.

A displacement reaction will happen when a metal higher in the electrochemical series is added to a solution containing a metal lower in the electrochemical series
The metal which is lower in the series is displaced or pushed out of its compound which is in solution.

When copper metal is added to colourless silver(I) nitrate solution, the colour changes will be the loss of the brown copper, the appearance of silver metal and the solution turning blue as the copper dissolves.

Copper has displaced silver from a compound containing silver ions.

When the nitrate spectator ions are removed, the reaction is between copper atoms and silver(I) ions.

The reactions of metals with acids (which contain the hydrogen ion), used to place hydrogen in the reactivity series is a further example of a displacement reaction.

• In this reaction, hydrogen is displaced from an acid by a metal.

• As with other displacements, only metals above hydrogen can cause this to happen.
  e.g. magnesium will react but copper will not.

• As lead will displace hydrogen from an acid, hydrogen is placed between lead and copper in the electrochemical series.

 
Oxidation

An oxidation reaction is when a reactant loses electrons in a chemical reaction.

An example of this is when magnesium metal is used to displace zinc from zinc sulphate solution. The magnesium loses electrons during the displacement reaction to form magnesium ions. This means the magnesium has been oxidised.

 
Reduction

An reduction reaction is when a reactant gains electrons in a chemical reaction.

An example of this is when copper ions are displaced from solution by a more reactive metal. The copper ions gain electrons during the displacement reaction to form copper metal. This means the copper ions have been reduced.

 
Redox Reactions

When both oxidation and reduction occur together, the complete reaction is called a REDOX reaction.

If magnesium metal was added to copper sulphate solution, the magnesium metal would be oxidised, while the copper ions were being reduced. This is an example of a redox reaction.

Both of the equations (known as ion-electron equations) can be written and combined, by eliminating the electrons involved, to produce a redox equation:

Example 1: Magnesium metal displacing copper metal from copper(II) sulphate solution

Example 2: aluminium metal displacing silver metal from silver(I) nitrate solution

Example 3: zinc metal displacing iron metal from iron(III) sulphate solution

 
The use of metals

Metals need to be recycled because they will not last forever.

How long a metal can last can be found from the previous bar graphs.
Metals are not finite.

Large quantities of metals are thrown away and the need for recycling is shown below:

 
Reactions of metals

a) With water

Metal + Water ----------> Metal hydroxide + Hydrogen

Potassium reacts vigorously, sodium very quickly, calcium quickly and magnesium slowly.

 

Potassium	+	Water	---------->	Potassium hydroxide	+	Hydrogen
K	+	H2O	---------->	KOH	+	H2
Sodium	+	Water	---------->	Sodium hydroxide	+	Hydrogen
Na	+	H2O	---------->	NaOH	+	H2
Calcium	+	Water	---------->	Calcium hydroxide	+	Hydrogen
Ca	+	H2O	---------->	Ca(OH)2	+	H2
Magnesium	+	Water	---------->	Magnesium hydroxide	+	Hydrogen
Mg	+	H2O	---------->	Mg(OH)2	+	H2

The order of metals reacting with water (most reactive first) is :

Potassium, sodium, calcium and magnesium.

b) Metal reacting with Acid

Metal	+	Acid	---------->	Salt	+	Hydrogen
Magnesium	+	Hydrochloric acid	---------->	Magnesium chloride	+	Hydrogen
Mg	+	HCl	---------->	MgCl2	+	H2

The order of reaction can be obtained by observation of the rate at which gas is given off.

All metals above hydrogen in the electrochemical series react with acids to displace hydrogen gas.

The order of metals reacting with acid (most reactive first) is

Magnesium, aluminium, zinc, iron, tin and lead.

Copper, mercury, silver and gold do not react, while potassium, sodium and calcium are too reactive to add to acid.

c) Metals reacting with oxygen

Oxygen can be made by heating potassium permanganate in a test tube and allowing the gas to pass through the preheated metal as shown.

A glow spreads through the metal (exothermic reaction), and the speed is related to the relative activity of the metal.

Metal	+	Oxygen	---------->	Metal oxide
Magnesium	+	Oxygen	---------->	Magnesium oxide
Mg	+	O2	---------->	MgO

The order of metals reacting with oxygen (most reactive first) is

Magnesium, aluminium, zinc, iron, tin, lead, copper and mercury.

Silver and gold do not react, while potassium, sodium and calcium are too reactive to react with oxygen in this way.

These reactions give an indication of the reactivity of the metal and are summarised below:

 
Metal ores

Ores are naturally-occuring compounds of metals from which metals can be extracted.
The three main types of ore are metal carbonates, metal oxide and metal sulphides.

 
Elements on earth

Metals such as gold and silver occur uncombined on earth because they are unreactive and because of this these elements were among the first to be discovered.

Other metals, such as those in the table above are found in compounds and have to be extracted.
You can quickly test your knowledge of the above information.

 
Extraction of metals from ores

The demand for metals is high and methods are now available to extract all metals from their ores.
Methods using carbon (coke) are cheaper and have been used longer than methods which use electricity.

 
Methods of extraction

a) Heating metal oxides

Silver oxide	---------->	Silver	+	Oxygen
Ag2O	---------->	Ag	+	O2

Few metals can be obtained in this way

b) Heating metal oxides with carbon

The main reaction is :

Metal oxide	+	Carbon	---------->	Metal	+	Carbon dioxide
Iron oxide	+	Carbon	---------->	Iron	+	Carbon dioxide
Fe2O3	+	C	---------->	Fe	+	CO2

This method is used to extract metals below aluminium in the reactivity series.

c) Heating with carbon monoxide

Iron is extracted from its ore in the blast furnace by heating with carbon (coke) in the presence of air.

At the bottom of the furnace the reaction makes carbon dioxide (Zone 1)

C + O₂ ----------> CO₂

Higher up, the carbon dioxide reacts with carbon to make carbon monoxide (Zone 2)

CO₂ + C ----------> CO

Further up the carbon monoxide reacts with iron oxide to make iron and carbon dioxide. (Zone 3)

The formation of a metal from a compound is known as reduction.

Fe₂O₃ + CO ----------> Fe + CO₂

d) Using electricity

Electricity can be used to split ionic compounds into their elements in a process called electrolysis.
The method is used to extract reactive metals above zinc in the reactivity series.

A large electric current is passed through the molten compound, and metal appears at the negative electrode. At the negative electrode, reduction is taking place (oxidation is taking place at the positive electrode).

Aluminium oxide ----------> Aluminium + Oxygen

The method used to extract a metal depends on the reactivity of the metal.

• The more reactive the metal, the more difficult it is to extract.
• The less reactive the metal, the easier it is to extract.

Electricity is used to extract the most reactive metals such as potassium, sodium, calcium, magnesium and aluminium.

 
Corrosion

Corrosion is a chemical reaction which involves the surface of a metal changing from an element to a compound. This natural change of metals into compounds is very costly.

 
Speed of Corrosion

Most metals corrode, but the speed at which they corrode is related to the chemical activity series. Metals high in the reactivity series (such as potassium) corrode very quickly while those lower in the series corrode much more slowly (such as silver and gold)

 
Rusting

Rusting is the special name given to the corrosion of iron
As iron, in the form of steel, is the most commonly used metal in the world, the corrosion of iron is important.

 
The Cause of Rusting

The experiment to the right shows that oxygen and water are both needed for rusting to occur.

Further proof that oxygen is needed is seen in the experiment on the left.

As about 80% of the air in the cylinder is left after rusting - it means that oxygen is used up during rusting and water rises to take its place.

 
Detecting Rusting

• The typical brown colour of rust is the end result of rusting.
• The first stage of the rusting process can be detected by ferroxyl indicator.
• As a blue colour is only made with iron(II) chloride solution, and with a rusting nail, it proves that iron(II) ions (Fe²⁺) are made during rusting.
• Fe²⁺ are produced when iron atoms lose 2 electrons.

Corrosion is an example of oxidation because it involves a loss of electrons.

The rusting process continues when iron(II) ions lose another electron to form iron(III) ions.

The iron(III) ions can be shown using a colourless solution of ammonium thiocyanate. The solution will turn blood-red.

The electrons 'lost' when iron is oxidised during the rusting process are accepted by water and oxygen (the requirements for rusting) and are shown in the following equation (this equation is given in the data booklet).

These hydroxide ions can also be detected by ferroxy indicator, which turns pink to indicate their presence.

As iron atoms lose electrons in rusting and oxygen/water molecules gain these electrons, rusting is described as a Redox reaction.

 
Increasing the rate of corrosion

Corrosion requires an electrolyte such as dissolved salt or acid rain. Acid can speed up corrosion in two ways:

(a) by acting as an electrolyte

(b) by reacting with the metal.

Electrolytes increase the speed of rusting, and cars rust faster in winter when salt is spread on the roads.
You can quickly test your knowledge of the above information.

The electrolyte helps to carry ions away from the rusting iron and this speeds up the oxidation (corrosion).

 
Using a battery to prevent the rusting of iron

The battery causes the nail connected to the positive terminal to rust rapidly, but the nail connected to the negative terminal does not rust.

• Iron has to lose electrons in order to rust.
• The negative terminal of the battery is pushing electrons onto this nail and this prevents this nail from losing any electrons. This nail cannot rust.
• Ferroxyl indicator turns pink because hydroxide ions are made here.
• Electrons flowing to the nail stop rusting.
• The positive terminal of the battery is removing electrons from the nail connected to it. This nail rusts rapidly, changing into iron(II) ions which turn ferroxyl indicator blue.
• Electrons flowing from the nail increase rusting.

You can quickly test your knowledge of the above information.

 
Connecting different metals to iron

Metals that push electrons onto iron stop rusting, but metals that let electrons flow from iron increase the speed of rusting.

• Metals that push electrons onto iron are higher than iron in the reactivity series (such as magnesium)
• Metals that let electrons flow from iron are lower than iron in the reactivity series (such as copper)

 
Preventing Corrosion

A. Physical Protection

This is where a metal is given a coating to stop it coming in contact with air and water and thus prevents corrosion.

Methods available for physical protection

• painting e.g. the Forth Rail Bridge.
• greasing or oiling - protects moving parts of machinery.
• coating with plastic - dish drainers have a metal core and a plastic coating.
• coating with other metals such as tin, zinc, silver, gold.

Tin-plating - metals can be coated with other metals which are less likely to corrode. Food cans are steel cans dipped into molten tin giving a layer of tin.

Electroplating - e.g. chromium-plating of car bumpers and the silver-plating of cutlery are done using this process to give an attractive appearance which provides protection against corrosion.

 
B. Chemical Protection

Tin-plating works well provided the layer of tin remains unbroken. If the tin layer to scratched, the iron corrodes quickly because electrons travel to tin from iron.

Zinc-plating works well if the layer of zinc remains unbroken and also when scratched because then the zinc corrodes quickly and electrons are pushed onto the iron.
This is an example of sacrificial protection - where a more reactive metal is allowed to corrode in order to protect a less reactive metal.
You can quickly test your knowledge of the above information.

New words and their meanings

BATTERY - a device containing chemicals that react to produce electricity.

CELL - the correct term for devices called batteries.

ELECTROLYTE - a substance that contains ions which can move (either molten ionic compounds or ionic compounds in solution or a watery paste as in a dry cell).

RECHARGEABLE CELL - a cell in which electricity can regenerate chemicals to allow the cell to produce electricity many times.

ELECTROCHEMICAL SERIES - a list of reactions in the data book (page 7) which can be used to determine which reaction of a pair is better at pushing electrons onto the other reaction.

VOLTAGE - a measure of the push of electrons between two reactions.

ION BRIDGE - a link containing ions which completes a circuit by allowing ions to travel through it.

DISPLACEMENT REACTION - where one metal higher in the electrochemical series displaces (pushes out) another metal from a solution of the other metal.

OXIDATION - the loss of electrons by a reactant in any reaction.

REDUCTION - the gain of electrons by a reactant in any reaction.

REDOX REACTION - a reaction in which reduction and oxidation occur together.

ION-ELECTRON EQUATION - an equation that shows ions and electrons involved in a reduction, oxidation or redox reaction.

OILRIG - This will help you remember oxidation and reduction. Oxidation Is Loss, Reduction Is Gain (of electrons).

ORE - a naturally-occuring compound of metals.

ELECTROLYSIS - splitting a substance into its elements using electricity.

CORROSION - the changing of the surface of a metal from an element into a compound.

RUSTING - the special name for the corrosion of iron.

FERROXYL INDICATOR - turns blue in the presence of iron(II) ions and turns pink in the presence of hydroxide ions.

PHYSICAL PROTECTION - stopping corrosion by keeping out air and/or water

CHEMICAL PROTECTION - using more reactive metals to push electrons onto iron

GALVANISING - coating iron with a layer of zinc

TIN PLATING - coating iron with a layer of tin

ELECTROPLATING - coating a metal with a layer of another metal using electricity

SACRIFICIAL PROTECTION - where a more reactive metal sacrifices itself to protect the less reactive metal