This booklet contains notes for guidance,
intended for students who will sit for the
Form III National Assessment in
Chemistry. It is a free booklet meant to
help Mauritian students in Chemistry and
it is not for sale.
Chemistry
National Assessment 2016
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Dushan BOODHENA
Table of Contents
CHAPTER ONE: CHEMICAL SUBSTANCES ....................................................................................... 4
Elements and symbols ........................................................................................................................ 4
Classifying elements............................................................................................................................ 4
As solids, liquids and gases ............................................................................................................. 4
As metals and non-metals............................................................................................................... 4
In a table called the Periodic Table ................................................................................................. 4
The reactivity series ............................................................................................................................ 6
Compounds and formulae .................................................................................................................. 6
Atoms and molecules.......................................................................................................................... 7
Classifying substances as elements, compounds or mixtures ............................................................ 7
Acids .................................................................................................................................................... 8
Importance of selected acids .......................................................................................................... 9
Bases and alkalis ................................................................................................................................. 9
Importance of selected bases and alkalis ....................................................................................... 9
Indicators ............................................................................................................................................ 9
The pH (potentiometric hydrogen ion concentration) scale .............................................................. 9
CHAPTER TWO: THE LANGUAGE OF CHEMISTRY.......................................................................... 12
Physical and chemical changes ......................................................................................................... 12
Shorthand representation of names of elements ............................................................................ 12
Shorthand representation of names of compounds ........................................................................ 13
Deducing the formulae of simple compounds using symbols and valencies ................................... 13
Chemical reactions, reactants and products .................................................................................... 15
Word equations and chemical equations ......................................................................................... 16
CHAPTER THREE: CHEMICAL REACTIONS IN GENERAL.................................................................. 18
Simple equipment and glassware used in chemistry........................................................................ 18
Making different compounds by chemical reactions ....................................................................... 19
The reactivity series .......................................................................................................................... 20
The reactions of some metals with the oxygen of the air ............................................................ 20
The reactions of some metals with water .................................................................................... 21
The reactions of some metals with dilute mineral acids .............................................................. 22
Displacement reactions ................................................................................................................ 22
Action of heat on metallic carbonates .......................................................................................... 23
Summary of reactivity series of metals......................................................................................... 23
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Laboratory preparation of hydrogen ................................................................................................ 23
Laboratory preparation of oxygen .................................................................................................... 24
Laboratory preparation of carbon dioxide ....................................................................................... 25
The rusting of iron and its prevention .............................................................................................. 25
CHAPTER FOUR: IMPORTANT CHEMICAL REACTIONS .................................................................. 28
Characteristic reactions of acids ....................................................................................................... 28
Characteristic reactions of bases and alkalis .................................................................................... 28
Neutralisation reactions and their applications ............................................................................... 29
Salts ................................................................................................................................................... 29
Solubility of salts ........................................................................................................................... 30
Preparation of soluble salts .......................................................................................................... 30
Uses of selected salts .................................................................................................................... 32
Combustion and its importance........................................................................................................ 32
Respiration and its importance ......................................................................................................... 32
Photosynthesis and its importance................................................................................................... 33
Percentage composition of the gases in the air................................................................................ 33
How the process of respiration and photosynthesis maintains the composition of air ................... 33
Air pollution and its effects ............................................................................................................... 33
Burning of fossil fuels, global warming and acid rain ....................................................................... 34
CHAPTER FIVE: EXPERIMENTAL TECHNIQUES IN CHEMISTRY ....................................................... 36
Dissolving substances ....................................................................................................................... 36
Different types of mixtures ............................................................................................................... 36
Changes of states .............................................................................................................................. 36
Melting and melting point ............................................................................................................ 36
Boiling and boiling point ............................................................................................................... 37
Other changes of states ................................................................................................................ 37
Importance of pure substances ........................................................................................................ 37
Separation of mixtures and their importance in everyday life ......................................................... 38
Use of magnet ............................................................................................................................... 38
Decantation................................................................................................................................... 38
Filtration ........................................................................................................................................ 38
Crystallisation ................................................................................................................................ 39
Simple distillation.......................................................................................................................... 39
Fractional distillation .................................................................................................................... 40
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Sublimation ................................................................................................................................... 40
Chromatography ........................................................................................................................... 41
INDEX.…………………………………………………………………………………………………………………………………………43
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CHAPTER ONE: CHEMICAL SUBSTANCES
Elements and symbols
Experiment 1: Observing and classifying some metallic elements.
Experiment 2: Observing and classifying some non-metallic elements.
An element is a pure substance that cannot be broken down into simpler substances by any
ordinary chemical process.
It is made up of its own types of atoms and each element has a name and a chemical symbol, for
example hydrogen (H) and zinc (Zn). A symbol is a shorthand way of writing the name of an element.
All known substances are made from elements.
The diagram below shows how we can represent an element in the gaseous state:
Classifying elements (Optional)
As solids, liquids and gases
Elements can be classified as solids, liquids or gases. For example, iron and carbon are solids.
Mercury and bromine are liquids. Oxygen and nitrogen are gases.
As metals and non-metals
They can also be classified as metals and non-metals. Examples of metals or metallic elements are
iron, mercury, copper, gold, aluminium, etc. Examples of non-metals or non-metallic elements are
carbon, sulfur (also written as ‘sulphur’), bromine, chlorine, oxygen, etc.
In a table called the Periodic Table
Similar elements can also be grouped together in a table called the Periodic Table. A colour copy of
the periodic table is shown on the next page.
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Periodic Table of the Elements
Alkali Metals
Alkaline Earth Metals
Transition Metals
Other Metals
Nonmetals
Noble Gases
Lanthanoids
Actinoids
hydrogen
1
H
1.00794
lithium
beryllium
3
4
C
Br
He
solid
liquid
gas
Tc
synthetic
helium
2
He
key
element name
boron
carbon
nitrogen
oxygen
fluorine
4.002602
neon
atomic number
5
6
7
8
9
10
Li
Be
symbol
B
C
N
O
F
Ne
6.941
sodium
9.012182
magnesium
atomic mass
10.811
aluminium
12.0107
silicon
14.00674
phosphorus
15.9994
sulfur
18.9984
chlorine
20.1797
argon
11
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
22.98977
potassium
24.3050
calcium
scandium
titanium
vanadium
chromium
manganese
iron
cobalt
nickel
copper
zinc
26.981538
gallium
28.0855
germanium
30.97376
arsenic
32.065
selenium
35.453
bromine
39.984
krypton
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
Cr
Mn
K
Ca
Sc
Ti
V
39.0983
rubidium
40.078
strontium
44.95591
yttrium
47.867
zirconium
50.9415
niobium
37
38
39
40
41
42
43
51.9961
54.93805
molybdenum technetium
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
55.845
ruthenium
58.9332
rhodium
58.6934
palladium
63.546
silver
65.409
cadmium
69.723
indium
72.64
tin
74.9216
antimony
78.96
tellurium
79.904
iodine
83.798
xenon
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
85.4678
caesium
87.62
barium
88.90585
lutetium
91.225
hafnium
92.90638
tantalum
95.94
tungsten
[98]
rhenium
101.07
osmium
102.9055
iridium
106.42
platinum
107.8682
gold
112.411
mercury
114.818
thallium
118.710
lead
121.760
bismuth
127.60
polonium
126.9045
astatine
131.293
radon
55
56
71
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Lu
Hf
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
200.59
ununbium
204.3833
207.2
ununquadium
208.980
[209]
[210]
[222]
Cs
132.90545
francium
Ba
137.327
radium
174.967
178.49
lawrencium rutherfordium
Ta
W
Re
Os
180.9479
dubnium
183.84
seaborgium
186.207
bohrium
190.23
hassium
192.217
195.078
196.96655
meitnerium darmstadtium roentgenium
87
88
103
104
105
106
107
108
109
110
111
112
114
Fr
Ra
Lr
Rf
Db
Sg
Bh
Hs
Mt
Ds
Rg
Uub
Uuq
[223]
[226]
[262]
[261]
[262]
[266]
[264]
[269]
[268]
[271]
[272]
[285]
[289]
lanthanum
cerium
57
58
praseodymium neodymium promethium
59
samarium
europium
gadolinium
terbium
dysprosium
holmium
erbium
thulium
ytterbium
60
61
62
63
64
65
66
67
68
69
70
Dy
Ho
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
138.9055
actinium
140.116
thorium
140.90765
protactinium
144.24
uranium
[145]
neptunium
150.36
plutonium
151.964
americium
157.25
curium
158.9253
berkelium
89
90
91
92
93
94
95
96
97
162.50
164.930
californium einsteinium
Er
Tm
Yb
167.259
fermium
168.934
mendelevium
173.04
nobelium
98
99
100
101
102
Ac
Th
Pa
U
Np
Pu
Am
Cm
Bk
Cf
Es
Fm
Md
No
[227]
232.038
231.0359
238.0289
[237]
[244]
[243]
[247]
[247]
[251]
[252]
[257]
[258]
[259]
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The vertical columns in a Periodic Table are called ‘Groups’. Group numbers are written in capital
Roman numbers (I, II, III, IV, V, VI, VII and 0). Thus the Group IV elements consist of carbon, silicon,
germanium, tin and lead. The group number of an element can be used to deduce its valency.
The horizontal rows in the Periodic Table are called ‘periods’. Period numbers are written in HinduArabic numbers (1, 2, 3, etc.). Thus, the Period 2 elements consist of lithium, beryllium, boron,
carbon, nitrogen, oxygen, fluorine and neon.
The reactivity series
Metals show chemical properties which are different from those of non-metals. But even between
metals themselves the reactions can be different: some metals will react faster than others. When
metals are placed in order of decreasing reactivity, we obtain a reactivity series.
The order of reactivity of some metals, with decreasing reactivity, is as follows:
potassium
K
sodium
Na
calcium
Ca
magnesium
Mg
aluminium
Al
zinc
Zn
iron
Fe
lead
Pb
copper
Cu
mercury
Hg
silver
Ag
gold
Au
more reactive
less reactive
Mnemonic to remember reactivity series: “PSC MAZIL? He Can Make Solid Gold!”
Compounds and formulae
Experiment 3: Observing some common compounds.
A compound is a pure substance which contains two or more elements chemically combined
together in a fixed proportion.
A compound is made up of more than one type of atoms and each compound has a name and a
chemical formula. A formula is a way of representing a chemical compound using symbols for the
atoms present. For example, water is a compound and has formula H2O. Carbon dioxide is another
compound of formula CO2.
A compound does not retain the properties of the elements it contains. Thus, the compound
magnesium oxide is a white solid. Yet, the magnesium from which it is made is a grey, shiny metal.
The oxygen from which it is also made is a colourless, non-metallic gas.
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The diagram below represents a compound in the gaseous state. It is made up of two different
elements chemically combined together.
Atoms and molecules
An atom is the smallest, indivisible particle of an element that can take part in a chemical reaction.
An atom is itself made of smaller, subatomic particles called protons, neutrons and electrons. The
protons and neutrons are situated in the central part or the nucleus of an atom. Electrons move
round the nucleus in circular orbits and they are responsible for the valency and the chemical
properties of an element.
A molecule is the smallest part of a chemical compound that can take part in a chemical reaction.
A molecule is one of the fundamental units forming a chemical compound. Many molecules, like
those of water or carbon dioxide, consist of groups of atoms held together by special forces of
attraction called ‘covalent bonds’.
Classifying substances as elements, compounds or mixtures
Experiment 4: To separate a mixture of sand and water by filtration.
Substances can be classified as elements, compounds or mixtures.
A mixture is an impure substance which contains two or more elements or compounds in
proportions which may vary because they are not chemically combined together.
Examples of mixtures are tap water, petroleum, etc. These constituents can easily be removed or
separated by physical means like filtration, distillation, etc.
A mixture retains the original properties of each element present in it. For example, air will support
combustion because oxygen is present in it. Air will also turn limewater milky because it contains
carbon dioxide.
Note: Air is a mixture containing by volume 78% nitrogen, 21% oxygen, 1% noble gases (argon and
neon, which exist as single atoms), 0.034% carbon dioxide and variable amounts of water
vapour and pollutants.
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The diagram shown below represents a mixture in the gaseous state. This mixture contains two
different elements and one compound:
The table below summarises the differences between elements, compounds and mixtures:
Element
Compound
Mixture
Impure substance in
which components
(elements and/or
compounds) are mixed
in varying proportions.
Have a range of
melting and boiling
points.
Pure substance.
Pure substance in
which elements in it
are combined in a
fixed ratio.
Specific melting and
boiling point for a
given element.
Elements have a
chemical name and a
symbol, for example,
hydrogen (H) and zinc
(Zn).
Specific melting and
boiling points for a
given compound.
Compounds have a
name and a formula,
for example, water
(H2O) and carbon
dioxide (CO2).
Breaking down or
separation
Cannot be broken
down by chemical
means.
Can be broken down
by thermal
decomposition,
electrolysis or by
other chemical means.
Can be easily
separated into its
components by
physical means like
filtration, distillation,
chromatography, etc.
Examples
All metals and nonmetals in the Periodic
table.
Water (H2O), carbon
dioxide (CO2), kitchen
salt (NaCl), etc.
Air, tap water, blood,
etc.
Purity
Melting and boiling
points
Naming and
shorthand notation
Mixtures do not have
a specific naming
system nor a specific
shorthand notation.
Acids
Experiment 5: To demonstrate the effect of mineral acids on litmus, phenolphthalein and methyl
orange indicators.
An acid is a hydrogen containing compound which, when dissolved in water, can produce a
solution of pH less than 7. We can also define an acid as a compound, which turns damp blue litmus
paper red.
Examples of acids are hydrochloric acid, sulfuric acid, nitric acid, ethanoic acid and carbonic acid.
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Importance of selected acids
Hydrochloric acid is naturally present in the stomach and it helps to digest food. Ethanoic acid, which
is the acid present in vinegar, can be used to neutralise wasp stings. Sulfuric acid is present in car
batteries.
Bases and alkalis
Experiment 6: To demonstrate the effect of alkalis on litmus, phenolphthalein and methyl orange
indicators.
A base is a compound, which reacts with an acid to form a salt and water only. Examples of bases
are magnesium oxide and copper (II) oxide, which are oxides of metals.
If a base is readily soluble in water, we call it an alkali. An alkali is defined as a compound, which
produces an aqueous solution of pH greater than 7. We can also define an alkali as a compound,
which turns red litmus paper blue.
Examples of alkalis are potassium hydroxide, sodium hydroxide, calcium hydroxide and aqueous
ammonia.
Note: If a solution is neither acidic nor alkaline we call it a neutral solution. Neutral solutions have
pH=7.
Importance of selected bases and alkalis
Toothpastes contain bases which neutralise the acids released by bacteria in the mouth, which
would otherwise cause tooth decay. Magnesium hydroxide is used as an antacid to neutralise excess
acidity in the stomach.
Indicators
An indicator is a substance which has one colour in very acidic solutions and another colour in very
alkaline solutions.
Colour in strongly
Colour in neutral
Colour in strongly
Name of indicator
acidic solution
solution
alkaline solution
litmus
red
purple
blue
phenolphthalein
colourless
colourless
pink
methyl orange
red
orange
yellow
The pH (potentiometric hydrogen ion concentration) scale
The pH scale, which measures the acidity or alkalinity of a solution, ranges from 0 (very acidic) to 14
(very alkaline). A pH meter can be used to measure the pH of a given solution accurately. The pH
indicator, also known as the Universal Indicator, is a mixture of several indicators and it can be used
to estimate the pH or acidity of a solution. It is normally red in strongly acidic solutions (pH=0 to 2),
yellowish-green in neutral solutions (pH=7) and violet in strongly alkaline solutions (pH=12 to 14).
0
1
2
3
acidic
4
5
6
7
neutral
8
9
← more acidic
10
11
12
13
14
alkaline
more alkaline →
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Acids have pH values less than 7. A strong acid is a solution that has a very low pH of 0 to 2.
Examples of strong acids are hydrochloric acid, sulfuric acid and nitric acid. A weak acid is a solution
that has a pH of 3 to 6. Examples of weak acids are ethanoic acid and carbonic acid.
Alkalis have pH values greater than 7. A strong alkali is a solution that has a very high pH of 12 to 14.
Examples of strong alkalis are potassium hydroxide and sodium hydroxide. A weak alkali is a solution
that has a pH of 8 to 11. Examples of weak alkalis are calcium hydroxide and aqueous ammonia.
Neutral solutions have pH value equal to 7.
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Ranges from 0 (very
acidic) to 14 (very
alkaline)
Substances which
have one colour in
very acidic solutions
and another colour in
very alkaline
solutions
Measures the acidity
or alkalinity of a
solution
made up of only one
type of atom
has a chemical
symbol, e.g Zn
a pure substance that
cannot be broken
down into simpler
substances by any
ordinary chemical
process
Estimated using
Universal or pH
indicator
As solids (iron,
carbon), liquids
(mercury, bromine)
and gases (nitrogen,
oxygen)
has a name e.g zinc
Measured accurately
using a pH meter
shorthand way of
writing name of
element
by physical state
by metallic nature
Indicators
Classifying
elements
Elements and
Symbols
By placing them in
Groups and Periods in
a Periodic Table
Smallest, indivisible
particle of an element
that can take part in a
chemical reaction
definition of element
Acids have pH range 0
to 6
As metals
(aluminium, copper)
and non-metals
(hydrogen, chlorine)
definition of atom
Alkalis have pH range
8 to 14
Smallest part of a
chemical compound
that can take part in a
chemical reaction
definition of molecule
The pH scale
Neutral solutions like
water have pH=7
Atoms and
molecules
Pure substance
containing two or
more elements
chemically combined
together in a fixed
proportion
Chemical Substances
Compound which
reacts with an acid to
form a salt and water
only
Compound producing
aqueous solution of
pH>7
Alkali is a base which
dissolves readily in
water
definition of compound
Base
definition
Compounds and
formulae
Base and alkali
Compounds have a
name and a chemical
formula
Mixtures
Acid
Alkali
definition
Compounds are
different from the
elements they
contain
definition
definition
Hydrogen-containing
compound producing
aqueous solution of
pH<7
use of symbols to
represent atoms in a
compound
definition
Turns damp litmus
paper blue
examples
separation
examples
Turns damp litmus
paper red
chemical properties
Impure substance
which contains two or
more elements or
compounds in
proportions which
may vary
Turns
phenolphthalein pink
Turns methyl orange
yellow
Made up of more
than one type of
atoms
potassium hydroxide,
sodium hydroxide,
calcium hydroxide,
aqueous ammonia
Turns
phenolphthalein
colourless
Turns methyl orange
red
hydrochloric acid,
sulphuric acid, nitric
acid
Constituents can be
separated by physical
means
Mixtures exhibit
chemical properties
of each constituent
present
Constituents not
chemically combined
together
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CHAPTER TWO: THE LANGUAGE OF CHEMISTRY
Physical and chemical changes
Experiment 7: To investigate whether the filament of a bulb changes mass when an electric current
is passed through it.
Experiment 8: To compare the mass of the magnesium ribbon to the mass of magnesium oxide that
it subsequently produced by burning it.
Experiment 9: Action of water on a small piece of sodium.
The following table summarises the main differences between a physical and a chemical change:
Physical change
Chemical change
Produces no new substance. For example,
melting an ice cube will not change the mass of
the water.
Is generally reversible. For example, the steam
that we obtain from boiling water can be
cooled down to obtain the liquid water back.
Is not accompanied by energy changes, except
for those involving changes of states. When a
piece of iron is magnetised it does not absorb
or give out heat or light.
Produces no change in mass. Thus, if we heat a
filament by passing an electric current through
it, the filament does not become heavier or
lighter.
Always produces a new substance. For
example, burning a piece of magnesium (grey,
shiny solid) in oxygen (colourless gas) will
produce magnesium oxide (white solid).
Is generally irreversible. When iron rusts in the
presence of air and moisture, the rust formed
cannot be easily converted back into iron. This
is why rusting is considered as a wasteful
process.
Is accompanied by considerable heat changes.
If a small piece of sodium is dropped into water
there is so much heat evolved that a flame is
observed.
Produces substances whose masses are
different from those of the original substances.
Magnesium oxide will be heavier than the piece
of magnesium which originally produced it by
burning.
Shorthand representation of names of elements
Elements have a name and a chemical symbol. When we write the first letter in the symbol of an
element, we use a capital letter. For example, the symbol of carbon is C.
The symbols of many elements also contain a second letter. In this case, the second letter is written
as small letters, keeping its first letter a capital letter. For example, the symbol of chlorine is Cl.
Notice how the letter ‘l’ is written so that the examiner does not think it is the capital letter ‘I’.
The symbol of iodine should also be written as ‘I’ and not as ‘I’ to avoid confusion. The symbol of
aluminium should be written as ‘Al’ and not as ‘Al’.
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Shorthand representation of names of compounds
Compounds have a symbol and a chemical formula. The formula is not only a shorthand
representation of its name; it even tells us how many atoms are chemically combined together to
form the compound. Thus, the formula of water is H2O, meaning that a water molecule is formed
when two atoms of hydrogen and one atom of oxygen are chemically combined together. Similarly,
the formula of carbon dioxide is CO2, meaning that carbon dioxide is made when one atom of carbon
chemically combines with two atoms of oxygen.
*Deducing
the formulae of simple compounds using symbols and valencies
The valency of a chemical species is its combining power with other chemical species. It helps us to
determine the chemical formula of a compound. The table below gives the symbols and the
valencies of selected elements.
Symbol
Valency
aluminium
Al
3
Non-Metal
argon
Symbol
Valency
Ar
0
barium
Ba
2
bromine
Br
1
calcium
Ca
2
carbon
C
2 or 4
copper
Cu
1 or 2
chlorine
Cl
1
gold
Au
1 or 3
fluorine
F
1
iron
Fe
2 or 3
helium
He
0
lead
Pb
2 or 4
hydrogen
H
1
magnesium
Mg
2
iodine
I
1
mercury
Hg
1 or 2
neon
Ne
0
potassium
K
1
nitrogen
N
3
silver
Ag
1
oxygen
O
2
sodium
Na
1
phosphorus
P
3 or 5
tin
Sn
2 or 4
silicon
Si
4
zinc
Zn
2
Sulfur or sulphur
S
2, 4 or 6
Metal
To write down the formula of a simple compound, containing only two elements, proceed as follows:
1.
2.
3.
4.
Recall and write down the symbols of the two elements it contains side by side.
Recall and write down the valency of each element below the symbol of that element.
Cross-arrange the symbol of one element with the valency of the other element.
Simplify the formula of the compound to its lowest terms, if necessary.
Worked example:
Write down the formula of the compound aluminium oxide.
Aluminium has symbol Al and its valency is 3. Oxygen has symbol O and its valency is 2.
Al
O
3
2
The formula of aluminium oxide is Al2O3.
The above formula contains 2 aluminium atoms and 3 oxygen atoms. The total number of atoms in it
is 2+3=5.
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A radical is a group of atoms that has a valency left over, although the atoms in it are chemically
combined together. If we know the formula and the valency of a radical, we can deduce the formula
of a compound containing this radical. The following table gives the names, symbols and valencies of
selected radicals:
Name
Formula
Valency
ammonium
NH4
1
carbonate
CO3
2
hydroxide
OH
1
nitrate
NO3
1
sulfate or sulphate
SO4
2
Worked example:
Deduce the formula of calcium nitrate.
Calcium has symbol Ca and its valency is 2. The nitrate radical has formula NO3 and its valency is 1.
Ca
NO3
2
1
The formula of calcium nitrate is Ca(NO3)2. If we do not put brackets in the formula of the
compound, someone will think that there are 32 oxygen atoms in the formula when, in fact, there
are only 6!!!
The above formula contains 1 calcium atom, (1x2) or 2 nitrogen atoms and (2x3) or 6 oxygen atoms.
The total number of atoms in it is 1+2+6=9.
The table below gives the names and formulae of selected compounds whose formulae cannot be
deduced by the ‘cross-arrangement’ method described above:
Compound
Formula
alcohol (ethanol)
C2H5OH
ammonia
NH3
carbon dioxide
CO2
carbon monoxide
CO
carbonic acid
H2CO3
dinitrogen oxide (laughing gas)
N2O
ethanoic acid (vinegar)
CH3COOH
glucose (cane sugar)
C6H12O6
hydrochloric acid
HCl
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Compound
Formula
hydrogen bromide
HBr
hydrogen chloride
HCl
hydrogen fluoride
HF
hydrogen iodide
HI
hydrogen peroxide
H2O2
hydrogen sulfide
H2S
methane
CH4
nitric acid
HNO3
nitric oxide (nitrogen monoxide)
NO
nitrogen dioxide
NO2
nitrous acid
HNO2
steam or water or ice
H2O
sucrose
C12H22O11
sulfur dioxide
SO2
sulfur trioxide
SO3
sulfuric acid
H2SO4
sulfurous acid
H2SO3
Worked example:
What is the total number of atoms in a molecule of sulfuric acid?
Sulfuric acid has formula H2SO4
The total number of atoms in it = 2 ‘hydrogens’ + 1 ‘sulfur’ + 4 ‘oxygens’ = 7←
Chemical reactions, reactants and products
A chemical reaction is a change in which one or more chemical elements or compounds (the
reactants) form new compounds (the products).
A chemical reaction involves a chemical change during which there is a rearrangement of atoms to
form new substances. It is this rearrangement of atoms that causes energy to be absorbed or given
out.
Reactants are those original substances whose atoms have not yet combined. Products are those
new substances which are formed as a result of a chemical change.
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Word equations and chemical equations
A word equation is a conventional way of representing a chemical reaction by writing the names of
the reactants on the left and the names of the products on the right. An arrow is used to indicate the
direction of the reaction.
A chemical or molecular equation is written in by using symbols for elements and by using formulae
for compounds (which are assumed to be made up of molecules). But we need to balance the
chemical equation to make it chemically correct. We may also write the state symbols to indicate
whether each reactant or product was a solid (s), a liquid (l), a gas (g) or dissolved in water (aq).
To write down a balanced chemical reaction, proceed as follows:
1. Make sure that the reaction does occur!!!
2. Write down the word equation.
3. Recall, deduce and write the symbols or formulae of all the reactants and the products. (This
has already been explained above).
4. Balance the equation by putting numbers in front. We should not change the formula of a
compound to balance the equation. Check, by counting, if all the particles are balanced.
Worked example:
Write down balanced equations for the following reactions:
(a) Magnesium burns in oxygen to form magnesium oxide.
Word equation: magnesium + oxygen ⟶ magnesium oxide
Molecular equation: 2 Mg +
O2 ⟶
2 MgO
(b) Zinc reacts in dilute hydrochloric acid to form zinc chloride and hydrogen gas.
Word equation:
zinc + hydrochloric acid ⟶ zinc chloride + hydrogen
Molecular equation: Zn
+
2 HCl
⟶
ZnCl2
+
H2
(c) When a solution of silver nitrate is mixed with a solution of calcium chloride, the products
are silver chloride and calcium nitrate.
Word equation: silver nitrate + calcium chloride ⟶ silver chloride + calcium nitrate
Molecular equation: 2 AgNO3 +
CaCl2
⟶
2 AgCl
+
Ca(NO3)2
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Physical change
Produces no new
substance
Always produces a
new substance
Is generally reversible
Is generally
irreversible
Is not accompanied
by energy changes
Is accompanied by
considerable heat
changes
Produces no change
in mass
Produces substances
whose masses are
different from those
of the original
substances
Chemical change
The language of
chemistry
Recall and write
down the symbols of
the two elements it
contains side by side
Make sure that the
reaction does occur!!!
Step 1
To write the
formula of a
compound
Before starting
Write down the word
equation
Step 1
To write a
balanced
chemical
equation
definition of radical
Step 2
Recall, deduce and
write the symbols
and/or formulae of all
the reactants and the
products
group of atoms with
valency left over
Step 2
Step 3
Step 4
Step 3
to use it to determine formula of a comound
Balance the equation
by putting numbers in
front
Recall and write
down the valency of
each element below
the symbol of that
element
write down valency of
radical below its
formula
Cross-arrange the
symbol of one
element with the
valency of the other
element
Simplify the formula
of the compound to
its lowest terms, if
necessary
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CHAPTER THREE: CHEMICAL REACTIONS IN GENERAL
Simple equipment and glassware used in chemistry (Optional)
Experiment 10: To demonstrate some common laboratory apparatus.
The following pieces of apparatus are often used in chemistry:
test-tube
beaker
measuring cylinder
pipette
round-bottomed flask
flat-bottomed flask
conical flask
burette
funnel
evaporating dish
A test-tube is used to hold small amounts of liquids which are to be analysed. Boiling tubes have the
same shape but they are slightly larger and they can withstand higher temperatures. Boiling tubes
are used to heat solids or liquids during analysis.
A beaker is used to hold large volumes of liquids but it cannot be used for accurate measurement.
Beakers exist in different sizes.
A conical flask can be used during a volumetric analysis. A volumetric analysis can be used, for
example, to determine the exact concentration of an acid or of an alkali. Conical flasks also exist in
different sizes.
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A measuring cylinder is more accurate than a beaker or a conical flask. Measuring cylinders also exist
in different sizes. The zero mark of a measuring cylinder is situated at its bottom.
A pipette is the most accurate piece of apparatus for measuring volumes of liquids. It is also used in
volumetric analysis. However, it can only hold a fixed volume of liquid. For example, a 25 cm3 pipette
will measure 25.00 cm3 very accurately but it cannot be used to measure a volume like 22.13 cm3.
A burette can measure any volume of liquid ranging from 0.0 to 50.0 cm3. It is also used in
volumetric analysis. When a burette tap is opened, we can run the desired volume of liquid into
another container, usually a conical flask. This is why the zero mark of this instrument is at the top.
Both round-bottomed and flat-bottomed flasks can be used to prepare gases in the laboratory. A
round-bottomed flask is preferred when the reactants need to be heated.
A funnel is used for filtering and to fill a burette with a liquid. Funnels are made of glass or plastic.
An evaporating dish is made of porcelain or similar material. It can be used to dry a wet solid by
evaporation. An evaporating dish can also be used to prepare crystals of a certain salt.
Making different compounds by chemical reactions (Optional)
Compounds can usually be made in one of the following ways:
1. By direct combination of their constituent elements, for example, sodium chloride (common
salt) can be made by mixing sodium and chlorine directly:
sodium + chlorine → sodium chloride
2 Na + Cl2
→
2 NaCl
2. By the reaction of an element with another compound, for example, the compound sodium
hydroxide is made by adding the element sodium to the compound water:
sodium + water → sodium hydroxide + hydrogen
2 Na + 2 H2O →
2 NaOH
+
H2
3. By the reaction of a compound with other compounds, for example, the compound
potassium chloride can be made by mixing hydrochloric acid with potassium hydroxide:
hydrochloric acid + potassium hydroxide → potassium chloride + water
HCl
+
KOH
→
KCl
+ H2O
4. By decomposing other compounds by heat (thermal decomposition), for example, the
compound calcium oxide (lime) can be made by heating the compound calcium carbonate
(limestone) strongly:
heat
calcium carbonate →
CaCO3
∆
→
calcium oxide + carbon dioxide
CaO
+
CO2
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The reactivity series
Metals show chemical properties which are different from those of non-metals. But even between
metals themselves the reactions can be different: some metals will react faster than others. When
metals are placed in order of decreasing reactivity, we obtain a reactivity series.
The order of reactivity of some metals, with decreasing reactivity, is as follows:
potassium
K
sodium
Na
calcium
Ca
magnesium
Mg
aluminium
Al
zinc
Zn
iron
Fe
lead
Pb
(hydrogen)
(H)
copper
Cu
mercury
Hg
silver
Ag
gold
Au
more reactive
less reactive
Mnemonic to remember reactivity series: “PSC MAZIL? He Can Make Solid Gold!”
Note: Although hydrogen is a non-metal, it is placed in the reactivity series as a reference, for
predicting certain reactions.
The reactions of some metals with the oxygen of the air
Experiment 11: Heating a magnesium ribbon strongly in air.
Experiment 12: Heating iron wool strongly in air.
Most metals will react with the oxygen of the air according to the general equation:
metal + oxygen → metal oxide
1. Very reactive metals like potassium or sodium become easily tarnished (covered with a dull
oxide layer) in air:
potassium + oxygen → potassium oxide
4K
+ O2 →
2 K2O
sodium + oxygen → sodium oxide
4 Na + O2 →
2 Na2O
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2. Other reactive metals have to be heated in air to produce their oxides. A magnesium ribbon
will burn in the oxygen of the air with a bright flame to produce magnesium oxide which is a
white powder:
heat
magnesium + oxygen → magnesium oxide
2 Mg
+ O2 →
2 MgO
Iron wool will burn in air to produce iron oxide which is a brown solid:
heat
iron + oxygen → iron (III) oxide
4 Fe + 3 O2 →
2 Fe2O3
3. Less reactive metals like copper do not burn when heated in air but on strong heating they
become coated with an oxide layer. With copper, a black oxide layer is formed:
heat
copper + oxygen →
copper (II) oxide
∆
2 Cu + O2 →
2 CuO
4. Unreactive metals like silver and gold do not react with the oxygen of the air even when they
are strongly heated.
The reactions of some metals with water
Experiment 13: The reaction of calcium with water.
1. Potassium and sodium react vigorously with cold water to form their hydroxides, with the
evolution of hydrogen gas and large amounts of energy:
potassium + water → potassium hydroxide + hydrogen
2 K + 2 H2O →
2 KOH
+
H2
2. Calcium also reacts with cold water, although less vigorously:
calcium + water → calcium hydroxide + hydrogen
Ca + 2 H2O →
Ca(OH)2
+
H2
3. Magnesium, zinc and iron do not react with cold water but with steam to form their
corresponding metal oxides and hydrogen gas:
heat
magnesium + water →
Mg
magnesium oxide + hydrogen
Δ
+ H2O →
heat
iron + water →
MgO
+
H2
triirontetraoxide + hydrogen
Δ
3 Fe + 4 H2O → Fe3O4 + 4 H2
Note: Aluminium has an inert, non-porous oxide layer on its surface which prevents
reaction.
4. Lead shows very slight reaction with the steam in a Bunsen flame which is at a very high
temperature.
5. Copper, mercury, silver and gold do not react with water nor steam, even under extreme
conditions.
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The reactions of some metals with dilute mineral acids
Experiment 14: To observe the reaction between zinc and dilute sulfuric acid.
Theoretically, all metals above hydrogen in the reactivity series react with dilute acids to give a salt
of the acid and hydrogen gas:
metal + dilute acid → salt + hydrogen
1. The reaction of potassium and sodium with a dilute acid is explosive.
2. Other metals like calcium, magnesium, zinc and iron will produce their corresponding salts,
liberating the colourless and odourless gas hydrogen.
magnesium + dilute hydrochloric acid → magnesium chloride + hydrogen
Mg
+
2 HCl
→
MgCl2
+
H2
zinc + dilute sulfuric acid → zinc sulfate + hydrogen
Zn +
H2SO4
→
ZnSO4 +
H2
iron + dilute nitric acid → iron (II) nitrate + hydrogen
Fe +
2 HNO3 → Fe(NO3)2
+
H2
3. The reaction between lead and dilute acids is too slow to be observed.
4. Metals like copper, mercury, silver and gold do not react with dilute acids.
Displacement reactions
A metal higher in the reactivity series will displace a metal which is lower in the series, from a
solution of its salt.
reactive metal + salt of less reactive metal ⟶ salt of reactive metal + less reactive metal
Displacement reactions take place because a more reactive metal will take the place of a less
reactive one to form new and more stable compound in which the chemical bonds are stronger.
zinc
zinc
blue copper
(II) sulfate
solution
after one
hour
colour of blue
solution
gradually fades
reddish-brown
shiny deposit of
copper appears
Zinc will displace copper from a solution of copper (II) sulfate, as shown above:
zinc + copper (II) sulfate → zinc sulfate + copper
Zn +
CuSO4
→
ZnSO4
+ Cu
The blue colour of the solution slowly disappears, while a reddish-brown, shiny deposit of copper
appears. If the zinc is replaced by iron, the reaction is similar except that the solution turns green.
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Action of heat on metallic carbonates
1. Sodium carbonate and potassium carbonate will only melt, but will not decompose on
strong heating.
2. All other metallic carbonates decompose on heating to produce their corresponding oxides
and carbon dioxide gas:
heat
unstable metallic carbonate →
metallic oxide + carbon dioxide
Δ
CaCO3(s) → CaO (s) + CO2(g)
Summary of reactivity series of metals
REACTION WITH
REACTION WITH
ELEMENT
AIR
WATER
potassium (K)
easily coated with
hydrogen
an oxide layer
liberated and
sodium (Na)
(tarnished) under
hydroxides
normal conditions
formed readily
calcium (Ca)
magnesium (Mg)
with steam
pure metal
aluminium (Al)
hydrogen formed,
produces oxide
zinc (Zn)
leaving oxides of
on heating
metals
iron (Fe)
very strong
lead (Pb)
heating is
HYDROGEN (H)
required to
produce metal
copper (Cu)
no reaction
oxide
mercury (Hg)
no reaction even
silver (Ag)
on
strong heating
gold (Au)
REACTION WITH
DILUTE ACIDS
hydrogen and salt
of acid formed
violently
hydrogen and salt
of acid formed
no visible
reaction
no reaction
METALLIC
CARBONATES
Not
decomposed
by strong
heating
Decomposed
on heating to
produce their
corresponding
oxides and
carbon dioxide
gas
Note: An alloy is a mixture consisting of two or more metals (e.g. bronze is an alloy containing
copper and tin) or a metal and a non-metal (e.g. stainless steel is an alloy of iron, carbon,
chromium and nickel).
Laboratory preparation of hydrogen
Experiment 15: Laboratory preparation, collection and test for hydrogen gas.
Hydrogen gas is prepared by adding a dilute acid like hydrochloric acid or sulfuric acid to a metal like
magnesium or zinc:
magnesium + hydrochloric acid → magnesium chloride + hydrogen
Mg
+
2 HCl
→
MgCl2
+
H2
zinc + sulfuric acid → zinc sulfate + hydrogen
Zn +
H2SO4
→
ZnSO4
+
H2
The apparatus shown below is used to prepare and collect hydrogen gas in the laboratory. The
hydrogen is collected over water by upward delivery (downwards displacement of water).
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sulfuric acid
tap
funnel
delivery
tube
gas
jar
hydrogen
conical
flask
zinc
beehive
shelf
glass
trough
water
When the sulfuric acid (colourless liquid) comes into contact with the zinc (grey, shiny solid) there is
effervescence and a colourless, odourless gas that burns in air with a pop sound is evolved (the gas is
hydrogen). A colourless solution containing zinc sulfate is left behind.
Laboratory preparation of oxygen
Experiment 16: Laboratory preparation, collection and test for oxygen gas.
Hydrogen peroxide is a colourless liquid that slowly decomposes into water, liberating oxygen gas.
The reaction can be represented by the equations shown below:
hydrogen peroxide ⟶ water + oxygen
2 H2O2
→ 2 H2O + O2
This reaction can be made to occur faster by adding a small amount of manganese (IV) oxide to the
hydrogen peroxide. The manganese (IV) oxide does not itself react but it only makes the reaction
occur faster: it acts as a catalyst for this reaction. The manganese (IV) oxide does not become a new
substance at the end of the reaction.
tap
funnel
hydrogen
peroxide
delivery
tube
gas
jar
oxygen
conical
flask
manganese (IV)
oxide
beehive
shelf
glass
trough
water
The apparatus shown above is used to prepare and collect oxygen gas in the laboratory. The oxygen
is collected over water by upward delivery (downwards displacement of water).
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When the colourless hydrogen peroxide is run into the conical flask, there is effervescence and a
colourless and odourless gas that rekindles a glowing splinter of wood is produced (the gas is
oxygen). The manganese (IV) oxide acts as a catalyst in this reaction causing the reaction to occur
faster: its chemical composition remains exactly the same at the end of the reaction.
Laboratory preparation of carbon dioxide
Experiment 17: Laboratory preparation and collection of carbon dioxide gas.
Carbon dioxide gas is prepared by adding a dilute acid like hydrochloric acid or nitric acid to a metal
carbonate like calcium carbonate or sodium carbonate:
calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
CaCO3
+
2 HCl
→
CaCl2
+ H2O +
CO2
sodium carbonate + nitric acid → sodium nitrate + water + carbon dioxide
Na2CO3
+ 2 HNO3 →
2 NaNO3 + H2O +
CO2
tap
funnel
hydrochloric
acid
delivery
tube
gas
jar
carbon dioxide
conical
flask
calcium
carbonate
beehive
shelf
glass
trough
water
The apparatus shown above is used to prepare and collect carbon dioxide gas in the laboratory. The
carbon dioxide is collected over water by upward delivery (downwards displacement of water).
The green solid copper (II) carbonate can also be added to dilute sulfuric acid. There is effervescence
and the green solid becomes smaller with the evolution of a colourless and odourless gas that turns
lime water milky (the gas is carbon dioxide). A blue solution containing copper (II) sulfate is left
behind:
copper (II) carbonate + sulfuric acid → copper (II) sulfate + water + carbon dioxide
CuCO3
+
H2SO4
→
CuSO4
+ H2O +
CO2
The rusting of iron and its prevention
Experiment 18: To show that water and air are necessary for rusting to occur.
Rusting is the formation of a hydrated form of iron (III) oxide. Water (moisture) and air are essential
for rusting to occur. Rust is a dull brown solid which falls off the surface of iron or steel, exposing
new layers to this wasteful process.
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Prevention of rusting can be achieved by placing a barrier around the metal. This can be done by
painting, oiling, greasing, plastic-coating, galvanising (covering with a layer of zinc) or by alloying.
Zinc and magnesium are higher than iron in the reactivity series. If blocks of these metals are
attached to underground steel pipes or iron surfaces in a ship’s body, they corrode in preference to
the iron, protecting the latter. This is known as the sacrificial protection of iron.
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Pipettes are very
accurate but measure
only fix volumes
Colourless, odourless
gas
Slightly soluble in
water (soft drinks)
Measuring cylinders
used to measure
volumes of liquids
Prepared by the
action of dilute acids
on carbonates
Turns lime water
milky
Burettes have a tap at
the bottom and their
zero marks are at the
top
Laboratory
equipment and
glassware
preparation
Beakers hold large
volumes of liquids
Carbon dioxide
By reaction of an
element with another
compound, e.g
sodium+water→sodiu
m
hydroxide+hydrogen
Flat-bottomed flasks
are used to contain
reaction mixtures.
Round-bottomed
flasks withstand heat
Conical flasks used in
volumetric analyses
test
By direct combination
of elements, e.g
sodium+chlorine→so
dium chloride
Test-tubes used for
analysis. Boiling tubes
withstand heat
By thermal
decomposition, e.g
calcium carbonate
→calcium
oxide+carbon dioxide
Funnels used for
filtering and filling
burettes
Metals placed in
order of decreasing
reactivity
How compounds
are made
Evaporating dishes
used for drying
powdered solids and
to prepare crystals
definition
Colourless, odourless
gas
metals in the series
Slightly soluble in
water (making
aquatic life possible)
Prepared by
decomposing
hydrogen peroxide by
manganese (IV) oxide
catalyst
The reactivity
series
Chemical reactions in
general
Oxygen
By reaction of a
compound with other
compounds, e.g
hydrochloric
acid+potassium
hydroxide→potassiu
m chloride+water
potassium,sodium,cal
cium,magnesium,alu
minium,zinc,iron,lead
,(hydrogen),copper,m
ercury,silver,gold
mnemonic
PSCMAZIL Can Make
Solid Gold!!!
preparation
test
Reactions of
metals
Relights a glowing
splinter of wood
Displacement
reactions
Hydrogen
Hydrated form of iron
(III) oxide
composition
A metal higher in the
reactivity series will
displace a metal
which is lower in the
series, from a
solution of its salt
preparation
test
Dull, reddish-brown
solid that falls off iron
surfaces
conditions
Prepared by action of
dilute acids on
reactive metals
Burns in air with a
pop sound
Water and oxygen of
the air needed for
rusting to occur
Wasteful process
With dilute mineral
acids
Heating of metallic
carbonates
prevention
Insoluble in water
Aluminium, zinc, iron
produce their oxides
on heating
Rusting
Colourless, odourless
gas
Lighter than air
With oxygen of the
air
potassium and
sodium carbonates
do not decompose
Prevented by
painting, oiling,
greasing, plasticcoating, galvanising,
alloying
remaining carbonates
decompose to their
oxides, liberating
carbon dioxide
Potassium, sodium,
calcium easily
tarnished under room
conditions
With water
Lead, copper,
mercury produce
their oxides on strong
heating
Silver and gold do not
react even on strong
heating
Potassium, sodium
react violently
Calcium, magnesium
aluminium, zinc, iron
liberate hydrogen
Lead reacts too
slowly
Copper, mercury,
silver, gold do not
react
Potassium, sodium,
calcium liberate
hydrogen with cold
water
Zinc and iron liberate
hydrogen with steam
Lead, copper,
mercury, silver and
gold do not react
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CHAPTER FOUR: IMPORTANT CHEMICAL REACTIONS
Characteristic reactions of acids
Experiment 19: To show the reaction between calcium carbonate and nitric acid.
1. Acids turn litmus paper red. They will also turn phenolphthalein indicator colourless and turn
methyl orange indicator red.
2. Acids react with bases and alkalis to produce a salt and water only. This type of reaction is
called a neutralisation.
acid + base or alkali ⟶ salt + water
For example, hydrochloric acid will react with sodium hydroxide to form sodium chloride
and water. The word equation for this reaction is:
hydrochloric acid + sodium hydroxide ⟶ sodium chloride + water
HCl
+
NaOH
→
NaCl
+ H2O
3. With reactive metals, acids produce a salt and hydrogen gas.
acid + reactive metal → salt + hydrogen
For example, zinc will react with sulfuric acid to produce zinc sulfate and hydrogen gas:
zinc + sulfuric acid → zinc sulfate + hydrogen
Zn +
H2SO4
→
ZnSO4 +
H2
4. Acids react with carbonates to produce a salt, water and carbon dioxide gas.
acid + carbonate → salt + water + carbon dioxide
For example, nitric acid will react with calcium carbonate to form calcium nitrate, water and
carbon dioxide gas:
nitric acid + calcium carbonate → calcium nitrate + water + carbon dioxide
2 HNO3 +
CaCO3
→
Ca(NO3)2 + H2O +
CO2
Characteristic reactions of bases and alkalis
Experiment 20: To demonstrate the reaction between sodium hydroxide and ammonium chloride.
1. Alkalis turn litmus paper blue. They will also turn phenolphthalein indicator pink and methyl
orange indicator yellow.
2. Bases react with acids to form a salt and water only:
copper (II) oxide + sulfuric acid →copper (II) sulfate + water
CuO
+
H2SO4 ⟶
CuSO4
+ H2O
3. Bases react with ammonium salts, on warming, to form a salt, water and ammonia gas:
∆
base + ammonium salt → salt + water + ammonia
∆
NaOH + NH4Cl → NaCl + H2O + NH3
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Neutralisation reactions and their applications
Neutralisation is the reaction between an acid and an alkali to produce a salt and water only. Thus,
the right amounts of nitric acid and sodium hydroxide will neutralise each other to produce sodium
nitrate and water:
nitric acid + sodium hydroxide → sodium nitrate + water
HNO3 +
NaOH
→
NaNO3
+ H2O
Hydrochloric acid and calcium hydroxide will also neutralise each other to produce calcium chloride
and water:
hydrochloric acid + calcium hydroxide → calcium chloride + water
2 HCl
+
Ca(OH)2
→
CaCl2
+ H2O
Neutralisation is also used in everyday life to:
1. neutralise excess acidity in the digestive system with antacids like calcium carbonate,
magnesium hydroxide (magnesia) or sodium hydrogencarbonate (also known as sodium
bicarbonate and found in baking soda):
sodium hydrogencarbonate + hydrochloric acid → sodium chloride + water + carbon dioxide
NaHCO3
+
HCl
→
NaCl
+ H2O +
CO2
2. neutralise bee stings with weak alkalis like sodium bicarbonate (baking soda),
3. neutralise wasp stings with weak acids like ethanoic acid (vinegar),
4. neutralise the excess acid in the soil with lime (calcium oxide). High soil acidity is often a
consequence of acid rains,
5. neutralise the oxides emitted from factory chimneys and which are responsible for acid
rains. Calcium carbonate is used to reduce the effect of acid rain,
6. neutralise the acids released by bacteria in the mouth with toothpastes.
Salts
Salts are compounds formed when an acid neutralises a base or an alkali. A salt is formed when the
hydrogen atoms of an acid are partly or completely replaced by a metal or by an ammonium radical.
The first name of a salt tells us from which metal or metallic compound they have been prepared:
for example, sodium chloride (found in common salt or kitchen salt) contains sodium. The second
name of a salt tells us from which acid it has been prepared: for example, calcium sulfate (found in
‘plaster of Paris’) has been prepared from sulfuric acid.
Some acids contain only one hydrogen atom per molecule and they are said to be monobasic acids.
Hydrochloric acid (HCl) and nitric acid (HNO3) are examples of monobasic acids. Other acids may
contain more than one hydrogen atoms per molecule and they are said to be polybasic acids.
Sulfuric acid (H2SO4) is an example of a polybasic acid. An acid salt is the salt of a polybasic acid in
which not all the hydrogen atoms have been replaced by metals. Examples of acid salts are sodium
hydrogen carbonate (NaHCO3) and sodium hydrogensulfate (NaHSO4).
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Solubility of salts
The method of preparing a particular salt will depend on its solubility in water. The following table
summarises the solubilities of different salts in water:
SOLUBLE SALTS
All nitrates
All potassium, sodium and ammonium salts
All chlorides
All bromides
All iodides
All sulfates
Except potassium, sodium and ammonium
carbonates.
INSOLUBLE SALTS
—
—
except silver and lead chlorides
except silver and lead bromides
except silver and lead iodides
except barium, lead and calcium sulfates
All carbonates
Preparation of soluble salts
1. By titration (Optional)
This method can be used for soluble salts that can be prepared by neutralisation. The acid
may be run from a burette to neutralise the alkali which has been pipetted into a conical
flask. The salt is then obtained from the solution by crystallisation.
exact
amount
of acid
exact
amount
of alkali
titration
determine
with
indicator end point
repeat without
indicator
solution of salt
(filtrate)
heat gently
concentrated
(saturated) solution
allow to cool
crystals of the salt +
uncrystallised
solution
filter and rinse with
distilled
water
pure crystals of the
salt
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2. By the action of an acid on a metal (e.g. zinc sulfate from zinc and dilute sulfuric
acid).
This method can be used for metals above hydrogen in the reactivity series (except
potassium and sodium) with dilute acids.
filter to remove
add excess metal
dilute
excess metal +
solution of salt
unreacted metal
till reaction stops
solution of salt
(residue)
acid
(filtrate)
heat gently
concentrated
(saturated) solution
allow to cool
crystals of the salt +
uncrystallised
solution
filter and rinse with
distilled
water
pure crystals of the
salt
The crystals obtained are usually dried between folds of filter paper. If we heat the hydrated
crystals strongly they will decompose to form the anhydrous salt.
3. By the action of an acid on an insoluble base or carbonate (e.g. copper (II) sulfate
from copper (II) oxide and dilute sulfuric acid)
This general method is suitable for metals which do not react directly with acids.
dilute
acid
add excess
base/carbonate
and warm
filter to remove excess
excess
base/carbonate
base/carbonate
+ solution of salt
solution of salt
(filtrate)
warm
saturated solution
allow to cool
crystals of the salt +
uncrystallised
solution
filter and rinse with
distilled water
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pure crystals of the
salt
31
If the above steps are followed for copper (II) oxide or copper (II) carbonate with dilute
sulfuric acid, the blue hydrated salt CuSO4.5H2O will be obtained. It should be dried between
folds of filter paper.
If the white, anhydrous sample of the salt is to be prepared, we can heat directly. To test for
water, add a liquid sample to anhydrous copper (II) sulfate: if the solid changes from white
to blue, the liquid contains water.
Note: Insoluble salts can also be prepared by precipitation, that is, a chemical reaction in which the
salt is formed as a suspension of small solid particles when two soluble reactants are mixed.
Uses of selected salts
1. Ammonium phosphate which has formula (NH4)3PO4 is used as a fertiliser that helps plants
produce strong roots.
2. Ammonium sulfate which has formula (NH4)2SO4 is a salt used as a nitrogen fertiliser to help
plants grow well.
3. Calcium sulfate (CaSO4) is used in plaster of Paris which is used as a cast for setting broken
bones.
4. Potassium chloride (KCl) helps plants carry out the process of photosynthesis. One problem
that may arise with these salts is that they are washed by heavy rains and they pollute our
rivers and lakes.
5. Sodium bicarbonate (NaHCO3) is used in baking and in the treatment of mild indigestion.
6. Sodium chloride is used for food preservation and to enhance the taste of foods.
7. Sodium fluoride is used in toothpaste to prevent cavities.
Combustion and its importance
Combustion is a chemical reaction in which a substance reacts rapidly with oxygen, producing heat
and light. For example, carbon glows in air to form carbon dioxide gas, while magnesium burns with
a bright flame to form magnesium oxide:
carbon + oxygen → carbon dioxide
C + O2 →
CO2
magnesium + oxygen → magnesium oxide
2 Mg
+ O2 →
2 MgO
Combustion of fuels is used to produce heat and to generate electricity.
Respiration and its importance
Respiration is a slower chemical process which takes place in living organisms, by which food
substances combine with the oxygen of the air to release energy, carbon dioxide and water vapour:
glucose + oxygen → carbon dioxide + water
C6H12O6 + 6 O2 →
6 CO2
+ 6 H2O
This process provides energy for the normal functioning of all tissues and cells in living matter.
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Photosynthesis and its importance
Photosynthesis is the chemical process by which green plants synthesize organic compounds from
carbon dioxide and water in the presence of sunlight:
sunlight
carbon dioxide + water →
6 CO2
sunlight
+ 6 H2O →
glucose + oxygen
C6H12O6 + 6 O2
Photosynthesis allows green plants to manufacture food.
Percentage composition of the gases in the air
Dry air is a mixture of gases whose composition, by volume is:
CONSTITUENT
PERCENTAGE BY VOLUME
nitrogen
78%
oxygen
21%
noble gases
less than 1%
carbon dioxide
0.034%
Air may also contain a variable amount of water vapour and air pollutants.
How the process of respiration and photosynthesis maintains the
composition of air
During respiration, living organisms use up oxygen from the atmosphere and they release carbon
dioxide.
During photosynthesis, green plants absorb this carbon dioxide from the atmosphere and release
oxygen.
In this way, the composition of the air is kept constant, as far as the levels of oxygen and carbon
dioxide are concerned.
Air pollution and its effects
Air Pollution is the addition of harmful substances to the atmosphere, resulting in damage to the
environment, human health, and quality of life. Air pollutants may consist of gases like carbon
monoxide, nitrogen oxides or sulfur dioxide. They may also consist of solid particles like ash, dust,
smoke or soot (a form of carbon).
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The table below gives more details about common gaseous pollutants.
POLLUTANTS
SOURCES
EFFECTS
carbon monoxide
incomplete
combustion of
carbon-containing
substances
aerosol propellants,
refrigerants and
solvents
poisonous gas which
cuts off oxygen supply
to the body
chlorofluorocarbons
(CFCs)
nitrogen monoxide
and nitrogen dioxide
lightning activity and
internal combustion
engines
smoke
produced by burning
materials in air
sulfur dioxide
volcanoes and
combustion of fossil
fuels
deplete atmospheric
ozone layer causing
harmful ultraviolet
rays to reach the
Earth’s surface
responsible for acid
rain and smog
causes pulmonary
irritation or poisoning
MEASURES THAT CAN
BE TAKEN
use catalytic
converters to oxidise
the carbon monoxide
to carbon dioxide
use aerosol
propellants,
refrigerants and
solvents which do not
react with ozone
fit catalytic converters
in motor vehicles to
reduce the nitrogen
oxides to nitrogen
do not burn materials
in open air if they
emit harmful smokes
use calcium carbonate
to reduce the effect of
‘acid rain’
leads to the formation
of acid rain which
causes damage to
animal life, plant life,
metal structures and
buildings
Note: Acid rain is a form of precipitation containing a heavy concentration of sulfuric and nitric acids.
Burning of fossil fuels, global warming and acid rain
Global warming refers to the measurable increases in the average temperature of the Earth’s
atmosphere, oceans and landmasses. Gases like carbon dioxide, methane or nitrogen dioxide (which
are referred to as Greenhouse gases) can trap too much of the Sun’s heat and cause global warming.
Carbon dioxide and nitrogen dioxide come from the burning of fossil fuels like petroleum, coal and
natural gas (methane). Global warming will cause a rise in the sea level. It will also be responsible
for destructive climatic conditions. This will eventually cause the extinction of many plant and animal
species.
Acid rain is a form of precipitation containing a heavy concentration of sulfuric and nitric acids. Acid
rains cause damage to plants, animals and man-made buildings and structures. The sulfuric acid
comes from sulfur dioxide (SO2) which itself originates from volcanoes and from the combustion of
fossil fuels. The nitric acid comes from oxides of nitrogen (NO and NO2) which themselves originate
from lightning activity and internal combustion engines.
Note: Global warming and acid rains are not the only threats to our atmosphere and our
environment. The use of chlorofluorocarbons (CFCs), which are compounds containing the
elements chlorine, fluorine and carbon, can deplete our ozone layer, causing more ultraviolet
(UV) radiation to reach the Earth’s surface and increasing the incidence of skin cancers and
cataract.
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Protects us from
Sun's harmful
ultraviolet radiation
Acid + carbonate→
salt + water + carbon
dioxide
Base + ammonium
salt→ salt + water +
ammonia
Acid + base or alkali→
salt + water
To neutralise excess
stomach acidity using
antacids
To neutralise bee
stings with baking
soda
To neutralise excess
soil acidity with lime
Acid + reactive
metal→ salt +
hydrogen
Depleted by
chlorofluocarbons
(CFCs)
with carbonates
Ozone layer depletion
can lead to increased
incidence of skin
cancers
To neutralise wasp
stings with vinegar
To neutralise acidic
exhaust chimney
gases with calcium
carbonate
with acids
with bases
with metals
Applications
importance
To neutralise
bacteria-produced
acids in the mouth
with toothpastes
Reactions of
bases and alkalis
Reactions of acids
dangers
Causes damage to
plants, animals and
man-made buildings
and structures
compound formed
when an acid
neutralises an alkali
Earth's ozone
layer
Both acids originate
from burning fossil
fuels
Neutralisation
consequnces
source
Precipitation
containing a heavy
concentration of
sulphuric and nitric
acids
definition
definition
Acid rain
definition
Important chemical
reactions
soluble salts
Salt
Increase in the
average temperature
of the Earth’s
atmosphere, oceans
and landmasses
definition
Global warming
preparation
acid salt formed
when hydrogen of
acid partially replaced
by metal
all nitrates,
ammonium salts and
sodium salts are
soluble
most chlorides,
bromides, iodides
and sulphates are
soluble
most carbonates are
insoluble
causes
Greenhouse gases
like carbon dioxide,
methane or nitrogen
dioxide responsible
for global warming
by titration or by the
action of acids on
metals, metal oxides
or carbonates
sources
Combustion
Air pollutants
Respiration
Photosynthesis
Carbon dioxide and
nitogen dioxide
released by burning
fossil fuels
definition
Air composition
definition
definition
carbon monoxide
78% nitrogen
21% oxygen
importance
CFCs
Chemical process by
which green plants
synthesize organic
compounds from
carbon dioxide and
water in the presence
of sunlight
nitrogen oxides
smoke
sulphur dioxide
<1% noble gases
carbon dioxide
0.034%
Allows green plants
to manufacture food
importance
Slower chemical
process by which
food substances
combine with the
oxygen of the air to
release energy,
carbon dioxide and
water vapour
chemical reaction in
which a substance
reacts rapidly with
oxygen, producing
heat and light
Examples are burning
of substances in air
Provides energy for
the normal
functioning of all
tissues and cells in
living matter
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CHAPTER FIVE: EXPERIMENTAL TECHNIQUES IN CHEMISTRY
Dissolving substances
A solute is a substance dissolved in a solvent to form a solution.
A solvent is a liquid that dissolves another substance or substances to form a solution.
A solution is a homogeneous mixture of a liquid (the solvent) with a gas or solid (the solute).
Thus, if we dissolve a spoonful of sugar completely in water we obtain a sugar solution. The water is
said to be the solvent and the sugar is the solute. Acetone is the solvent for nail polish.
A suspension is a mixture in which small solid or liquid particles are suspended in a liquid or gas.
Thus, if we shake a spoonful of flour in water we obtain a suspension of flour in water.
Different types of mixtures (Optional)
A mixture is an impure substance which contains two or more elements or compounds in
proportions which may vary. A mixture is a system of two or more distinct chemical substances.
Homogeneous mixtures have their atoms or molecules interspersed, that is, their components do
not appear separately. Examples of homogeneous mixtures are air, a sugar solution, a dilute solution
of hydrochloric acid, a mixture of alcohol and water, etc.
Heterogeneous mixtures have distinguishable phases, that is, its components appear separately.
Examples of heterogeneous mixture are a mixture of iron filings and sulfur, a mixture of oil and
water, a mixture of sand and water, etc.
In a mixture the components in it are not chemically combined together and the components retain
their individual chemical properties. For example, air will support combustion because oxygen is
present in it. Air will also turn limewater milky because it contains carbon dioxide.
Unlike compounds, mixtures can be separated by physical means like filtration, distillation,
crystallization, etc.
Changes of states
Experiment 21: To show that during boiling the temperature of water remains the same.
Experiment 22: To show that ammonium chloride sublimes on heating.
During any change of state, the temperature does not change.
Melting and melting point
Melting is the change of state from solid to liquid at constant temperature. This occurs when the
solid particles gain heat energy and they vibrate so fast that they can break their orderly
arrangement.
The melting point of a solid is the temperature at which the solid turns into a liquid at constant
temperature.
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Boiling and boiling point
Boiling is the change of state from liquid to gas at constant temperature. The temperature at which
boiling occurs is called the boiling point of the liquid. Boiling occurs when the liquid particles gain
heat energy and they move so fast that they can escape into the air.
The boiling point of a liquid is the temperature at which the liquid turns into a gas at constant
temperature.
Evaporation is also the change of state from liquid to gas but it occurs at any temperature.
Other changes of states
Freezing or solidification is the change of state from liquid to solid at constant temperature. The
freezing point of a liquid is the temperature at which the liquid turns into a solid at constant
temperature.
Condensation is the change of state from gas to liquid (or solid) at constant temperature.
Certain solids, like ammonium chloride, solid carbon dioxide (dry ice) or iodine do not melt when
heated—they change directly from solid to gas: these substances sublime on heating. Sublimation is
the change of state from solid to gas at constant temperature.
SOLID
Freezing
(solidification)
melting
condensation
(reverse sublimation)
sublimation
boiling
GAS
LIQUID
condensation
Note: Water evaporates from the sun-heated lakes and the sea to produce vapours which rise and
form clouds. When these clouds cool and condense they reach the ground as rain. This rain
water flows as rivers, which carry the liquid water back to lakes and the sea. This cycle then
repeats itself and form what is known as the ‘water cycle’.
Importance of pure substances
A pure substance is one that contains only one kind of matter. Elements and compounds are pure
substances.
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If impurities are present in an element or compound, we then have a mixture. Substances need to
be very pure if they are to be used as a medicine; the presence of impurities in a prescribed drug
may cause the person to suffer from other complications. In the laboratory, substances have to be
very pure so that we obtain the expected result when we perform experiments.
Separation of mixtures and their importance in everyday life
Use of magnet
Experiment 23: To separate a mixture of iron and sulfur using a magnet.
A mixture of two solids, among which one is magnetic, can be separated by using a strong magnet.
One example is a mixture of sulfur powder and iron filings. If we plunge the magnet into the mixture
and move it through, the iron filings will remain on the magnet when the latter is withdrawn. The
sulfur powder is left behind.
Examples of magnetic materials are iron, steel, cobalt and nickel.
Decantation
Decantation is the process of separating a liquid from a settled solid suspension or from a denser
immiscible liquid, by carefully pouring it into a different container.
A mixture of cooking oil and water can be partially separated by decanting the oil into another
container, although the use of a separating funnel will be preferred for such a separation.
For example a mixture of lead and water can be partially separated by decantation, simply by
pouring off the water in another container. The lead will remain at the bottom. A better method of
separation will, however, be filtration.
Filtration
Filtration is used to separate insoluble solids from a liquid, for example, powdered chalk and water.
It is performed in the laboratory using a funnel lined with a filter paper. Clear water will be collected
as the filtrate and chalk will be left on the filter paper as the residue.
filter
paper
chalk
(residue)
mixture of chalk
and water
clear water
(filtrate)
It can also be used to separate a mixture of two solids among which one is soluble, for example,
sodium chloride and calcium carbonate. Water should first be added and the mixture well stirred
before filtering.
In industry, filtration is used to purify water or beer. In the home, we use a tea strainer to remove
the solid particles; this is also a method of filtration.
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Crystallisation
Experiment 24: To separate crystals of copper (II) sulfate from an aqueous solution of copper (II)
sulfate.
Crystallisation is used to obtain a pure solid from an aqueous solution, for example, crystals of
copper (II) sulfate from aqueous copper (II) sulfate solution.
crystals
appear
copper (II) sulfate
solution
boiling
water
heat
We do not heat the solution directly so as not to decompose the crystals.
Simple distillation
Experiment 25: To separate a mixture of water and ink by simple distillation.
Simple distillation is used to separate a mixture of two miscible liquids which have different boiling
points, for example a mixture of ethanol and water. Ethanol is collected as the distillate, while water
is left in the boiling tube as the residue. Simple distillation is used to obtain fermented liquor.
thermometer
water out
condenser
mixture of
ethanol and
water
water in
boiling
stones
heat
distillate
Simple distillation is performed using a simple distillation apparatus. Water is made to flow from the
bottom to the top of the condenser. Boiling stones are also used to make the boiling smooth.
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Fractional distillation
Fractional distillation is more efficient than simple distillation and it can be used to separate and
purify a mixture of several miscible liquids, for example, crude oil or liquid air.
water out
Fractionating
column
packed with
glass beads
water in
very pure
distillate
liquid
mixture
As the vapour rises up the fractionating column, it gets purer and purer. The liquid with the lowest
boiling point is collected first.
Sublimation
This is the conversion of a solid into a vapour without the solid first melting. Examples of solids that
sublime are iodine, ammonium chloride, anhydrous aluminium chloride and anhydrous iron (III)
chloride.
glass funnel
mixture of
ammonium
chloride and
sodium chloride
watch
glass
gentle heating
pure ammonium
chloride
(sublimate)
sodium chloride
(residue)
A mixture of sodium chloride and ammonium chloride can be separated by the set-up shown above.
The ammonium chloride that collects over the inverted glass funnel can be scraped off with a knife
and collected.
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Chromatography
Experiment 26: To separate the colours present in black ink using chromatography.
Chromatography is a sensitive test which can be used to separate and identify small amounts of
substances. We use it especially for coloured substances or for very complicated ones.
most soluble
component
chromatography
paper
solvent
front
A
after some hours
spot of
black ink
pencil line
least soluble
component
B
C
D
solvent
front
original
spot
solvent
If a spot of black ink is applied on chromatography paper, which is then dipped into a suitable
solvent, it will be separated into its components as shown. Each component or spot represents one
substance present in the black ink.
The solvent front is the highest level reached by the solvent.
solvent
front
B
A
urine
sample
drug X
drug Y
drug Z
pencil
line
We can use chromatography to test for the presence of illegal substances in the urine sample of a
suspect. From the chromatogram above, we can see that the urine sample contained drug X and
drug Z, because they match horizontally with those of spots A and B respectively.
Chromatography is sometimes used for colourless substances. To make the spots visible, a locating
agent has to be sprayed on the chromatography paper.
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Used for coloured
substances
To separate and
identify small
amounts of
substances
substance dissolved
homogeneous
mixture of solid and
liquid
Used for complicated
substances
During change of
state temperature
does not change
solution
mixture of solid
particles suspended
in a liquid
liquid that dissolves
solute
Melting: change from
solid to liquid
solvent
Example of such
mixture: black ink
Melting point: fixed
temperature at which
solid turns into liquid
Freezing: change
from liquid to solid
Boiling point: fixed
temperature at which
liquid turns into gas
suspension
Chromatography
Boiling: change from
liquid to gas
Dissolving
substances
Used to identify
illegal substances in
urine samples of
suspects
Freezing point: fixed
temperature at which
liquid turns into solid
Changes of states
Condensation:
change from gas to
liquid or solid
Sublimation: change
from solid to gas
Used to separate
mixture of solids
among which one
sublimes on heating
Experimental
techniques in
chemistry
sublimation
Solids that sublime
are: iodine,
ammonium chloride,
anhydrous aluminium
chloride, anhydrous
iron (III) chloride
To separate mixture
of solids among
which one is
magnetic
Use of magnet
Magnetic materials:
iron, steel, cobalt,
nickel
Example of such
mixture: sulphur and
iron filings
Similar to but more
efficient than simple
distillation
To separate a mixture
of more than two
miscible liquids
To obtain a very pure
distillate
Fractional
distillation
Decantation
To separate a liquid
from a dense solid,
e.g water and sand
Simple distillation
Crystallisation
To separate a mixture
of two miscible
liquids which have
different boiling
points
Higher boiling point
liquid remains in
distillation flask as
the residue
Example of such
mixture: liquid air
Lower boiling point
liquid is collected as
the distillate
To separate two
immiscible liquids of
different densities,
e.g cooking oil and
water
Filtration
Example of such
mixture: ethanol and
water
Simply pour off less
dense liquid at the
top
Used to obtain a solid
from a solution
Example of such
mixture: copper (II)
sulphate solution
Do not heat crystals
directly with flame
To separate insoluble
solids from a liquid
Solid remains on filter
paper as the residue
Clear liquid is
collected downward
as the filtrate
Example of such
mixtures: sulphur and
water or
fermentation
mixtures like beer
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A
D
Acid ................................................................................. 8
Acid rain .................................................................. 29, 34
Acid salt ........................................................................ 29
Acids ............................................................................. 10
Aerosol propellants....................................................... 34
Air ................................................................................... 7
Air (composition of) ...................................................... 33
Air Pollution .................................................................. 33
Alkali ................................................................... 9, 10, 28
Alloy .............................................................................. 23
Aluminium .................................................................... 21
Ammonia ...................................................................... 28
Ammonium phosphate ................................................. 32
Ammonium salts ........................................................... 28
Ammonium sulphate .................................................... 32
Antacids ........................................................................ 29
Atom ............................................................................... 7
B
Decantation .................................................................. 38
Direct combination ....................................................... 19
Displacement reactions ................................................ 22
Distillate ........................................................................ 39
Distillation ..................................................................... 39
E
Element ........................................................................... 4
Elements, compounds and mixtures compared.............. 8
Energy ........................................................................... 12
Ethanoic acid ................................................................... 9
Evaporating dish ........................................................... 18
Evaporation ................................................................... 37
F
Baking soda ................................................................... 29
Balanced equations....................................................... 16
Base ................................................................................ 9
Beaker ........................................................................... 18
Bee stings ...................................................................... 29
Boiling ........................................................................... 37
Boiling point .................................................................. 37
Boiling stones ................................................................ 39
Boiling tubes ................................................................. 18
Fertiliser ........................................................................ 32
Filtrate........................................................................... 38
Filtration........................................................................ 38
Flat-bottomed flask ....................................................... 18
Formula ........................................................................... 6
Formulae of traditional compounds ............................. 14
Fossil fuels ..................................................................... 34
Fractional distillation .................................................... 40
Freezing......................................................................... 37
Freezing point ............................................................... 37
Funnel ........................................................................... 18
C
G
Calcium ................................................................... 21, 22
Calcium carbonate .................................................. 29, 34
Carbon dioxide preparation .......................................... 25
Carbon monoxide ......................................................... 34
Catalytic converters ...................................................... 34
Cataract ........................................................................ 34
Changes of states .......................................................... 36
Chemical change ........................................................... 12
Chemical reaction ......................................................... 15
Chlorofluorocarbons (CFCs) .......................................... 34
Chromatography ........................................................... 41
Combustion................................................................... 32
Compound ................................................................ 6, 13
Condensation ................................................................ 37
Condenser ..................................................................... 39
Copper .................................................................... 21, 22
Copper (II) sulfate ......................................................... 22
Crystallisation ............................................................... 39
Gases ............................................................................... 4
Global warming ............................................................. 34
Gold......................................................................... 21, 22
H
Heat changes................................................................. 12
Heterogeneous mixtures .............................................. 36
Homogeneous mixtures ................................................ 36
Hydrochloric acid ...................................................... 9, 28
Hydrogen .......................................................... 20, 22, 28
Hydrogen peroxide ....................................................... 24
Hydrogen preparation................................................... 23
I
Impurities ...................................................................... 38
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Indicators ........................................................................ 9
Ink ................................................................................. 41
Insoluble salts ............................................................... 30
Internal combustion engines ........................................ 34
Iron ......................................................................... 21, 22
Iron wool....................................................................... 21
Irreversible changes ...................................................... 12
L
Laboratory apparatus ................................................... 18
Lead ........................................................................ 21, 22
Liquids ............................................................................. 4
Litmus ....................................................................... 8, 28
pH scale ........................................................................... 9
Phenolphthalein ............................................................ 28
Photosynthesis ........................................................ 32, 33
physical and a chemical change compared ................... 12
Physical change ............................................................. 12
Pipette........................................................................... 18
Plaster of Paris .............................................................. 29
Pollutants ...................................................................... 33
Polybasic acids .............................................................. 29
Potassium................................................................ 21, 22
Potassium carbonate .................................................... 23
Potassium chloride........................................................ 32
Prevention of rusting .................................................... 26
Products ........................................................................ 15
Pure substances ............................................................ 37
M
R
Magnesia ...................................................................... 29
Magnesium ............................................................. 21, 22
Magnesium hydroxide .................................................... 9
Magnet ......................................................................... 38
Magnetic materials ....................................................... 38
Making compounds ...................................................... 19
Manganese (IV) oxide ................................................... 24
Mass.............................................................................. 12
Measuring cylinder ....................................................... 18
Melting.......................................................................... 36
Melting point ................................................................ 36
Mercury ........................................................................ 22
Metallic carbonates ...................................................... 23
Metals ............................................................................. 4
Methyl orange .............................................................. 28
Mixture ..................................................................... 7, 36
Molecular equation ...................................................... 16
Molecule ......................................................................... 7
Monobasic acids ........................................................... 29
N
Neutral solutions ...................................................... 9, 10
Neutralisation ......................................................... 28, 29
Neutralisation (applications in life) ............................... 29
Nitrogen fertiliser ......................................................... 32
Nitrogen monoxide and nitrogen dioxide ..................... 34
Non-metals ..................................................................... 4
O
Oxygen preparation ...................................................... 24
P
Periodic Table ................................................................. 5
Radicals ......................................................................... 14
Reactants ...................................................................... 15
Reactions of metals with dilute acids............................ 22
Reactions of metals with oxygen .................................. 20
Reactions of metals with water .................................... 21
Reactivity series ........................................................ 6, 20
Residue ................................................................... 38, 39
Respiration .................................................................... 33
Respiration .................................................................... 32
Reversible changes........................................................ 12
Round-bottomed flask .................................................. 18
Rusting of iron ............................................................... 25
S
Salt .......................................................................... 28, 29
Salts (uses) .................................................................... 32
Silver ....................................................................... 21, 22
Skin cancers .................................................................. 34
Smog ............................................................................. 34
Smoke ........................................................................... 34
Sodium .................................................................... 21, 22
Sodium bicarbonate ...................................................... 29
Sodium carbonate ......................................................... 23
Sodium chloride ............................................................ 32
Sodium fluoride............................................................. 32
Sodium hydrogencarbonate ......................................... 29
Sodium hydroxide ......................................................... 28
Solids ............................................................................... 4
Solubility of salts ........................................................... 30
Soluble salts .................................................................. 30
Soluble salts (preparation) ............................................ 30
Solute ............................................................................ 36
Solution ......................................................................... 36
Solvent .......................................................................... 36
Solvent front ................................................................. 41
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Sublimation ............................................................. 37, 40
sulfur dioxide ................................................................ 34
Sulphuric acid.................................................................. 9
Summary of reactivity series......................................... 23
suspension .................................................................... 36
symbol ............................................................................ 4
symbols of elements ..................................................... 13
Upward delivery ............................................................ 23
T
W
Tarnishing of metals ..................................................... 20
Test for carbon dioxide ................................................. 25
Test for hydrogen ......................................................... 24
Test for oxygen ............................................................. 25
Test-tube....................................................................... 18
Thermal decomposition ................................................ 19
Titration ........................................................................ 30
Toothpastes .............................................................. 9, 29
Wasp stings ................................................................... 29
Water cycle’ .................................................................. 37
Water pollution ............................................................. 32
Word equations ............................................................ 16
Writing down formula of compounds ........................... 13
V
Valencies of elements ................................................... 13
Valency.......................................................................... 13
Volcanoes...................................................................... 34
Z
Zinc.......................................................................... 21, 22
U
Ultraviolet rays ............................................................. 34
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