This revision guide is designed to help you study for the

chemistry part of the IGCSE Coordinated Science course.

The guide contains everything that the syllabus says you

need you need to know, and nothing extra.

The material that is only covered in the supplementary

part of the course (which can be ignored by core

candidates) is highlighted in dashed boxes:

Some very useful websites to help you further your

understanding include:

•http://www.docbrown.info/ - whilst not the

prettiest site this contains a lot of very useful and

nicely explained information.

•http://www.bbc.co.uk/schools/gcsebitesize/scienc

e/ - well presented with many clear diagrams,

animations and quizzes. Can occasionally lack

depth.

•http://www.chemguide.co.uk/ - whilst mostly

targeted at A-Levels this site contains very detailed

CHEMISTRY REVISION GUIDE

for IGCSE Coordinated Science

CHEMISTRY REVISION GUIDE

for IGCSE Coordinated Science

Whilst this guide is intended to help with your revision, it

should not be your only revision. It is intended as a

starting point but only a starting point. You should make

sure that you also read your text books and use the

internet to supplement your study in conjunction with

your syllabus document.

Whilst this guide does contain the entire syllabus, it just

has the bare minimum and is not in itself sufficient for

those candidates aiming for the highest grades. If that is

you, you should make sure you read around a range of

sources to get a deeper knowledge and understanding.

targeted at A-Levels this site contains very detailed

information suitable for those looking to deepen

their knowledge and hit the highest grades.

Finally, remember revision is not just reading but should

be an active process and could involve:

•Making notes

•Condensing class notes

•Drawing Mind-maps

•Practicing past exam questions

•Making flashcards

The golden rule is that what makes you think makes you

learn.

Happy studying, Mr Field.

C1: THE ELEMENTS OF

CHEMISTRY

Atom: The smallest

particle of matter

An atom: Some atoms:

Molecule: A small

particle made from

more than one atom

bonded together

Molecules of an

element:

Molecules of a

compound:

Element: A

substance made of

only one type of

atom

A solid element: A gaseous element:

Compound: A

substance made

from two or more

A solid compound A gaseous

compound:

CHEMICAL FORMULAS

Formulas tell you the atoms

that make up a compound

Eg 1. H2O – two H, one O

Eg 2. C2H6O – two C, six H, one

O

Eg 3. Mg(OH)2 – one Mg, two

O, two H*

Eg 4. CH2(CH3)2 – three C, 8 H*

*In this case everything in

brackets is doubled

CHEMICAL EQUATIONS

•Show the reactants you start with and the products you make

•Must contain an arrow (�) NOT an equals sign (=)

•Must be balanced – same number of atoms on each side

Eg. CH4 + O2 � CO2 + H2O

This is unbalanced as there are 4 ‘H’ on the left but only 2 ‘H’ on the

right. This must be corrected by adding a ‘2’ in front of the ‘H2O’:

CH4+ O2 � CO2 +2H2O

Now the H balances but there 4 O on the right and only 2 on the left.

This must be balance by placing a ‘2’ in front of the ‘O2’

CH4 + 2O2 � CO2 + 2H2O

Now there is 1 ‘C’, 4 ‘H’ and 4 ‘O’ on each side so it balances.

MOLES AND MOLAR MASS

•A ‘mole’ is the name we give to the number

6.02x1023 – it is used to talk about the numbers of

particles involved in chemical reactions.

•It is the number of atoms such that one mole has

the same mass in grams as an atom’s atomic mass.

•Eg 1: Carbon has an atomic mass of 12.0,

so a mole of carbon has a mass of 12.0g

•Eg 2: Iron has an atomic mass of 55.8, so

masses of the elements in its formula:

•Eg 1: C2H6O (C=12.0, H=1.0, O=16.0)

Molar mass = 12.0 x 2 + 1.0 x 6 + 16.0 x 1 = 46.0g

The number of moles of a substance present in a

given mass is given by:

Moles = mass used

molar mass

Eg: How many moles of ethanol (C2H6O) are

Solids, Liquids and

Gases

from two or more

different elements

bonded together

Mixture: A

substance made

from two or more

elements or

compounds mixed

but not joined

A mixture of compounds and elements:

•Eg 2: Iron has an atomic mass of 55.8, so

a mole of iron has a mass of 55.8g

•The molar mass of a compound is the sum of the

Eg: How many moles of ethanol (C2H6O) are

present in 69.0 g?

Moles = mass used = 69.0 = 1.5 mol

molar mass 46.0

MOLE CALCULATIONS

What mass of carbon dioxide (the unknown) is needed to produce

when 50g iron (the known) oxide is reduced to iron.

•Balanced Equation: 2Fe2O3 + 3C � 4Fe + 3CO2

•Moles of Known (Fe2O3)= mass / molar mass

= 50.0 / (55.8 x 2 + 16.0 x 3) = 0.313

•Moles of Unknown (CO2) = (moles of known /knowns) x unknowns

= (0.313 / 2) x 3 = 0.470 mol

•Mass of Unknown (CO2) = moles x molar mass

= 0.470 x (12.0 x 1 + 16.0 x 2) = 20.7g

COMBINING POWERS

This is the number of ‘bonds’ an element forms

The combining power is given by the periodic table.

•Groups I and VII form 1 bond

•Groups II and VI form 2 bonds

•Groups III and V from 3 bonds

•Group IV forms 4 bonds

•Group VIII forms 0 bonds

•Eg: NH3 – N (Gp V) has three bonds to Hs, each of the

three Hs (Gp I) has one bond to N.

ATOMIC STRUCTURE

Atoms are made of:

Protons: mass = 1, charge = +1

Neutrons: mass = 1, charge = 0

Electrons: mass = 0, charge = -1

In a square on the periodic table

the smaller number, the proton

number gives the number of

protons or electrons and the

bigger number, the nucleon

number the number of protons

and neutrons together.

Eg 1: Boron has 5 protons,

6 neutrons, 5 electrons

Eg 2: Phosphorus has 15

protons, 16 neutrons 15

electrons

C2: CLASSIFYING THE

ELEMENTS

Structure of the Periodic Table (PT on last page!)

Elements arranged in order of increasing proton number.

PERIODS:

•The rows in the periodic table.

•For example lithium, carbon and chlorine are all in period 2.

GROUPS:

•The columns in the periodic table.

•Use roman numbers: I, II, III, IV, V, VI, VII, VIII

•For example, F, Cl, Br, I are all in Group VII (the halogens)

•Elements in the same group have similar properties and react

in similar ways.

•Eg. The halogens all react in the same way with sodium to

form sodium fluoride (NaF), sodium chloride (NaCl) and sodium

bromide (NaBr)

Periodic Patterns

Across a each period (row) you see the same patterns repeated. For

example in each period:

•The size of atoms decreases from Group I to Group VIII

•The number of electrons in the outer-shell increase by one as

Non-metals

Transition Metals

Gro

up

VIII: N

ob

le G

ase

s

Gro

up

I: Alk

ali M

eta

ls

Gro

up

II: Alk

ali-E

arth

Lanthanides and Actinides (metals)

Other

Metals

H

ME

TALS

Conduct

electricity,

conduct heat,

higher density,

malleable

NO

N-

ME

TALS

Insulate

electricity,

insulate heat,

lower density,

brittle

Gro

up

VII: H

alo

ge

ns

Group I: The Alkali Metals (Li, Na, K ...)

Oxides

Formed when an element

reacts with oxygen for

example: lithium oxide

(Li2O), calcium oxide (CaO),

carbon dioxide (CO2),

sulphur trioxide (SO3)

ACIDIC OXIDES •The number of electrons in the outer-shell increase by one as

you move across each group

•The melting point increases from Group I to Group IV and

then decreases from Group IV to Group VIII

Group I: The Alkali Metals (Li, Na, K ...)

As you go down Group I, the alkali metals get:

•More reactive

•More dense

•Harder

•Higher melting point

The alkali metals react with water in the same way:

Metal + water � metal hydroxide + hydrogen

2Li + 2H2O � 2LiOH + H2

The metal hydroxide is an alkali – it makes a pH greater than 7

when it dissolves in water (hence the name alkali metals)

Lithium is high in Group I so reacts much more slowly than

potassium which is lower in the group.

ACIDIC OXIDES

Many non-metal oxides

dissolve in water to make

acids: carbon dioxide makes

carbonic acid, sulphur

trioxide makes sulphuric acid

BASIC OXIDES

Some metal oxides are bases

(such as CaO): they

neutralise acids. Those that

also dissolve in water are

called alkalis (such as Na2O).

Group VII: The Halogens (F, Cl, Br, I)

As you go down Group VII, the halogens get:

•Less reactive

•More dense

•Higher melting point (F/Cl – gases, Br – liquid, I – solid)

•Darker coloured (pale green� dark brown)

Chlorine was used as a weapon because it’s very reactive,

fluorine is so reactive it corrodes the bottles it is stored in!

The Challenge

You need to be able to use an element’s position in the periodic table

to predict its properties. This means being familiar with the properties

of groups I, II, VII and VIII and understanding them in depth.

You may wish to research groups III, IV and VI in more detail using

textbooks or the internet)

Group III (B, Al.....)

Does not follow simple patterns – B

and Al react in very different ways.

The oxide of aluminium (Al2O3) is

amphoteric – this means

sometimes it acts like an acid and

sometimes like a base.

Group IV (C, Si, Ge...)

Carbon exists in different forms

(allotropes) – diamond, graphite,

Buckminster fullerene, nanotubes.

Si and Ge are semiconductors –

sometimes they conduct electricity

and sometimes not.

Group VI (O, S, Se...)

Main interesting point is the oxides

of sulphur: sulphur dioxide (SO2)

and sulphur trioxide (SO3) both

exist and dissolve to form

sulphurous- (H2SO3) and sulphuric-

(H2SO4) acids respectively.

C3: PETROCHEMICALS –

Refining Crude Oil

Hydrocarbons

Hydrocarbons are compounds of hydrogen and carbon only.

The carbons are linked together with hydrogen atoms attached

to them. They carbons can be arranged in straight chains (1),

branched chains (2) or even rings (3).

Refining Oil – Fractional Distillation

Oil is a mixture of hundreds of hydrocarbons. This mixture has to be separated into its useful

components using fractional distillation. Very hot crude oil is pumped into the fractionating column

where the hydrocarbons separate out by their boiling points, rising through the column until they get

cold enough to condense. The compounds that condense at a particular temperature are called a

FRACTION.

1

23

fuel gas, 1-4 carbons

petrol, 5-9 carbons

naptha, 6-11 carbons

kerosene,

11-18 carbons

diesel, 15-21 carbons

COOLER

Bubble Caps: the

gaseous fractions

bubble up through

these until they get cool

enough when they then

condense.

ALKANES

The simplest

hydrocarbons are the

alkanes, they are

saturated hydrocarbons

which means they only

contain single bonds. They

are unreactive and make

good fuels and solvents.

ALKENES

These are unsaturated which means they contain at least one

double bond. They are very valuable as a starting point for

making lots of other compounds....more on the next page!

As you move down the column, the fractions have longer

carbon chains. This increases the attractive forces between

molecules which leads to:

•Higher boiling points

•Higher viscosity

•Lower flammability

Fuel Gas – used for fuel, and to make other chemicals

Naptha – used mostly to make other useful compounds

Kerosene – fuel for aeroplanes

Fuel oil – fuel for large ships

Bitumen – used to surface roads

fuel oil, 20-27 carbons

Greases and wax,

25-30 carbons

bitumen, 35+ carbonsHOTTER

FORMULAS

Molecular formula: tells you all the

atoms present in a molecule. Quick to

write but little information.

Graphical formula: drawing showing

how all the atoms in a molecule are

connected – takes longer but tells you

much more information.

C3: PETROCHEMICALS –

Using the Products

Ethene

The gas ethene can be made into many other

compounds so is too valuable to burn.

The double bond in ethene and other alkenes has a cloud of

electrons around it which makes them very reactive. The

reactions of alkenes involve adding things to the double bond

Reaction with Steam

Ethene reacts with steam in the presence of a phosphoric acid

catalyst to make ethanol which can be used as a solvent or to

make other useful compounds.

C2H4(g) + H2O(g) ���� C2H5OH(g)

Testing for unsaturated hydrocarbons

When an orange solution of bromine is added to alkenes in the

Cracking

Because there is a greater need for hydrocarbons with shorter carbon chains we sometimes need to

cut longer chains into shorter ones using the process of cracking.

A long alkane is heated, vaporised and passed over a ceramic catalyst produce a shorter alkane and

an alkene.

Eg. 1: C8H18 � C4H10 + C4H8

Eg. 2: C10H22 � C7H16 + C3H6

Note:

•The with the alkenes for each carbon there are 2 H (CnH2n); with the alkanes, for each C there are 2 H

plus 2 extra (CnH2n+2).

•Any combination of alkene and alkane can be made, including straight and branched chains, so long

as the numbers of atoms balance.

Polymers (plastics)

Polymers are very large molecules made from lots of smaller ones (monomers) joined together.

Polymers can be many thousands of monomers long.

Addition Polymers – eg polythene, polystyrene, polyvinylchloride (PVC)

These are formed by monomers containing a C=C double bond. The double bonds link together to

form a continuous chain.

When an orange solution of bromine is added to alkenes in the

presence of UV light, the bromine reacts with the double bond

on the alkene to make a bromoalkane. The bromine water

loses its colour so this makes it a good test for alkenes:

C2H4 + Br2 ���� C2H4Br2

Homologous series

The alkenes are a homologous series, this means they are all

similar (in this case containing a C=C double bond) but differ

only in the length of their

carbon chain. The

alkanes are also a

homologous series.

The beginning of the

name tells you the

number of carbons and

the end part the type of

compound.

Condensation Polymers – eg nylon, polyester, Kevlar

These are formed from monomers that contain a carboxylic acid group (-COOH) and either an –OH or

an –NH2 group. The ‘acid‘ end of one monomer joins with the –OH/-NH2 of the other, spitting out

water.

Thermoplastics

The polymer chains are only weakly attracted to

each other so these can be continuously melted

and re-moulded. Easy to recycle.

Thermosets

The polymer chains are joined with cross links,

this means they decompose when heated instead

of melting. Can’t be recycled (easily) or re-

moulded.

polymer chains free to move when hot Cross-links prevent chains from moving

C4: CHEMICALS FROM

PLANTS

Carbohydrates

Carbohydrates are compounds of carbon, hydrogen and oxygen

such as glucose, starch and cellulose.

Glucose:

One of the most important carbohydrates

(to humans) is glucose, C6H12O6. There are

two important forms of glucose – alpha

and beta. In alpha glucose (pictured) the

OH circled in red points down, in beta

glucose it points up; this seemingly small

difference has big consequences.

Starch:

Starch is a polymer made of thousands of alpha glucose units

joined together. Plants use it to store energy and animals

Amino Acids and Proteins

Proteins are polymers made of many amino acid

monomers joined together

Amino Acids:

Amino acids are compounds of

carbon, hydrogen, oxygen,

nitrogen (and sometimes

sulphur) have the general

structure shown left. The ‘NH2’

is the ‘amino’ part and the

‘COOH’ is the ‘acid’ part.

The ‘R’ means residue and can be any atom or

group of atoms from something as simple as a

hydrogen atom to something more complex like a

benzene ring.

Proteins

Proteins are long chains of amino acids whose

properties are decided by the ‘R’ group on each

amino acid. Proteins are condensation polymers

forming water each time two amino acids join.

Proteins are extremely important in biology – life

could not exist without them.

Useful Natural Products:

Cellulose, rubber and wood all have a wide range

of uses. This is likely to increase in the future as

they are renewable resources and we are more

Rubber:

Rubber is a natural polymer with chains that are

able to move past each other when stretched and

then spring back. It has many uses including joined together. Plants use it to store energy and animals

(including humans) can easily digest it to get at that energy.

Food such as bread, rice, noodles, pasta and potatoes contain a

lot of starch.

Cellulose:

Cellulose is a polymer made of many thousands of beta glucose

units joined together. Plants use it to build their cell walls and

give them strength. It can only be digested by bacteria and not

animals.

Starch and cellulose are both condensation polymers – each

time two glucoses join, one water molecule is produced.

Although they are very large molecules, the bonding in

carbohydrates is just ordinary covalent bonding (see Unit

C17).

Semi-permeable membranes

Semi-permeable membranes are membranes with

tiny holes in them. Small molecules such as

glucose can move through these holes whereas

large ones like starch can’t. The wall of our

intestine is a semi-permeable so when we eat

something containing starch – like rice – the

starch molecules must first be digested into

glucose molecules so they are small enough to be

able to pass into our blood.

they are renewable resources and we are more

aware of the need to live sustainably.

Cellulose:

Comes from wood and has many uses, by far the

most important of which is making paper and

cardboard. The long fibres of cellulose are tangled

into a fine, flexible web.

then spring back. It has many uses including

making car tyres, rubber gloves and balloons.

Wood:

Wood is strong, cheap and readily available and

finds many uses especially for construction and

furniture making.

����

����Too big!!!

Starch

GlucoseSemi-permeable

membrane

C5: MATERIALS AND

STRUCTURES

Properties of materials

Materials can be described using words such as:

•Strength – how much force it can resist?

•Elasticity – how stretchy is it?

•Hardness – how difficult is it to scratch?

•Porosity – can air/water pass through it?

•Transparency – does light pass through it?

•Conductivity – does it conduct electricity or heat?

•Biodegradability – does it break down naturally outside?

Molecules

A molecule is a small particle made from a few

non-metal atoms bonded together – often fewer

than 10 but sometimes much more (think

polymers).

covalent bonds.

Molecular compounds have low melting points

due to the weak intermolecular forces and do not

conduct electricity as all electrons are stuck in

Ionic Compounds

Most compounds of a metal and a non-metal a

made of ions – atoms that have gained or lost

electrons. Usually the metal atom loses electrons

to make a positive ion (cation) and the non-metal

gains electrons to make negative ion (anion).

The positive and negative ions strongly attract

each other – this is an ionic bond.

Giant Ionic Structures

Ionic compounds don’t form molecules, they form

crystals made of alternating positive and negative

ions repeating millions of times in all directions.

This is called a giant ionic lattice.

Properties of Ionic Compounds

When you melt or dissolve an ionic compound it

conducts electricity because the ions are free to

move towards the positive and negative

electrodes. When

solid the ions are

stuck in position

and there are no

free electrons so

they don’t conduct.

Glass

Glass is made of silicon (IV) oxide - aka silica, SiO2 – with various

metal oxides (such as sodium oxide or calcium oxide) added to

it. The biggest source of silicon (IV) oxide is sand.

The metal ions cause the glass to have an amorphous giant polymers).

The atoms in a molecule are joined by strong

covalent bonds. In a solid each molecule is held

close to its neighbour by weak intermolecular

forces. When a substance melts, it is these weak

intermolecular forces that break NOT the strong

conduct electricity as all electrons are stuck in

bonds and so unable to move.

Giant Covalent Lattices

A crystal made of a repeating pattern of atoms

joined with covalent bonds that repeats millions

of times in all directions.

Examples include silica (SiO2) diamond (C) and

graphite (C). They have high melting points

because melting requires the breaking of strong

covalent bonds. The don’t conduct electricity

(except graphite) as there are no electrons free to

move – they are stuck in bonds.

Graphite

Diamond

Silica

The metal ions cause the glass to have an amorphous giant

structure, this is different to other giant structures because

the atoms are disordered and do not form regular patterns.

Metal ions can be added to glass to give it colour for example:

•Cobalt – blue

•Iron (II) oxide – blue-green

•Manganese – pale violet

•Copper oxide – turquoise

•Titanium – yellowish-brown

Recycling glass is beneficial for the environment as it uses less

energy and resources. However it is hard to control the quality

and consistency so is unsuitable for specialised applications.

Ceramics have a similar structure to glass and are made from

clay that is fired at high temperature causing a chemical

reaction that fuses its particles together.

C6: OXIDATION AND

REDUCTION

Oxidation and Reduction

Oxidation is when something gains oxygen. Reduction is when

something loses oxygen. Whenever one thing gets oxidised,

another thing must get reduced (and vice versa).

2Fe2O3 + 4Al � 2Al2O3 + 4Fe

In this reaction, the iron (in iron oxide) is reduced and the

aluminium is oxidised (to aluminium oxide).

You can describe aluminium as a reducing agent because it

reduces the iron. Reducing agents must be more reactive than

the element they are reducing – in this case we had aluminium

which is more reactive than iron.

Calculating % Metal Content

Factories processing ores need to know the

percentage metal content so they know

whether they can make enough money from it

and how much metal to expect to produce.

% Metal Content = mass metal in ore x 100

formula mass of ore

Eg. What percentage of iron is present in iron

ore, Fe2O3? (Atomic masses: Fe=55.8, O = 16.0)

% Iron = 2 x 55.8 x 100 = 69.9%

2 x 55.8 + 3 x 16.0

To calculate how much ore is needed to make a

given amount of iron, divide the amount you

want by the percentage (expressed as a decimal)

Eg: To make 100 kg iron you need:

Mass Iron needed = 100 = 143 kg

0.699

Reactivity Series Extracting Metals From Their Ores The limestone (CaCO ) reacts with impurities such as silicon to form

Extracting Minerals from the Environment

In order to extract metals from their ores, we

must first extract their ores from the earth. This

can be done by open-cast mining (just dig a big

hole) or shaft-mining (mining underground).

There are a number of issues associated with

both processes:

•Dangerous – many workers killed each year

•Polluting – can cause the release of heavy

metals and other poisons into the environment

•Habitat destruction – caused not just by the

mine but all the roads etc needed to service it

•Waste Disposal – vast mounds of spoil made.

•Dusty

•Increased heavy traffic

•Noisy

•Creates jobs – but can make an area over

dependent on one income source

•Ugly – destroys the natural beauty of places

Reactivity Series

MOST REACTIVE

Potassium, K

Sodium, Na

Calcium, Ca

Magnesium, Mg

Aluminium, Al

(Carbon, C)

Zinc, Zn

Iron, Fe

Tin, Sn

Lead, Pb

Copper, Cu

Silver, Ag

Gold, Au

Platinum, Pt

LEAST REACTIVE

RE

AC

TIV

ITY

Extracting Metals From Their Ores

Rocks that contain a significant amount of a metal are called ores.

The metals in an ore are not present in their pure form but are

bonded to other elements to form compounds – often oxides or

sulphides. For example iron can be extracted from iron ore (Fe2O3,

iron (III) oxide) and lead can be extracted from an ore called galena

(PbS, lead sulphide).

Metals that are less reactive than carbon can be extracted by using

carbon as a reducing agent (to steal the oxygen/sulphur). Metals that

are more reactive than carbon must be produced by electrolysis.

Iron is less reactive than carbon so can be reduced by it. This is done

in a blast furnace. Study the diagram then read the following:

•Step 1: Carbon (coke) reacts with oxygen (from the hot air blast)

C (s)+ O2(g) � CO2(g)

•Step 2: Carbon dioxide reacts with more carbon to make carbon

monoxide

CO2(g) + C(s) � 2CO(g)

•Step 3: Carbon monoxide reduces the iron oxide (iron ore) to make

molten liquid iron.

Fe2O3(s) + CO(g) � Fe(l) + CO2(g)

The limestone (CaCO3) reacts with impurities such as silicon to form

an easy-to-collect waste called slag (calcium silicate, CaSiO3):

CaCO3 +SiO2 � CaSiO 3+ CO2

Step 1 happens here

Step 2 happens here

Step 3 happens here

C7: IONS AND ELECTROLYSIS

Products of Electrolysis

Electrolysis

Electrolysis is a process in which electricity is

used to break compounds down into their

elements. The mixture being electrolysed is

called an electrolyte and must be liquid (either

melted or dissolved) to allow the ions to move.

Cations (positive ions – remember they are

’puss-itive’) ions move to the cathode (the

negative electrode) where they gain electrons,

usually forming a metal.

Anions (negative ions) move to the anode (the

positive electrode) where they lose electrons,

usually forming a non-metal.

In the electrolysis of copper chloride (CuCl2)

(right) positive copper ions move to the cathode

and form copper metal. Negative chloride ions

more to the anode and form chlorine gas.

Cu2+

Cu2+

Cu2+

Cu

Cu

CuCl- Cl-

Cl-Cl-Cl-

Cl-

Cl Cl

Cl Cl

Cathode

(negative

electrode)

Anode

(positive

electrode)

Anions move

to anode

Cations move

to cathode

Layer of metal

formed

Bubbles of gas

formed

Molten

SaltSalt Solution

Cathode Metal

Metal, except with reactive

metals (K, Na, Li Ca, Mg) in which

case H2 gas is produced plus a

solution of metal hydroxide

Anode Non-metalNon Metal, except sulphates in

which case O2

Electrolysis of Aluminium

Aluminium can’t be extracted by reduction of

aluminium oxide (Al2O3) using carbon as carbon

is less reactive than aluminium. Instead

Molten aluminium oxide ( the electrolyte) is

placed in a large carbon lined vessel which acts

as the cathode. A large anode made of carbon is

lowered into the electrolyte. The processes that

Purification of Copper

When copper is made it contains lots of impurities. The copper

is purified by electrolysis. A large lump of impure copper is is less reactive than aluminium. Instead

aluminium is produced by electrolysis.

lowered into the electrolyte. The processes that

take place are:

At the cathode:

Aluminium ions gain electrons

to make liquid aluminium

Al3+ + 3e- � Al(l)

At the anode:

Oxide ions lose electrons to

make oxygen gas

O2- � ½ O2(g) + 2e-

The oxygen reacts with the

carbon anode so it has to be

replaced regularly

is purified by electrolysis. A large lump of impure copper is

used as the anode, the electrolyte is copper sulphate solution

and the cathode is made of pure copper.

At the anode, instead of anions losing electrons, neutral

copper atoms lose electrons to become copper ions .

Cu(s) ���� Cu2+(aq) + 2e-

These then move through the electrolyte to the cathode

where they become copper atoms again.

Cu2+(aq) + 2e- ���� Cu(s)

The anode loses mass

as copper atoms leave

it and the cathode

gains mass as copper

atoms join it. The

impurities sink to the

bottom as a pile of

sludge.

Some Tests

You need to know two tests for elements

that can be made during electrolysis

Chlorine gas – bleaches damp litmus paper

Oxygen – can relight a glowing wooden

splint

The Electrolysis of Sodium Chloride (NaCl)

The electrolysis of brine (sodium chloride solution)

makes sodium hydroxide (many uses in industry),

chlorine gas (used for many things including

hydrochloric acid) and hydrogen gas (also used for

many things including hydrochloric acid).

C8: SOLVENTS AND

SOLUTIONS

Solutions

A solution is a mixture in which a solute is dissolved in a liquid

solvent. When you dissolve something, it is still ‘there’ but it

has been broken down into individual molecules or ions that

are too small to see. If you added 10g of salt to 100g water, the

solution will weigh 110g NOT 100g because the salt is still

present, but just well mixed with the water.

Solubility and Concentration

Some substances are more soluble in water than others,

which means that more of the substance is able to dissolve.

Sodium chloride is very soluble in water but silica (SiO2) is

insoluble.

Generally ionic compounds, such as copper sulphate,

dissolve in water whereas covalent compounds dissolve in

non-aqueous solvents such as ethanol, acetone or hexane.

The ‘strength’ of a solution is called concentration. It is

measured in units of ‘mol dm-3’ (pronounced ‘moles per

decimetre cubed). 1.0 mol dm-3 means that 1.0 mole of

solute is dissolved in 1.0 litres (dm3) of solution. In general:

Concentration = moles of solute .

volume of solution in litres

Eg. 75.0g of glucose (C6H12O6) is dissolved in 250 cm3 of

water, what is the concentration of this solution?

Moles of solute = mass used ÷ molar mass

= 75.0 ÷ (6 x 12.0 + 12 x 1.0 + 6 x 16.0)

= 0.42 mol

Concentration = moles solute ÷ volume in litres

= 0.42 ÷ (250 ÷ 1000*) = 1.68 mol dm-3

Some More Tests

You need to remember the chemical

tests for the following ions:

Chloride ions:

•Add acidified silver nitrate solution

•See a white precipitate of insoluble

silver chloride

•Cl-(aq) + AgNO3(aq) � AgCl(s) + NO3-(aq)

Sulphate ions:

•Add acidified barium nitrate solution

•See a white precipitate of insoluble

barium sulphate

•SO42-

(aq) + Ba(NO3)2(aq) � BaSO4(s) +

. 2NO3-(aq)

Both these reactions rely on solid

particles of an insoluble product being

made, this precipitates out of the = 0.42 ÷ (250 ÷ 1000*) = 1.68 mol dm-3

*There are 1000 cm3 in 1 litre so this turns cm3 into litres

made, this precipitates out of the

solution as ‘cloudy powder’.

Hard and Soft Water

Hard water contains small amounts of dissolved calcium and magnesium minerals

that can slowly form scale (deposits of calcium carbonate, magnesium hydroxide and

calcium sulphate) which clogs pipes.

eg: Ca(HCO3)2(aq) � CaCO3(s) + H2O(l) + CO2(g)

You can often tell water is hard by the behaviour of soap: in soft water it forms a

bubbly lather and in hard water it leaves behind a grey scum. Softening water

involves converting the minerals to insoluble compounds that settle out of the water.

Temporary hardness caused by magnesium- or calcium- hydrogen carbonate can

be removed by boiling:

Mg(HCO3)2(aq) � MgCO3(s) + H20(l) + CO2(g)

Permanent hardness caused by calcium sulphate can only be removed by sodium

carbonate (washing soda):

CaSO4(aq) + Na2CO3(aq) � CaCO3(s) + Na2SO4(aq)

Or by ion exchange. The water is passed through a column containing Na+ ions, these

get swapped over with Ca2+ ions: Ca2+(aq) + 2Na+

(s) � Ca2+(s) + 2Na+

(aq)

Drinking Water

Water drawn from rivers can

contain pollutants such as

fertilizers, dissolved organic

matter, harmful bacteria and

industrial waste that make it unfit

to drink. At treatment plants, two

main processes are used to make

water safe:

Filtration – the water is passed

through a series of increasingly

fine filters that trap suspended

particles. Activated carbon is used

to filter out dissolved pollutants.

Chlorination – chlorine is added

to the water which destroys

bacteria.

Cleaning

Often non-aqueous (not water) solvents are

used in cleaning as they can dissolve the dirt,

for example acetone can dissolve nail

varnish.

Detergents are used to clean up oils and fats,

for example in laundry powder or washing-

up liquid.

Detergent molecules have two ends , one

end dissolves in water and the other in oil

which allows oil and water to mix

C9: ACIDS AND ALKALIS

Reactions of Acids

You need to memorise these reactions, each one shows the

general word equation then a specific example with symbols.

Acids and Metals

Acid + Metal � Salt + Hydrogen

•Hydrochloric acid + lithium � lithium chloride + hydrogen

• 2HCl(aq) + 2Li(s) � 2LiCl(aq) + H2(g)

Acids and Base (like alkali but not always soluble)

Acid + Base � Salt + Water

•Sulphuric acid + sodium hydroxide� sodium sulphate + water

• H2SO4(aq) + NaOH(aq) � Na2SO4(aq) + H2O(l)

Acids and Carbonates

Acid + Carbonate � Salt + Water + Carbon Dioxide

•Nitric acid + calcium carbonate� calcium nitrate + water +

What’s the salt?

To work out which salt is formed during neutralisation reactions you need to know the ions formed

by the acid or alkali when it dissolves.

Working out the name is

easy, you just combine

the name of the cation

from the alkali with the

anion from the acid.

For example potassium

sulphate and sulphuric

acid makes potassium

sulphate.

Magnesium hydroxide

and phosphoric acid

makes magnesium

phosphate

Working out the formula of the salt is a little more complicated, the key is to make sure the

positive and negative charges on the cancel each other out to zero.

Substance Cation(s) Formed Anion(s) Formed

Hydrochloric acid, HCl 1 H+ Cl- , chloride

Nitric acid, HNO3 1 H+ NO3- , nitrate

Sulphuric acid, H2SO4 2H+ SO42- , sulphate

Phosphoric acid, H3PO4 3 H+ PO43- , phosphate

Sodium hydroxide, NaOH Na+ , sodium 1 OH-

Potassium hydroxide, KOH K+ , potassium 1 OH-

Magnesium hydroxide, Mg(OH)2 Mg2+ , magnesium 2 OH-

Ammonium hydroxide, NH4OH NH4+ , ammonium 1 OH-

Eg 1. Potassium nitrate

K+ has one plus charge

Eg 2. Magnesium phosphate

Mg2+ has two plus charges•Nitric acid + calcium carbonate� calcium nitrate + water +

. carbon dioxide

• HNO3(aq) + CaCO3(s) � Ca(NO3)2(aq) + H2O(l) + CO2(g)

Neutralisation Reactions

All acids form hydrogen ions (H+ )

when they dissolve, all alkalis form

hydroxide ions (OH-). During

neutralisation, the H+ and OH- react to

form water:

H+(aq) + OH-

(aq) � H2O(l)

This reaction is exothermic, which

means it gives out heat and gets hot.

Finally, to write a balanced equation, you need to get the right number of waters, the simplest way

is to remember that each ‘H+’ from an acid makes one water.

K+ has one plus charge

SO42- has two minus charges

You need two K+ to balance out one

NO3- so the formula is K2SO4

Mg2+ has two plus charges

PO43- has three minus charges

So you need three Mg2+ to balance out

two PO43- so the formula is Mg3(PO4)2

Eg 1. Potassium hydroxide and sulphuric acid

As we have seen it makes K2SO4 which requires

one H2SO4 and two KOH. Two H2O are made

since the one H2SO4 produces two H+ ions

H2SO4 + 2KOH � K2SO4 + 2H2O

Eg 2. Magnesium phosphate

As we have seen it makes Mg3(PO4)2 which

requires two H3PO4 and three Mg(OH)2. Six H2O

are made since each of the two H3PO4 produces

three H+ ions.

2H3PO4 + 3Mg(OH)2 � Mg3(PO4)2 + 6H20

The pH Scale

•Acids have a pH of

less than 7

•Alkalis have a pH

greater than 7

•pH can be

measured with

colour changing

indicators or digital

pH meters

Some uses of Bases

Antacids, used to cure indigestion, are basic salts – such as

carbonates – that react with and neutralise acids.

Lime (calcium oxide, CaO) is used on a large scale to neutralise

acidic industrial waste.

Testing Carbonates

To test for carbonates, add a

sample to some acid and

bubble the gas collected

through limewater. If the

limewater goes cloudy, the

sample contained a carbonate.

How much energy?

Carry out a neutralisation reaction in an insulated container

such as a polystyrene cup. By measuring the temperature

change and the volumes you can work out how much heat was

given out by the reaction (H = m.c.ΔT). You can then divide this

by the number of moles of acid you had to work out how much

energy one mole of acid produces.

C10: SOIL, ROCKS AND RATES

Rates of Reaction

For a chemical reaction to happen, the reacting particles need

to collide with enough energy. Anything that increases the

number of collisions or their energy will increase the rate.

Temperature

Increasing temperature increases the rate of a reaction

because particles are moving faster which means more

collisions and higher energy collisions.

Concentration

Increasing the concentration of a solution increases the rate of

a reaction because it means there are more particles available

to react which leads to more collisions.

Surface Area/Particle size

Increasing the total

surface area of

The Rock Cycle

The rocks that make up the

Earth’s surface are in a constant

state of slow change that takes

place on a timescale of millions

of years. Igneous rocks are

formed by magma from the

mantle that comes out through

volcanoes or moves to near the

surface and cools before

erupting. Sedimentary rocks

are formed by small particles of

rock that get eroded ,

transported and built up in a

layer thick layer that squashes

the particles at the bottom

together forming new rock.

Metamorphic rocks are formed

when sedimentary rocks get hot enough to partially melt, changing their structure. This process of

constant change is called the rock cycle.

Weathering Some Uses of Rockssurface area of

particles (by using

finer powder)

increases the rate of a

reaction because it

means more particles

at the surface are

exposed to collisions.

How fast?

On a graph showing the change in concentration of reactants

or products, the gradient of the line tells you the reaction rate:

steeper = faster,

flat = stopped

Weathering

This is the process whereby rocks are broken into ever

smaller pieces by exposure to the environment. There are

three classes of weathering:

•Physical: for example water (the force of waves and

rivers knocking bits off), exfoliation (caused by the

day/night heating/cooling cycle leading to cracks that

gradually expand over time) or freeze-thaw (water

seeps into cracks, freezes, expands and enlarges the

crack).

•Chemical: for example hydrolysis (when rocks like

feldspar react with acidic rainwater to form kaolin

(china-clay)) or carbonation (naturally occurring

carbonic acid in rain water (dissolved CO2) reacts with

limestone to form soluble calcium hydrogen carbonate

(Ca(HCO3)2))

•Biological: the force of plants roots growing into

cracks and forcing them apart.

Weathering releases nutrients present in the rocks and so

is vital for making soils fertile.

Some Uses of Rocks

Limestone:

•Used to remove impurities during iron

production

•Lime (CaO, produced by thermal

decomposition of limestone) used to

raise pH of acid soils

Sand:

•Used in glass production

Yet More Tests

You need to remember these chemical

tests:

•Oxygen (see Unit C7)

•Hydrogen – lighting a test-tube of H2

with a splint gives a squeaky pop

•Carbon dioxide – when bubbled

through limewater it turns it cloudy.

C11: FERTILISERS

Ammonia, NH3

The ammonia is alkaline gas – forming

ammonium hydroxide (NH4OH) when it dissolves

in water. Ammonia is very important since it is

used to turn unreactive nitrogen gas (N2) into

important nitrate (-NO2/3) containing compounds

such as fertilisers and explosives, this is known as

nitrogen fixation.

The Haber Process

This the process used to produce ammonia from nitrogen and

hydrogen gases.

N2(g) + 3H2(g) � 2NH3

The reaction is reversible which means some of the products

turn back to reactants as soon as they are made, this means it

Sulphuric Acid, H2SO4

Sulphuric acid is another important chemical used in huge

range of industrial processes. It is produced by the Contact

Process. There are three chemical reactions. First sulphur is

burnt in air to produce sulphur dioxide (SO2):

S + O2 � SO2

Secondly SO2 is reacted with further oxygen to make sulphur

trioxide (SO3):

2SO2 + O2 � SO3

This reaction is reversible, so to maximise the amount of SO3

made, they use a high temperature (425OC), medium-high

pressure (1-2 times atmospheric pressure) and a catalyst

(vanadium (V) oxide, V2O5). Finally, the sulphur trioxide is

produced by first dissolving it in sulphuric acid to make oleum

(H2S2O7) which then makes more sulphuric acid on the addition

of water:

SO3 + H2SO4 � H2S2O7

H2S2O7 + H2O � 2H2SO4

Fertilisers

Fertilisers are chemicals applied

to plants to improve their growth

and increase the amounts of

products such as fruits, nuts,

leaves, roots and flowers that

they produce for us. They work by

supplying plants with the vital

elements they need including

nitrogen - in the form or nitrate

(NO3- containing) salts,

phosphorous – in the form of

phosphate (PO43- containing) salts

and potassium (K+ containing)

salts.

Salts containing suitable ions can

be prepared by reacting various

combinations of potassium

hydroxide, ammonia, nitric acid

and phosphoric acid (see Unit C9).Nitric Acid (HNO3)

Nitric acid is prepared via a number of steps starting with the turn back to reactants as soon as they are made, this means it

takes a long time to make an economical amount of ammonia.

To speed it up, the reaction is done at high temperature

(~450OC), high pressure (~200 times atmospheric pressure) and

with a catalyst (iron oxide).

oxidation of ammonia to nitric oxide (NO):

4NH3 + 5O2 � 4NO + 6H2O

This reaction is quite slow so a platinum catalyst is used to

speed it up. Next, the nitric oxide is oxidised to nitrogen

dioxide (NO2):

2NO + O2 � 2NO2

Finally the nitrogen dioxide is reacted with water to produce

nitric acid (HNO3):

3NO2 + H2O � 2HNO3 + NO

More Tests....ayooohhhh!!

You need to remember the

following tests:

•Ammonium ion (NH4+) – add a

few drops of cold sodium

hydroxide. If ammonium is

present it will produce ammonia

which you can smell and the

fumes will turn damp red litmus

blue.

•Nitrate ion (NO3-) – boil the

sample with sodium hydroxide

and aluminium foil. If nitrate is

present, ammonia will be

produced so the fumes will turn

damp red litmus blue

Eutrophication

When it rains on fields that have been treated with nitrate

fertilisers, they can dissolve in the rain water and be washed

through the soil into streams, rivers and lakes. The nitrates

then fertilise the growth of lots of algae in the water. When

this dies, it sinks to the bottom and is rapidly decomposed by

bacteria which use up most of the oxygen dissolved in the

water, causing most fish and other aquatic life to suffocate.

This process if called eutrophication and is a major problem.

This is not such a problem with phosphate salts since they are

much less soluble so do not make it to the water in such large

amounts and potassium on salts on their own can not cause

such an effect.

Ammonia and Ammonium

Ammonia (NH3) is a gaseous

compound the forms an alkali

when it dissolves in water.

The similarly named ammonium

(NH4+) is an ion formed when

ammonia reacts with acids

forming ammonium salts such

as ammonium nitrate (NH4NO3) or ammonium sulphate

((NH4)2SO4).

C12: DYES AND DRUGS

Dyes

Dyes are compounds used to colour fabrics.

Initially many dyes were produced from natural substances, for

example ‘tyrian purple’ was produced from sea-snails, ‘red

carthamine’ from safflowers and turmeric was used to dye

things yellow. More recently, synthetic dyes have been

invented, the first of which was the mauve coloured dye

‘mauvine’ invented by William Perkin in 1856. Synthetic dyes

have replaced natural ones for most uses.

Substances called mordants are often added to help fix dyes to

their fabrics. Before modern chemicals, one of the most widely

used mordants was urine!

Melting/Boiling Point

The melting and boiling point of substance

depend on how pure they are. For example Drugs

Paper Chromatography

Paper chromatography is a technique that can be

used to separate mixtures of dyes or pigments

such as chlorophyll. Firstly a very strong solution

of the dye-mixture is prepared, this is then used to

build up a small and intensely coloured spot on a

piece of absorbent paper. This is then placed in a

jar of solvent (with a lid). As the solvent soaks up

the paper, it dissolves the coloured spot, causing it

move up the paper as the solvent does. However

because dyes have different levels of solubility,

they move up the paper at different speeds

causing the individual colours to separate out. The

solvent or combination of solvents can be changed

to get the best possible separation of spots.

Purity

It is important for chemists to be able to produce

pure substances, this is especially the case for

things like food additive or drugs that are used in depend on how pure they are. For example

water boils at 100OC but if you ‘pollute’ it with

some salt, the melting point decreases to minus

3-4OC and can increase the boiling point to 103-

4OC.

This effect is often used by chemists to judge the

purity of the compounds they have made.

Drugs

Drugs are chemicals that are used to change (generally to

improve) physical or mental wellbeing. Many drugs have come

from the study of plants, which is one of many reasons we

should do all we can to protect the rainforests. Example include

quinine used to treat malaria from the cinchona tree,

vincristine from the rosy periwinkle flower used to treat

childhood leukaemia and procaine from the coca plant used as

a local anaesthetic.

Nb. You do not need to

remember these

vincristine

procaine

quinine

Analgesics

Analgesics are drugs that relieve pain – better

known as painkillers. Common examples include:

Aspirin

Paracetamol

things like food additive or drugs that are used in

the human body. Whilst the drug or food

additive itself may not be harmful to the body,

there is no guarantee that any impurities won’t

be. Because of this, chemists use a wide range of

techniques including recrystallisation,

chromatography and measuring the melting

point to test the purity of their products.

Chemotherapy

Chemotherapy refers to taking specific drugs

such as ‘cis-Platin’ that are designed to kill

cancerous tumours. Most chemotherapy drugs

generally target cells that divide rapidly – like

cancer cells but also healthy cells like those

found in bone marrow and the digestive tract.

This means chemotherapy, whilst effective,

takes a very heavy toll on the body.

C13: COLLOIDS

Colloids

A colloid is a mixture of one very fine particles of one phase of

matter evenly spread throughout another. There are three

major types of colloid:

•Sol – fine particles of solid suspended in a liquid, for example

blood (red blood cells suspended in water) or paint (pigment

particles suspended in water)

•Gel – small droplets of liquid trapped in a solid matrix, for

example strawberry jelly (water trapped in solid gelatine)

•Emulsion – fine droplets of one liquid suspended in another,

for example mayonnaise (oil droplets suspended in vinegar) or

milk (fat droplets suspended in water)

Diagram shows droplets of

oil (blue) suspended in water

Emulsifiers

An emulsifier is a substance used to help make an emulsion by allowing the two liquids to mix

without dissolving. For example you can’t make mayonnaise without egg yolk, the lecithin in the egg

yolk is an emulsifier and allows the particles of oil to stay suspended in the vinegar. Without egg

yolk the oil and vinegar would separate out straight away.

Emulsifiers work in a similar way to detergents. They are large molecules with one end able to

dissolve in one type of substance and the other end able to dissolve in another type. So for example,

in an oil/water emulsion, one end of the emulsifier can dissolve in water and the other in oil

allowing small droplets of oil to be suspended in the water (or vice versa).

oil (blue) suspended in water

(red)Colloids and Light

Colloids are always opaque (not transparent),

light can’t travel through them.

The reason for this is that as photons of light hit

the suspended particles they bounce off in a

random direction rather than passing straight

through, this is called scattering. You can see this

effect when you bend a plastic ruler too much.

Normally light passes straight through it but

when you bend it just enough, you cause small

areas of plastic molecules to rearrange

themselves whilst others stay as they were.

When light hits these areas it bounces off rather

than passing straight through so the ruler

appears white rather than transparent.

Longer (redder) wavelengths of light get scattered less than shorter (bluer) ones. This is what

causes the colours of a sunset – the bluer colours of sunlight get scattered by dust suspended in

the air whereas the redder colours are able to pass straight through.

This box is blank because there’s just not that much in this unit!

C14: FUELS

Combustion (burning)

In combustion, a fuel reacts with oxygen

releasing heat. There are three requirements

for fire, described by the fire-triangle. If the

fuel is a hydrocarbon or carbohydrate the

products of combustion are water and carbon

dioxide (carbon monoxide if oxygen is limited).Fuel

A fuel is any substance that is burnt for the heat energy it

releases. Good fuels need to have a number of properties, they

must: burn easily, release a lot of heat on burning, burn

without producing harmful products, be easy to handle and

burn steadily (i.e. without exploding).

Solid fuels (such as wood and coal)

Easy to handle and so safe to use but uses are limited by the

inability to flow. Also tends to burn ‘dirtily’ producing the most

pollution such as carbon dioxide, carbon monoxide and soot.

Finely powdered solid fuels can explode due to the high surface

area causing an extremely high reaction rate (see unit C10)

Liquid fuels (such as oil, ethanol, petrol)

Harder to handle as they flow (so they don’t stay in one place)

Fossil Fuels

Fossil fuels are formed from the remains of dead

animals (oil and gas) and plant (coal) that

collected over millions of years and were

transformed by heat and pressure.

Most of the energy that human society uses

comes from burning fossil fuels in power stations,

factories and motor vehicles.

There are a number of problems with fossil fuels

•They are running out quickly (especially oil)

•They release carbon dioxide when burnt leading

to global warming

Oxygen, O2

Oxygen gas has many uses:

•Medicine – helps patients

with impaired respiratory

systems to breathe.

•Welding – burns with ethyne

gas to give a super hot flame.

Reducing Pollution

It is important to reduce the pollution and

carbon dioxide that we produce. The only real

way to do this is to burn fewer fossil fuels by:

•Using less energy – by using less things that

require energy such as computers and cars that

require or using more energy efficient ones and

avoiding energy waste (leaving lights on, using

air-con with an open window etc.).

•Using cleaner energy sources – renewable

energy such as solar, wind and hydroelectric

In cars, the carbon monoxide (CO), nitrogen

oxides (NO ) and unburnt hydrocarbons (HC) that but this makes them easy to store and to feed into engines.

Gas fuels(natural gas, methane)

The hardest to handle and store requiring high pressure

canisters. Flowing means they are well suited to use in engines

and they also burn cleanly producing least pollution.

to global warming

•The soot released when they burn damages

wildlife and coats buildings in unsightly black

Pollution

Burning fuels often produces pollution in the form of

things like carbon monoxide, nitrogen oxides, un-

burnt hydrocarbons and products of impurities such as

lead and sulphur.

Carbon monoxide (CO)

Formed when fuels burn in too little oxygen. Binds

permanently to haemoglobin in your blood preventing

your blood from carrying oxygen.

Lead (Pb)

Highly toxic causing a wide range of symptoms,

including affecting the brain development of children.

Acidic Oxides

The oxides of sulphur (SO2, SO3) and nitrogen (many

including NO and NO2) are acidic which means when they

dissolve in water they form acids. These gases are

produced from burning fossil fuels, especially coal which is

high in sulphur.

When these gases dissolve in the water in clouds they

form acid rain which damages wildlife (by releasing toxic

aluminium ions) and corrodes the stonework on buildings.

NO2 is also a key component in the formation of smog

which is a major health risk in large cities.

oxides (NOX) and unburnt hydrocarbons (HC) that

they produce can be converted to less harmful

things with a catalytic converter.:

2CO + O2 � 2CO2

2NOx � N2 + xO2

2HC + 2O2 � 2CO2 + H2O

The catalytic converter is part of a car’s exhaust

pipe and uses a catalyst of platinum, rhodium or

palladium to speed up these reactions.

Green Methane

Methane gas, a good fuel can be made from

decomposing plant and animal waste (this is

often called ‘biogas’). This makes fuel from a

waste product so is environmentally better.

Hot Words

Exothermic reactions get hotter (like burning

fuels) whereas endothermic reactions get colder

(for example dissolving salt).

C15: BATTERIES

Reactivity of Metals (check Unit

The reactivity of metals can be seen by the way they react with

steam or with acid (see Unit C6 for the reactivity series).

Reaction with water (see Unit C2 for details of this reaction):

The most reactive metals (K-Ca) react with cold water, fairly

reactive metals (Mg-Fe) will only react with steam whereas the

least reactive metals (Sn-Pt) don’t react at all.

Reaction with dilute acids (see Unit C9 for details)

The reaction of metals with acids shows a similar patter with

the most reactive metals (K-Ca) reacting violently, the fairly

reactive metals (Mg-Pb) reacting gradually more slowly and the

least reactive metals (Cu-Pt) not reacting at all.

The reactivity of metals relates to how easily they form ions,

more reactive metals like K form K+ ions much more easily

Voltaic Cells

Voltaic cells produce electricity and comprise an

electrolyte solution with electrodes made of two

different metals dipping into it. The voltage of a

cell can be changed by changing the metals used

for the electrodes.

A chemical reaction takes place where the more

reactive metal forms ions forcing ions of the less

reactive metal to become atoms. This causes

electrons to flow around the circuit from the more

reactive metal to the less reactive metal. When

one of the reactants (metal or electrolyte) runs

out, the reactions stop so the cell no longer pro

The voltage of the cell is related to the position

of the two metals on the reactivity series; the

further apart they are the higher the voltage and

the closer together they are the lower the

voltage. For example a cell using of potassium

and platinum (furthest apart) would have a very

high voltage, a cell using iron and tin (close

together) would have a very low voltage.

V

Electrode

(metal 1)

Electrolyte

Solution

voltmeterElectrode

(metal 2)

Rusting

Rust (hydrated iron (III) oxide) affects most structures made of iron (or steel (see Unit C16)) and more reactive metals like K form K+ ions much more easily

than less reactive metals like Cu can form Cu+ ions.

Types of Cell

Simple Cells (Dry Cells): These are ‘normal’ batteries that can

only be used once before they go flat. The reactions that

generate the electricity are irreversible (they only go one way).

Rechargeable Cells: Like the batteries found in a your mobile

phone or a car. The reactions are can be reversed when you

put electricity back through the battery allowing you to charge

it up.

Fuel Cells: Fuel cells ‘burn’ a continuous

supply of hydrogen and oxygen to make

electricity. The gases are separated by a

special membrane that only lets hydrogen

pass through it in the form of H+ ions. This

means that when the H2 and O2 react,

energy is released as electricity instead of

heat.

Rust (hydrated iron (III) oxide) affects most structures made of iron (or steel (see Unit C16)) and

causes huge damage:

Iron + oxygen + water ���� hydrated iron (III) oxide

Rust can be prevented by taking steps making sure either oxygen or water can’t reach the iron. The

main ways to do this involve covering the metal with, paint (bridges and other structures), oil/grease

(moving machine parts) or another metal such as zinc (galvanising). Rust is also sometimes prevented

by using sacrificial prevention where a more reactive metal (Mg, Zn) is used which corrodes in

preference to the iron, this is often used to protect the hulls of ships.

When zinc and aluminium oxidise they form a waterproof layer of oxide that protects the metal from

further damage. When iron rusts, the rust flakes off exposing more iron, this is why rust is so

damaging.

Which Battery?

Dry Cells:

Dry Cells are cheap, convenient

and can store a lot of energy but

can only be used once and so are

environmentally unfriendly. They

last a long time, gradually giving

lower and lower voltage.

Rechargeable Batteries:

Can be used many times and the

ones containing lithium are light.

They are also expensive and

older versions lose their ability to

store energy after a number or

recharges. They give the same

voltage until they go flat.

Fuel Cells:

Fuel cells run on a continuous

supply of hydrogen and oxygen

so don’t run out as such. They

are very expensive and new

technology and there is little

infrastructure for producing and

distributing hydrogen gas.

C16: METALS AND ALLOYS

Metals and their Structure

Metals have a giant structure of atoms arranged in tightly

packed rows in three dimensions.

Metals are joined together

by metallic bonding in

which metal ions are

attracted to a ‘sea of

electrons’ that surrounds

them. This can happen

because the outer-shell

electrons of the metal

atoms delocalise (meaning

they are able to move

around many atoms rather

than just one) leaving

Some Commonly Used Metals and Alloys

Metal Properties

Iron Soft, magnetic, conducts

electricity

Copper Easy to shape, ductile,

excellent heat /electrical

conductor

Aluminium Low density, strong

resists corrosion

Titanium Low density, very strong,

high melting point

Gold Shiny colour, very low

reactivity, tarnish

resistant

Alloy Metals Properties

Amalgam Mercury and

Cu/Ag/Zn/Sn

Soft when formed,

rapidly hardens – used

for dental fillings

Brass Copper and

Zinc

Malleable, soft, shiny

golden colour, good

conductivity – often

used for decorative

purposes on door

handles

Solder Tin and Lead Good electrical

conductor, very low

melting point – used

to join electrical

components

Alloys With Numbers

Given the % of metal in a

mass or alloy, the amount

Steel – a very useful alloy

Iron is too soft to have many uses, but when carbon

(and sometimes a few other things) is added to make

Transition

Metals

These metals than just one) leaving

behind positive ions.

Alloys

Alloys are ‘mixtures of metals’ (although sometimes they can contain a non-metal) that are

made by mixing molten metals.

Alloys often have very different properties to the metals they are made from and by varying

the metals they can be tailored to have specific desirable properties – this is called metallurgy.

Alloys are often harder than the metals they are made from. In a pure metal the rows of

atoms are neatly lined up meaning they can slip past each easily when the metal is hit leaving

a dent. Because alloys are made of atoms of different sizes, the atoms don’t line up neatly so

can’t slip past each other so easily when hit.

Arrangement

of atoms in an

alloy

Arrangement

of atoms in a

pure metal

mass or alloy, the amount

can be found similar to

with ores (see Unit C6).

(and sometimes a few other things) is added to make

steel it becomes the most widely used metal of all.

Mild Steel (~99.75% Fe, ~0.25% C): Strong, cheap

and easy to shape – used in construction

High Carbon Steel (~99% Fe, ~1% C): Very strong but

brittle – used to make tools

Stainless Steel (80% Fe, 15% Cr, 4% Ni, 1% C):

Corrosion resistant – Cr and Ni form waterproof

oxide layer preventing further corrosion

Manganese Steel (84% Fe, 15% Mn, 1% C): Extremely

hard – used for things like railways

The Last Tests!!!!

Flame tests are used to indentify

metals in compounds by the colour

of flame they produce when burnt:

Lithium - red

Sodium - yellow

Potassium - lilac

Calcium – orangey-red

Barium – pale green

Copper – blue-green

Cu2+ – NaOH(aq) produces a blue

precipitate which dissolves in

NH3(aq)

Fe2+ - NaOH(aq) produces a dark

green precipitate insoluble in

NH3(aq)

Fe3+ - NaOH(aq) produces a brown

precipitate insoluble in NH3(aq)

Zn2+ - NaOH(aq) forms a white

precipitate soluble in NH3(aq)

These metals

have high

densities and

boiling points

and tend to

form coloured

compounds

such as the

bright blue

copper (II)

sulphate

C17: ATOMS, BONDING AND

THE PERIODIC TABLE

Electron Arrangement/Configuration

Electrons are arranged around atoms in specific shells. The

most important shell is the outer one as this controls an

atom’s chemistry. We call the electrons in the outer shell

‘valence electrons’ The number of electrons in the outer shell

is the same an element’s group number.

The number of electrons around an atom is given by the

atom’s proton number. They are arranged in shells as follows:

•1st Shell – Holds two electrons

•2nd/3rd/4th Shells – Hold 8 electrons

•Example 1: Carbon. Proton

number is 6 which means

there are 6 electrons: 2 in the

1st shell and 4 in the second

•Example 2: Chlorine. Proton

number is 17 which means

there are 17 electrons: 2 in

the 1st shell, 8 in the second

and 7 in the 3rd.

Ionic Bonding

An ionic bond is the attraction between two oppositely charged ions. Cations (positive) are formed

when atoms (usually metals) lose electrons. Anions (negative) are formed when atoms (usually non-

metals) gain electrons.

Atoms will lose or gain electrons until they have a complete outer shell: elements in Groups I, II and

III will lose 1, 2 and 3 electrons respectively to form 1+, 2+ and 3+ ions. Atoms in Groups V, VI and

VII gain 3, 2 and 1 electrons to form 3-, 2- and 1- ions. In an ionic compound the number of positive

and negative and charges must cancel out to neutral.

Example: NaF, sodium in Group I forms a 1+ ion

and fluorine in group VII forms a 1- ion so one

Na+ is needed to balance out one F-

Example: Li2O, lithium in Group I forms a 1+ ion

but oxygen in Group VI forms a 2- ion so two Li+

are needed to balance out one O2-

Li+O2-Li+Na+F-

Covalent Bonding

A covalent bond is the attraction of two atoms (usually non-metals)

Isotopes:

Isotopes are atoms with the

Checking Your Answer: To check

you are right, the period gives the

number of shells and the group gives the number of electrons

in the outer shell. For example chlorine is in Period 3 and

Group VII so it has 3 shells and 7 electrons in the outer shell.

Ions: The configuration of ions is the same as for atoms but

you have to take electrons away from positive ions and add

extra for negative ions. For example oxygen and lithium.

and 7 in the 3 .

C

Cl

Li Li+O2-O

A covalent bond is the attraction of two atoms (usually non-metals)

to a shared pair of electrons. Small groups of covalent bonded

atoms can join together to form molecules.

The atoms share enough electrons to complete their outer shells.

*Nb: In bonding diagrams you only draw the outer shell and you use

different shapes/colours to show where electrons have come from.

Example: H2O*, hydrogen is has

one valence electron and needs

one more to complete the 1st

shell, oxygen has six valence

electrons electrons so needs

two more. Thus one oxygen will

react with two hydrogens:

Example: CO2*, carbon is has

four valence electrons so needs

four more to complete its outer

shell, oxygen needs two more.

Thus each carbon will react with

two oxygens, sharing two

electrons with each one. A bond

involving two shared pairs is a

double bond. H HO

O OC

Isotopes are atoms with the

same proton number but

different nucleon number.

For example carbon has two

main isotopes – C-12 and

C-13. Carbon has a proton

number of 6 so they both

contain 6 protons and 6

electrons but C-12 has 6

neutrons and C-13 has 7.

A Noble Matter

Atoms strive to have either 2

electrons (H, He) or 8

electrons (everything else) in

their outer shells as this is

very energetically stable –

just like the Noble Gases.

When bonding atoms gain

and lose electrons to do this.