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A level CHEMISTRY
PRIMER
The purpose of this booklet is to bridge the gap between GCSE Double Award Science
and Separate Science Chemistry and to prepare for A level.
By the start of the course you should be confident with the content, much of which should
already be familiar. So by the start of term:

read through the notes (pages 2 to 9).

learn the ionic charges given in the table on page 7.

complete all the even-numbered questions in the exercises1 that follow page 10,
(writing your answers in the spaces provided.) If there are any exercises that you
are finding difficult then try the odd numbers as well.
1
Taken from UA008883 GCE Chem Moles wkbk Iss3
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What is it all about?
Chemistry is the creative central science, the science of changing matter.
What is the matter?
Nearly all real stuff / matter is a mixture of substances. A pure substance has characteristic
properties such as colour, solubility, melting point (Tm) and boiling point (Tb).
Elements are the simplest pure substances. There are about one hundred elements, divided into
metals and nonmetals although the division is not always clear-cut. Metals conduct electricity
whereas nonmetals do not; graphite is a metalloid as it does conduct but generally behaves as a
non-metal.
Compounds, of which there are millions, are more complicated pure substances, being
composed of two or more elements bonded (not mixed!) together.
In chemical changes, chemical reactions, new substances are formed. Elements cannot be
broken down, they can only build up (synthesise) compounds. Compounds can be converted
into other compounds and may be broken down (decomposed) into simpler compounds or into
elements.
There are three common states of matter: solid, liquid and gas. It is possible to have one state
spread through another as in colloids (such as foam, suspensions, gels).
Atomic Theory (formulated by Dalton in 1808, updated with comments in italics).
1.
All matter is made of minute indivisible particles, atoms. Atoms are themselves made of
smaller particles but an atom is the smallest, electrically neutral particle of an element that
can participate in a chemical reaction.
2.
Atoms cannot be created or destroyed except in nuclear reactions.
3.
The atoms of an element are identical. Atoms of an element have the same atomic
number but not necessarily the same mass number (isotopes).
4.
Atoms of an element are different from the atoms of all other elements.
5.
When elements react to form compounds, the atoms of the elements combine, bond, in
simple whole number ratios. In chemical reactions, new substances are formed as atoms
change their partners.
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Relative atomic mass scale
Hydrogen atoms are the lightest.
The relative atomic mass, r.a.m. or Ar, of an element is the number of times heavier an atom of
the element is than one hydrogen atom. This definition will be revised at the start of the AS
course.
Since a carbon atom is twelve times heavier than a hydrogen atom
Ar (C) =
12
For water particles, made of three atoms bonded together, the relative formula mass, Mr, is
calculated thus
Mr (H2O)
=
2*1.0 + 16.0 =
18.0
and for hydrated aluminium sulfate, the relative formula mass
Mr (Al2(SO4)3.16H2O =
2*27.0 + 3(32.1 + 4*16.0) + 16*18.0
=
630.3
Note that the “.16H2O” is the water of crystallization, water that is part of the crystals.
Relative formula mass is perhaps a better term for what is more commonly known as relative
molecular mass.
Mr (NaCl)
=
23.0 + 35.5 =
58.5
Note that relative atomic and relative formula (molecular) masses have no units.
Atomic Structure
Model of an atom
subatomic
mass on Ar charge
particle
scale
proton
1
+1
neutron
1
0
electron
almost 0
-1
volume occupied by
electrons
nucleus holding
protons & neutrons
Numbers
atomic number = proton number
mass number
=
( =
electron number
)
neutron number + proton number
Points

Both atomic number and mass number really refer to the nucleus of the atom.

The mass number of an atom is effectively the same as its Ar, because the mass of the electrons is negligible.
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Electron arrangement
Electrons are arranged in shells (energy levels) around the nucleus.
The first shell (nearest the nucleus – the lowest energy level) can hold only two electrons.
The second shell can hold only eight electrons.
The third shell can hold eight electrons.
group
1
2
period
1
2
3
4
3
4
5
6
7
0
F
2.7
Cl
2.8.7
He
2
Ne
2.8
Ar
2.8.8
H
1
Li
2.1
Na
2.8.1
K
2.8.8.1
Be
2.2
Mg
2.8.2
Ca
2.8.8.2
B
2.3
Al
2.8.3
C
2.4
Si
2.8.4
N
2.5
P
2.8.5
O
2.6
S
2.8.6
Periodic Table Elements arranged in order of atomic number.
Period
Elements having the same electron shell filled.
Group
Elements have the same number of electrons in the outer shell (valence
electrons) and have similar chemical properties.
Noble gases
Elements have a full (or empty) outer shell and are chemically unreactive.
Atoms of the other elements bond to get this arrangement.
Atomic size
Across a period the increasing nuclear charge causes the atoms to get smaller.
The chemistry of an element is governed by its outer valence electrons.
When showing the valence electrons (in a dot & cross diagram), set out in the order north, south,
east, west before pairing in the same sequence, e.g.
15
4
3
SYMBOL
8
7
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Covalent bonding
Nonmetal atoms can effectively gain electrons by sharing pairs forming molecules. Covalent
bonding then results from the attraction of the nuclei for the shared pair of electrons.
Using the dot & cross approach: show just one atom of each element before bonding (to reduce
clutter); bonding involves the pairing of unpaired electrons.
Examples:
structural formula
molecular formula
Points

Covalent bonds are formed by non-metal atoms sharing pairs of electrons.

The number of covalent bonds formed by an atom = number of unpaired valence
electrons

Particles of atoms bonded covalently are called MOLECULES.

Inside molecules the bonds are strong.

Between molecules there are weak forces.
Note
Use the terms bonds & forces to differentiate the relative order of magnitude of
interactions.
It is sensible to learn the formulae of the most common covalent substances.
elements: hydrogen H2, oxygen O2, halogen (i.e. chlorine/bromine/iodine) Hal2, nitrogen N2
compounds with trivial names: water H2O, ammonia NH3, methane CH4, ethane C2H6,
propane C3H8, butane C4H10.
Others have systematic names providing an indication of their formulae: carbon monoxide CO,
carbon dioxide CO2, phosphorus trichloride PCl3, silicon tetrachloride SiCl4, phosphorus
pentachloride PCl5.
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Ionic bonding
When a metal combines with a nonmetal, electrons are transferred between the atoms forming
ions. Ionic bonding then results from attractions between the oppositely charged ions.
Using the dot & cross approach: show just one atom of each element before bonding (to reduce
clutter); ions are charged as a result of having unequal numbers of protons and electrons.
Examples
formula
Points

Ionic bonds are formed when electrons are transferred between atoms.

Metal atoms lose electrons, becoming positive IONS, ca+ions.

Non-metal atoms gain electrons, becoming negative IONS, anions.

Ionic bonding is the strong electrostatic attraction between oppositely charged ions..

Ions are arranged in a giant lattice, with ions of the opposite charge nearest, e.g.
sodium chloride structure:
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Ionic Formulae
The charges of the most commonly encountered ions are:
Cations
Anions
gp 1 metals M+
halides (gp 7) Hal-
hydrogen H+
1+
ammonium
hydrogencarbonate
1-
NH4+
HCO3-hydroxide OH-
silver Ag+
manganate(VII) MnO4nitrate NO3nitrite NO2-
most metals M2+
carbonate, CO32chromate(VI) CrO42dichromate(VI) Cr2O72-
2+
2-
oxide O2sulfide S2sulfite SO32sulfate SO42-
3+
3-
3+
aluminium Al
chromium(III) Cr
3-
phosphate PO4
3+
iron(III) Fe3+
Ions containing two or more elements are called radical ions.
The common radical ions can be remembered from parent acids (compounds containing
replaceable hydrogen).
Sulfuric acid (hydrogen sulfate) H2SO4
=
2 H+ + SO42-
Sulfurous acid (hydrogen sulfite) H2SO3
=
2 H+ + SO32-
Nitric acid (hydrogen nitrate)
HNO3
=
H+ + NO3-
Nitrous acid (hydrogen nitrite)
HNO2
=
H+ + NO2-
Phosphoric acid (hydrogen phosphate) H3PO4
=
3 H+ + PO43-
Carbonic acid (hydrogen carbonate)
=
2 H+ + CO32-
or
H+ + HCO3-
=
H+ + OH-
H2CO3
(hydrogen hydrogencarbonate)
Water (hydrogen hydroxide)
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Rules for naming ionic compounds
These will be refined during the course
1. the cation is stated before the anion.
2. binary (two element) compounds end in -ide, e.g. Na2O sodium oxide, MgCl2 magnesium
chloride, LiH lithium hydride
3. exceptions to rule 2 are hydroxides and cyanides, e.g. Ca(OH)2 calcium hydroxide and KCN
potassium cyanide
4. with metals other than those in groups 1, 2 or 3, e.g. the transition metals which can have
more than one charge, the charge must be shown using Roman numerals, e.g. FeCl2 iron(II)
chloride & FeCl3 iron(III) chloride and similarly PbO lead(II) oxide & PbO2 lead(IV) oxide.
5. oxyanions end in -ate or –ite, the latter indicating a lower oxygen content than the former,
2-
2-
-
-
e.g. sulfate SO4 & sulfite SO3 and nitrate NO3 & nitrite NO2 .
Method for deducing the formulae of ionic compounds
In an ionic compound the charge of the cations must be balanced by the charge of the anions.
Knowing the charge of the ions therefore gives a simple way of deducing the chemical formula.
Ions
Substance
formula
(see page 7)
sodium chloride
Na
1+
Cl
1-
NaCl
sodium oxide
Na
1+
O
2-
Na2O
magnesium oxide
Mg 2+ O 2-
MgO
calcium chloride
Ca 2+ Cl 1-
CaCl2
aluminium oxide
Al 3+ O 2-
Al2O3
sodium hydroxide
Na 1+ OH 1-
NaOH
calcium hydroxide
Ca 2+ OH 1-
Ca(OH)2
calcium carbonate
Ca 2+ CO3 2-
CaCO3
Ca 2+ HCO3 1-
Ca(HCO3)2
calcium
hydrogencarbonate
Note that in the formula

the figure 1 is not written (but is assumed if no other number present)

a bracket must be placed around a radical if it is multiplied by 2 or more

numbers are cancelled / simplified if possible.
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Writing equations
Tackle in four steps.
1.
Write word equation.
2.
Put formulae below words.
3.

Metal
just symbol

Covalent stuff (compound & non-metal)
you need to know formula

Ionic compound
use ionic charges
Balance.
The numbers of atoms on each side of the equation must be the same. Extra numbers are only
allowed in front of a formula.
4.
Add state symbols. This is an optional extra in most cases but is essential in some situations (e.g.
precipitation reactions).
(s)
solid
(l)
liquid
(g)
gas
(aq)
aqueous, i.e. dissolved in water
Examples These have been set out step by step to show the working; in practice this is unnecessary!
magnesium
+
→
oxygen
magnesium
oxide
Mg
+
O2
→
MgO
2 Mg
+
O2
→
2 MgO
2 Mg (s)
+
O2 (g)
→
2 MgO (s)
sodium
+
→
sodium
water (hydrogen
hydroxide)
+
hydrogen
hydroxide
Na
+
H2O
→
NaOH
+
H2
2 Na
+
2 H2O
→
2 NaOH
+
H2
2 Na (s)
+
2 H2O (l)
→
2 NaOH (aq)
+
H2 (g)
sulfuric (hydrogen +
acid
sulfate)
→
iron(III)
hydroxide
iron(III)
+
water
sulfate
H2SO4
+
Fe(OH)3
→
Fe2(SO4)3
+
H2O
3 H2SO4
+
2 Fe(OH)3
→
Fe2(SO4)3
+
6 H2O
3 H2SO4 (aq)
+
2 Fe(OH)3 (s) →
Fe2(SO4)3 (aq)
+
6 H2O (l)
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Exercise 1:
Calculating the relative formula mass of compounds
See page 3 of the notes. You will find relative atomic masses on the periodic table on page 10.
1
H2O
2
CO2
3
NH3
4`
SO2
5
C2H4
6
H2SO4
7
NaCl
8
CuSO4
9
C2H5OH
10
HNO3
11
CaCl2
12
FeSO4
13
Na2CO3
14
Pb3O4
15
Na2SO4
16
KMnO4
17
K2Cr2O7
18
KHCO3
19
CH3CO2H
20
CH3COCH3
21
Ca(OH)2
22
Mg(NO3)2
23
Ca(HCO3)2
24
Pb(NO3)2
25
Al(NO3)3
26
Fe2(SO4)3
27
(NH4)2SO4
28
CuSO4.5H2O
29
(COOH)2.2H2O
30
(NH4)2SO4.Fe2(SO4)3.24H2O
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Exercise 2:
Writing formulae from their names
See pages 5, 7 & 8 of the notes.
1
Carbon Dioxide
2
Carbon Monoxide
3
Phosphorus Trichloride
4
Phosphorus Pentachloride
5
Silicon Tetrachloride
6
Silicon Dioxide
7
Sulfur Dioxide
8
Sulfur Trioxide
9
Nitrogen Dioxide
10
Dinitrogen tetraoxide
11
Sodium Chloride
12
Potassium Bromide
13
Magnesium Chloride
14
Silver Chloride
15
Calcium Iodide
16
Aluminium Chloride
17
Sodium Hydroxide
18
Potassium Nitrate
19
Ammonium Chloride
20
Magnesium Oxide
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21
Calcium Carbonate
22
Sodium Carbonate
23
Lithium Sulfate
24
Sodium Phosphate
25
Aluminium Oxide
26
Magnesium Hydroxide
27
Aluminium Hydroxide
28
Barium Nitrate
29
Ammonium Carbonate
30
Aluminium Sulfate
31
Copper(II) Oxide
32
Copper(II) Sulfate
33
Iron(II) Chloride
34
Copper(I) Chloride
35
Lead(II) Carbonate
36
Lead(IV) Oxide
37
Tin(IV) Chloride
38
Iron(III) Chloride
39
Iron(III) Sulfate
40.
Silver(I)Oxide
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Exercise 3:
Deducing names from formulae
See pages 5, 7 & 8 of the notes.
Questions 1 - 10
Some of these you just need to know, so look them up if you don’t.
1
H2O
2
CO2
3
NH3
4
O2
5
SO3
6
HCl
7
CH4
8
H2SO4
9
HNO3
10
C8H18
11
NaCl
12
Ca(NO3)2
13
Al2(SO4)3
14
KMnO4
15
(NH4)2CO3
16
KClO3
17
KHCO3
18
CsAt
19.
Sr(OH)2
20
NH4VO3 (V is Vanadium)
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Questions 21 - 30
21
FeSO4
22
FeCl3
23
PbO
24
MnO2
25
Cu(NO3)2
26
CuCl
27
AgNO3
28
Co(NO3)2
29.
PbCl4
30.
Cr(OH)3
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Exercise 4
Balancing equations
Balance the following equations. All the formulae are correct.
1
H2
+
O2
→
H2O
2
BaCl2
+
NaOH
→
Ba(OH)2
+ NaCl
3
H2SO4
+
KOH
→
K2SO4
+
H2O
4
K2CO3
+
HCl
→
KCl
+
H2O
+
CO2
5
CaCO3
+
HNO3
→
Ca(NO3)2
+
H2O
+
CO2
6
Ca
+
H2O
→
Ca(OH)2
+
H2
7
Pb(NO3)2
+
NaI
→
PbI2
+
NaNO3
8
Al2(SO4)3
+
NaOH
→
Al(OH)3
+
Na2SO4
9
N2
+
H2
→
NH3
10
H3PO4
+
NaOH
→
Na3PO4
11
NaNO3
→
NaNO2
+
O2
12
CH4
+
O2
→
CO2
+
H2O
13
C4H10
+
O2
→
CO2
+
H2O
14
H3PO4
+
NaOH
→
NaH2PO4
+
H2O
15
6NaOH
+
Cl2
→
NaClO3
+
NaCl +
16
Fe2O3
+
CO
→
Fe
+
CO2
17
C2H5OH
+
PCl3
→
C2H5Cl
18
2KMnO4
+
HCl
→
MnCl2
19
Al(OH)3
+
NaOH
→
NaAlO2
20
Pb(NO3)2
→
PbO
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+
+
+
+
Cl2
H2O
H2O
H3PO3
+ 8H2O + KCl
+
H2O
NO2
+
O2
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Exercise 5
Writing equations in symbols from equations in words
In the following examples you will need to convert the names of the materials into formulae and
then balance the resulting equation.
1.
Hydrogen gas reacts with oxygen gas to make water.
2
Liquid silicon tetrachloride reacts with water to produce solid silicon dioxide and
hydrogen chloride gas.
3.
Zinc metal reacts with copper sulfate solution to produce solid copper metal and zinc
sulfate solution.
4.
When octane (C8H18) vapour is burnt with excess air in a car engine, carbon dioxide and
water vapour are produced.
5
When magnesium is added to dilute nitric acid, a solution of magnesium nitrate is
produced and bubbles of hydrogen gas.
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6.
When lead(II) nitrate crystals are heated in a dry tube lead(II) oxide, nitrogen dioxide gas
and oxygen are produced.
7
When a solution of calcium hydrogencarbonate is heated, a precipitate of calcium
carbonate is produced together with carbon dioxide gas and water.
8.
Solid calcium hydroxide reacts with solid ammonium chloride on heating to produce solid
calcium chloride, steam and ammonia gas.
9.
A solution of ammonium hydroxide will neutralize sulfuric acid to make ammonium
sulfate solution and water.
10
When solutions of silver nitrate and calcium chloride are mixed a white precipitate of
silver chloride is formed.
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Exercise 6
What’s wrong here?
The following equations have one or more mistakes. Are the formulae correct? Is the balancing
correct? Are the state symbols correct and, most importantly, does the reaction actually happen?
Identify the error(s) and then rewrite the equation correctly.
1
Na (s) + H2O (l) → NaOH (aq) + H (g)
2
PbNO3 (aq) + NaCl (aq) → PbCl (s) + NaNO3 (aq)
3
CaOH2 (aq) + 2 HCl (aq) → CaCl2 (aq) + 2H2O (l)
4
C2H4 (g) + 2 O2 (g) → 2 CO2 (g) + 2 H2 (g)
5
MgSO4 (aq) + 2 NaOH → Ca(OH)2 (s) + Na2SO4 (aq)
6
Cu(NO3)2 (s) + CuO (s) → 2 NO (g) + O3 (g)
7
Cu (s) + H2SO4 (aq) → CuSO4 (aq) + H2 (g)
8
AlCl2 (s) + 2 KOH (aq) → Al(OH)2 (s) + 2 KCl (aq)
9
NaCO3 (s) + 2 HCl (aq) → NaCl2 (aq) + CO2 (g) + H2O (l)
10
2 AgNO3 (aq) + MgCl2 (aq) → Mg(NO3)2 (s) + 2 AgCl (aq)
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