• Mendeleev created the
periodic table in 1871.
• He arranged elements in
order of their
• -groups = down.
• -period= across
• Metals are on the left
hand side of the
He arranged them by the
• Atomic mass
• Melting Point
• Formula of the oxide
Patterns emerged down
the groups and across the
periods which confirmed
C2.2 Structure of the atom
• Atoms are made up of
PROTONS, NEUTRONS and
• They are found in the positions
shown on the diagram.
• Atom is the smallest of particles.
protons and neutrons
Proton 1 +1
Neutron 1 0
Electron 1/1840 -1
Electrons fill shells from the middle in the order of
2, 8, 8, 18 (how many elements are in each period)
The group number
corresponds to the
electrons in the
C2.3 The Modern Periodic Table• Top number – MASS NUMBER
The number of protons and neutrons
• Bottom number – ATOMIC NUMBER
The number of protons
(also the same as electrons)
Protons = 13
Electrons = 13
Neutrons = 27-13 = 14
Protons and neutrons are packed together
tightly in the nucleus (high density)
Electrons are spread out in shells (low
C2.4 Electron shells
• Electron located in the shell.
• Electronic configuration = how the electrons are
arranged in an atom.
C2.5 Ionic Bonding
• Ionic bonds form between METALS and NON-METALS.
• Ionic bonding involves the transfer of ELECTRONS.
• Metallic Ions are POSITIVELY charged (ANIONS).
• Non-metallic elements are NEGATIVELY charged
an atom must lose
charged and atom
must gain electrons.
C2.5 Ionic compounds- non metal and
• Conduct electricity when MOLTEN (melted) and in an AQUEOUS SOLUTION
(dissolved in water)
• DO NOT conduct electricity as a SOLID
• Have high MELTING and BOILING points
• Usually SOLID at ROOM TEMPERATURE
Ion = an atom with a positive or negative charge.
Cations = metal atoms lose electrons to form positively charged ions called cations.
Anions = Non-metal atoms gain electrons to form negatively ions called anions.
Transferring electrons = ionic bonds.
Ionic compounds have a lattice structure, with a regular arrangement of ions, held
together by electrostatic forces between oppositely charged ions.
C2.6 Ionic compounds
• Ionic compounds contain ions.
Below are some of the common
Compound ions = contain
more than one element.
Names of ionic compounds =
If the compound contains
oxygen it ends with an –
ate…..eg….. Iron sulfate = Fe
If no oxygen then they end in
an –ide…….Iron Sulfide….FeS
C2.7 Properties of ionic
• Solid = does not conduct electricity
• Liquid = conduct electricity…WHY….ions are free to
move around and they are charged.
• Melting points = when a substance changes from
solid to liquid.
• Boiling point = changes from a liquid to a gas.
• Ionic substances = high boiling points and high
• Ionic bonds = strong and lots of them to break so
needs lots of energy.
C2.8 SolubilitySoluble = substance dissolves.
Insoluble = does not dissolve
Salt = Acid + alkali
Table explains which salts are soluble and which are not.
silver nitrate (soluble) + sodium chloride (soluble) → silver chloride (insoluble) + sodium nitrate (soluble)
AgNO3 (aq) + NaCl (aq)→ AgCl (s) + NaNO3 (aq)
reaction - An
Barium sulfate is used as a ‘barium meal’ to x-ray
patients because it is
1. Opaque to x-rays
2. Safe to use
3. Insoluble so doesn’t enter the blood
• Precipitates are SALTS that are formed in chemical reactions that DO NOT
DISSOLVE in the solvent used in the reaction.
C2.11 Ion Tests
• Ion tests for metals
(ANIONS) are usually done
by FLAME TESTS.
• Each ION produces a certain
Testing for CATIONS is done through
1. Chloride – add nitric acid and silver
nitrate, if a white precipitate forms
then chloride ions are present.
2. Sulphate – add hydrochloric acid
and barium chloride, if a white
precipitate forms then sulphate ions
3. Carbonate – add an acid and if the
gas produced turns limewater MILKY
the carbonate ions were present.
(carbonate release carbon dioxide)
C2.12 Covalent Compounds
Learn these covalent molecules. Be able to
draw them using dot and cross diagrams.
Covalent bond is when a pair of electrons
are shared between two atoms forming a
Carbon Dioxide CO2
Hydrogen Chloride HCl
C2.13 Properties of Covalent
Simple Covalent Substances
• Low melting point
• Low boiling point
• Poor conductor of
• Because there are weak
Giant Covalent Substances
• High melting point
• High boiling point
• Because there are
strong forces between
2.13 Diamond vs. Graphite
• Used to make cutting tools.
• Very hard with strong bonds
• Does not conduct electricity
because there are no free
• Used to make electrodes and
• There are strong forces between
molecules within the layers – but
there are weak forces between
• There is one free electron for
each carbon atom meaning it can
C2.15 Miscible or Immiscible?
Miscible – Liquids that completely mix to form a solution.
Immiscible – Liquids that do not form a solution but separate into two layers e.g.
oil and water.
Separating two immiscible liquids
Use a separating funnel.
1. Open the tap and let the lower liquid flow out and collect in a beaker.
2. Close the tap before the upper liquid starts to run out.
3. Change the beaker and then let the other liquid flow out.
Separating two miscible liquids
Use fractional distillation to separate liquid air to form nitrogen and oxygen.
Exploits the fact they have different boiling points.
1. Cool air to -200oC to liquefy it.
2. Warm up to -185oC to evaporate the nitrogen and keep the oxygen liquid.
Used to separate mixtures of coloured compounds including inks, dyes and
colouring agents in food.
1. A spot of the mixture is placed near the bottom of a piece of chromatography
paper and the paper is then placed upright in a suitable solvent, e.g. water.
2. As the solvent soaks up the paper, it carries the mixtures with it.
3. Different components of the mixture will move at different rates.
4. This separates the mixture out.
Different chromatograms and the separated components of the
mixtures can be identified by calculating the Rf value using the
The Rf value of a particular compound is always the same - if the
chromatography has been carried out in the same way. This
allows industry to use chromatography to identify compounds in
Rf = distance moved by the compound ÷
distance moved by the solvent
C2.17 Chemical Classification
Melting Point Boiling Point Solubility in
Ionic High High Most Dissolve Conduct
or in solution,
not when solid
low low Some dissolve
High High Insoluble Do not
Metallic High except
C2.14 Classifying Substances
Substance Melting Point Boiling Point Electrical
High High Yes in solution
High High Yes in solution
Hexane Low Low No Insoluble SIMPLE
Paraffin Low Low No Insoluble SIMPLE
Silicon dioxide High High No Insoluble GIANT
High High Yes in solution
Sucrose High High Soluble GIANT
Could you classify these substances based on their properties? Know these examples!
C2.18 Metallic Bonding and
The structure of metals – a regular
arrangement of positive ions
surrounded by a sea of delocalised
Metals are malleable because ions can slide over each other if a large enough
force is applied. The electrons hold the ions together meaning the metal spreads
out rather than breaks.
Metals conduct electricity because the delocalised electrons are free to move. If a
voltage is applied, the electrons all move in the same direction and this is called a
Most metals are transition metals and they typically have high melting points and
produce coloured compounds.
C2.19 Alkali Metals
Alkali metals are soft with relatively low
Lithium, Sodium and Potassium react
with water to form hydroxides (which are
alkaline) and hydrogen gas.
Alkali metals get more reactive as you move down
the groups because the atoms lower down have more
electron shells. This means the outer electron is
further away from the nucleus, not held in place as
strongly, and easier to lose – making it more reactive.
Lithium + water Lithium Hydroxide + Hydrogen
Sodium + water Sodium Hydroxide + Hydrogen
Potassium + water Potassium Hydroxide + Hydrogen
Halogen Colour State at room
Fluorine Pale Yellow Gas
Chlorine Green/yellow Gas
Bromine Brown Liquid
Iodine Grey Solid
Metal + Halogen Metal Halide
Potassium + Bromine Potassium Bromide
2K(s) + Br2(l) 2KBr2(S)
Calcium + Chlorine Calcium Chloride
2Ca(s) + Cl2(l) 2KCl2(S)
They all follow
C2.21 Displacement of the
• Displacement reactions – a more reactive element
will displace (take the place of) a less reactive
Potassium Bromide + Chlorine Potassium Chloride
Potassium Chloride + Bromine No reaction
Chlorine is more
bromine so will
Sodium Iodide + Chlorine Potassium Chloride
Sodium Chloride + Iodine No reaction
Chlorine is more
reactive than iodine
so will displace the
C2.22 More halogen reactions
Halogens react with hydrogen to make hydrogen halides
which will dissolve in water to make acidic solutions.
Hydrogen + Fluorine Hydrogen Fluoride
H2(g) + F2(g) 2HF(g)
Hydrogen + Chlorine Hydrogen Chloride
H2(g) + Cl2(g) 2HCl(g)
Fluorine Reaction Reaction Reaction
be used to
work out the
C2.23 Noble gases
Noble gases are
they have a full outer
shell of electrons.
Discovery of Noble Gases
1. Chemists noticed the
density of nitrogen made
in a reaction was different
from that of nitrogen in
2. Chemists made a
hypothesis about the
composition of the air.
3. Chemists performed tests
to show the presence of
the Noble gases in the air.
Uses of Noble Gases Reason
Xenon and Argon inside
To stop the hot filament
reacting with oxygen and
Argon and Helium used in
To form a blanket over the
hot metal to stop it reacting
with oxygen in the air
Argon used in fire
It is non flammable and can
be used to fill a computer
server room (for example) if
a fire breaks out.
Helium used to fill balloons
and air ships
It has low density
Neon in fluorescent lamps Produces a red light when
electric current is passed
through it under low
C2.24 and C2.25 Temperature
• Endothermic – reaction that takes heat energy in, decreasing the temperature of the reaction mixture
and its surroundings
• Exothermic – reaction that releases heat energy, increasing the temperature of the reaction mixture
and its surroundings
• During a chemical reaction there is usually a
transfer of energy between the reactant and
• Not many reactions are endothermic
Sodium hydrogen carbonate + Hydrochloric acid Sodium Chloride + Water + Carbon dioxide
• Other endothermic reactions include:
– Dissolving ammonium nitrate
• Takes in heat energy
• Gives out heat
• Most reactions are exothermic
• Temperature increase
• All combustion reactions are exothermic
• E.g. Methane + Oxygen
• Explosions – release a lot of heat and gases very quickly
• Energy needed to break bond
• Energy released when bonds made
C2.26, 2.27 and 2.28 Rates of Reactions and Collision Theory
• Rate of reaction – The speed at which a reaction takes place
• Concentration – A measure of how much solute is dissolved in
a fixed volume of solvent.
• Surface area– The total area of all the surfaces of an object or
Factors affecting Rate
• More concentrated =
• More particles = more
collisions = faster reaction
• Higher temperature = faster reaction
e.g. And egg cooks faster in boiling water than warm water
• Particles have more energy = move faster
– More effective collisions (collide with more energy)
– Collide more frequently
3. Surface area (SA)
• Solid broken up into smaller pieces = larger SA
• Greater surface area = faster reaction
• More surface area = more particles on the surface
therefore more frequent collisions
• A = Smaller SA (block)
• B = Larger SA (powder)
• Fast reactions = Burning,
• Slow reaction = Rusting, apple
• Many chemical
processes use catalysts
to increase rate of
production of products
• Catalysts help to lower
the temperature and
pressure needed = less
energy needed = saves
• Catalyst – A substance that speeds up the rate of a reaction without being used
up in the reaction
• Catalytic converter – Device fitted to car exhausts with a thin layer of transition
metal catalyst on a honeycomb structure giving a large surface area.. The
catalyst speeds up the reaction to combine Carbon monoxide and unburned
petrol into carbon dioxide and water.
• Reduce pollutants in exhaust gases
• Combine Carbon monoxide (CO) with oxygen
• Carbon dioxide and water released instead
• Contain transition metals (platinum, rhodium or palladium)
• Expensive metals so small amounts used
• Catalyst spread over a honeycomb structure = large Surface Area
• Works faster with the hot gases from the engine heat it up
C2.30 and 2.31 Relative Mass
• Compares mass to that of carbon 12
• E.g. Hydrogen is 12X lighter that
• Sum of all masses in the formula
• Simplest whole number ratio of
• E.g. H2O there are double the number
of hydrogens to oxygen
• Relative Atomic mass (RAM) – The
mean mass of an atom relative to the
mass of an atom of carbon-12, which
has a mass of 12.
• Empirical formula (EF) – The simplest
whole number ration of atoms of each
element in a compound
• Molecular formula (MF) - The actual of
atoms of each element that combine to
make a molecule of a compound.
• Relative formula mass (RFM)– The sum
of the relative atomic masses of all the
atoms in a formula
• No atoms are lost or
made in a chemical
• They are rearranged into
• You can use relative mass
and balanced equations
to calculate the mass of a
reactant or product.
mass of an element
in a compound
= X X
Products need to be make as cheaply as possible. Chemists need to make sure the
reaction creates as much product as possible.
1.Maximum calculated amount of a product that could be formed from a given
amount of reactants.
2.Can be calculated from the balanced equation
3.Assumes all reactant are turned into products and all products harvested
1.The actual amount of product obtained from a chemical reaction.
1.The actual yield divided by the theoretical yield as a percentage
Yield usually less than expected – 3 reasons:
1. Reaction may be incomplete
2. Some product is lost
Yield – the amount of useful product obtained from a
C2.34 Waste and Profit
Cost Effective Process
• All factories must consider environmental impacts including disposal of waste.
• Companies want the most profit
• Reactions need to be
• High in percentage yield
• All products are useful so there are no waste products (or uses found for waste
products – econimically viable)
• Reactions are fast (quick to make lots of product)
Disposal of Waste Products
• Industry make many useful substances (cement, pesticides, plastics)
• Many of the reactions make waste products:
Must be disposed of following strict laws
Expensive to dispose
Can cause environmental problems
Can cause social problems (house prices drop or unpleasant smells)
• By-products – Any product formed in a reaction in
addition to the required product
• Waste products – By-products that have no uses