IGCSE Chemistry
The particulate nature of matter
Definition
Atom: A basic unit of matter that consists of a dense central nucleus surrounded by a
cloud of negatively charged electrons.
Molecule: A group of two or more atoms held together by covalent chemical bonds.
Matter: The physical substance having mass and occupying space
Constituent particles: Particles that cannot be broken into smaller pieces of energy
Kinetic theory
All types of matter are composed of tiny particles which are in constant motion, the physical properties of
matter are dependent on the movement of these constituent particles
Solids, Liquids and Gas properties
Properties
Solid
Liquid
Bonds
Strong
intermolecular Relatively strong
bonds
intermolecular bonds; slightly
weaker than those of a solid
Fixed shape
There is a fixed shape
There is no fixed shape; liquids
take the shape of their
container
Fixed Volume
There is a fixed volume
Liquids have fixed volume
Arrangement of
particles
Density
Gas
Very weak intermolecular
bonds
There is no fixed shape; gases
take the shape of their
container
Gases do not have a fixed
volume`
Particles are close together Particles are close together Particles are far apart and
and vibrate around a fixed but not as close as solids
bounce randomly around in
point
the container
Fixed and quite high
Fixed and quite high
Variable and quite low
Inter-particle forces
Forces of attraction greater Forces
of
attraction Forces of attraction much less
than average energy of comparable to average energy than average energy of
particles
of particles
particles
Particle separation
Particles
together
Illustration
held
close Close but slightly separated
Widely sepearated
Different states of matter in terms of kinetic particle theory
The kinetic theory can be used as a scientific model to explain the different states of matter and the
interrelationship between them:
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Solids. The constituent particles are held close together by inter-particle attractive forces, the
particles have little freedom of movement and they can only vibrate around a fixed position
Liquids. The constituent particles are held close together but the forces are weaker and the
particles have more freedom of movement. The particles have more energy and they move around
in a random manner, often colliding with one another.
Gases. The forces of attraction between the constituent particles are so weak and their individual
energy is so high that forces of attraction become ineffective and the individual particles are free to
move around, restricted only by the walls of the container they are kept in. The particles move
randomly at very high speeds, colliding with each other and with the walls of the container in the
process.
Interconversion of different states of matter in terms of kinetic particle theory
The kinetic theory can be used to explain how a substance undergoes change from one state to another
If energy in the particles of a solid is increased (e.g. by heating), the attractive forces binding the particles
together are weakened. The particles are able to move around more freely. This results in a flexible shape
The solid shape gradually changes into a liquid one. The temperature at which this transition from solid to
liquid is the melting point of the solid.
If the liquid is further heated, the energy of the particles is increased even further. This causes faster
movement and the subsequent thinning of the liquid. A stage is reached when the particles at the surface
acquire enough energy to overcome the forces of attraction holding them together. They escape to form a
gas. The temperature at which this transition from a liquid to gas is known as the boiling point. The process
of conversion from liquid  gas is called evaporation
The process from solid  liquid  gas can be reversed by decreasing the temperature of the gas. This
results in the decrease of average energy of the gas particles, making the forces of attraction stronger. As a
result, the gas starts to condense into a liquid. If this liquid is further cooled, the energy of the particles is
decreased to such an extent that it is fully overcome by the attractive forces and the liquid freezes to
become a solid.
Experiment techniques
Paper chromatography
 a method used to separate and distinguish the substances within a mixture.
 used to separate the colouring agents and dye in products like food and ink.
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The mixture you want to separate is placed on the edge of strip of paper, right above the water.
In this case, it is a black ink spot.
The strip of paper is then dipped in the solvent (water).
The water level slowly rises, taking the dye molecules with it.
Depending on how strong the attraction is between the dye molecules and the water, it will
move more slowly or quickly as the water level rises.
If the ink is a mixture of two or more dyes, then the different colours will move at different
rates.
This separates the different components of ink, allowing us to identify the colours used.
Filtration
 a method of separation used to remove undissolved solids from liquids.
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Here, filter paper and a glass funnel is used to separate the solid from the liquid.
The filter works by allowing the liquid to flow through the paper, but not the solid.
This is because the particles of the solid are too big to pass through the minute holes in the filter
paper.
At the end of the filtration, you are left the residue of the solid on the filter paper and the liquid in
another container.
Crystallization
 Used to separate dissolved solids from liquids, unlike filtration.
How it works:
 The solution is heated in an open container.
 This allows the solvent to evaporate, leaving us with a saturated solution.
 The saturated solution is a solvent which contains as much of the solid dissolved within it.
 The saturated solution is left to cool.
 At this point, crystals of the solid will grow in the solution.
 When the solution has completely evaporated, you are left with the crystallized solids.
Distillation
 a method of separation used to extract a pure liquid from a mixture.
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In this example, we will use a mixture of both ethanol and water.
The mixture is first heated in a round bottomed flask.
The fact that ethanol has lower boiling point is crucial.
At 78oc, the ethanol will evaporate into vapour.
At this point, the vapour will travel down the condenser filled with cool, running water and be
cooled down into a liquid.
The ethanol liquid (distillate) will be collected into a beaker whilst the water is left in the flask.
The ethanol is now separated from the water.
Fractional Distillation
 Is similar to the normal distillation method above.
 Separates a liquid mixture into its individual components
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The above diagram uses the example of crude oil.
Like distillation, the mixture is heated and evaporated.
The substances within the crude oil have different boiling points and they each evaporate at
different temperatures.
The temperature decreases as the gases go up the fractioning column.
The smaller molecules condense at the higher portions of the column at lower temperatures.
Conversely, the larger molecules sink to the bottom and condense at higher temperatures
Importance of Purification
 Purifying and separation methods allow us to obtain specific substances and chemicals needed for our
everyday life. Like above, the separation of crude oil into its components is essential, as the petrol
allows us to keep our cars running and the naphtha is used for making plastics.
 Purifying chemicals and substances is important, especially if we ingest it. Impurities in things like drugs
and food could be potentially hazardous and detrimental to our health if swallowed.
Atomic structure and the Periodic table
Definition
Element: A pure chemical substance consisting of one type of atom
Compound: A pure chemical substance consisting of two or more different chemical
Elements
Mixture: The physical combination of two or more substances that are not combined
chemically
Atomic number: The number of protons/electrons in an atom.
Mass number: The total number of protons and neutrons in an atom. Also known as
nucleon number.
Key ideas
 All elements are made up from atoms
 An atom consists of subatomic particles. They include:
o Proton – overall charge of +1
o Electron – overall charge of -1
o Neutron – no charge; it is neutral
 Because an atom has no overall charge, the number of protons in any atom is equal to the number
of electrons
An atom
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The small circles within the large circle is located in a region called the nucleus.
The nucleus is home to the protons and neutrons.
Orbiting the nucleus are the electrons. These are called “Electron Orbitals”
Atomic number and mass number
Atomic no.
Element symbol
Mass no.
Therefore a carbon-12 contains 6 protons (ie. the atomic number = 6), 6 electrons and 6 neutrons. This is
sometimes written as:
...where the atomic is written under the mass number.
General structure of an atom.
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2 electrons go in the first shell
8 electrons go in the second shell
8 electrons go in the third shell
2 electrons go in the fourth shell
Some notes
 10 more electrons can go in the third shell (only) after the fourth shell has been completed
 8 more electrons can go in the fourth shell (only) after the third shell has been completed
Structure of the periodic table
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The vertical columns in the periodic table are called groups. The elements in the groups have
similar properties
Each group indicate the number of electrons in its outside shell…with the exception of group 8
where it has a full outer shell.
The horizontal rows of elements are called periods.
The periods state number of electron shells
The main block (in dark blue) are the transition metals
The line between the green and purple block divide metals on the left and side and non-metals on
the right.
Energy levels:
o 1st energy level holds up to 2 electrons
o 2nd energy level holds up to 8 electrons
o 3rd energy level holds up to 8 electrons.
The groups:
o Group 1: Alkali Metals
 1+ charge
o Group 2: Alkali Earth Metals
 2+ charge
o
o
o
o
o
o
o
Group 3 (not important)
 3+ charge
Transition metals charges are indicated
Group 4:
 4+ / 4- charge
Group 5:
 3- charge
Group 2:
 2- charge
Group 7: Halogens
 1- charge
Group 0: Noble Gases
 0 charge (inert)
Ions and Ionic bonding
Definition
Cation: an ion with fewer electrons than protons, giving it a positive charge. It is a
usually a metal.
Anion: An ion with more electrons than protons, giving it a net negative charge. It is
usually a non-metal
Ion: An atom or molecule in which the total number of electrons is not equal to the
total number of protons, giving it a net positive or negative electrical charge.
Ionic bonding: A type of chemical bond formed through an electrostatic attraction
between two oppositely charged ions. Ionic bonds are formed between a
cation and an anion.
Noble gases: The elements in Group 18 of the periodic table. They comprise the
elements helium, neon, argon, krypton, xenon, and radioactive radon.
These gases are found as single atoms as they do not like to interact with
other atoms on the periodic table. Their outer-most electrons shells
(valence shells) are completely full with eight electrons with the exception
of helium having two
Electron configuration: The distribution of electrons of an atom or molecule (or other
physical structure) in atomic or molecular orbitals
Lattice: A regular repeated three-dimensional arrangement of atoms, ions, or
molecules in a metal or other crystalline solid
Examples
Sodium chloride
 A sodium atom has an electron arrangement of 2,8,1
o This means that it has one more electron than the stable electron arrangement of 2,8
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A chlorine atom has an electron arrange of 2,8,7
o This means that it is one electron short of a stable electron arrangement of 2,8,8
If the chlorine atom “takes” an atom from the sodium atom, they both will have full electron
shell.
The sodium atom now loses one electron. This means that there is an overall positive charge as
there is there is more protons (positive charge) than electrons (negative charge)
o The sodium atom is an cation
The chlorine atom now gains one electron. This means that there is an overall negative charge
as there is there is less protons (positive charge) than electrons (negative charge)
o The sodium atom is an anion
One electron
transferred from
Na to Cl
Na+Cl- is formed
Sodium chloride lattice
Magnesium oxide
 Electron arrangement in magnesium atom 2,8,2
o Two electrons more for a full shell of 2,8
 Electron arrangement in oxygen atom 2,6
o Two electron short for a full shell of 2,8
 This means the magnesium atom has to give 2 electrons to the oxygen atom
Two electrons
transferred from
Mg to O
Mg2+Cl2- is formed
Some common mistakes and tips
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It is incorrect to speak of a ‘sodium chloride molecule’, for example.’ This would assume that one
sodium ion joins with one chloride
A frequent mistake is to use terms such as atoms “swapping” electrons. THIS IS WRONG
It is important to stress in your answers that there is a complete transfer of electrons in ionic
bonding. Electrons go from the metal (e.g. sodium) to the non-metal (e.g. chlorine)
Covalent bonding
Definition
Volatilitiy: The tendency of a substance to vaporize.
Solubility: Degree to which a substance dissolves in a solvent to make a solution
Molecule: An electrically neutral group of two or more atoms held together by
covalent chemical bonds.
Covalent bonding: A form of chemical bonding that is characterized by the sharing of
pairs of electrons between atoms. The stable balance of attractive
and repulsive forces between atoms when they share electrons is
known as covalent bonding.
Some key ideas
 Covalent bonding involves the sharing of electrons rather than the complete transfer.
 When a pair of electrons is shared, it is called a single covalent bond or single bond.
 When 2 pairs of electrons are shared, it is called a double covalent bond or double bond.
 When 3 pairs of electrons are shared, it is called a triple covalent bond or triple bond.
Examples
Chlorine molecule (Cl2)
 A chlorine atom has an electron arrangement of 2,8,7.
 When two chlorine atoms bond together they form a chlorine molecule.
 If one electron was transferred from one chlorine atom to the other, only one atom could achieve a
stable electron arrangement.
 Instaed, one electron from each atom is donated to form a pair of electrons which is shared
between both atoms
o This is known as single covalent bond.
This holds them together.
 Often shown as Cl–Cl
Two electrons “shared”
between both Chlorine
atoms.
Cl2 is formed
Oxygen molecule (O2)
 Oxygen atoms have a electron arrangement of 2,6.
 Each oxygen atom donates two electron
 The four electrons (2 pairs) are shared between both atoms
o This is known as double covalent bond
 Often shown as O=O
Two electrons “shared”
between both Oxygen
atoms.
O2 is formed
Other examples
Water
Methane
H2O –Single covalent bond
CH4 – Single covalent bond
Nitrogen
N2 – Triple covalent bond
Comparisons between covalent compounds and ionic compounds
Differences and similarities of covalent and ionic compounds
Similarities
No full outer shell
Ended up with full outer shell
Ended up with compounds
Stable at the end
Differences
Covalent
Atoms combines to another atom
Non-metals only
Sharing pairs of electrons
Ionic
Has electrostatic attraction as one atom is positive
when it loses electrons and one atom is negative
when it gain electrons
Metal and non-metals
Transfer of electrons
Properties
Examples
Solubility
Ionic
Salt, sugar, copper sulphate
Soluble in water
Covalent
Sand, sulphur
Insoluble (sometimes soluble in organic
solvents
Conductivity
Do not conduct electricity when solid Do not conduct electricity
but do when dissolved
Macromolecules
Metallic bonding
Definition
Lattice: A molecular compound is simply 2 or more different types of molecules/atoms
put together.
Macromolecules: A molecule containing a very large number of atoms.
Giant covalent structure: An element made with very strong bonds between the atoms
to create various materials.
Allotropes: The property of some chemical elements to exist in two or more different
forms, known as allotropes of these elements.
Key ideas
 Found only in metals
 A metal consists of close-packed together regular arrangement of positive ions
o It is surrounded by a ‘sea’ of electrons that bends the ions together
An example
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Bonding
Around any one ion in each layer, there is six ions in a hexagonal manner
Metals are malleable = they can be hammered into any shape.
o This is due to the mobility of the electrons; the layers of atoms can slide over each other
without fracturing the structure.
Structure
Bonding, structure and properties
Properties
Ionic
Giant structure
Sodium chloride (NaCl)
Magnesium oxide (MgO)
Covalent
Molecular
Chlorine (Cl2)
Iodine (I2)
Methane (CH4)
Macromolecules (large molecules)
Polyethene [(C2H4)nH2]
Starch [(C6H12O6)n]
Giant structure
Silicon dioxide (SiO2/Si2O4)
Metallic
Giant structure
Copper (Cu)
High melting and boiling point
Usually soluble in water
Insoluble in organic solvents
Conduct electricity when molten or dissolved in water
(electrolyte)
Usually gases or low boiling point liquids
Some (iodine/sulphur) are low melting point solids
Usually insoluble in water
Soluble in organic solvents
Do not conduct electricity
Solids.
Usually insoluble in water
More soluble in organic solvents
Do not conduct electricity
Solids
High melting points
Insoluble in water and organic solvents
Do not conduct electricity
Solids
High density (ions closely packed)
Good electrical conductors (free electrons)
Stoichiometry
Definition:
Stoichiometry: It deals with the relative quantities of reactants and products in chemical
Reactions
Reactants: substance or compound that is added to a system in order to bring about a
chemical reaction, or added to see if a reaction occurs.
Products: formed during chemical reactions as reagents are consumed
Relative atomic mass: The mass of an atom of a chemical element expressed in atomic
mass units. Written as Ar
Relative molecular mass: The sum of the relative atomic masses of the constituent
atoms of a molecule. Written as Mr
Balancing equations
Key ideas
 Reactions can been summarised using chemical symbols
 All reactions have to be balanced
 Law of conservation of mass : Mass is neither created nor destroyed in ordinary chemical and
physical changes.
Examples
Reacting calcium hydroxide with hydrochloric acid
 Two reactants
o Calcium hydroxide symbol = Ca(OH)2
o Hydrochloric Acid = HCl
 Produces two products:
o Calcium chloride = CaCl2
o Water = H2O
The equation is written as:
Ca(OH)2 + HCl CaCl2 + H2O
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You cannot alter the formula because of the law of conservation of mass but they can be
rearranged to balance the equation.
The equation is unbalanced because there are different numbers of atoms on each side:
Left-hand side
Right-hand side
1 Ca
1 Ca
2O
1O
3H
2H
1 Cl
2 Cl
So by “balancing” the equation it will become:
Ca(OH)2 + 2HCl CaCl2 + 2H2O
Thus you will have:
Left-hand side
Right-hand side
1 Ca
1 Ca
2O
2O
4H
4H
2 Cl
2 Cl
State symbols
Key ideas
 Sometimes state symbols are added to symbol equations to show wheather the substance is sold,
liquid or gas or whether it is in a solution
 The state symbols
o (s) solid
o (l) liquid
o (g) gas
o (aq) in aqueous solution, ie. where the solvent is water
Calculating RAM and Relative Molecular Mass
Key ideas
 The relative atomic mass (Ar) is simply a number and has no units
 For compounds, you can use the relative atomic mass to find the relative formula mass.
Examples
1) Find the mass of H2O, given Hydrogen = 1, Oxygen = 16
Mr = (2x1)+16 = 2+16 = 18
The mole concept
Definition
Molecular formula: The simplest way of representing the actual number of atoms in a
molecule (e.g. H2O for water)
Mole: A quantity or amount of particles (6 x 1023) which is fixed for any substance (unit
= mol)
Avogadro’s Constant: The number of particles in 1 mole of any substance.
Molar mass: The mass (in grams) of 1 mole of any substance.
Mass number: The number of nucleons in an atom of an element (protons + neutrons)
Relative atomic mass: The mass of an atom relative to the mass of a carbon atom (12) –
sometimes abbreviated to Ar
Relative molecular mass: The mass of a formula calculated by taking the sum of the
relative atomic masses. Sometimes abbreviated to Mr.
Empirical formula: The simplest ratio of atoms of an element in a chemical formula, as
worked out by experiment.
Some formulae
Moles
Mass / g
No. of moles
x
Molar mass / g/mol
Volume of a gas
Volume / dm3
No. of moles * 24mol/dm-3 (At RTP and STP)
Concentration of a solution
Moles
Concentration (M) * Volume (dm3)
Some word equations
1) Metal + acid → salt + hydrogen
2) Metal carbonate + acid → salt + carbon dioxide + water
3) Acid + alkali → salt + water
4)
Metal + salt solution → (When metal is more reactive) Displacement reaction
Metals
Definition
Displacement reaction: A reaction in which a more reactive element displaces a less
reactive element from solution.
Malleable: Metals are very malleable as they can be beaten into thin sheets or different
shapes.
Metal: An element that is shiny, conducts heat and electricity, is malleable and is ductile
is probably a metal. Metals usually have high melting and boiling points, and high
densities.
Metallic bond: An electrostatic force of attraction between the mobile ‘sea’ of electrons
and the regular array of positive metal ions within a solid metal.
Alloy: A metal made by mixing two or more molten metals together e.g. brass
The presence of different sized metal atoms in the structure of an alloy increases
its strength, as the atoms are no longer able to slide over each other. In a pure
metal all atoms are of the same size.
Reactivity series: An order of reactivity, giving the most reactive metal first, based on
results from experiments with oxygen, water and dilute HCl.
Rusting: Iron (III) oxide.
Sea of mobile electrons: The positive metal ions exist in a giant structure surrounded by
valence electrons that are not attached to any particular ion.
They (the electrons) are able to move freely and allow the
metal to conduct electricity, and also allow layers of ions to be
malleable.
Steel: An alloy between metallic iron and non-metallic carbon.
Flame tests: Different metal compounds produce characteristic colours in a flame.
Cation: A positive ion
Precipitate: A solid formed when two solutions are mixed. Different metals form
characteristic coloured precipitates in sodium hydroxide and ammonia.
General properties of metals
 Solid
 Hard
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Strong
Shiny
Malleable
Tough (endures)
Feel cold
Conduct electricity and heat
Dense
Magnetic - Iron, Nickel and Cobalt.
Sonorous
Expand on heating
React with oxygen
React with acids
Formula
Alkali metal + water = Hydrogen + Metal hydroxide
Aluminium
 Very reactive, very quickly reacts with Oxygen. However, Aluminium Oxide is quite unreactive.
 Used in food containers because it is resistant to corrosion
 Alloys used for fuselage of planes, wings etc. because they are low in density.
 Is strong, light and malleable.
Galvanize
 To coat iron or steel with a protective coat of zinc.
 Zinc is more reactive - therefore it reacts instead of iron / steel.
 Sacrificial protection
The Reactivity Series
METAL
ION
Cesium
Cs+
Rubidium
Rb+
REACTIVITY
reacts with water
Potassium
K+
Sodium
Na+
EXTRACTION
electrolysis
Lithium
Li+
Barium
Ba2+
Strontium
Sr2+
Calcium
Ca2+
Magnesium Mg2+
reacts with acids
Aluminium Al3+
Carbon
included for comparison
Manganese Mn2+
Zinc
Zn2+
Chromium
Cr2+
reacts with acids
Iron
Fe2+
Cadium
Cd2+
Cobalt
Co2+
smelting
with coke
Nickel
Ni2+
Tin
Sn2+
Lead
Pb2+
Hydrogen
H+
Antimony
Sb3+
Bismuth
Bi3+
Copper
Cu2+
Mercury
Hg2+
Silver
Ag+
Gold
Au3+
Platinum
Pt2+
Included for comparison
heat or
may react with some strongly oxidizing acids physical
extraction
The blast furnace
The formulas.
1) Carbon + Oxygen → Carbon Dioxide
C + O2 → CO2
2) Carbon Dioxide + Carbon → Carbon Monoxide
CO2 + C → 2CO
3) Iron Oxide + Carbon Monoxide → Iron + Carbon Dioxide
Fe2O3 + 3CO → 2Fe + 3CO2
4) Carbon + Iron Oxide → Iron + Carbon Dioxide
3C + 2Fe2O3 → 4Fe + 3CO2
5) Calcium Carbonate → Calcium Oxide + Carbon Dioxide
CaCO3 → CaO + CO2
6) Calcium Oxide + Silicon Oxide → Calcium Sillicate
CaO + SiO2 → CaSiO3
Redox Reaction
Definition
Mineral: A naturally occurring substance of which rocks are made.
Melt: A solid changes to a liquid at the boiling point.
Ore: A naturally occurring mineral from which a metal can be extracted.
Oxidation: A reaction in which oxygen is added to an element or compound
A reaction involving the loss of electrons from an atom, molecule or ion
A reaction in which the oxidation number of an element is increased.
Brought about by oxidation agent.
Reduction: A reaction in which oxygen is removed from a compound
A reaction involving the gain of electrons by an atom, molecule or ion
A reaction in which the oxidation number of an element is decreased.
Brought about by reducing agent.
Reactivity Series: Order of reactivities of different elements.
Extraction: Metals are extracted from their ores by a process of reduction. Carbon can
be used to do this if the metal is below carbon in the reactivity series.
Electrolysis
Definition
Electron: Sub-atomic particles which pass the electric current in a circuit
Ion: An atom which has become charged, by either losing or gaining electrons
Electrolysis: Using electrical energy to separate ions from an ionic compound, and turn them into neutral
elements. Often used as a method for obtaining a reactive metal from its ore
Electrode: A conducting material which passes an electrical current into the substance being electrolysed
Inert: Unreactive
Anode: The electrode connected to the + terminal in electrolysis (The site of electron
loss / oxidation)
Cathode: An electrode through which electric current flows out of a polarized
electrical device
Ore: A rock that contains reacted metal
Key ideas
 Electrolysis is an important process. It was developed by Michael Faraday in the early nineteenth
century when suitable supplies of electricity became available. He discovered a number of new
elements, e.g. the alkali metals and established the basic laws of electrolysis.
 The common examples of solids conducting electricity are metals and graphite (a form of carbon).
 Electricity passes through metals and graphite because of freely moving (delocalized) electrons.
 Electricity passes through metals and graphite with no chemical change taking place.
 Melting an electrolyte or dissolving it in water breaks up the structure and the ions are free to move.
 Electrolysis is the splitting up of an electrolyte when molten or in solution.
 Ions move (or migrate) towards the electrode of opposite charge
 A cell consists of two different metals dipping in a solution that conducts electricity (electrolyte). In the
cell, chemical energy is converted into electrical energy. The more reactive metal becomes the negative
pole of the cell from which electrons flow.
 The further apart the metals are in the reactivity series the higher will be the voltage of the cell.
 The most reactive or a halogen would get made if the electrolyte is an aqueous solution at the anode
 In the cathode, the least reactive gets made instead.
Electrolysis of molten lead(II) bromide
Key ideas
 The positive electrode has a shortage of electrons and the negative electrode has a surplus of electrons.
Electrons are constantly flowing through the wire from the positive electrode to the negative electrode
 At the negative electrode (cathode) electrons are transferred from the cathode to the ion and the ion is
changed to a metal atom.
 At the positive electrode (anode), electrons are transferred from the ion to the anode and the ion is
changed to a bromine atom. Two bromine atoms combine to form a bromine molecule.
What you will see
 Lead starts to from around the negative electrode
 Bromine gas is produced at the positive electrode
 The overall reaction is:
PbBr2  Pb + Br2
Migration of ions
When lead bromide is melted, positive lead ions, Pb2+, moves towards the negative electrode.
The negative bromide is, Br-, move towards the positive electrode
Discharging of ions
When the ions arrive at the electrodes, they are discharged.
At the cathode
Pb2+ + 2e- → Pb ⇐ oxidisation
At the anode
Br- → Br2 + 2e- ⇐ reduction
How to remember
Cathode
 Remember RED METAL CAT
o Reduction occurs
o The metal is made
Anode
 Remember AN OX
o Oxidation occurs
Energetics of reactions
During all chemical reactions, an energy change occurs. In the reaction, heat is either released or absorbed.
When a reaction releases heat to the surroundings, we call that reaction an Exothermic Reaction. The
reaction that absorbs energy from the surroundings are called Endothermic Reactions.
Exothermic Reaction
The reactants have more energy than the products here, so a small amount of energy is required to activate
the reaction.




Release of heat
Energy needed for the reaction to occur is less than the total energy released.
Extra energy is released, usually in the form of heat.
The release of heat means that an exothermic reaction increases temperature of the surroundings.
Endothermic
 Heat absorbs energy from the surroundings.
 Temperature of surroundings decreases during an endothermic reaction because energy from
surroundings is required to drive the reaction, hence decreasing the temperature of the
surroundings.
Demonstrate understanding that exothermic and endothermic changes relate to the
transformation of chemical energy to heat (thermal energy).
In order to actually start a reaction, a certain amount of energy will be provided to the reactants; We often
call this the Energy of Activation because this energy is essentially required to start the reaction.
The energy here is used to break the bonds between the molecules of the atoms of the reactants. The
bonds then subsequently rearrange and bond again, which releases energy.
However, if the energy provided to activate the energy is less than the energy released when the bonds
form together, the reaction gave out more than it took/absorbed, which makes this a exothermic reaction.
If the energy given to activate is more than the energy released during the bond formation, the reaction is
endothermic.
The total energy change is called enthalpy.
Kinetics
Definition
Catalyst: A substance that alters the rate of reaction but remains chemically unchanged
at the end of the reaction.
Activation energy: The minimal amount of energy needed for a chemical reaction to
take place.
Key ideas
 The theory used to explain why chemical reactions occur is called the collision theory
 For a reaction to take place, the particles of the substances that are reacting have to collide. If they
collide, with enough energy then they will react.
 The minimum amount of kinetic (movement) energy that two particles need if they are going to
react when they collide is called the activation energy.
 There are therefore two main ways of increasing the rate of a reaction:
1. Increase the number of collisions
2. Increase the amount of kinetic energy so that more collisions lead to a reaction.
Calculating rates of reactions
The rate of reactions can be calculated in the following raction
Rate of reaction ∝ 1/time
For practical reasons, reactions used in the laboratory for studying rate of reaction must not be too fast or
too slow.
Having selected a suitable reaction, it is necessary to find a change that can be observed during the reaction.
An estimate of the rate of reaction can be found from the time for a measurable change to take place.
Factors affecting rate of reaction
Factor
Reactions affected
Temperature
All
Concentration
Pressure
Light
Particle size
Change made in conditions
Effect on rate of reactions
Increase 10°C
Approx. doubles rate
Decrease 10°C
Approx. halves rate
All
Increase in concentration of one Increases the rate of reaction
of the reactants
Reactions
involving Increase the pressure
Greatly increases the rate of
mixtures of gases
reaction.
Wide
variety
of Reaction in sunlight or UV Greatly increases the rate of
reactions
including sunlight
reaction
reactions with mixtures
of
gases
including
chlorine or bromine
Reactions
involving Using one or more solids in a Greatly increases the rate of
solids and liquids, solids powdered form
reaction
and gases, or mixtures
Using a catalyst
of solids
Adding a substance to a A specific substance which Increases the rate of reaction.
reaction mixture
speeds up the reaction without
being used up
 Surface Area
Increasing the surface area will subsequently increase the rate of reaction, as increasing surface area will
increase the chances of particles colliding with each other and will hence increase the rate of reaction.
In the exam, they might ask you questions like, “which reacts faster, magnesium or magnesium powder”.
The obvious answer is the powder because the powder has a much larger surface area, hence increasing
the rate of reaction.
Concentration
Increasing the concentration increases the rate of reaction, as there will be collisions per second per unit
volume.
The reason for this is increasing the concentration results in there being more particles in each cm 3 of space,
so there will be more frequent collisions between particles.
As the reaction occurs and the reactants get slowly used up, the concentration of the substance then
decreases. This explains for a slower rate of reaction as the reaction proceeds for a period of time.
Temperature
Increasing the temperature will increase the rate of reaction. There are two important reasons for this:
1. Particles will move faster and have more kinetic energy so there will be more collisions per second.
2. More colliding particles will have the necessary activation energy required, hence allowing more
successful collisions.
The second reason is a more important factor in explaining the increased rate of reaction than that of the
first.
Catalysts
Adding a catalyst increases the rate of reaction, but it itself is not used up in the reaction. Catalyst speed up
the reaction by lowering the activation energy or providing an alternative pathway for the reacting particles.
Practical method
There are two ways to measure the rate of reaction. First, let’s define what a reaction is.
Reactant  Product
Above is a reaction summed up to its essence. A reactant forms a product. The reactant is what you start
off with. It could be an acid and alkali or a metal and an acid. As you start the reaction, the reactants
slowly get used up. The reactants slowly become products.
Therefore, the process of where reactants become the products is the “reaction process”.
Hence, we can measure speed of reaction through two factors:
1. Speed at which reactants are used up
2. Speed at which products are formed.
When a gas is produced during a reaction, we can easily measure the reaction by measuring the “Volume of
the gas produced”.
An example of such reaction is:
Magnesium + Hydrochloric –> Magnesium Chloride + Hydrogen
We can measure the volume of hydrogen produced. However, to do this, we need to devise an experiment
suitable for high school students.
The experiment proposed to measure the volume of gas produced is described below:
Apparatus
 Gas syringe
 Excess Dilute Hydrochloric Acid
 Magnesium
 Stopwatch
 Conical Flask
Diagram of Experiment
Method
1. Clean Magnesium with sandpaper so ensure any impurities are cleaned off.
2. Put the dilute hydrochloric acid into the flask.
3. Add the magnesium into the flask.
4. Simultaneously, add the stopper + gas syringe.
5. Start stopwatch.
6. Measure the volume moved by the plunger every minute.
We will talk about the interpretation of this data in Point 4. Obviously, you are going to produce hydrogen
in this reaction. The hydrogen will show itself in the form of bubbles, and these bubbles will rise up the flask
and into the gas syringe, hence pushing the plunger.
Practical method 2
Concentration
Remember, we did the experiment on Point 2. Repeat the experiment again, however, this time use two
different types of concentrations of HCl.
For one experiment use, “x” concentration of HCl
For the second experiment, use “2x” concentration of HCl
Temperature
Again, repeat the experiment but this time with two different types of temperatures of HCl, and compare
the differences of Volume of Gas produced.
As you use a higher temperature, we see a steeper graph, hence concluding that a higher temperature
leads to a higher rate of reaction.
Surface Area
Repeat the experiment again but this time use two different sizes of magnesium. The magnesium’s you
should use are:
1) Normal magnesium chips
2) Magnesium chips of same mass, but smaller pieces.
Catalysts
Again, repeat the experiment two times again, once with a catalyst, and once without.
Interpreting data
We did the experiment in Step 2. Now we have to process and collect our results.
Below are a possible set of results we might have received from the experiment.
Time (minutes)
Volume of Hydrogen
1
13
2
27
3
34
4
39
5
41
6
41
7
41
Let’s
graph our results:
How can we calculate rate of reaction?
We use this simply formula:
Rate of Reaction = Volume of Gas Produced / Time.
Let’s notice a few things:
 The rate of reaction is greatest when the graph is the steepest, notably at the start of the reaction
 The reaction is finished when the graph becomes horizontal, when there is no further increase in
the volume of gas.
We can’t really calculate the instantaneous change in rate of reaction, but we can calculate
the average rate of reaction.
Average Rate of Reaction = Total Volume of Hydrogen / Total Time
= 41cm3/ 7 minutes
= 41/7 cm3 / per minute
Effects of temperature and concentration to collision
Temperature:
Concentration:
Explaining different rates using particle model
Particles in solids, liquids and gases are moving. This movement is much greater in gases than in liquids and
in liquids more than solids.
In a reaction mixture, the particles of the reactants collide. Not every collision leads to a reaction. Before a
reaction occurs, the particles must have a sufficient amount of energy, known as the activation energy. If a
collision between particles can produce sufficient energy, if the collide fast enough and in the right
direction, a reaction will take place. Not all collisions will result in a reaction.
A reaction is speeded up if the number of collisions is increased.
Increasing the concentration
If concentration is increased, there are more collisions between particles and so there are more collisions
leading to reaction and the reaction is faster. Increasing concentration increases rate of reaction. The
reason for this is increasing the concentration results in there being more particles in each cm 3 of space, so
there will be more frequent collisions between particles.
Increasing the pressure can be explained in the same way, because increasing the pressure of a mixture of
gases increases the concentration by forcing the particles closer together.
Increasing the temperature
The temperature of will give the particles more kinetic energy, so the frequency of collisions between
particles will increase and the number of successful collisions will also increase. Using sunlight or ultraviolet
light has the same effect as increasing temperature.
Using smaller pieces of solid
When one of the reactants is a solid, the reaction must take place on the surface of the solid. By breaking
the solid into smaller pieces, the surface area is increased, giving a greater area for collisions to take place
and hence causing an increase in the rate of reaction.
Effect of light
There are certain chemical reactions which can take place only in the presence of light. The most common
examples are photosynthesis and photographic films.
Photosynthesis is the process in which water absorbed by roots of the plants reacts with carbon dioxide
taken in by green plant leaves, in the presence of light, to form glucose and oxygen. The rate of
photosynthesis increases with the increase in the intensity of light and vice versa. The green pigment
chlorophyll, present in plant leaves, acts as a catalyst in the process.
6CO2 + 6H2O → C6H12O6 + 6O2
Carbon dioxide + Water (+ Light energy) → Glucose + Oxygen
In conventional photography, the photographic film used is a transparent plastic strip coated with an
emulsion containing tiny crystals of silver bromide (AgBr). When exposed to light, the silver cations present
in silver bromide crystal accept an electron from the bromide ions (Br) and silver atoms are produced
Ag+ + e- →Ag
Silver ion + electron → Silver atom
The amount of silver deposited depends on the intensity of light falling on the photographic film and thus,
depending on the light reflected by various parts of the subject being photographed, the amount of silver
deposited in different portions of the film is different. This is how the details of the subject are captured on
film and that is why the lighter shades appear darker and vice versa in negatives of photographs.
Catalysts
Catalyst
Catalysts usually speed up reactions. A catalyst which slows down a reaction is called a negative catalyst or
inhibitor.
Manganese(IV) oxide catalyzes the decomposition of hydrogen peroxide into water and oxygen:
2H2O2(aq) → 2H2O(l) + O2(g)
Catalysts are often transition metals or transition metal compounds. The catalyst provides a surface where
the reaction can take place
Using a catalyst lowers the activation energy for the reaction. This implies that more collisions have
sufficient energy for reactions to take place.
Application
We know that increasing the temperature, concentration and surface area of a substance can increase the
rate of reaction. However, it’s important that you don’t overdo it. Overdoing any one of these factors can
lead to some severe consequences, including explosions.
Here are examples of places where adding too much of any of these factors can lead to some serious
consequences:

Flour Mills: Flour particles are very tiny, just like all particles. Hence, flour particles tend to have a
large surface area. If there is a lot of flour dust in the air, a spark from the machine is enough to
cause a reaction between the flour and the spark to form an explosion.

Coal Mines: In a coal mine, you have all sorts of flammable gases in the air. At the right
concentration, these gases form an explosive mix with the air, and this is enough to set off an
explosion.
Enzymes
Enzymes are biological catalysts. Hydrogen peroxide is decomposed into water and oxygen by an enzyme
in fruits and vegetables. Catalase is a protein. Unlike chemical catalysts, such as manganese (IV) oxide,
catalase works only under particular conditions. It works best at 37°C. At higher temperatures the protein
structure is permanently changed (denatured): it no longer decomposes hydrogen peroxide.
Uses
Enzymes are used in many industrial processes like:
 Fermentation of solutions of starch and sugar using enzymes in yeast to produce beer and wine
 Making cheese and yogurt by the action of enzymes on milk
 Enzymes (proteases and lipases) in washing powders break down protein stains in cold or warm
water


Soft-centered chocolates are made by injecting hard-centered chocolates with the enzyme
invertase
Isomerase is used to turn glucose syrup into fructose syrup. This is sweeter and can be used in
smaller quantities in slimming products.
Organic chemistry
Names of compounds and homologous series
Organic compounds
 Organic compounds are those with molecules containing carbon and one or more elements
 There is a wide variety of organic compounds, depending on the length of the carbon chain, the
elements it is attached with and their structural arrangement in space.
 All living things are made of organic compounds
Homologous series
 Homologous is derived from Greek homos (same) and logos (ratio).
 It is a series of chemical compounds having the same functional group but differing in composition
by a fixed group of atoms.
 Functional groups determine most of the chemical properties of compounds
 Compounds in a homologous group have the following points in common:
o Similar name endings
o Similar chemical structure
o Similar chemical behaviour
o Can be represented by a general formula
 The properties in a homologous series change with the change in number of carbon atoms per
molecule.
 The more the number of carbon atoms present, the higher the melting and boiling points
Most of the simple organic compounds can be broadly classified into the following homologous groups.
1. Alkanes: The names of chemicals in this group ends with ‘-ane’.
2. Alkenes: The names of in this group end with ‘-ene’
3. Alcohols: The names in this group end with ‘-ol’
4. Carboxylic acids: The names in this group end with ‘-oic acid’
Most of the man-made organic compounds are derived from crude oil, which along with coal and natural
gas, is the most widely used fuel at present.
Isomerism
Definition: The existence of two or more compounds with the same molecular formula
but different structural formulae.
With alkanes containing up to three carbon atoms, there is only one possible structure for each molecular
formula.
However when there are four or more carbon atoms in an alkane, it is possible to have different structures.
Two isomers of butane are:
Butane
2-methylpropane
These two isomers have similar chemical properties but different physical properties. In each structure,
every carbon atom is joined to 4 other atoms and every hydrogen atom is joined to one other atom.
Isomerism becomes more common with higher alkanes, e.g. 75 isomers of decane C10H22. Isomers can also
occur with other homologous series (families) e.g. alkenes and also with compounds in different
homologous series e.g. C2H5OH (an alcohol) and CH3OCH3 (an ether).
Separating crude oil
Some key ideas
 Crude oil is a mixture of hydrocarbons
 Hydrocarbons are compounds of carbon and hydrogen only.
 Most hydrocarbons belong to a family called alkanes.
The alkane family
Name
Formula
Structure
State at room Melting point
temp.
Gas
Increases down
the family
Methane
CH4
Ethane
C2H6
Gas
Propane
C3H8
Gas
Butane
C4H10
Gas
Pentane
C5H12
Liquid
Boiling point
Increases
down the
family
A diagram
What happens
 When you heat the crude oil, the compounds start to evaporate as particles will more kinetic
energy and therefore will more likely be able to break bonds. The compounds which are smaller
and lighter evaporate first as it takes less energy to evaporate these.
 The hot vapour rises and the vapour then condenses in the cool test tube.
 When the thermometer reaches 100 degrees, the first test tube is then replaced with an empty one.
The liquid in the first test tube is the first fraction from the distillation.
 Repeat, replacing the test tube at 150 degrees, 200 degrees, and 250 degrees.
Alkanes and Alkenes
Some key ideas
 If a hydrocarbon has 1 carbon atom in its chemical makeup, its name begins with a “meth-“
 If a hydrocarbon has 2 carbon atoms in its chemical makeup, its name begins with a “eth-“
 If there are single bonds between carbons in it chemical makeup, its name ends with “-ane”
 If there are double bonds between carbons in it chemical makeup, its name ends with “-ene”.
 Alkanes all fit a formula of CnH2n+2
 Alkanes burn in air or oxygen
Some alkanes and alkenes
CH4
C2H4
Simple alkanes and alkenes
Alkenes
Key ideas
 Alkenes are a series of homologous organic compounds
 The general formula is: CnH2n
 They are hydrocarbons like alkanes, but are more reactive than them.
 They possess a double covalent i.e. unsaturated bond.
One of the two bonds in double covalent can be broken to add extra atoms in the molecule i.e. the bond is
not fully saturated and there is scope for other atoms to be added, hence it is called unsaturated.
Name
Formula
Molecular Structure
Ethene
C2H4
Propene
C3H6
Butene
C4H8
Pentene
C5H10
Testing for alkenes
A bromine test for alkenes is used to distinguish from alkanes. Alkanes are unreactive except for
combustion and do not react with bromine water. Alkenes, however, will readily decolorize bromine
water.
The equation is as follows:
Alkene + Br2 (aq)  bromoalkane
Note: Alkenes decolourises bromine water. Alkanes don’t.
Alcohol
Key ideas
 Alcohols are derived from an alkane by removing an –H and adding an –OH
 Ethene is mixed with steam and passes over a phosphoric acid catalyst at 600°C and at high
pressure
 Ethanol can be prepared by the fermentation of sugar solutions using enzymes in yeast
 When perfume is sprayed on the skin, the cooling felt is due to the evaporation of ethanol.
Some alcohols
Name
Formula
Methanol
CH3OH
Ethanol
C2H5OH
Propanol
C3H7OH
Butanol
C4H9OH
Pentanol
C5H11OH
Molecular structure
Manufacture of ethanol
Large quantities of ethanol are manufactured for industrial use.
There are two methods of producing ethanol – from ethane or from sugar
From ethene
Large amounts of ethane are produced from cracking factions from crude oil. Much of this is used to make
poly(ethane) and ethanol. Formation of ethanol involves an addition reaction
 Cracking breaks longer-chain alkanes into short-chain alkenes, which are useful as a starting point
for plastics.
 This is done by:
o Heat - 400°C to 700°C
o Catalyst – Aluminium oxide
C2H4(g) + H2O (g)  C2H5(g)
Ethene + Steam Ethanol
From ethane
The solution is kept in a warm place for several days. It is actually a dilute solution of ethanol.
C6H12O6(aq)  2C2H5OH(aq) + 2CO2 (g)
Glucose  Ethanol + Carbon dioxide
A more concentrated solution of ethanol is produced by fractional distillation.
The method used to manufacture ethanol depends upon the materials available.
1. In developed countries, such as the United States and in Europe, there are large amounts of ethane
available. Making ethanol from ethene would be preferred.
2. In countries such as Mauritius, which do not have crude oil but do have sugar produced by sugar
cane in large amounts, fermentation would be preferred.
Ethanol made from fermentation process is a batch process. Ethanol is produced from ethane by a
continuous process.
Advantages and disadvantages
Ethanol by fermentation
Advantages
Uses renewable resources
e.g. sugar cane
Uses waste materials
Disadvantages
Ethanol from ethane
Advantages
Fast reaction
Continuous process
Does not need large reaction vessels
Produces pure ethanol
Disadvantages
Large volume needed to produce small amount of
ethanol
Needs large reaction vessels
Fractional distillation is expensive
Fermentation is slow
When ethanol reaches a certain concentration, the
reaction stops.
Uses a non-renewable resource
Energy is needed to produce steam
High percentage of ethane remains unreacted and
must be recycled
Uses of ethanol
1. Ethanol is widely used as a solvent. It is also used in paints, varnishes, perfumes, etc.
2. Ethanol is used as a fuel. In Brazil, either pure ethanol or a mixture of petrol and ethanol is used as
a fuel in cars. Ethanol burns in excess air to form carbon dioxide and water.
C2H5OH +3O2  2CO2 + 3H2O
3. Ethanol is used for making other organic chemicals e.g. ethanoic acid, esters.
4. Ethanol is used in alcoholic drinks.
Different drinks contain different percentages of alcohol
Drink
Approximate percentage of ethanol
Beers
4
Wine
12
Fortified wine
18
(e.g. sherry)
Spirits (whisky)
35
However there are some harmful effects of ethanol including:
 Impaired coordination and judgement
 Slower reaction
 Promotes aggression
 Causes depression and other mental disorders
 Causes ulcers, high blood pressure, brain and liver damage
Pure ethanol cannot be purchased in shops. We can buy methylated spirits. This is ethanol with added
methanol. Methanol is highly toxic. Other substances are added to make it undrinkable and a purple dye is
added as a warning.
Carboxylic acid
Key ideas
 Carboxylic acid form as series of homologous organic compounds
 The general formula is: CnH2n+1COOH
 They are weak acids, with the functional group – COOH
Some carboxylic acid
Name
Molecular formula
Structural Formula
Molecular Structure
Methanoic acid
Ethanoic acid
CH2O2
C2H4O2
HCOOH
CH3COOH
Propanoic acid
C3H6O2
CH3CH2COOH
Butanoic acid
C4H8O2
CH3CH2 CH2 CH2COOH
Pentanoic acid
C5H10O2
CH3CH2CH2CH2COOH
Ethanoic acid, which is the main constituent of vinegar, is also the most well-known carboxylic acid.
Ethanoic acid
Ethanoic acid is a weak acid with te molecular formula of C2H4O2 and it has the structural formula given
below.
It is called ethanoic acid because, like ethane, it contains two carbon atoms. Ethanoi acid is prepared in
industry by passing ethanol and air over a heated catalyst.
C2H5OH + O2  CH3COOH + H2O
Properties of ethanoic acid
Ethanoic acid is a weak acid. It has a similar reactions to other acids.
1. Indicators. Ethanoic acid turns litmus paper red. Solutions of ethanoic acid have a pH value of about
4.
2. Metals. Ethanoic acid reacts with reactive metals
3. Metal oxides. Ethanoic acid reacts with a base to form a salt and water only.
4. Metal carbonates. Ethanoic acid reacts with a carbonate to produce salt and water and carbon
dioxide.
Some formulas
Alkanes
combustion
Alkane + oxygen  carbon dioxide + water
Incomplete combustion
Alkane + little oxygen  carbon monoxide + water
Cracking
Long chain alkane  short chain alkane + alkene
C10H22  C5H12 + C5H10
Alkenes
Bromination
C2H4 + Br2  C2H4Br2
Hydrogenation
C2H4 + H2  C2H6
Hydration
C2H4 + H2O  C2H4OH
Addition polymerisation
Alcohols
Combustion
Macromolecules
Key points
 Cracking involves passing the vapour of the high boiling point fraction over a catalyst at high
pressure.
 Small ethane molecules, produced by cracking, are joined together by a process called addition
polymerisation
 Alkanes are saturated hydrocarbons used as fuels and chemical feedstock. Alkanes can be cracked
to produce alkenes and alkenes can polymerised to form addition polymers
 Thermoplastic polymers e.g. poly(ethene) melt easily when heated. Thermosetting polymers e.g.
Bakelite do not melt when heated. On stronger heating the decompose
 Some polymers can produce toxic gases: hydrogen cyanide from polymers containing nitrogen,
hydrogen chloride from polyemers containing chlorine
Making addition polymers
Higher boiling point fractions are more difficult to sell as there is less demand for them. The petrochemical
industry breaks up these long chains to produce short molecules. This decomposition reaction is called
cracking. Compounds such as ethene are produced.
Ethene




Belongs to a family of alkenes
Is an unsaturated hydrocarbon with a formula C2H4
Is a gas at room temperature
Molecules contain a carbon-carbon double bond
There is a simple test to distinguish between ethane and ethene. If ethene gas is bubbled through a
solution of bromine, the solution changes from red-brown to colourless.
Ethene + Br2 (aq)  1,2-dibromoaethane
This is an example of addition reaction. Two reactants react to form a single product and the double bond
in ethene becomes a single bond. There is no colour change when ethane is added to a solution of bromine.
Ethene is called the monomer and poly(ethene) is called the addition polymer. In order to produce this
polymer, the ethene vapour is passed over a heated catalyst. A series of addition reactions occur.
Uses of addition polymers
Addition polymers such as poly(ethene) and poly(vinyl chloride) have many uses. They have replaced
traditional materials such as metals, paper, cardboard and rubber.
Common uses include:
Poly(ethene) – wrappings for food, storage containers, milk crates
Poly(vinyl chloride) – wellington boots, insulation for electrical wiring
Advantages
Do not absorb water
Can be moulded into shape
Can be coloured
Low density
Strong
Disadvantages
Do not rot away and can cause litter problems
Not easy to recycle as there are many types
Burn to form poisonous fumes
Hardening natural oils
Alkenes undergo addition reactions.
If a mixture of ethene and hydrogen is passed over a heated nickel catalyst, an addition reaction can take
place, the product of which is ethane.
Natural oils such as sunflower oil are liquid and unsaturated i.e. they contain carbon-carbon double bonds.
These oils can be hardened by addition reactions with hydrogen. The oil and hydrogen are passed over a
nickel catalyst at 170°C. The resulting fat is used as margarine.
Further polymers
Condensation polymers
Some general characteristics of this class of compounds:
 Are more complex than addition polymers
 Are formed by reaction of two different chemicals
 A water molecule is last during each reaction
o Hence, the process the is called condensation polymerisation
 Some common types are Nylon, Terylene (polyester) etc.
Nylon
Nylon is used to produce yarns which are used in the manufacture of cloths, ropes, racket strings. It is
formed by the following reaction:
(1,6-Diaminohexane) + (Hexanedioc acid)  Nylon + Water
H2N(CH2)6NH2 + HOOC(CH2)4COOH  H2N(CH2)6NHOC(CH2)4COOH + H2O
The structure consists of a series of two different molecules joined together by an amide link, found in
proteins.
What happens
Double bonds break, which allows monomers molecules to ultimately join together.
However, in condensation polymer, no double bonds break. Alternatively:
 Two different monomers join.
 The monomers join at their function groups by eliminating small molecules.
Structure of polyamides
Terylene
Terylene, like nylon, is used to produce yarns which are also used in the manufacture of various fabrics, etc.
It is formed by the following reaction:
(Ethane-1,2 diol) + (benezene-1,4-dicarboxylic acid)  Terylene + Water
HO(CH2)2OH + HOOC(C6H4)COOH  HON(CH2)2OCO(C6H4)4COOH + H2O
The structure consists of a series of two different molecules joined together by an ester link, found in fats.
Thermoplastic and thermosetting polymers
Polymers can be classified as thermoplastic or thermosetting. There is a difference in structure between
thermoplastic and thermosetting polymers. A simple representation of the two types of polymer of the two
types of polymer is shown below. In a thermoplastic polymer the chains are not linked. On melting, the
chains are able to move freely over each other. In a thermosetting polymer, there are strong links between
the polymer chains. The rigid structure is not easily broken down.
Natural rubber consists of chains of polymer molecules. It is naturally soft and sticky. It can be hardened by
a process of vulcanisation where sulphur atoms link the chains by cross-linking.
Disposal of polymers
Unlike materials such as paper, cardboard and wood, polymers do not rot away when tipped in landfill sites.
They are said to be non-biodegradable.
Recycling polymers is not economic as the costs of collection and sorting are greater than the costs of
making new polymer.
Polymers can be incinerated. Carbon dioxide and water vapour are produced.
Monomer Polymer
Ethene  Poly(ethene)
Fuels
The fossil fuels are:
 Crude oil
o Petrol
o Diesel
o Kerosene
o Butane
o Propane
 Natural gas
o Methane
o Ethane
 Coal
The reason these are called fossils are because they are the remains of plants and animals that lived a few
million years ago.
 Crude oil: These are formed from the remains of dead organisms that fell to the ocean floor and
were then buried by the thick sediment. The high pressure in which the dead organisms are buried
eventually converts the dead organisms into crude oil, but this is a process that takes millions of
years.
 Natural Gas: This is composed mainly of methane and is often found with petroleum. High
temperatures and pressure causes the compounds to break down into gas.
 Coal: This is the remains of lush vegetation that grew in ancient swamps. Over the millions of years,
high pressure and heat eventually converted the vegetation into coal.
When they are burnt with oxygen, they form carbon dioxide and water.
Combustion reactions
Definition
Complete combustion reaction: When the (usually) hydrocarbon burns/reacts
completely with oxygen gas (O2) it forms form only CO2
(carbon dioxide) and water (H2O). There are no
byproducts to a complete combustion reaction
Incomplete combustion reaction: A reaction or process which entails only partial
burning of a fuel. This may be due to a lack of oxygen
or low temperature, preventing the complete
chemical reaction. If, hydrocarbons burn with
imperfect efficiency, it will also produce carbon
monoxide and things like nitrogen oxides.
Hydrocarbons: An organic compound consisting entirely of hydrogen and carbon.
Complete combustion reactions (some examples)
When alkanes are burned in excess air or oxygen, carbon dioxide and water are produced.
The formula
CnH2n+2 + (1.5n+0.5)O2 → (n+1)H2O + nCO2
Methane
CH4 + 2O2  CO2 + 2H2O
Octane (petrol)
2C8H18 + 25O2  16CO2 + 18H2O
Incomplete combustion
Alkanes burn in a limited of a supply of air to produce carbon monoxide and water vapour. Carbon
monoxide is very poisonous.
C8H18 + O2  CO2 + CO + C + H2O
The formula
CnH(2n+2) + (n+0.5)O2 → (n+1)H2O + nCO
CnH(2n+2) + (0.5n+0.5)O2 → (n+1)H2O + nC
Sulphur
Key ideas
 Sulphur is a non-metallic, yellow coloured solid.
 It is very useful element.
 It is found in crude oil and in metal ores, e.g. copper pyrite, zinc blende, lead sulphide, etc.
 It is also found in elemental state in some places.
 It is extracted mostly by contact process.
Uses of sulphur
 Sulphur is used to produce sulphuric acid, a very important chemical.
 It is used to produce sulphur dioxide, which is used as a bleach in the manufacture of wood pulp for
paper and also as a preservative because of its antibacterial properties.
 Sulphur is used to manufacture several other compounds, such as sulphur dioxide, sulphites and
sulphates.
 It is used in the manufacture of matches, fireworks, vulcanized rubber, dyes etc.
 Being a fungicide and a sterilizing agent, it is used as a food preservative and in medicines, e.g.
sulpha drugs and skin ointments
 Its salt, sodium thiosulphate (commonly called hypo), is used in photography for fixing negatives
and prints.
Contact process
The contact process is a method used to produce Sulphuric Acid. Sulphuric acid is made in industry in a
three-stage process.
[In practice, this is done in two stages – the sulphur trioxide is dissolved in concentrated sulphuric acid and
then diluted with the required amount of water to make concentrated suphuric acid.]
Only stage 2-3 is reversible. Controlling this reaction is the secret for getting the maximum yield in the
whole process.
To get the best yield, low temperatures are desirable. However, low temperatures will slow down the
reactions. There has to be a compromise between getting a good yield and getting the yield quickly.
Vanadium (V) oxide can be used as a catalyst to speed up the reaction.
Conditions for Contact Process to occur
 450°C
 2-9 atm (pressure)

Vanadium Pent-oxide (V2O5) is used as a catalyst
Properties of dilute sulphuric acid
 Colored
 Corrosive liquid
 Strong oxidizing agent
 Reacts violently with bases.
 Doesn’t conduct electricity
 Brittle
 Insoluble in Water
Acids, Bases and Salts
Definition
Acid: A chemical species that donates protons or hydrogen ions and/or accepts
electrons.
Base: A chemical species that donates electrons or hydroxide ions or that accepts
protons.
pH: A measure of hydrogen ion concentration; a measure of the acidity or alkalinity of a
solution.
Properties
Key points
 Acids:
o Are compounds containing hydrogen that can be replaced by a metal
o Dissolve in water to form hydrogen, H+ ions
o Are proton donors
 Bases
o Are chemical opposites of acids
o
o
Dissolve in water to form hydroxide, OH- ions
Are proton acceptors
Acids
Acid
Hydrochloric acid
Sulphuric acid
Nitric acid
Ethanoic acid
Ethanedioic acid
Citric acid
Formula
HCl
H2SO4
HNO3
CH3COOH
C2O4H2
C3O8H7
Mineral salts
Contained in vinegar
Contained in rhubarb leaves
Contained in lemon juice
Only one hydrogen atom in ethanoic acid can be replaced, forming sodium ethanoate CH3COONa.
Properties of acids
Although there are a large number of different acids, there are a number of general chemical reactions
common to all acids.
Indicators
Acids turn indicators to their characteristic colours, e.g. litmus turns red.
Universal Indicator
This is a substance that changes color when added to another substance depending on its pH. The indicator
and the substance should be in aqueous form.
Litmus Paper or Solution
This indicator is present in two colors: red and blue. We use blue litmus if we want test a substance for
acidity. We use red litmus if we want to test a substance for alkalinity. Its results are:
 Acids: Turns blue litmus paper/ solution red,
 Bases: Turns red litmus paper/ solution blue,
 Neutral: if it is used as paper the color doesn’t change. If it is used as solution it turns purple.
Note: use damp litmus paper if testing gases.
Phenolphthalein
This is an indicator that is used to test for alkalinity because it is colorless if used with an acidic or neutral
substance and it is pink if it is used with a basic substance.
Methyl Orange
This indicator gives fire colors: Red with acids, yellow with neutrals and orange with bases.
Fairly reactive metals
Acids react with fairly reactive metals (e.g. magnesium and zinc) to form a salt and hydrogen gas.
Metal + Acid  Salt + Hydrogen
Characteristics of the reaction
 Bubbles are given out
 Temperature rises (the reaction is exothermic, heat is released)
 Metal disappears
Metal oxides/hydroxides
3Acids react with metal oxides or hydroxides to form a salt and water only.
Acid + Metal (Hydr)Oxide  Salt + Water
Characteristics of reaction
 Amount of metal (hydr)oxide decreases
 Temperature increases (exothermic reaction)
 Solution changes colour.
Metal carbonates
Acids react with carbonates (or hydrogencarbonates) to form carbon dioxide, a salt and water.
Acid + Metal Carbonate  Salt + Water + Carbon Dioxide
Characteristics of reaction
 Metal carbonate starts to disappear
 Temperature rises (exothermic reaction)
 Colour Change
Strong and weak acids
Some acids completely ionise when they dissolve in water. These are called strong acids. A solution of a
strong acid will have a high concentration of hydrogen ions, H+.
e.g. sulphuric acid
H2SO4  2H+ + SO42Other acids do not ionise completely on dissolving in water. Some of the molecules remain un-ionised in
the solution. These are called weak acids.
CH3COOH  CH3COO- + H+
-3
In solution of ethanoic acid (1 moldm ), there are about four, in every thousand molecules, which get
ionised
Bases
Base
Sodium Hydroxide
Potassium Hydroxide
Ammonia
Calcium Carbonate
Formula
NaOH
KOH
NH3
CaCO3
Properties of bases
There are a number of general chemical reactions common to all bases.
Indicators
Bases turn indicators to their characteristic colours e.g. red litmus turns blue.
Acid
Bases react with acids to form salt and water.
KOH + HCl  KCl + H2O
In gaseous form, the reaction is different.
NH3(g) + HCl(g)  NH4Cl (s)
Strong and weak bases
Bases which ionize completely, when dissolved in water, are called strong bases.
A solution of a strong base will have a high concentration of hydroxide ions, OH-.
e.g. sodium hydroxide
NaOH(s) + H2O(l)  KCl (s) + H2O(l)
Bases which ionize only partially when dissolved in water are called weak bases.
e.g. ammonia
NH3(g) + H2O(l)  NH4+ (aq) + OH- (l)
Titration
Steps to figuring out titration questions
1. Work out moles
2. Use ratio from equation
3. Work out concentration
In application
In a titration experiment, 23.1 cm3 of H2SO4 with a concentration of 1 moldm-3 neutralizes 25 cm3 KOH.
What is the concentration of KOH?
1. Work out the moles
1 x 0.0231 moles
2. Use ratio from equation
2KOH + H2SO4  K2SO4 + 2H2O
2
: 1
0.0462
: 0.0231
3. Work out concentration
0.0462/0.025 = 1.8 moldm-3
Neutralisation
Neutralisation is the reaction of acid and alkali, in the correct proportions, to produce a neutral substance.
Examples of neutralisation
1. Soil that is too acidic is not as fit to grow crops as neutral soil.
Slaked lime (calcium hydroxide) or limestone (calcium carbonate) can be added to the soil to neutralize
it.
2. Hydrochloric acid in the stomach helps in the digestion of food.
Indigestion is caused by excess acid. The pain can be relieved by taking a weak alkali such as sodium
hydrogencarbonate or magnesium hydroxide.
3. Coal-fired power stations produce sulphur dioxide, which can produce acid rain. The sulphur dioxide
can be removed from the waste gases before they escape into the atmosphere. Limestone removes
sulphure dioxide from the waste gases.
4. Insect bites and stings involve an injection of a small amount of chemical into the skin. These chemicals
cause irritation. Nettle stings, bee stings and ant bites involve methanoic acid being injected into the
skin. Wasp stings involve the injection of an alkali into the skin. The irritation can be removed by
neutralisation of the acid or alkali
5. Acid rain, caused by sulphur dioxide escaping into the atmosphere, can make lakes and rivers acidic.
This can affect organisms in the water such as killing fish. Blocks of limestone, put into the water, can
reduce the acidity.
Relative acidity and alkalinity
The degree to which acids and bases ionise in solution determines the relative strength of acids and bases.
To measure the acidity/alkalinity of a substance, a scale called pH scale has been developed. The pH scale
runs from 0 to 14.
If the value of pH is less than 7, the solution is acidic. Lesser the pH value, the more acidic is the solution i.e.
more is its relative acidic strength.
If the value of pH is more than 7, the solution is basic. More the pH value, the more basic is the solution i.e.
more is its relative basic strength.
To check whether a substance is acidic or alkaline, substances called indicators are used. Indicators change
colour on coming in contact with acids or alkalis.
To obtain an idea of how acidic or alkaline a substance is, an indicator called universal indicator is used.
Universal indicator is a mixture of several indicators and its colour changes with change in acidity or
alkalinity of the solution, in which it is kept. The colour shown by this indicator can be matched against the
pH scale to determine the relative degree of acidity/alkalinity of a substance.
Carbonates
Key points
 The rocks of the Earth are the source of a wide range of materials.
 Limestone is the mineral extracted from the Earth in the largest amounts. Often, it is found in
beautiful areas and its mining can damage the environment
 Limestone is used to make quicklime (calcium oxide) and slaked lime (calcium hydroxide).
 An ore is a rock that contains enough of a metal compound for it to be worth extracting the metal.
Rocks as building materials
Rocks such as limestone, sandstone and slate are used as building materials. As transport costs are very
high, rocks are often used as building materials close to where they are dug out of the ground, known as
quarrying.
Building materials made from rocks
As natural rocks are expensive, new materials have been developed to replace them.
Material
Bricks
Mortar
How it is made
By baking clay to a high temperature
Mixture of calcium hydroxide, sand,
and water made into a thick paste
Cement
Heating limestone with clay
(containing aluminium silicates)
Concrete
Made by mixing cement with sand
and small stones
Glass
Mixing limestone, sand (silicon
dioxide) and sodium carbonate
together and melting the mixture
More information
Hard and brittle – regular shape
It sets by losing water and by absorbing
carbon dioxide from the air. Long crystals
of calcium carbonate are formed
It consists of a complex mixture of calcium
and aluminium silicates. When it is mixed
with water, a chemical reaction occurs
producing calcium hydroxide, and this sets
in a similar way to mortar
Used to make many objects such as drain
covers that were previously made from
cast iron. Concrete can be strengthened by
steel reinforcing rods.
The resulting mixture of calcium and
sodium silicates cools to produce glass.
Coloured glass is due to transition metal
oxides present in the mixture
Uses of limestone
Apart from its use in building construction, limestone has several other uses, some of which are listed
below.
 Neutralising acidic soil by mixing soil with powdered limestone.
 In the manufacture of sodium carbonate, the key ingredient for making soaps, detergents etc.
 Used in the manufacture of cement.
Quicklime and slaked lime have several uses, some of which are listed below.
 Neutralising acidic soil by mixing it with powdered lime.
 Improving drainage properties of clay soils – when mixed with powdered lime, their clay nature
decreases.
 Used as a water-softening agent and also to manufacture bleaching powder.
 Used for making mortar in building construction
 Used in the manufacture of paper and glass, in leather tanning and sugar refining.
 Limewater (solution of slaked lime in water) is used in medicine as an antacid as a neutralizer for
acid poisoning and treatment of burns.
 Used for neutralizing acidic industrial waste products
Metals from rocks
Most metals are found in the earth as deposits of ore.
Metal
Name of ore
Sodium
Rock salt (halite)
Magnesium
Magnesite
Compound of metal present
Sodium chloride
Magnesium chloride
Aluminium
Iron
Zinc
Mercury
Bauxite
Haematite
Zinc blende
Cinnabar
Aluminium oxide
Iron (III) oxide
Zinc sulphide
Mercury (II) sulphide
Some ores contain only small amounts of metal compounds. The metal compound in these ores may be
concentrated by froth flotation before the metal is extracted. The ore is added to a detergent bath and the
mixture agitated. By careful control of the conditions, it is possible to get the metal compound to float
while the impurities sink to the bottom.
Air and Water
Water
Key points
 Water is a very important as it is essential for so many fundamental processes.
 It is however, very difficult to get pure because it is very good at dissolving other substances.
 This ability to dissolve other things has important consequences.
 A clean water supply is essential for good health.
 Many diseases are a consequence of impure water supplies.
Chemical test for water
1. Anhydrous Cobalt Chloride = turns from blue to pink
CoCl2 + 6H2O → *Co(H2O)6]Cl2
2. Anhydrous Copper(II) Sulphate = turns from
CuSO4·5H2O(s) → CuSO4(s) + 5H2O(g)
to blue
Physical tests
1. Boil to 100°C
2. Freeze at 0°C
3. Measure its mass
a. 1g = 1 ml = 1cm3
b. I kg = 1 dm3
Treating Water
Water extracted from the earth is always infested with impurities. This water might be contaminated with
disease and bacteria. That is why it is crucial to “purify” the water before it is drunk. This is done by two
processes, Filtration and Chlorination.
Here is how it works:
 Water is extracted from reservoirs and sent to be “treated”
 The water is first passed through a filter to filter out large objects such as rocks or mud.
 Smaller particles in the water is removed by adding Aluminium Sulfate which causes the smaller
particles to stick together in large pieces and settle down the filter.
 Water is now passed through sand and gravel filters which continue to filter off the smaller
particles and kills bacteria.
 Now its time for chlorination
 Chlorine gas is first bubbled through the water to kill the bacteria that exists in the water.


Sodium Hydroxide may be added in the water to prevent the water from being acidic from the
chlorine.
Water is delivered to the people that need them.
Producing tap water
There are various essential steps in producing tap water:
 Water is taken from a clean river or reservoir
 The water is passed through a screen (sieve) to remove solid objects
 It is left to stand to allow solid material to settle out
 It is then filtered through a gravel bed to remove impurities in suspension
 It is then chlorinated. Chlorine is added in small amounts to kill bacteria and other harmful microorganism
Other forms of treatment may be used in special cases.
These include:
 Adding aluminium sulphate to coagulate colloidal clay in water
 Using a mixture of carbon and water to remove tastes and odours from the water
 Using lime to correct acidity of the water
 Adding sulphur dioxide to remove excess chlorine
The water produced at the end of all these processes is not pure water. It is water that is safe to drink.
Recycling waste water
After water has been used by a home or a factory, it must be cleaned up before being returned to rivers
and possibly re-used for water supply.
This is done in a sewage treatment plant. The treatment of waste water involves:
 Filtering it to remove solid materials
 Using bacteria to break down the waste.
Nitrogen
Key points
 Nitrogen in the air is very useful
 It dilutes the air and is inert
Uses




Making ammonia
Combining with hydrogen to make fertilizers and explosives
Crisp packets
Liquid nitrogen is used for freezing = wart removing
Ammonia
Key points
 Ammonia, NH2, is a compound of nitrogen and hydrogen.
 It is produced in large quantities by the Haber process using nitrogen from the air as a raw material
Uses of Ammonia
 Explosives
 Fertilisers
Haber Process
 Nitrogen is gotten from the air by fractional distillation
 Hydrogen is gotten from water (H2O) or natural gas (CH4)
 The Haber Process is a reversible reaction
By choosing the best conditions, chemists attempt to produce the highest yield of ammonia economically.
The best conditions are:
1. One part of nitrogen to 3 parts of hydrogen by volume
a. Increase exothermic reaction and shift equilibrium to the right increases yield
2. A high pressure
a. 200 atm
b. Higher pressures favour the side with less moles of gas
c. It is dangerous and expensive
3. A low temperature.
a. There is more need of “forwards reaction” or exothermic reaction
i. Exothermic reactions are favored by lowered temperatures
b. There is less need of “backwards reaction” or endothermic reaction
i. Endothermic reactions are favored by higher temperatures
c. So lower temperatures produce higher yields of ammonia
d. However, using a low temperature reduces the rate of reaction.
e. Using an iron catalyst speeds up the reaction.
i. Has no effect on yield.
4. Keep adding reactants
5. Keep removing products and reusing them
Composition of air
Common air pollutants
Carbon Monoxide
 Poisonous pollutant of air
 Main source is factories that burning Carbon-containing fossil fuels as carbon is one of the products
of incomplete combustion of fossil fuels.
Sulphur Dioxide
 Contributes to acidic rain
 Main Source comes from two products:

1. Combustion of sulphur
 2. Extraction of metals from their sulfide ores
 Mixes with water vapour of cloud and air.
 This forms Sulphuric Acid (H2SO4)
 When it rains, the rain water becomes acidic .
 Acidic water is dangerous because it causes the death of sea creatures, acidifies soil which can
cause death to plants and cause deforestation.
 May also cause lung cancer
Oxides of Nitrogen
 Formed in high temperatures when nitrogen and oxygen react.
 In Cars, the engine operates at a high temperature, giving the nitrogen and the oxygen in the air
and engine a chance to react, hence forming nitrogen monoxide. Nitrogen monoxide further reacts
with the oxygen in the air to form Nitrogen Oxide.
 Nitrogen oxide is dangerous in that it also rises in the air and mixes with rain water to form nitric
acid. This can also cause acid rain
 Additionally, Nitrogen oxygen can cause certain respiratory problems.
Car exhausts and catalytic converter
As we mentioned earlier, oxides of nitrogen are present in car exhausts, and these can cause problems both
to the environment and us humans. Therefore, scientists need to find a way to remove the “oxides of
nitrogen” in cars.
This is done through a catalytic converter


The catalytic converter is fitted at the end of the car exhaust.
The purpose of the catalytic converter which catalyzes the reaction between the Nitrogen Oxide
and Carbon Monoxide, which in turn produces two harmless separate gases, nitrogen and carbon
dioxide. The carbon dioxide comes from the fact that carbon is already present in the car’s engine.
Equation of the Reactions
1. 2NO + 2CO —> 2CO2 + N2
Nitrogen Oxide + Carbon Monoxide —-> Carbon Dioxide + Nitrogen
2. 2NO2 + 4CO —> 4CO2 + N2
Nitrogen Dioxide + Carbon Monoxide —> Carbon Dioxide + Nitrogen
Greenhouse effect and global warming
 The Sun sends energy to the earth in two discrete forms, heat and light.
 Some of the heat is reflected back to the sun/space, but some is trapped in the earth.
 This is caused by the existence of some gases and we call this the greenhouse effect.
 The primary greenhouse gases are carbon dioxide and methane.
The greenhouse effect is a serious threat to our world. The reason for this can be described by the
proliferation of greenhouse gases which causes the greenhouse effect. Increased combustion of carbon in
industries which mass produce Carbon Dioxide as a side product and the cutting down of trees which
release CO2 via respiration are two major reasons why the greenhouse effect is becoming more serious.
The increase of heat trapped in the earth causes an average rise in sea level and global average
temperatures, and we call this effect global warming.
Formation of carbon dioxide
 Formed in power stations by the complete combustion of Carbon containing fuels.
 Formed as a product as respiration.
 When an acid reacts with a carbonate, Carbon Dioxide is usually formed.
Rust prevention
Painting
 Iron or steel object is painted
 The paint creates a barrier which prevents the air or water from coming in contact with the
iron/steel object. This is commonly done in car bodies and bridges.
 Disadvantages
o Oil can rub off, the paint can flake off, the plastic can chip off.
o Oiling and painting has to keep doing it all the time.
Electroplating
 Electroplated with another metal that doesn’t corrode.
 Metals commonly used are chromium and tin since they are very unreactive.
 Used in food cans.
Galvanization
 Covering the whole object by a layer of zinc
 Done either dipping the object into some molten zinc or by electroplating the object with zinc.
 The zinc provides a barrier to prevent the air and water from coming in contact with the iron/steel,
as the zinc is corroded instead of the iron
 Advantage
 Disadvantages:
o Expensive
o Not very energy-efficient
Sacrificial Protection
 This method revolves on the idea that metals higher up in the reactivity series will react in
preference, so a metal in the higher in the reactivity series (e.g. zinc and magnesium) is used as a
“protection”, and is corroded instead of the iron
 Used often in ships, or bridge columns
 Disadvantages
o Must be replaced from time to time as when the metal finishes corroding, the iron/steel
starts to rust again.
o
Cathodic protection
 A technique used to inhibit corrosion on buried or immersed structures by supplying an electrical
charge that suppresses the electro-chemical reaction.
 Connect the metal to be protected with a piece of another more easily corroded "sacrificial metal"
to act as the anode of the electrochemical cell.
 The sacrificial metal then corrodes instead of the protected metal.
 Used to protect a wide range of metallic structures in various environments
 Common applications are:
o steel water or fuel pipelines
o storage tanks such as home water heaters
o steel pier piles; ship and boat hulls
o offshore oil platforms and onshore oil well casings
o metal reinforcement bars in concrete buildings and structures
 Disadvantages
o A side effect of improperly applied cathodic protection is the production of hydrogen ions,
leading to its absorption in the protected metal and subsequent hydrogen embrittlement of
welds and materials with high hardness.
o A process of disbondment of protective coatings from the protected structure (cathode)
due to the formation of hydrogen ions over the surface of the protected material (cathode).
Miscellaneous
Oxidation Numbers
Oxidation numbers are used to determine whether, in a chemical reaction, which s oxidised or reduced.
The rules are:
1) Any element has an oxidisation no. of 0
a. E.g. H2 = 0 , O2 = 0, Na=0, Fe=0
2) Any ion has the oxidisation no. of its charge
a. E.g. Fe3+ = 3+, Cl- = 13)
4)
Hydrogen is +1
Oxygen is -2
Example
The oxidation number of an element indicates the number of electrons lost, gained, or shared as a result of
chemical bonding. The change in the oxidation state of a species lets you know if it has undergone oxidation
or reduction.
Oxidation can be defined as "an increase in oxidation number".
In other words, if a species starts out at one oxidation state and ends up at a higher oxidation state it has
undergone oxidation.
Conversely,
Reduction can be defined as "a decrease in oxidation number".
Any species whose oxidation number is lowered during the course of a reaction has undergone reduction.
Na + Cl2 -----> 2NaCl
The Na starts out with an oxidation number of zero (0) and ends up having an oxidation number of 1+. It
has been oxidized from a sodium atom to a positive sodium ion.
The Cl2 also starts out with an oxidation number of zero (0), but it ends up with an oxidation number of 1-.
It, therefore, has been reduced from chlorine atoms to negative chloride ions.
The substance bringing about the oxidation of the sodium atoms is the chlorine, thus the chlorine is called
an oxidizing agent. In other words, the oxidizing agent is being reduced (undergoing reduction).
The substance bringing about the reduction of the chlorine is the sodium, thus the sodium is called a
reducing agent. Or in other words, the reducing agent is being oxidized (undergoing oxidation).
Oxidation is always accompanied by reduction. Reactions in which oxidation and reduction are occurring
are usually called redox reactions.
Organic equations
1. alkane + oxygen
heat
methane + oxygen
Balancing equations
6HNO3 + Al2(CO3)3  2Al(NO3 )3 + 3CO2 + H2O
Start here
Note:
However many carbons there are in the reactants, use that number on CO2 and H2O
1. H2SO42+ + Na2CO3  Na2SO4 + CO2 + H2O
There is only one carbon in the reactants, so the number of CO2 and H2O is 1
2. H3PO43+ + 3KHCO3 K3PO4 + 3CO2 + 3H2O
There are three carbons in the reactants, so the number of CO2 and H2O is 3
3. 6HNO3 + Al2(CO3)3 2Al(NO3)3 + 3CO2 + 3H2O
There are three carbons in the reactants, so the number of CO2 and H2O is 3
Results of some electrolysis experiments
Solution
Electrodes Ion discharged
at positive
electrode
(Anode)
Dilute sulphuric
Carbon
OH- (aq)
acid
Dilute sodium
Carbon
OH- (aq)
hydroxide
Copper(II) sulphate
Carbon
OH- (aq)
Copper(II) sulphate
Copper
None
Copper(II) chloride
Carbon
Cl- (aq)
Concentrated
Carbon
Cl- (aq)
sodium chloride
Molten sodium
Mercury
Cl- (aq)
chloride
cathode
Molten Al2O3
O2
Aqueous HCl
Cl2
Molten PbBr2
Br2
Formulae of acids and their salts
Acid
Formula
Hydrochloric acid
HCl
Nitric acid
HNO3
Sulphuric acid
H2SO4
Carbonic acid
H2CO3
Phosphoric acid
H3PO4
Product at
positive
electrode
Oxygen
Ion discharged
at negative
electrode
(Cathode)
H+ (aq)
Oxygen
H+ (aq)
Hydrogen
Oxygen
None
Chlorine
Chlorine
Cu2+ (aq)
Cu2+ (aq)
Cu2+ (aq)
H+ (aq)
Copper
Copper
Copper
Hydrogen
Chlorine
Na+
Sodium
amalgam
Oxygen
Al
H2
Pb
Salts
Chloride
Nitrate
Sulphate
Carbonate
Phosphate
Product at
negative
electrode
Hydrogen
Formula
ClNO3SO42CO32PO43-
Acids and salts
Some common alkalis and bases
Type
Name
Formula
Alkalis
Sodium hydroxide
NaOH
Bases
Strong or weak?
Strong
Where found or used
In oven cleaners; in making soap and
paper; other industrial uses
Salts in common use
Salt
Parent acid
Ammonium chloride
Hydrochloric acid
Ammonium nitrate
Ammonium sulphate
Calcium carbonate
Nitric acid
Sulphuric acid
Carbonic acid
Calcium sulphate
Sodium carbonate
Sulphuric acid
Carbonic acid
Magnesium sulphate
Copper(II) sulphate
Calcium phosphate
Sulphuric acid
Sulphuric acid
Phosphoric acid
Colour and other Uses
characteristics
White crystals
Fertilisers;
dry
cells
(batteries)
White crystals
Fertilisers; explosives
White crystals
Fertilisers
White
Decorative stones; making
lime and cement and
extracting iron
White crystals
Wall plaster; plaster casts
White crystals or In cleaning; water softening;
powder
making glass
White crystals
Health salts (laxative)
Blue crystals
Fungicides
White
Making fertilisers
Pattern of solubility for various types of salts
Salts
Soluble
Sodium salts
All are soluble
Potassium salts
All are soluble
Ammonium salts
All are soluble
Nitrates
All are soluble
Ethanoates
All are soluble
Chlorides
Most are soluble
Sulfates
Most are soluble
Carbonates
Sodium
potassium
ammonium carbonates
Hydrated salts
Hydrated salts
Copper(II) sulphate
Cobalt(II) chloride
Iron(II) sulphate
Magnesium sulphate
Sodium carbonate
Calcium sulphate
Formula
CuSO4 ∙ 5H2O
CoCl2 ∙ 6H2O
FeSO4∙ 6H2O
MgSO4∙ 7H2O
Na2CO3∙ 10H2O
CaSO4∙ 2H2O
Insoluble
None
None
None
None
None
Silver chloride
Lead(II) chloride
Barium sulphate
Lead(II) sulphate
Calcium sulphate
and Most are insoluble
Colour
Blue
Pink
Green
White
White
White
Chemical tests and results
Cations
Cation
Ammonium
Effect of adding sodium hydroxide
Ammonia produced on warming
Effect of adding ammonia solution
test with damp red litmus)
NH4+
Copper(II)
Cu2+
Iron(II)
Fe2+
Iron(III)
Fe3+
Calcium
Ca2+
Magnesium
Mg2+
Zinc
Zn2+
Aluminium
Al3+
Anions
Anion
carbonate
(CO32-)
chloride (Cl-)
(in solution)
bromide (Br-)
(in solution)
iodide (I-)
(in solution)
sulphate (SO42-)
(in solution)
nitrate (NO32-)
(in solution)
Test for gases
Gas
Ammonia
NH3
Carbon dioxide
CO2
Light blue gelatinous ppt. of copper
hydroxide, insoluble in excess
sodium hydroxide
Green gelatinous ppt. of iron(II)
hydroxide; insoluble in excess
Rust-brown
gelatinous
ppt.;
insoluble in excess
White ppt. of calcium hydroxide;
insoluble in excess
White ppt. of magnesium hydroxide;
insoluble in excess
White ppt. of zinc hydroxide; soluble
in excess, giving a colourless solution
White ppt. of aluminium hydroxide;
soluble in excess, giving a colourless
solution
Light blue gelatinous ppt.; dissolve in
excess ammonia, giving deep blue
solution
Green gelatinous ppt.; insoluble in excess
Test
add dilute hydrochloric acid to
solution
acidify solution with dilute nitric
acid, then add aqueous silver nitrate
acidify solution with dilute nitric
acid, then add aqueous
acidify solution with dilute nitric
acid, then add aqueous silver nitrate
acidify
solution
with
dilute
hydrochloric acid, then add barium
chloride solution
OR
acidify solution with dilute nitric
acid, then add barium nitrate
solution
make solution alkaline with sodium
hydroxide solution, then add
aluminium foil (or Devarda’s alloy)
and warm carefully
Result
effervescence, carbon dioxide produced
(test with limewater)
white ppt. of silver chloride formed: ppt.
soluble in ammonia solution
cream ppt. of silver bromide formed:
only slightly soluble in ammonia solution
yellow ppt. of silver iodide formed: only
slightly soluble in ammonia solution
white ppt. of barium sulphate formed
Colour and smell
Colourless, pungent
smell
Colourless, odourless
Rust-brown gelatinous ppt.; insoluble in
excess
No ppt. (or only a very slight ppt.)
White ppt.; soluble in excess
White ppt.; soluble in excess
White ppt.; insoluble in excess
white ppt. of barium sulphate formed
ammonia gas given off (test with moist
red litmus.
Test
Hold damp red litmus paper
( or U.I. paper) in gas
Bubble
gas
through
limewater
(calcium
hydroxide solution
Test result
Indicator paper turns blue
White
ppt.
of
carbonate formed
turns milky)
calcium
(solution
Chlorine
ClHydrogen
H2
Oxygen
O2
Pale green, choking Hold damp litmus paper (or
smell
Universal Indicator paper) in
gas
Colourless, odourless Hold a lighted splint in gas
Colourless, odourless
Indicator paper is bleached
white (blue litmus will turn red
first)
Hydrogen burns with a squeaky
‘pop’
Hold a ‘glowing’ wooden The splint re-lights
splint in gas
Flame test colours
Metal ion
Sodium
Formula
Na+
Colour of flame
Yellow
Potassium
Calcium
K+
Ca2+
Lithium
Copper
Barium
Li+
Cu2+
Ba2+
Lilac
Brick red
(orange-red)
Crimson
Blue-green
Apple green
How to write the evaluation for coursework
Evaluation
1. What did the results show?
2. Did you expect this/ is it consistent with theory?
a. What is the theory?
b. Why?
3. How reliable are the results?
a. Why?
4. How valid are the results?
a. How well did you control the variables?
b. How well did the result reflect what you want to test?
5. Explain the current theory for predicted results and compare yours
a. Propose sources of error/variation
b. Propose improvements
Some GCSE Questions
1. List 5 metals and 5 properties they have in common.
Iron, zinc, copper silver, gold; hard, shiny, good conductors of heat and electricity.
2. Write a symbol equation of the reaction of carbon with iron oxide
2Fe2O3 + 3C  4Fe + 3CO2
a. What is oxidized (reducing agent)?
Carbon
b. What is reduced?
Iron
3. Write a symbol equation for the reaction of magnesium with oxygen.
2Mg + O2  2MgO
a. Is the product acidic or alkaline? Explain.
Alkaline because it is a metal oxide.
4. What is an alloy?
An alloy is a mixture of metals.
a. Name 3 alloys and a use for each.
Brass is used for instruments
Steel is used for buildings
Bronze is used for statues
b. Why are they more useful than pure metals?
They are stronger and less reactive.
5. Work out these formulae.
a. Sodium chloride
NaCl
b. Magnesium chloride
MgCl2
c. Zinc(II) sulfate
ZnS
d. Lithium oxide
Li2O
6. Draw a labeled diagram of the electrolysis of molten sodium chloride
Na+ + e-  Na (Reduction)
2Cl  Cl2 + 2e- (Oxidation)
7. What are the products at the anode and cathode for aqueous sodium chloride? Write half equations.
Cathode: Hydrogen Anode: Chlorine
Na+ + e-  Na (Reduction)
2Cl  Cl2 + 2e- (Oxidation)
8. CaCO3 + 2HCL CaCl2 + H2O + CO2
a. 100g of calcium carbonate reacts. What mass of calcium chloride is produced?
1 CaCO3  1 CaCl2
100g
1 mole of CaCl2 (110)
40+12+(16x3)
1 x 110 = 110g
b. What volume of CO2 is produced?
CaCO3  CO2
100g  1 mole of CO2
100
Volume (dm3) = 1 x 24 = 24dm3
9. CuO + H2  Cu + H2O
a. What is the reducing agent?
Hydrogen
b. 55g of copper oxide reacts. What mass of copper is made?
55  0.6875 moles
80
Mass = 0.6875 x 64 = 44g
10. What do the terms endothermic and exothermic mean?
Exothermic refers to reactions or processes that release energy by either heat or light
Endo thermic refers to a process or reaction that absorbs energy from surrounding in the form of heat.
11. Draw apparatus that could be used to electroplate a spoon with silver.