CfE Chemistry
Summary Notes
S2/3
Pupil Name: _______________________
You must keep this booklet safe as it will also be used in
National 4/5 Chemistry.
Unit 1 – Solutions, Compounds and Mixtures
Elements and the Periodic Table
•
•
•
Everything is made from elements in the Periodic Table.
The Periodic Table contains around 109 different natural
elements.
Man-made elements are shown with a * next to them in the
Periodic Table.
Each element has its own symbol.
E.g. Silver - Ag, Iron - Fe and Sodium - Na.
Each element has its own atomic number.
E.g. Aluminium is number 13.
Each element is made from only ONE type of atom.
E.g. Silver is only made from silver atoms.
Compounds
•
Compounds are made by chemically joining two or
more different elements together.
•
When the element sodium reacts together with
the element chlorine we get the compound sodium
chloride.
Naming Compounds
If a compound has the ending…
-IDE…
-ITE…
-ATE…
this means there are only 2 elements present.
this means the compound has 2 elements PLUS oxygen.
this means the compound also has 2 elements PLUS oxygen.
Name of Compound
Lithium bromide
Sodium nitrite
Calcium carbonate
Lithium
Sodium
Calcium
Elements Present
Bromine
Nitrogen
Carbon
Oxygen
Oxygen
2
Word Equations
Sodium metal can react with chlorine gas to produce the compound sodium
chloride. This can be written in a shorter format called a word equation:
The ‘
REACTANTS
PRODUCTS
Sodium + Chlorine
Sodium Chloride
‘ means produces, makes or gives.
Chemical Reactions
All chemical reactions produce
NEW substances.
Signs of a chemical reaction include:
1. A GAS forming.
This can be seen as fizzing or bubbling.
This is often known as effervescence.
2. A COLOUR CHANGE.
3. A PRECIPITATE formed.
This is when solutions react to form a solid.
4. During the reaction there is a change in ENERGY.
This can be exothermic – where heat is released – or
Endothermic – reaction gets cold.
Chemical reactions do NOT involve:
Melting
Freezing
boiling
condensing
evaporating
dissolving
The above are known as physical changes.
3
Elements, Compounds and Mixtures
Mixture
Compound
Chemically Joined
Easily separated
4
Solute, Solvent and Solutions
A solute is a substance which
can be dissolved in a liquid.
A solvent is a liquid in which a
solute can be dissolved.
A solution is what is produced
when a solute is dissolved in a
solvent.



A saturated solution is one in which no more solute can be dissolved.
A concentrate solution is one which contains a large amount of dissolved solute.
A dilute solution is one which contains a small amount of dissolved solute.
Solubility
If a substance dissolves in water it can be called soluble in water. Eg salt
If a substance does not dissolve in water it is called insoluble in water. Eg sand
•
The higher the temperature of the solvent the faster the
substance will dissolve.
E.g. Sugar will dissolve faster in water at 80ºC than water
at 25ºC
•
The smaller the particle size of the substance the quicker
the substance will dissolve.
E.g. Sugar granules will dissolve quicker than sugar lumps.
Some substances that do not dissolve in water can be dissolved in
other solvents.
•
Example: Nail varnish does not dissolve in water but does
dissolve in the solvent acetone.
5
Separating Mixtures
There are various methods used in the science lab to separate mixtures;
1. Filtration is used to separate a solid
from a liquid. The solid collected in
the filter paper is called the residue
and the liquid collected is called the
filtrate.
Eg. Used to separate sand and water.
2. Distillation can be used to separate
two different liquids. As different
liquids can have different boiling
points, one liquid can be boiled to
produce a gas which is then condensed
back into a liquid again.
Eg. Used to separate alcohol and water.
3. Chromatography can be used to find out which liquids are
contained within a mixture of liquids. Chromatography can
separate the liquids.
Eg. Used to separate the colours in ink.
6
Unit 2 – Chemical Reactions
Rate of Reaction
How fast a new substance is made is called the ‘speed/rate of reaction’.
Reactions take place when particles in the chemicals collide with each other.
The rate of a reaction depends on:
1. The frequency or how often the collisions between the reactant particles
happen.
2. The energy with which reactant particles collide.
As the reaction progresses, the concentration of reactants decrease, and the concentration
of the products increases.
Concentration
•
•
•
The units of concentration can be shown as M (eg. 1M) or mol/l (eg. 0.5mol/l).
The higher the number, the greater the concentration of the solution.
At a higher concentration, there are more particles in the same volume of space.
This means that the particles are more likely to collide and therefore more likely to
react.
Low Concentration
High concentration
concentra
concentra
7
Particle Size



Any reaction involving a solid can only take place on the surface of the solid.
If the solid is split into smaller pieces, the surface area increases.
This means that there is an increased area for the reactant particles to collide with.

The smaller the pieces, the larger the surface area. This means more collisions
and a greater chance of reaction.
Low Surface Area
High Surface Area
•
The following graph shows the
difference in reaction rate.
•
Powder reacts faster than lumps and
therefore the powder graph has a
steeper gradient than the ‘lump’ graph.
This explains why potatoes chopped into
smaller pieces cook faster than larger lumps.
Temperature
•
•
•
If we increase the temperature of a reaction we give the particles
more energy.
This means they will move faster and therefore are more likely to
collide with other particles.
When the particles collide, they do so with more energy, and so the
number of successful collisions increases.
8
Catalysts
•
Another way we can speed up the rate of a chemical reaction is to add a catalyst.
•
Unlike changing particle size, concentration or temperature, a catalyst is a substance
we add to a reaction to speed the reaction up.
•
An advantage of using a catalyst is that it is not used up during the reaction and can
be used again.
•
Enzymes are ‘biological catalysts’ and can be found in animals and plants.
E.g. Amylase, which is found in saliva.
Hazard Symbols
Corrosive
Toxic
Flammable
Irritant
9
Unit 3 – Elements, The Periodic Table and Bonding
The Periodic Table
•
The Periodic Table is
split into groups which
run vertically up and
down the table.
•
The Table is also split
into periods which go
horizontally across.
Metals & Non-Metals
Group
Name
Properties
Examples
1
Alkali Metals
Very reactive
Sodium, Lithium
7
Halogens
Reactive, used for
killing bacteria
Chlorine, Fluorine
8/0
Noble Gases
Unreactive
Neon, Xenon
Middle section
Transition Metals
Can vary
Copper, Silver
The thick zig zag line shown in the
diagram separates the metals and
non-metals.
Metals are found to the left of
the line and non-metals to the
right.
This information is also found on
page 3 of your databook.
10
Metals v. Non-Metals
Metals
Non-Metals
Solid.
(except MERCURY which is LIQUID)
Can be solid, liquid or gas.
Conduct electricity.
Do NOT conduct.
(except CARBON in the form of GRAPHITE)
High melting and boiling point.
Lower melting and boiling point.
Higher density.
Lower density.
Metals have many uses; copper for wires, aluminium for cans and planes and iron for bridges
and fences.
The Atom
ELECTRON
The atom is made from 3 different
subatomic particles.
PROTON
NEUTRON
Particle Name
Mass
Charge
Location
PROTON
1
Positive
Inside nucleus
NEUTRON
1
No charge
Inside nucleus
ELECTRON
0
Negative
Outside nucleus
For every element:
The Atomic Number
The Mass Number
= The number of protons.
= The number of protons + neutrons.
To find the number of neutrons in an element we use:
Neutrons
=
Mass Number
-
Atomic Number
11
**In a NEUTRAL atom the number of protons is EQUAL to the number of electrons.**
Nuclide Notation
No. of Protons
No of Electrons
No of Neutrons
=
=
=
17
17
35 – 17= 18
Electron Arrangements
As the number of electrons increases they arrange themselves in a particular order in
energy levels/shells.
Energy level
Maximum number of electrons
1
2
2
8
3
8 or 18
12
The electron arrangement (Page 3 databook) shows how electrons are arranged in atoms.
E.g.
Sodium
2 electrons in its 1st shell
8 electrons in its 2nd shell
1 electron in its 3rd shell
The Periodic Table group number tells us the number of outer electrons the element has
Eg. Group 3 elements have 3 electrons in the outer level.
Ions
•
•
•
Noble Gases have a stable electron arrangement.
Noble Gases are like the rockstars of the Periodic Table, atoms want to be like them.
In order to do that they LOSE or GAIN electrons to have the same electron
arrangement.
•
When this happens atoms form IONS, which are charged particles.
13
E.g.
Electron arrangement:
Na
2,8,1
Na+
2,8
+
e1
Sodium now has the electron arrangement of Neon.
Rule:
Metals lose electrons to form POSITIVE IONS.
Non-metals gain electrons to form NEGATIVE IONS.
The Covalent Bond
Remember atoms want to have the electron arrangement of a noble gas so one way they can
do this is by sharing electrons.

A COVALENT bond is formed when 2 non-metal atoms share outer electrons.

The covalent bond is very strong and is difficult to break.

Covalent bonds can be single bond (Cl-Cl), double bond (O=O) or triple bond (N≡N).
14
Diatomic Molecules
•
A diatomic molecule contains 2 atoms. There are SEVEN elements that exist as
diatomic molecules on the periodic table.
Hydrogen
Oxygen
Nitrogen
Fluorine
Chlorine
Bromine
Iodine
-
H2
O2
N2
F2
Cl2
Br2
I2
Shapes of Molecules
When atoms bond together the molecules produced can form different shapes.
The four main shapes are…
Linear
V-Shaped/Bent
Pyramidal
Tetrahedral
Chemical Formula

Chemical formula is a shorthand way of showing elements and compounds using
symbols from the Periodic Table.

Some covalent formula (non-metals) can be found by using the symbols given in the
name of the compound.
Eg. Silicon Tetrachloride - SiCl4
Formula using prefixes
If there is a prefix present in the name of the compound it tells you the number of atoms
present.
E.g. Carbon Dioxide – the carbon does not have a prefix so there must be only one C, the
oxygen has a DI prefix which means there are two O’s. The finished formula is CO 2.
15
Prefix
Number
Mon or mono
1
Di
2
Tri
3
Tetra
4
Pent
5
Formula using the Crossover Rule
If the name of the compound does not contain any prefixes we can use the Crossover Rule.
The valency number of an element is used; this is the number of bonds an atom can form.
Group
1
2
3
4
5
6
7
8/0
Valency
1
2
3
4
3
2
1
0
Chemical formula is written using the following steps
1.
2.
3.
4.
5.
Symbols
Valency number
Swap - crossover
Divide - if there is a common number*
Write the final formula – remember not to show
the 1 e.g C1H4 is shown as CH4
Symbols
C
H
Valency
4
1
Swap
1
4
Divide*
Formula
CH4
*The divide step is not always required.
As the Transition Metals do not have a group number their valency is given using Roman
Numerals
1 = (l) , 2 = (ll), 3 = (lll), 4 = (lV), 5 = (V), 6 =(Vl)
Eg.
Nickel (II) bromide
Iron (III) oxide
-
Nickel has a valency of 2
Iron has a valency of 3
16
Examples
Iron (II) Oxide
Symbols
Valency
Swap
Divide
Formula
Fe
2
2
1
FeO
Hydrogen Oxide
O
2
2
1
Carbon Tetrachloride*
Symbols
Valency
Swap
Divide
Formula
C
Cl
CCl4
Aluminium Oxide
Symbols
Valency
Swap
Divide
Formula
Al
3
2
O
2
3
Al2O3
Symbols
Valency
Swap
Divide
Formula
H
1
2
O
2
1
H 2O
Silicon Hydride
Symbols
Valency
Swap
Divide
Formula
Si
4
1
H
1
4
SiH4
Nitrogen Trihydride*
Symbols
Valency
Swap
Divide
Formula
N
H
NH3
Carbon Sulphide
Symbols
Valency
Swap
Divide
Formula
C
4
2
1
CS2
S
2
4
2
* do not use valency rules when the compound has a prefix
17
Unit 4 - Fuels
Fossil Fuels
•
Coal, oil and natural gas were created millions of years ago and
are known as the Fossil Fuels.
•
Coal was formed from dead plants that sank to the bottom of
swampy water and over millions of years were buried with layers
of mud.
•
These layers of mud were then compressed by the pressure as more layers were
added and formed coal.
•
Oil and gas were formed in a similar way from tiny sea creatures that sank to the
bottom of the sea millions of years ago.
•
The pressure of more layers of sand caused the production of oil and natural gas.
Fuels

Fuels are substances which burn in oxygen to release energy.
E.g
Coal, oil, gas, wood, peat and sugar

However, not all substances which give out energy are fuels.
A battery is NOT a fuel as it does not burn.
18
Fractional Distillation
Crude Oil is a mixture of different chemical compounds.
•
By using fractional distillation we can separate the different compounds as each has a
different boiling point.
•
The fractional distillation tower is used to separate the mixture into groups with
similar boiling points called fractions.
•
The tower is hot at the bottom to collect fractions with a high boiling point, and cool
at the top to collect fractions with a low boiling point.
Fractional Distillation
Top of the Tower
Bottom of the Tower
Small molecules
Large molecules
Low boiling point
High boiling point
Low viscosity
High viscosity
High flammability
Low flammability
Easily evaporated
Difficult to evaporate
19
Composition of Air
Approximately only 20% of the air
around us is Oxygen. The majority is
made up of Nitrogen.
The ratio of Oxygen to Nitrogen in
air is 1 : 4.
Combustion
When a fuel is burned in oxygen it is called combustion. This is an example of an
exothermic reaction. There are 2 types of combustion.
1. Complete Combustion – A plentiful supply of oxygen.
FUEL + O2
CO2 (Carbon dioxide) +
H2O (Water)
2. Incomplete Combustion – A limited supply of oxygen.
FUEL + O2
C (Carbon/Soot) + CO (Carbon monoxide)
+
H 2O
**The test for oxygen is that it relights a glowing splint**
Alkanes
A homologous series is a set of compounds with similar chemical properties which can be
represented by a general formula.
Alkanes:
 are a subset of the set of hydrocarbons (contain only Hydrogen and Carbon)
 all end in the letters -ane
 are a homologous series with general formula CnH2n+2

contain single C-C bonds (saturated)
20

each alkane had a different name depending on how many carbon atoms are present.
Prefix
methethpropbutpenthexhept
oct-
Number of Carbons
1
2
3
4
5
6
7
8
E.g. The alkane with 3 carbons is called propane.
Note:
If you forget the prefixes then look at page 6
of the data booklet. The alkanes are listed in a
table, in order, so you can work out the number
of carbons from that.
Structure of the Alkanes
Name
Formula
Methane
CH4
Shortened Structural
Formula
CH4
Ethane
C2H6
CH3CH3
Propane
C3H8
CH3CH2CH3
Butane
C4H10
CH3CH2CH2CH3
Full Structural Formula
21
Pollution
The burning of fossil fuels causes POLLUTION!
1. Acid Rain
•
•
Fossil fuels can often contain sulphur, which when burned
produces sulphur dioxide SO2. This acidic gas dissolves in clouds
and falls as acid rain.
Nitrogen oxides NOx are made by lighting or by spark plugs in car
engines. This is another acidic gas which dissolves in clouds and
falls as acid rain.
2. Global Warming
•
•
•
When a fuel is burned in enough oxygen it produces CO2
– a “Greenhouse gas”.
This gas is building up in the atmosphere and causing the
world to heat up, also known as the “Greenhouse Effect”
This could mean the melting of the polar ice caps with a
rising of sea levels and more severe weather.
3. Transportation of Crude Oil
•
Oil tankers can crash spilling oil and causing damage to the sea.
Reducing Pollution
•
To reduce pollution we have to use RENEWABLE ENERGY to make our electricity and
run our cars and buses.
HYDROELECTRIC
SOLAR
WIND FARMS
WAVE POWER
BIO-FUELS
-----------
uses falling water to make electricity.
uses energy from the sun
uses the wind to make electricity.
uses the waves in the sea to make electricity.
from plants to help us run our cars and buses
Renewable Fuels
Ethanol, obtained from sugar cane, is a RENEWABLE fuel.
22
Unit 5 – Acids & Alkalis
Ionic Introduction
Ionic bonding occurs between a metal and a non-metal.
The metal gives electrons away and the non-metal accepts electrons.
Metals form positive ions
Non-metals form negative ions
Ionic compounds from an ionic lattice. The ions are held together by electrostatic
attraction.
Ionic Formula
The charge for each ion can be obtained from the Periodic Table as follows:
Group
1
2
3
4
5
6
7
0
Charge on ion
1+
2+ 3+ **
3-
2-
1-
**
** Groups 4 and 0 do NOT form ions**
Ionic formulae are worked out by writing the symbols or formulae for the positive and
negative ions. Then, the positive and negative charges must be "balanced" (if they are not
already the same) as below.
Eg.
sodium chloride Na+ Clcopper (II) sulphide Cu2+ S2-
If the positive and negative ions don't have the same number of charges, we have to work
out how many of the positive ions and how many of the negative ions would be needed to
make the whole compound neutral.
When we need to show more than one ion in a formula, we put brackets round the ion as
below.
Eg.
sodium oxide (Na+)2 O2calcium bromide Ca2+ (Br-)2
iron (III) chloride Fe3+ (Cl-)3
aluminium oxide (Al3+)2 (O2-)3
23
Examples
Calcium oxide
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
Sodium fluoride
Ca
2
2
1
CaO
Ca2+O2-
O
2
2
1
Magnesium chloride
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
Mg
2
1
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
Na
1
1
F
1
1
NaF
Na+F-
Potassium sulphide
Cl
1
2
MgCl2
Mg2+(Cl-)2
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
K
1
2
S
2
1
K2S
(K+)2S2-
Formula involving complex ions
•
•
•
•
Complex ions contain more than one kind of atom.
These are found on pg 6 of the databook.
E.g. CO32- , NO3-, NH4+ and OHThe valency of these ions is the same as their charge. The formula is worked out in
the same way, using the crossover method.
Examples
Potassium nitrate
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
K
1
1
KNO3
K+NO3-
Ammonium chloride
NO3
1
1
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
NH4
1
1
Cl
1
1
NH4Cl
NH4+Cl-
24
Calcium hydroxide
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
Sodium carbonate
Ca
2
1
OH
1
2
Symbols
Valency
Swap
Divide
Chem Formula
Ionic Formula
Ca(OH)2
Ca2+(OH-)2
Na
1
2
CO3
2
1
Na2CO3
(Na+)2CO32-
The pH Scale


pH Scale measures how acidic or alkaline a solution is.
The pH can be found using universal indicator, litmus paper or using a pH meter.
Acid
Common lab
Common household
Alkali
Hydrochloric acid HCl
Sulphuric acid
H2SO4
Nitric acid
HNO3
Sodium hydroxide
Potassium hydroxide
Coke
lemonade
Bleach
Oven cleaner
NaOH
KOH
All acids contain the hydrogen ion H+
All alkalis contain the hydroxide ion OH-
High
conc.
H+
ions
Acids –
Acid
s
pH less than 7
Low
conc.
H+
ions
Neutral
Neutral
1 2 3 4 5 6 7 8 9 10 11 12 13 14
–
Alkalis
Low
conc.
OHions
OHions
pH 7
High
conc.
Alkali -
pH more than 7
25
State Symbols
Making Acids

When we burn certain non-metals in the presence of O2,
a non-metal oxide is produced.
E.g.

C(s) + O2(g)
Solid
s
Liquid
l
Gas
g
Aqueous
aq
CO2(g)
The non-metal oxide, if soluble, can then be dissolved in water to produce an acid.
CO2(g) + H2O(l)
(H+)2 CO32-(aq) (Carbonic acid)
Acid Rain
•
Coal contains sulphur which when burned produces sulphur dioxide.
S(s) + O2(g)
•
SO2(g)
When the sulphur dioxide rises and is absorbed by the clouds sulphuric acid is
formed.
SO2(g) + H2O(l)
(H+)2 SO42- (aq) (Sulphuric acid)
Other causes of acid rain
•
The energy from lightening in thunderstorms and the energy from the spark plug in
an engine allows the nitrogen and oxygen in air to react together.
•
The nitrogen dioxide is then absorbed into the clouds to form nitric acid.
NO2(g) + H2O(l)
H+NO3- (aq) (Nitric Acid)
Making Alkalis

If a metal oxide or metal hydroxide dissolves in water, an alkali is formed.
i)
ii)
Sodium oxide dissolves to form sodium hydroxide solution.
Calcium hydroxide dissolves to form calcium hydroxide solution.
26
Diluting Acids & Alkalis
When water is added to an acid the pH moves
towards 7 as the concentration of H+ ions
decreases.
When water is added to an alkali the pH moves
towards 7 as the concentration of OHdecreases.
Neutralisation – Acid & Alkali
•
Acids can be neutralised by certain substances to make neutral compounds (pH 7).
The reaction is called neutralisation and the substance which reacts with the acid is
called a neutraliser or a base.
•
During neutralisation a SALT is always produced.
Everyday neutralisation reactions include:
a)
b)
c)
adding lime to rivers and lochs to reduce the effects of acid rain.
treating a bee sting (acid) with baking soda (alkaline).
treating a wasp sting (alkaline) with vinegar (acid).
27
Neutralisation Reactions
When a salt is produced in a neutralisation reaction part of the name comes from the acid
used in the reaction and the other part from the neutraliser/base.
Acid
Salt Produced
Hydrochloric
Chloride
Sulphuric
Sulphate
Nitric
Nitrate
1. Acid & Alkali
Acid + Alkali (neutraliser)
Salt (neutral) + Water (neutral)
The hydrogen (H+) from the acid reacts with the hydroxide (OH-) from the alkali to form
water.
Eg.
Full equation
H+Cl-(aq)
H+(aq)
+
+
NaOH-(aq)
OH-(aq)
Na+Cl-
+
H2O(l)
H2O(l) (pH 7)
2. Acid & Reactive metal
Reactive Metal + Acid
Eg.
Magnesium + nitric acid
Salt +
Hydrogen
Magnesium nitrate + Hydrogen
**The test for hydrogen gas is that is burns with a “pop”.**
3. Acid & Carbonate
Carbonate + Acid
Eg.
Calcium carbonate + Hydrochloric acid
Salt + Carbon dioxide + Water
Calcium chloride + Carbon dioxide + Water
**The test for carbon dioxide is that it turns limewater cloudy/milky**
28
The Mole
The mole is the formula mass of an element, compound or molecule, expressed in grams.
1. For Elements
The gram formula mass of Aluminium is 27 grams, this information is found in page 6 of your
databook. Therefore 1 Mole of Aluminium is 27 grams.
Note: For diatomic elements the databook value is doubled.
E.g. for chlorine Cl2 the mass of one mole is 71g not 35.5g.
2. For Molecules or Compounds
The chemical formula of a molecule or compound tells us which elements are present, and in
what quantity.
Calculating the mass of all the elements present will give the mass of one mole of the
molecule or compound.
E.g.
Lithium Oxide (Li2O).
The formula tells us there are 2 atoms of lithium and 1 atom of oxygen present.
Therefore:
1 Mole of Li2O = 30g
E.g.
Sodium Hydroxide (NaOH)
The formula tells us there is 1 atom of sodium, 1 atom of oxygen and 1 atom of
hydrogen present.
1 Mole of NaOH = 40g
29
Unit 6 – Metals, Reactivity Series, Electricity & Corrosion
Properties of metals
Iron is a metal that is strong and so it can be used to build bridges and
railway tracks.
Copper metal has many uses. It has good electrical conduction. This
means that it is very good at conducting electricity so it is used to make
electrical wires. Copper has good thermal conductivity. This means that
it is good at letting heat move through it and so it is sometimes used to
make cooking pots. Copper does not corrode so it is also good for making
water pipes.
Aluminium is a metal that has a very low density.
This means that it is light so it is used to make aeroplanes.
Metals that are used to make jewellery, have to be malleable.
This means that they are can be hammered easily into different shapes.
Gold, silver and platinum are very malleable metals. These metals are
also used because they do not corrode and stay shiny for a long time.
Tin metal does not corrode so it is a good metal to use to make food cans.
Alloys




E.g.
Pure metals do not always have the properties that we need for a particular job.
An alloy is a substance made by melting and mixing metals with other elements.
Alloys are often more useful than pure metals because they have different
properties from the pure metals.
This can make them more suitable for certain uses.
Solder is an alloy of TIN and LEAD
Stainless steel is an alloy of IRON, CARBON and CHROMIUM
(a transition metal)
30
Reactions of Metals
Metal + Water
Metal + Acid
Metal + Oxygen
Metal hydroxide + Hydrogen
Salt + Water
Metal oxide
Reactivity Series
By observing the reactions of metals we are able to build the “Reactivity Series” which
shows how reactive metals are relative to each other.
Extraction of Metals

Some metals are uncombined. This means we find them pure in the ground; not joined
to other elements. Examples of uncombined metals are silver and gold.

Most metals are not like this. We find them combined with other elements in a
compound known as an ORE. Most ores are metals joined to oxygen (metal oxide).

To obtain a metal from a metal oxide we need to separate the metal atoms from the
oxygen atoms. The more reactive the metal the tighter it is joined to the oxygen so
the harder it is to produce!
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Extraction of Metals
Batteries





In a battery electricity comes from a chemical reaction.
Batteries need to be replaced when the chemicals inside are used up.
Some batteries are rechargeable; e.g. the lead-acid battery.
All batteries contain electrolytes. (Usually ammonium chloride paste)
The purpose of the electrolyte is to complete the circuit.
Mains Power v. Batteries
Mains Power
Batteries
Advantages
Cheaper
Safe to use
Portable
Disadvantages
Danger of being electrocuted
Uses finite resources
Expensive
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The Electrochemical Series (ECS)
 Shows metals in order of their willingness to give up electrons.
 Shown on page 7 of the data booklet (Reactivity Series).
Cells


Electricity can be produced by connecting two different metals together (with an
electrolyte eg ammonium chloride) to form a cell.
Electricity can also be produced in a cell by connecting two different metals in
solutions of their metal ions.
Electrons flow from
the metal higher in
the ECS to the
metal lower i.e.
from magnesium to
lead.


The purpose of the “ion bridge” is to complete the circuit.
The greater the distance between the metals in the electrochemical series, the
higher the voltage produced.
Displacement Reactions

Reactions which occur when a metal higher up in the electrochemical series is added
to a solution containing ions of a metal lower down in the series.
E.g. Magnesium metal added to copper ions;
Mg(s) + Cu2+SO42-(aq)
Mg2+SO42-(aq)
+
Cu(s)
The higher magnesium metal will form ions in solution
The lower copper ions will form copper metal and come out of solution.
I.e. they will be displaced.
33

The ECS can be used to predict whether or not a displacement reaction will occur.
The metal being added must be higher than the ions in solution for displacement to
occur.
Corrosion




Corrosion is a chemical reaction which involves the surface of a metal changing from
an element to a compound (metal oxide).
This is an example of oxidation.
The corrosion of iron is known as rusting.
Water and oxygen are required for corrosion to occur.
Preventing Corrosion
A surface barrier to air (oxygen) and water can provide physical protection against
corrosion.
E.g.






Painting
Greasing
Coating with plastic
Galvanising - metal objects are dipped into molten zinc
Electroplating - silver, chromium and other metals can be deposited on the
surface of a metal
Tin-plating
Gas Tests - Summary
Gas
Test
Oxygen
Relights a glowing splint
Hydrogen
Burns with a ‘pop’
Carbon dioxide
Turns limewater milky/cloudy
34
Unit 7 – Bonding, Structure & Properties
Electrical Conductivity
Electric current is a flow of charged particles
Testing conductivity
1. electrons flow through
metals.
2. ions flow through solutions or
melts.
Testing a solid
Testing a solution

Metal elements and carbon (graphite) are electrical conductors.

Non-metal elements are non-conductors of electricity.

Covalent compounds do not conduct in any state as they only
contain non-metals.


Ionic compounds conduct when molten or in solution.
Ionic compounds do not conduct when solid.
State at Room
Temperature
 Covalent
compounds can be
solids, liquids or
gases.
 Ionic compounds
are all solids.
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Structure of Compounds
Ionic solids exist as lattices of
oppositely charged ions.
E.g. Na+Cl-
Some covalent solids are network structures
E.g. SiO2, diamond
Ionic lattices and covalent networks have strong bonds and
therefore high melting points.
Other covalent compounds exist as discrete molecules.
(see opposite)
Discrete covalent molecules have weak forces of attraction and
so these compounds have low melting points.
Solubility

Some covalent substances do not dissolve in water but will dissolve in other covalent
solvents.
Eg.


Nail varnish being dissolved in propanone.
Paint dissolving using turpentine.
Ionic compounds generally dissolve in water. Eg. Sodium chloride & copper sulphate.
Dissolving ionic compounds break up the ionic lattice.

Solid ionic lattice
Dissolved ionic compound
(ions are free to move)
36
Electrolysis

Electrolysis is used to break up a compound into its elements using electricity.

Positive metal ions are attracted to the negative electrode.

Negative non-metal ions are attracted to the positive electrode.

Electrodes are made from carbon in the form of GRAPHITE as it conducts electricity.

A d.c. (direct current) supply must be used if the products are to be identified.
Covalent compounds cannot be electrolysed as they do not conduct electricity and they
have no ions.
Coloured Ions

Some ions are coloured. Examples are shown in the table below.
Ion
copper
nickel
potassium
chromate
permanganate
sulphate
Colour
blue
green
colourless
yellow
purple
colourless
37
Migration of Ions
In the experiment below;


the positively charged blue copper ions move towards the negative electrode
the negatively charged orange dichromate ions move towards the positive electrode.
38
Glossary
Definition
Acid
Acid rain
Alkali
Alkali metals
Atomic number
Base
Blast furnace
Catalyst
Chemical
reaction
Combustion
Concentration
Concentrated
solution
Corrosion
Covalent bond
Crude oil
Diatomic
Dilute solution
Electrolysis
Electron
Electron
arrangement
Element
Endothermic
Exothermic
Fossil fuel
Substance with a pH less than 7.
Source of pollution caused by sulphur dioxide SO2 and nitrogen dioxide gas
NO2
Substance with a pH greater than 7.
Very reactive group 1 elements.
Number of protons in the nucleus of an atom.
Substance which will neutralise an acid.
A tall oven used to extract iron from iron ore by burning it with carbon at
high temperatures.
A substance that increases the rate of a chemical reaction without being
used up.
A change in which new substances are made and cannot easily be reversed.
Reaction which involves burning a substance in oxygen (O2).
The number of molecules of a substance in a given volume. Units are mol/l,
mol l-1 or M
A solution that contains a large amount of solid dissolved.
Surface of a metal changing from a metal to a compound (metal oxide).
A shared pair of electrons between two non-metal atoms.
A fossil fuel made up of a mixture of hydrocarbons.
Containing only 2 atoms. E.g. HCl, H2, NO, CO, O2
A solution containing a small mass of dissolved solute.
Breaking a compound into its elements using electricity.
Negatively charged subatomic particle found in electron shells.
Shows the number of electron shells and how many electrons they contain.
A substance made up of only one type of atom.
A reaction which releases heat to the surroundings ie gets cold.
A reaction which releases heat.
Fuels formed from fossils over millions of years at a high temperature and
pressure; coal, oil and gas.
39
Fractional
distillation
Group
Graphite
Halogen
Hydrocarbon
Method of separating crude oil depending on boiling point.
Insoluble
Ion
Ionic bond
A substance than cannot be dissolved in water.
A charged particle formed when an atom loses or gains electrons.
A bond formed between a metal and a non-metal.
Loam
A mixture of small pieces of rock or sand surrounded by decayed animal or
plant remains.
A compound containing metal, hydrogen and oxygen atoms which will
neutralise an acid and has a pH greater than 7.
A compound containing metal and oxygen atoms only.
Rock formed when heat and pressure cause changes in existing igneous or
sedimentary rocks over a long period of time.
Two or more substances brought together but are not chemically joined.
Eg. Air and crude oil.
Scale used to measure the ‘hardness’ of a mineral.
A small group of atoms that are held together by covalent bonds.
Group 8/0 elements in the Periodic Table. Unreactive elements with a full
outer shell of electrons.
The small, dense, postively charged centre of an atom, made up of protons
and neutrons.
Substance with a pH 7 which turns universal indicator green.
Horizontal line going across the Periodic Table.
Positive subatomic particle found within the nucleus.
When two solutions react and a solid is formed.
A substance that reacts quickly or easily. Eg. Alkali metals.
How quickly or easily a substance will react.
The specific name for the corrosion of iron.
A solution in which no more solute can be dissolved.
Metal
hydroxide
Metal oxide
Metamorphic
rock
Mixture
Mohs scale
Molecule
Noble gases
Nucleus
Neutral
Period
Proton
Precipitation
Reactive
Reactivity
Rusting
Saturated
solution
Soluble
Solute
Solution
Solvent
Vertical column of elements in the Periodic Table, with similar properties.
Form of carbon with is an electrical conductor. Used as electrodes.
Group 7 elements in the Periodic Table.
A molecule containing only hydrogen and carbon.
A
A
A
A
solute which can dissolve in a liquid.
solid which is dissolved in the solvent.
liquid containing a dissolved solid.
liquid which can dissolved a solute.
40
Transition
metals
Universal
Indicator
Unreactive
Valency
Metals that are found in the middle of the periodic table. They do not
have a group number given to them. Their valency is given in Roman
numerals. Eg. Titanium.
Solution used to test the pH of substances.
Acids – red, alkali – blue and neutral – green.
A substance that reacts very slowly or does not react at all. Eg. Noble
gases.
How many bonds an atom can form. Eg. Group 4 elements have a valency
(combining power) of 4 and therefore form 4 bonds.
41