MR.CHITEMBEYAS COPY 1.

1 The three states of matters

(C) 2017 STATES OF MATTER
There are different substances around us. All these substances are called matter.

Matter is any substance that occupies space and has weight.

Matter exists in different forms. These forms are called states of matter.
COTENT LEARNING OBJECTIVE (Pupils should be able to)
There are 3 states of matter: solid, liquid and gas.
The three states  Identify solids, liquids and gases
of matter  Describe the solid, liquid and gaseous states of mater and [Examples of each state]
explain their inter-conversion in terms of the kinetic
particle theory Solid – salt, wood and glass
Changes of state  Describe the changes of state occurring when substances
are heated or cooled Liquid – water, paraffin and oil Air is a mixture of gases
 Determine the temperature at which these changes occur
Gas – hydrogen, oxygen and water vapour (steam)

The table below shows the characteristics of these 3 states of matter

Solids Liquids Gases

NO fixed shape. No fixed shape.
Shape Fixed shape Takes the shape of Takes the shape of
the container the container
No fixed volume.
Volume Fixed volume Fixed volume Takes the volume of
the container
Very slightly
Compressibility Incompressible compressible Very compressible
negligible
Table: Characteristics of the 3 states of matter

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THE KINETIC THEORY [Experiment]

The kinetic Theory was proposed to explain the characteristics of the three A crystal of potassium dichromate (VI) is
states of matter. It states that all matter is made up of extremely small put in a beaker filled with water. Leave the
particles that are in constant motion These particles can be atoms, ions or beaker undisturbed and observe carefully.
molecules.
As the crystal dissolves the colour slowly
STATE SOLID LIQUID GAS spreads through the liquid, first covering the
bottom.

Diagram
of particles Eventually the colour distributes itself evenly
throughout the liquid.

It is the potassium dichromate (VI) particles

Arrangement
Packed closely Packed loosely Spaced widely which slowly move from an area of high
of particles
concentration to an area of low concentration. This is diffusion in liquid.
As well as
Move at very high Three factors which can affect the rate of diffusion
Movement Vibrate about a vibrating, can
speeds in the 1. The higher the temperature is, the faster the diffusion is.
of particles fixed position move rapidly over
space available 2. The bigger the size of particle is, the faster the diffusion is.
short distances
3. The larger the concentration gradient is, the faster the diffusion is.
Attractive forces
Forces Attractive and are not strong Forces between
between repulsive forces enough to hold particles are 1.2 Changes in state
particles counterbalance particles in a negligible
regular pattern
A change of state is a change where one state changes to another
DIFFUSION High
Low There are some types like Melting, Evaporation/Boiling, Freezing/Solidification,
Condensation and Sublimation.
Diffusion is the movement of
.
particles from an area of high
concentration to an area of low Physical change can easily reverse and produce no new substance
E.g. Melting, Evaporation, Condensation
concentration
Chemical change can not easily reverse and produce new substances
Direction of diffusion E.g. Combustion, Decomposition

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 As a substance is heated, it absorbs energy and its temperature rises. Then it
MELTING EVAPORATION
changes from a solid to a liquid and finally to a gas.
SOLID LIQUID GAS
As you can see in the graph above, there are 2 types of sections; Slope and Flat
FREEZING CONDENSATION

The flat sections on the graph indicate the melting and boiling points. Here the
SUBLIMATION temperature remains the same over a period of time, as the heat energy is being
used to change the state of the substance.
SUBLIMATION
Sublimation takes place when a solid changes into a gas directly, without Heat energy can be used either to raise the temperature
going through the liquid state. of a substance or to change the state of it
Substances which sublime are
・ iodine, ammonium chloride (NH4Cl),
・ ammonium sulphate ((NH4)2SO4)
・ carbon dioxide (CO2, called dry ice) BOILING AND EVAPORATION

Boiling and evaporation are both physical process that change a liquid into a gas.
HEATING CURVE
The liquid absorbs heat energy during these changes in state.
A heating curve is a graph showing changes in temperature with time for a
These must be differentiated with each other. The table below shows the
substance being heated
differences between these 2 process.
The graph below is a heating curve for a substance of water

Boiling Evaporation
Temperat
Occurs at boiling point Occurs at any temperature below
ure
boiling point
Boiling point 100 oC Occurs throughout the liquid Occurs only at the surface of the
liquid
Melting point 0 oC Bubbles observed No bubbles observed
Occurs rapidly Occurs slowly
Time Table: Differences between boiling and evaporation
Boling point is the
ice ice and water water and steam
temperature at which all of
water steam
a liquid changes into a gas.
Melting point is the
temperature at which a solid A pure substance has an exact boiling and
completely changes into a melting point. Impurities raise boiling
liquid. points and cause lower melting points.

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 2.1Measurement

VOLUMES OF LIQUID AND GASES

COTENT LEARNING OBJECTIVE (Pupils should be able to) There are some types of apparatus to measure the volumes of liquids.
 Name and use appropriate apparatus for the measurement They have differences in accuracy.
of time, temperature, mass and volume, including burettes,
Measurement pipettes and measuring cylinders
 Design arrangements of apparatus, given information about Least accurate Most accurate
the substances involved
 Describe and use methods of separation by the use of
suitable solvent, filtration, crystallisation, distillation.
 Suggest suitable purification techniques, given information
about the substances involved.
Method of
 Describe and use paper chromatography and interpret
purification
chromatograms.
 Identify substances and test their purity by melting point
and boiling point determination and by paper
chromatography.

Most of materials we meet in our environment are mixtures. Often, only one flask
substance from a mixture is needed, so it has to be separated from the mixture by
Measuring
physical means.
cylinder
There are many industries in Zambia which produce a variety of products. During
the production of any of these products the industry begins with impure raw
pipette
materials that are often mixtures. The final product has to be extracted from the
raw materials by using some of the techniques we are going to learn in this topic.
burette

A Burette is a long vertical graduated glass tube with a tap at one end which is
used to add controlled volumes of liquids with accuracy up to 0.1cm3.
A Pipette is a graduated tube which is filled by suction and used for
transferring exact volumes (10cm3, 25cm3) of liquids.
A Measuring cylinder is used to measure the approximate volume of liquids.
A flask/beaker is used only for estimating volumes of liquids.

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HEATING Carbon dioxide gas is sometimes collected by displacement of water. Carbon
dioxide is sparingly soluble in water to form carbonic acid. Thus, the volume of
In experiment, a Bunsen burner is usually gas collected will be less than expected because some carbon dioxide will
used for heating. You can identify the dissolve into the water.
temperature of the flame from its colour
2.2 Method of purification (Separation techniques)

Yellow flame (luminous flame)  Air hole is closed. It produces pollutant gas A Solution is a mixture in which the particles of
There are some techniques to
solute and solvent are evenly spread out.
like carbon monoxide. separate mixture. How to determine
A Solute is the substance which dissolves in the
Blue flame (non-luminous flame)  Air hole is half open. It is most generally used. the method to separate it depends on solvent to form a solution.
Blue-green flame  Air hole is completely open. It is used for strong heating. some physical properties such as A Solvent is a liquid that dissolves substances.
solubility, density and so on.
[e.g.] Salt water = Salt + Water
COLLECTION OF GASES
(Solution) (Solute) (Solvent)
2 factors determine the method used to collect a gas: the density of the gas and
the solubility of the gas in water. A mixture is a material formed by two or more different
substances which are physically combined together
Method of collection FILTRATION
Displacement of water Type of Gases to be collected
For gases that are insoluble in water, Used to separate out an insoluble solid from a liquid,
e.g. hydrogen, oxygen. e.g. separating sand form sand water and water mixture.

Mixture of solid
and lliquid
Method of collection Method of collection Residue is the solid trapped in the
Displacement of air –upward delivery Displacement of air -downward delivery filter during filtration.
Filter funnel Filtrate is the clear liquid that
Gas jar
Gas passes through the filter during
Filter paper residue filtration.
Delivery tube
Delivery tube

Gas Gas jar

filtrate
Very fine pores in the filter paper allow small particles
Type of Gases to be collected Type of Gases to be collected to flow through, but retain the large particles.
For gases that are less dense than air, For gases that are denser than air,
e.g. hydrogen, ammonia. e.g. hydrogen chloride, carbon dioxide.
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CRYSTALLISATION DISTILLATION

 used to separate out a pure solid from an impure solution, e.g. separating Distillation is conducted using EVAPOTATION and CONDENSTION for
copper (II) sulphate crystals from impure copper (II) sulphate solution. separation.
The impurities will remain dissolved in solution.
SIMPLE DISTILLATION
STEP 1 STEP 2
Stir to  used to separate a pure liquid from a solution containing dissolved solids,
dissolve Evaporating
solvent basin e.g. separating pure water from seawater.
dolid

Impure
solid Distillate is pure and the
thermometer
condensed liquid obtained by
The impure solid is dissolved in a solvent. The solution is heated to evaporate most distillation.
of the solvent.
STEP 3 STEP 4 flask

Sea water

crystals Filter paper

Boiling stones for
The hot solution is allowed to cool. The The cold solution is poured off to obtain smooth boiling
solid appears as pure crystals. the crystals. The crystals may be dried
by pressing them between sheets of Pure water
filter paper. (distillate)

Figure: crystallisation

A thermometer is placed at the mouth of the condenser to

Crystallisation must be different from evaporation. In crystallisation, the measure the temperature of the vapour entering it.
solvent is partially evaporated, leaving small amount of solution in which the This temperature is the boiling point of the distillate.
crystals form. Impurities are left behind in the solution when the crystals
are filtered off. In evaporation, all the solvent is removed. The crystals
formed may be impure.

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 FRACTIONAL DISTILLATION
Temperature
The graph shows the change in temperature as
 used to separate a pure liquid from miscible liquids. the mixture of ethanol and water is being heated in
e.g. ethanol from a mixture of ethanol and water. the flask. The temperature will remain at 78 C when 100 oC
the first ethanol is being collected. When all the Water distils
ethanol has evaporated over, the temperature will rise
Miscible describes two or more liquids which will 78 oC
again until it reaches 100 C. At this temperature,
diffuse together and form a single. Ethanol distils
water will be collected as the second distillate.
e.g.) alcohol and water
Immiscible describes two or more liquids that will Time
not mix together.
e.g.) oil and water Fractional distillation is also used to separate
1 the compounds of crude oil
2 fermented liquor to obtain alcoholic drinks of a higher
concentration.

Fractionating
column [EXAPMPLE] Separate a mixture of sand and salt.

-The common procedure is …

flask 1. Dissolve
Place the mixture of sand and salt in a beaker, add water and stir.
Ethanol
 Salt will dissolve while sand will not dissolve because of their solubilities.
(distillate)
Mixture of
ethanol Boiling stones
2. Filtration
and water for smooth Pour the liquid of the mixture along a glass rod in to the funnel with a filter paper.
boiling  Salt water will pass through the filter paper as a filtrate while Sand will be
trapped on the paper as a residue
3. Evaporating
Put a little amount of the filtrate into the evaporating dish. Heat the filtrate until
all the water is driven off.
Fractional distillation separates according to boiling points.  Salt(solute) will remain in the dish while water(solvent) will go away as steam.
The liquid with the lowest boiling point will be distilled first, followed by the liquid
with the next lowest boiling point. As a rough guide, the boiling points of the liquids Now, Steps to Separate Insoluble and Soluble Solids:
to be separated should be at least 20º C apart. Salt and sand have been separated 1. Dissolve 2. Filtration 3. Evaporating

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PAPER CHROMATOGRAPHY INTERPRETATION OF RESULTS

Substances in a mixture are separated according to their solubility in the same 1. From unknown dye, follow direction of solvent flow until you find a spot
solvent. The more soluble component in the mixture will tend to remain in the 2. Move across the chromatogram until you find a corresponding spot
solvent and travel further up the chromatogram, while the less soluble 3. Move in the other direction to identify the known dye
component will separate out onto the paper.

o
o o
cover o
solvent
flow o o
Black Red Blue Green Orange
solvent x------ x------ x------ x------ x
Drops of ink on the pencil line
U A B C D

PROCEDURE Unknown Dye U = Mixture of Dyes A + C

1 Use a pencil to draw the start line.

2 Use the black ink sample to make a small dot on the start line,
together with some other coloured ink to use as reference.
3 Fold the paper into a cylinder and place it into a beaker containing
the solvent, ensuring that the start line is above the solvent level. It is possible for different substance to travel the same
Cover the beaker while the chromatogram develops. distance in the same solvent in paper chromatography, i.e. have
4 Remove the chromatogram from the beaker just before the solvent the same solubility in the same solvent.
reaches the top of the paper.. To confirm the identity of a substance, another round of
chromatography is carried out using a different solvent. If the
spot still travel the same distance, then the spot must contain
 If the start line is drawn in ink, the components in the ink will also separate
the same substance.
out together with the sample dots when the chromatogram is run.
 If the start line is below the solvent level, the sample dots will dissolve into
the solvent in the beaker instead of travelling up the chromatogram paper.
 The beaker must be covered when the chromatogram is run to reduce
evaporation of the solvent from the beaker and to prevent the solvent from
evaporating off the paper as it moves up.
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ELECTROLYSIS PROCEDURE
Electrolysis is the process by which an electric current flowing through liquid A concentrated solution of sodium chloride (NaCl) is electrolysed. This is what happens:
containing ions causes the liquid to undergo chemical decomposition.

APPARATUS Figure shows the apparatus need for electrolysis.

Negative Positive
+ - + - + -
The cathode is a The anode is a positive
negative electrode to electrode to which the 1.The solution contains Na+ 2.At the cathode, it is the H+ 3. At the anode, the Cl- ions
which the cations are anions are attracted -
ions and Cl ions form the ions which accept electrons give up electrons more readily
attracted during + + during electrolysis. salt, and H+ and OH- ions since sodium is more reactive than the OH- ions do. Chlorine
- -
electrolysis. form water. The positive ions than hydrogen; gas bubbles off.
+ + - -
go to the cathode and the 2H+ + 2e-  H2 2Cl-  Cl2 + 2e-
+ + - - Anions are atoms or negative ions to the anode Hydrogen gas bubbles off while OH- ions remain in solution
Cations are atoms or + + + --- - molecules containing Na+ ions remain in solution.
molecules containing more electrons that
fewer electrons than When the hydrogen and chlorine bubble off, Na+ and OH- ions are left behind
protons, and so carrying
protons, and so carrying a solution of sodium hydroxde is formed.
a negative charge.
a positive charge. SUMMARY of SEPARATION TECHNIQUE
SEPARATION SUBSTANCES TO BE
EXAMPLE
TECHNIQUE SEPARATED
Electrolytes are liquids which conduct electricity. An Electrode is a piece of metal
Filtration  Insoluble solid and liquid Muddy water
 All ionic compounds when molten or in aqueous solution or carbon (graphite) placed in an
are electrolytes, as their ions are free to move electrolyte which allows electric Crystallisation /  Solute (soluble solid) from its
Salt solution
current to enter and leave Evaporation solution
during electrolysis. Distillation  Solvent from its solution Salt solution
Ethanol and water
Fractional  Miscible liquids with different
USES of electrolysis Crude oil
Distillation boiling points
Apart form decompositions, electrolysis is used as followings; Liquid air
Decantation /  Insoluble suspension settles to Mealie-meal and
 ELECTROPLATING This is the coating of a metal object with a thin layer of another Sedimentation form sediment water
metal by electrolysis. Separating Funnel  Immiscible liquids Oil and water
Charcoal dust and
 EXTRACTION OF METALS Metals which are high in the reactivity series must be Floatation  Less dense solid and liquid
water
extracted by electrolysis. A common example is the
Magnetic Iron filings and
extraction of aluminium.  Magnetic materials
Separation sulphur powder
 ELECTROREFINING This is a method of purifying metals such as copper by Paper Dyes and pigment
 Dissolved substances
electrolysis. Chromatography of ink

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 3.1 ATOM

Are you sure the meaning of terms like atoms, molecules, elements, compounds and
COTENT LEARNING OBJECTIVE (Pupils should be able to) mixtures. Can you distinguish them clearly? Here are definitions for them.
Atoms  State the relative charges and approximate relative masses
of protons, neutrons and electrons.
 Define proton number (atomic number) and nucleon number
An ATOM is the smallest particle of an element which can take
(mass number).
part in a chemical reaction and remain unchanged.
 Use and interpret such symbols as 6 C
12

A MOLECULE is the smallest particle of an elements or a

 Use proton (atomic) number and the simple structure of compound which exists independently
atoms to explain the periodic table, with special reference
to the elements of proton (atomic) number 1 to 20. An ELEMENT is a substance that cannot be broken down into two
 Define isotopes or more simpler substances by chemical means.
 Describe the build-up the electrons in ‘shells’ and explain
the significance of valency electrons and the noble gas A COMPOUND is a substance that consists of two or more
electronic structures. elements which are chemically combined in fixed proportions.
Chemical Bonding  Describe the formation of ionic bonding between metallic
and non-metallic elements, e.g. NaCl, CaCl2 A MIXTURE is a substance that consists of two or more
 Describe the formation of covalent molecules. substances which are not chemically combined
 Deduce the electron arrangement in other covalent
molecules. As you have seen they are considered in terms of substances while they are
 Construct ‘dots and cross’ diagrams to show the valency considered in terms of particles.
electrons in covalent molecules.
 Describe the differences in volatility, solubility and ATOMS and MOLECULES
electrical conductivity between ionic and covalent
compounds. In kinetic theory, we saw that matter consists of particles. What are these
Structures and  Describe the differences between elements, compounds and particles? It is very useful to know about them.
Properties of mixtures and between metals and non-metals.
materials Molecules can be thought of to be 3 types.
In kinetic theory we saw that matter consisted of particles. We looked at how 1. It consists of those elements in which a single atom forms the molecule.
these particles account for the differences in the physical properties of solids, - These molecules are called monatomic
liquids and gases. [e.g.] helium, neon and argon which are known as the noble gases
Now we shall consider all substances as chemical substances 2. It consists of atoms of the same element combined together.
- These molecules are called diatomic
[e.g.] oxygen, hydrogen, nitrogen and chlorine

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 3. It consists of atoms of different elements combined together. NUCLIDE NOTATION
Here the atoms form molecules of compounds. One convenient method of writing the names of elements is by applying a
[e.g.] carbon dioxide, water and sugar shorthand system in which each element is assigned a specific symbol.

Now we shall see more about an atom!

Mass Number (or Nucleon Number) is the total number
STURECTURE OF AN ATOM A of protons and neutrons found in the nucleus of an atom.

X Atomic Number (or Proton Number) is the number of

Atoms are made up of three fundamental particles. These are the proton, the
Z protons an element has in the nucleus of its atom.
neutron and the electron

Charge Mass Position in Atom A chemical symbol of an element is a letter or letters derived
PROTON + 1 from the name of the element.
nucleus
NEUTRON 0 1
ELECTRON - 1/2000 electron shells Information about the number of particle in an atom can be found from the
Properties of particles found in an atom periodic table.

- Most of an atom is empty space. The protons and neutrons cluster

- An atom is electrically neutral. together in the centre, forming the From Nuclide Notation many information of an atom can be obtained.
nucleus; this is the heavy part of
－
 Number of Neutrons = Mass Number – Atomic Number
the atom and positive charged.
 Number of Electrons = Number of Protons = Atomic Number
(So that Atoms should be electrically neutral)
＋
－ －
＋ ＋ [Example] What are the mass number, proton number and number of neutron in
＋ Aluminium, hydrogen and lead?
The electrons circle very fast around
the nucleus, at different levels from it
along electronic shells. Aluminium Hydrogen Lead
－
Form periodic table

Mass number 27 1 207

These particles are extremely small. If a golf Proton number 13 1 82
ball were magnified to the size of the Earth, No. of neutrons 27 – 13 = 14 1-1=0 207 – 82 =125
then an atom would be the size of a marble!
NOTICE from the above example that the number of neutrons is obtained from
They have a radius of around 10-10 m and a mass
the periodic table using the top number minus the bottom number, however, this is
of about 10-22g not true for all cases. For example, chlorine in the periodic table is represented as
shown left; this does not mean that the chlorine atom contains 35.5-17 = 18.5
neutrons! Chlorine contains 2 isotopes (see later in the chapter) and to account for
11 these 2 isotopes, its mass number is a calculated average value of 35.5
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ISOTOPES DRAWING of ELECTRONIC STRUCTURE
Many elements contain atoms that are slightly different from each other.
The number of electrons in each shell can also be shown by drawing as the
followings
ISOTOPES are different atoms of the same element Rules to draw electronic structure: 24
which have the same number of protons but different
[e.g.] 12 Mg
number of neutrons.  Use the symbol of the element to
represent the nucleus 12e- = 2, 8, 2
It is known that 3 isotopes of hydrogen exist, also isotopes of carbon are known.  Represent each e- with a cross
X X
 Each new circle represents another
X X
Element Isotopes No. of Protons No. of No. of electron shell X
Neutrons Electrons  Start filling up the shells from the first X X
X Mg X
Hydrogen 1
1
H 1 0 1 shell before going on to the next shell X
2
H 1 1 1  Group the electrons in pairs for easy X X
1
1 2 1 counting
3
1
H
Carbon 6C
12
6 6 6
 Nearly full shells want to get extra e- - lose or share (NON-METALS)
6 7 6
6C
13
 Nearly empty shells want to lose extra e- (METALS)
6 8 6
6C
14
-
 Valency = No. of e an atom wants to lose, gain or share
Table: Isotopes of some elements

Electrons found in the outer shell are called Valence electrons (v.e.)

ELECTRONIC SHELLS
Electrons are arranged in electronic shells around nucleus
The number of electrons in each shell is finite as shown below [Example 1] Lithium, Li [Example 2] Potassium, K

Shell(from a nucleus) 1st 2nd 3rd 4th ….. The lithium atom contains 3 electrons The potassium atom contains 19 electrons
Maximum number 2 8 8 ....... arranged in 2 shells. arranged in 4 shells
Table : Maximum No. of electrons that can occupy the shell
X X
X X
The number of electons in each shell is shown by the elctronic configuration
X X
ATOM Electronic Configuration X X X XX
Li X K X
Lithium 2:1 X X X
X
X
Potassium 2:8:8:1 X X
Table: Electronic Configuration of some atoms X X

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3.2 CHEMICAL BONDING IONI BONDING (ELECTROVALENT)

When atoms combine together in a chemical reaction, we say that a bond is FORMATION OF IONS
formed between the atoms during the reaction. A reactive atom will combine or Atoms can obtain a full outer shell and become stable when they lose or gain
form bonds with other atoms easily, while an unreactive atom will not. valence electrons. Charged particles called ions are formed.

WHY DO ATOMS FORM BONDS? 1. FORMATION of POSITIVE IONS Note that 1+ is written simply
as +. The 1 is not written.
Li
7
[Example 1] Lithium ３
 Atoms of noble gases possess the maximum number of electrons in their
outermost shell as shown in the diagram below. Lose 1 Valence +
× electron ×
X X
XX Li × Li
× X
XX × ×
He
X
Ne
X
XX X XX
X
X
X X X Ar X X Lithium atom, Li Lithium ion, Li+
We do not draw
× X 3protons 3protons empty shells
X X
XX 3electrons 3-1= 2electorns either.
XX 7-3 = 4 neutrons 4 neutrons
4 20 40
Helium, 2He Neon, Ne
10 Argon, Ar
18 Overall zero charge, i.e. neutral Overall 1+ charge

The lithium ion now carries a 1+ charge because it has an extra proton.
All the outermost shells (or valence shells) are completely filled. This type of This is represented by enclosing the ion in brackets and writing its charge
arrangement is very stable and highly unreactive. on the top right hand corner.

 Atoms of other elements have incompletely filled outermost shells, as a 27

[Example 2] Aluminium 1３ Al
result, these atoms are unstable and therefore, reactive. ●
Most of them tend to become stable like noble gases. ●● ●● 3+
●● Lose 3 Valence
● ●
● electrons ● ●
In order to achieve stable electronic structure, they share, gain ● Al ●●
● Al ●
● ●
or lose electrons in their outer electronic shells. ● ●●
● ● ●●

This is why atoms form bonds. Aluminium

● atom, Al Aluminium ion, Al3+
●13protons 13protons
There are 2types of chemical bonds that can be formed between 2 atoms: 13electrons 13-3= 10electorns
1. Ionic bonds – valence electrons are transferred from one atom to another 27-13 = 14 neutrons 14 neutrons
2. Covalent bonds – valence electrons are shared Overall zero charge, i.e. neutral Overall 3+ charge
In general, when an atom loses n valence electrons to form a stable ion,
the ion formed will carry n+ charge.
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 2. FORMATION of NEGATIVE IONS FORMATION OF IONIC BONDING
19
[Example 3] Fluorine, 9 F
・Ionic bonding occurs between metallic and non-metallic atoms.
- ・Valence electrons are transferred from the metallic atom to the non-metallic
●●
●●
Gain 1 Valence ● atom so that both atoms achieve a full outer shell and become stable.
● ●
● electron
F x ・Oppositely charged ions are formed. The metal ion carries positive charge,
● F ●
●
●
●
● while the non-metallic ion carries negative charge. These ions attract each
●●
●● other with strong electrostatic forces to form an ionic bond.

Fluorine atom, F Fluorine ion, F-

Rules for ‘Dot and Cross’ Diagrams for Ionic Bonding:
9protons 9protons
9electrons 9+1= 10electorns
 Calculate how many e- metal atom wants to lose
19-9 = 10 neutrons 10 neutrons
 Calculate how many e- non-metal wants to gain
Overall zero charge, i.e. neutral Overall - charge
 To find out how many metal and non-metal atoms combine, use:
Total No. of e- lost = Total No. of e- gained
The fluoride ion now carries a 1- charge because it has 1 extra electron.
 No. of Metal Atoms No. of Non Metal Atoms
Note: Non-metal name changes to end “-ide” x No. of e- lost x No. of e- gained
[e.g.] chlorine  chloride oxygen  oxide,
sulphur  sulphide Phosphorus  phosphide etc.  Draw each atom’s e- shells BEFORE losing or gaining
 Metal e- with a ‘cross’ x
32
[Example 4] Sulphur atom, 16 S  Non-metal e- with a ‘dot’ o
2-  Re-draw each atoms’ shells AFTER losing or gaining
XX XX
XX XX  REMEMBER that the metal e-‘s GAINED by non-metal are still a ‘x’
Gain 2 Valence  Check: Every ion should now have FULL OUTER e- shells
X X
X X X electrons X X X●
XX S XX S  INDICATE the CHARGE on each ion
XX XX
X X [EXAMPLE]
XX XX Magnesium fluoride contains magnesium and fluorine atoms. Draw the dot and
X X●
cross diagram to represent magnesium fluoride.
Sulphur atom, S Sulphur ion, S2-
ANSWER: Mg: 12 e- = 2, 8, 2 F: 9 e- = 2, 7 { Periodic Table }
16protons 16protons
BEFORE: X X O O
16electrons 16+2= 18electorns
X X O
32-16 = 16 neutrons 16 neutrons X O O
Overall zero charge, i.e. neutral Overall 2- charge X X O F
X Mg X O
X O O
In general, when an atom gains m electrons to form a stable ion, the ion will X X
 gain 1 e-  need 2 F atoms
carry m- charge. - 2+
 lose 2 e  Mg ion (to gain 2e-)  2F- (fluoride) ions

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AFTER: [EXAMPLE]
- 2+ - Water is a molecular compound containing hydrogen and oxygen atoms. Draw the
O O X X O O
O X O
dot and cross diagram to represent a molecule of water. Show all the electron
O O X X O O shells.
O F X X Mg X X F O
O X O
O O X X O O ANSWER: H: 1 e- = 1 O: 8 e- = 2, 6 { Periodic Table }
BEFORE: O O
H X O
O
・ The formula of the compound formed is written as MgF2. O O
 shares 1 e- O
・ Ionic compounds are electrically neutral, i.e. once the positive and negative ions
O O
combine, these charges ‘cancel’ each other out to give a neutral compound.  shares 2 e-
 need 2 x H atoms to share with 1 x O atom

It is not necessary to draw and show the movement of valence electrons from
one atom to another unless the question requires it. Most exam question will ask
AFTER: O O
for the final structure of the compound only.
O Reaction: H 2 + O2  H 2 O
X X H
HO
O O
O Structural formula: H–O–H
COVALENT BONDING (MOLECULAR) O O
Water molecule, H2O
FORMATION OF COVALENT BONDING
When drawing covalent structures, always draw the atom
・Covalent bonding usually takes place between non-metallic atoms.
that needs to form the most number of bonds in the
・Valence electrons are shared between these atoms.
centre, then add on the res of the atoms
・The molecules of the compound are held together by weak intermolecular
forces that are easily broken by heating.
DIFFERENCES in PROPERTIES of IONIC and COVALENT BONDING
Rules for ‘Dot and Cross’ Diagrams for Molecular Bonding:
IONIC (Electrovalent) COVALENT (Molecular)
 For a molecule of an element: Represent the e- of each atom with a ‘dot’ or  Ionic compounds can conduct electricity  Molecular compounds do not
a ‘cross’ when molten or aqueous because the ions conduct electricity in any form
 For a molecule of a compound: Represent the e- of each element with a ‘dot’ are free to move  Molecular compounds have low
or a ‘cross’  Ionic compounds have high MP and BP MP and BP due to weak inter-
 (Usually) Only draw the OUTER e- shell due to the strong electrostatic forces molecular forces
 Draw the SHARED e- FIRST between charged ions  Soluble in organic solvents (E.g.
 Then ADD the REMAINING e- to the non-shared section for EACH ATOM  Soluble in water, insoluble in organic ethanol, petrol)
 Check: Each atom should have the correct number of e- for their shell solvents Insoluble in water
 Double-check: Every atom should now have FULL OUTER e- shells
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Table below summarises properties (a) a monatomic gaseous element made up of atoms, e.g. helium
(b) an gaseous element made up of diatomic molecules, e.g. hydrogen, oxygen,
Type of Bonding nitrogen etc.
Property (c) a solid element, e.g. iron, copper, etc.
IONIC (Electrovalent) COVALENT (Molecular)

Conduct YES COMPOUNDS

NO
Electricity (when AQUEOUS or MOLTEN) A compound is made up of two or more types of atoms chemically combined
together. It can not be separated using physical means. Chemical means such as
MP and BP HIGH LOW electrolysis are needed
Volatility NON-VOLATILE VOLATILE
MIXTURES
Usual State
SOLID GASES OR LIQUIDS A mixture is made up of two or more elements or compounds physically
(at room temp)
combined together. The components can be separated easily form one another
Composition IONIC LATTICE MOLECULES using physical means such as filtration, a magnet, distillation, etc.

+ - + - + - + - ELEMENTS MIXTURES COMPOUNDS

Diagram
- + - + - + - +

+ - + - + - + -

+ + + + Diagram
Examples NaCl, CaO, MgF2 H2O, CO2, O2

Melting Fixed melting Melts and boils Fixed melting and

3.3 STRUCTURE AND PROPERTIES OF MATERIALS and Boling and boiling over a range of boiling point
Points point temperature
ELEMENTS,COMPOUNDS AND MIXTURES Easily separated Chemical means
Separation ----- using physical means such as electrolysis
ELEMENTS such as distillation are needed
Elements are made up of only one kind of atoms. The diagram below shows  Copper  Air  Carbon dioxide
examples of elements existing as atoms as well as molecules.  Iron  Sea water  Water
(a) (b) (c) Examples  Oxygen  Metal alloys  Common salt
 Nitrogen  Rock salt  Ethanol
 Carbon  Ziggy solution
Table: the differences of Elements, Mixtures and Compounds

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 4.1 FORMULAE and EQUATIONS

It is useful to know the names of elements and compounds, how these names can
be represented and how chemical changes involving elements and compounds
maybe described.

COTENT LEARNING OBJECTIVE (Pupils should be able to) CHEMICAL FORMULAE

 State the symbols of the elements and formulae of One convenient method of writing the names of compounds is by using chemical
compounds. formulae.
 Deduce the formula of a simple compound from the relative There are some types of chemical formulae. Here we talk about Molecular
numbers of atoms and vice versa. formulae. Later we will see Empirical formulae and Structural formulae

 Determine the formula of an ionic compound from the

Formulae and A chemical formula is a way of showing the proportions of elements present in a
charge on the ions present and vice versa
Equations chemical compound using symbols for the atoms present.
 Construct equations with sate symbols, including ionic
equations.
 Deduce, from experimental results of the identity of the Rules for Chemical Formulae: It shows the number of atoms
reactants and products, the balanced chemical equation for
a chemical reaction.  A SMALL number multiplies ONLY the elements or radicals to the LEFT
 Define relative atomic mass, Ar.
Stoichiometric  Define relative molecular mass, Mr.  A BIG number at the FRONT multiplies ALL the elements in the
calculations  Perform calculations concerning reading masses reacting formulae
masses using simple proportions. It shows the number of molecules

[EXAMPLE]

・ H2O represents 2 atoms of hydrogen and 1 atom of oxygen in 1 molecule of

water
・ 2CO2 represents 1 atom of carbon and 2 atoms of oxygen in 2 molecules of
carbon dioxide
・ 4Na2SO4 represents 2 atoms of sodium, 1atom of sulphur and 4 atoms of
oxygen in 4molecules of sodium sulphate.

[EXAMPLE 2]
How many atoms are represented by 2Al2(SO4)3 ?
ANS: 2 x (2 x Al + 3 x (S + 4 x O) ) = 2 x ( 2 + 3 x (1 + 4) ) = 34 atoms

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CHEMICAL FORMULAE FOR ELEMENTS VALENCY
If the Valencies of the elements which take part in the compound are known,
1 Metals exist as atoms. The chemical formula for a metal is its symbol. writing the chemical formula is simple.
[EXAMPLE] Sodium  Na Magnesium  Mg Iron  Fe

2 Most non-metals, with the exception of the noble gases, exist as molecules. Its
Valency is the combining power of an atom or radical.
chemical formula will show both the symbol as well as the number of atoms that
In ionic compounds it is the same as the charge on the ion.
make up the molecule. In covalent compounds it is equal to the number of bonds formed.
Where the subscript ‘2’ shows that the molecule
[EXAMPLE] Hydrogen  H2,
is made up of two hydrogen atoms joined
together.

+ METALS + - NON-METALS -
Noble gases exist as atoms. The chemical formula for a noble gas is thus its symbol.
Element Symbol of ion Element Symbol of ion VALENCY
[EXAMPLE] Helium  He Neon  Ne.
Hydrogen H-
+
Lithium Li Fluorine F-

Metallic Element Chemical Formula Non-Metallic Element Chemical Formula

Sodium
Potassium
Na+
K+
Chlorine
Bromine
Cl-
Br-
1
Calcium Ca Chlorine Cl2
Iodine I-
Zinc Zn Oxygen O2
Copper Cu Nitrogen N2
Lead Pb Carbon C Magnesium Mg2+
2
Oxygen O2-
Calcium Ba2+
Manganese Mn Sulphur S Sulphur S2-
Barium Ca2+
Mercury Hg Argon Ar
Table ; Chemical formulae for some common elements

3
Nitrogen N3-
Aluminium Al3+
Phosphorous P3-
CHEMICAL FORMULA for COMPOUNDS
List of Valency for common ions
IONIC COMPOUNDS
The formulae of both the positive ion and the negative ion must be determined
before the chemical formula of the ionic compound can be written.
for metal = the number of electron in the outermost shell
for non-metal = 8 – the number of electron in the outermost shell

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 Some metals can form positive ions with different charges, depending on the RADICALS
compound that they are found in. Some negative ions exist in groups with an overall charge.

Element Symbol Valency Element Symbol of Valency A radical is a group of atoms within a compound that
of ion ion maintains its identity throughout a chemical reaction.
Copper(I) C+ 1 Mercury(I) Hg+ 1 It can not exist by itself
Copper(II) C2+ 2 Mercury(II) Hg2+ 2
Iron(II) Fe2+ 2 Lead(II) Pb2+ 2
Iron(III) Fe3+ 3 Lead(IV) Pb4+ 4
Tin(II) Sn2+ 2 Cobalt(II) Co2+ 2 RADICAL Symbol of its ion VALENCY
Tin(III) Sn3+ 3 Cobalt(III) Co3+ 3
Hydroxide OH-
Chromium(II) Cr2+ 2 Nickel(II) Ni2+ 2
Nitrate NO3-
Chromium(III) Cr3+ 3 Nickel(IV) Ni4+ 4
Nitrite NO2-
Manganese(II) Mn2+ 2 Silver(I) Ag+ 1
Hydrogen carbonate HCO3-
1
Manganese(IV) Mn4+ 4 Zinc(II) Zn2+ 2
Hydrogen sulphate HSO4-
Chlorate ClO3-
List of valency for common ions with variable charges
Manganate(VII) MnO4-
Ethanoate CH3COO -
Note that all these ions are formed from
transition metals. All charges are positive
Ammonium NH4+

Sometimes the charges on silver and zinc ions are not represented. Assume then
Carbonate CO32-
that the silver ion is Ag+, and the zinc ion is Zn2+
2
Sulphate SO42-
Sulphite SO32-
Dichromate(VI) Cr2 O72

Phosphate PO43- 3
List of common radicals with valencies

When writing chemical formulae involving

radicals, never take apart with each
element – take it as a whole group

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How to Deduce Chemical Formulae from Valencies or Ions: How to Deduce the Valency or Ion from the Chemical Formulae:
(WORKING BACKWARDS)
1. Write the symbols for the combining elements and radicals  Write the chemical formula of the compound
 Include all subscripts even 1
Magnesium oxide  Mg (magnesium), O (oxygen)  Re-write the symbols for each element / radical, but swap the subscript
2. Write the valency of each on the top right-hand side (leave off the charges numbers and write at the top right-hand side
for ions)  MULTIPLY BOTH NUMBERS by the same number to get the CORRECT
Mg2 O2 VALENCY for any known element / radical
 Include the CHARGE of the ION if required
3. Re-write the symbols, but swap the valencies and write at the bottom right-
 1st ion is + ve and 2nd ion is - ve
hand side
[EXAMPLE]
Mg2 O2
What is the charge of the iron ion in Fe2(SO4)3 ?
Fe2 (SO4)3
Mg2 O2 The charge on the ion is
(SWAP)
4. Find the lowest ratio of the two numbers the same as the valency
Fe3 (SO4)2
2:2  1:1 ANS: The charge on the iron ion is +3
the chemical formula is MgO

[EXAMPLE] COVALENT COMPOUNDS

Ignore writing subscript 1
Many exceptions exist to the rules for writing the chemical formulae of covalent
Aluminium sulphate compounds, making them difficult to remember. Some general rules:
Al3 (SO4)2  Many gases are made up of diatomic molecules, i.e. H2, O2, N2 etc
(SWAP)  Group VII elements also exist as diatomic molecules, i.e. Cl 2, Br2, F2, etc.
Al2 (SO4)3
2:3 lowest ratio You can use the rule for some Note that when HCl is in gaseous
Make sure;
covalent compounds like water, form, it is called hydrogen chloride
ANS: Al2(SO4)3 Put BRACKETS around any carbon dioxide, etc. You can try it. gas; when it is dissolved in water,
RADICALS before swapping!!! it forms a solution
Calcium Carbonate Compound Formula Compound Formula called hydrochloric
Remove the brackets from the
Ca2 (CO3)2 Carbon monoxide CO Nitric acid HNO３ acid.
(SWAP) radical with subscript 1
Carbon dioxide CO2 Sulphuric acid H2SO４
Ca2 (CO3)2 Sulphur dioxide SO2 Hydrochloric acid HCl
2:2  1:1 lowest ratio
Sulphur trioxide SO３ Methane CH４
ANS: CaCO３
Ammonia NH３ Ozone O３
Hydrogen chloride HCl Ethanoic acid CH３COOH
VERY VERY IMPORTANT: Silicon dioxide SiO2 Ethanol C2H５OH
You must LEARN the VALENCIES of the List of common covalent compounds
COMMON ELEMENTS, RADICALS and IONS
if you wish to progress any further with your Chemistry!!!! 20
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CHEMICAL EQUATIONS STEPS FOR BALANCING THE EQUATION:
A Chemical equation is a way of summarizing a chemical reaction. Although it can  There must be the same number of atoms on both sides, so that all atoms
be written in words, an equation is often written using chemical symbols and a are accounted for and none are lost or gained.
chemical formula. When you write chemical equations, to start with, it is advisable  Only balance by putting a number IN-FRONT of the formulae where
to follow these steps: needed i.e. you can not change the chemical formulae of a given substance

Steps for making chemical equations

1. Identify reactants and products in reaction 1. Find the total number of atoms on each side of the equation
and write down the equations in words [EXAMPLE]
using either the information given Reactants are the chemical elements Na + O2  Na2O
or compounds that a chemical
or your own chemical knowledge.
reaction starts with.
Left-side Right-side
Products are the chemical elements
[EXAMPLE] or compounds that are produced Na 1x1=1 1x2=2
during a chemical reaction. O 1x2=2 1x1=1
Reactants: sulphuric acid, sodium hydroxide
Products : sodium sulphate, water 2. Find an element that doesn’t balance and pencil in a number
(IN-FRONT!!!) to try and correct it
sulphuric acid + sodium hydroxide  sodium sulphate + water  For odd and even in-balances, try swapping the numbers

2. Put every substance into the correct chemical formula. 2Na + O2  Na2O

H2SO4 + NaOH  Na2SO4 + H2O Left-side Right-side

3. Balance the equation. (Changing the proportions of reactants and products, in Na 2x1=2 1x2=2 OK!
such a way that the number of atoms of each element is the same on both sides.) O 1x2=2 1x1=1 NO

H2SO4 + 2NaOH  Na2SO4 + 2H2O

4. Finally, put state symbols in the equation for every reactant and product.
3. See if this works, it may create another in-balance but pencil in another
H2SO4(aq) + 2NaOH(aq) -> Na2SO4 (aq) + 2H2O(l) number (IN-FRONT) and see if this works

state symbols 2Na + O2  2Na2O

(s) solid / precipitate
(l) liquid / molten Left-side Right-side
(g) gas / vapour Na 2x1=2 2x2=4 NO
(aq) aqueous O 1x2=2 2 x1=2 OK!
Most pupils get injured at the balancing
/dissolved in water solution stage! Let’s look at it in details!! Balance just ONE type of atom at a time!!!
/ dilute
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 4. Continue chasing unbalanced elements and the equation will balance eventually GOING FROM A WORD EQUATION TO A BALANCED EQUATION
Remember, if the numbers are not working, rub them out and try again
2↓ IMPORTANT POINTS TO REMEMBER:
4Na + O2  2Na2O
 Most non-metals form molecules, e.g. H2, O2, Cl2, F2, N2 (not C, S or Si)
 Metals do not form molecules!!! e.g. Al, Cu, Mg, Fe
Left-side Right-side
 Better to memorize Formulae of common compounds [e.g.] H2O, CO2,
Na 4x1=4 4x2=4 OK!
ammonia NH3, methane CH4, hydrochloric acid HCl, sulphuric acid
O 1x2=2 2 x1=2 OK!
H2SO4, nitric acid HNO3
Congratulation!!
 Use the valency to find the formulae of (unknown) compounds

 Always double-check your answer that every element is balanced [EXAMPLE]

magnesium + hydrochloric acid  magnesium chloride + hydrogen
[EXAMPLE] 3H2 + N2  2NH3
1.Write the FORMULAE for each compound
H: 2x3=6 3x2=6 (try swap) Mg + HCl  MgCl2 + H2
N: 2 1x2=2 OK!
2.BALANCE the equation
Mg + 2HCl  MgCl2 + H2

Beginners often find it difficult to balance chemical Left-side Right-side

equations, especially the more complicated ones. A Mg 1x1=1 1x1=1 OK!
few rules to bear in mind H 2x1=2 2 x1=2 OK!
 If an equation cannot be balanced, it may be wrong. Cl 2x1=2 1x2=2 OK!
Either the formulae of one or more of the substances involved is /are It’s completed!
written wrongly or there may be missing/extra substances in the
equation.
 Never change the chemical formula of compounds when balancing Now let’s try whole the steps!!
equations. You can only add numbers in front of the chemical formula.
For example, 2NaOH and Na2OH has different meanings. 2NaOH means GOING FROM A “WORDY” QUESTION TO A BALANCED EQUATION:
you have 2 units of NaOH (=2 Na, 2O and 2H), while Na2OH means you
have 2 Na, 1O and 1 H  Identify ALL the REACTANTS and PRODUCTS from the question
(underline each compound referred to in question)
 Write the word equation: REACTANTS  PRODUCTS
The Second dangerous point is putting  Write the symbol equation (remember molecules, metals, common formulae
Remember: PRACTICE, PRACTICE,
chemical PRACTICE!!!
formulae!!! Let’s look at it! and to use the valency)
(find balanced equations in text books and try to balance them on your own)
 Balance the equation
 Put state symbols
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[EXAMPLE] IONIC EQUATIONS
Sodium chloride solution is formed by titrating aqueous sodium hydroxide with Ionic equations are used when a chemical reaction involves the coming together of
dilute hydrochloric acid. Water is also formed. ions in solution.
Write the balanced chemical equation for the reaction including state symbols. IONIC EQUATIONS show only the changes taking place in a chemical reaction

WORKING OUT: Formation of Ions from a chemical formula

REACTANTS: sodium hydroxide, hydrochloric acid
PRODUCTS: sodium chloride, water  Usually, only (aq) compounds split to form ions
Every reactant on the left-hand side
Every product on the right-hand side
 First ion is + ve (metal, H+ or NH4+)
Second ion is – ve (non-metals and radicals)
sodium hydroxide + hydrochloric acid  sodium chloride + water  Charge on Ion = Valency
 Radicals are unchanged i.e. stay together
NaOH + HCl  NaCl + H2O  SMALL number denoting number present moves to
Left-side Right-side a BIG number in-front
Na 1x1=1 1x1=1 OK! ve: a Valence Electron is an
 No. of + ve charges = No. of - ve charges electron found in the outermost
O 1x1= 1x1=1 OK!
electron shell of an atom.
H 1 x1 + 1 x 1 =2 1x2=2 OK! STEPS:
Cl 1x1=1 1x1=1 OK! 1 Make sure you have a BALANCED chemical equation
[EXAMPLE]
already balanced! Hydrochloric acid + calcium carbonate  calcium chloride + carbon dioxide + water

2HCl (aq) + CaCO3 (aq)  CaCl2 (aq) + CO2 (g) + H2O (l)
Note that the ‘liquid’ state and the ‘aqueous’ state is not the
same. The ‘liquid’ state of a substance is pure. For a solid
2 Split only soluble ionic compounds [(aq) compounds] into its ions.
substance, the liquid state is obtained by heating the substance
 Insoluble ionic compounds, elements and covalent compounds remain unchanged.
until it melts, while the ‘aqueous’ state of a substance is obtained
by dissolving it in water.
2H+(aq) + 2Cl- (aq) + Ca2+(aq) + CO32- (aq)  Ca2+ (aq) + 2Cl- (aq) + CO2 (g) + H2O (l)

ANS: NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l) 3 Cancel out spectator ions
Spectator ions are the ions that appear in both
the left and right side of the equation. electron
found in the outermost electron shell of an atom.
No score is given if you cannot give a balanced equation, even though 2H+(aq) + 2Cl- (aq) + Ca2+(aq) + CO32- (aq)  Ca2+(aq) + 2Cl- (aq) + CO2 (g) + H2O (l)
the formulae of the compounds in your equation are correct. State
symbols are not necessary in your balanced chemical equation unless the 4 Rewrite the equation without spectator ions.
question requires it.
2H+ (aq) + CO32- (aq)  CO2 (g) + H2O (l)
AGAIN: PRACTICE, PRACTICE, PRACTICE!!! This is IONIC EQUATION!!!
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[EXAMPLE] 4.2 Stoichiometric calculations
Copper metal is displaced from its solution when an iron nail is placed into a
solution of blue copper (II) sulphate. The clear solution which remains after the RELATIVE MASSES
reaction is complete is iron (II) sulphate.
Derive the ionic equation for this reaction. RELATIVE ATOMIC MASS

WORKING OUT:
Relative atomic mass is the average
REACTANTS: Copper (II) sulphate, iron
mass of a large number of atoms of a
PRODUCTS : copper, iron (II) sulphate The symbol for relative
particular element.
atomic mass is Ar.
Copper (II) sulphate + iron  copper + iron (II) sulphate
All naturally occurring elements are mixture of isotopes and therefore the
CuSO4 (aq) + Fe(s)  Cu (s) + FeSO4 (aq) BALANCED relative atomic mass of an element takes into account the percentage of various
isotopes that may be present.
Cu2+ + SO42- (aq) + Fe(s)  Cu (s) + Fe2+ + SO42- (aq) SPLIT / CANCEL
Ar is simply the average of the mass numbers for each of the isotopes
IONIC EQUATION: Cu2+(aq) + Fe(s)  Cu (s) + Fe2+(aq) present in the element.

Compounds that are sparingly soluble or very Ar = Sum for each Isotope { % Present x Mass Number }
sparingly soluble can be considered as insoluble
when writing ionic equations involving them.
[EXAMPLE]
When you construct ionic equations, the number of each Chlorine gas is 75% chlorine-35 atoms and 25% chlorine -37 atoms.
particle and the total charge must be the same on both sides
of the equation. Ar = (75% X35) + (25% X 37) = 35.5 You can find Ar for each
element in Periodic Table.
[EXAMPLE] The relative atomic mass is a ratio and therefore has no unit.
Cl2 + 2Br-  Cl- + Br2
is not a balanced ionic equation since the total charge on
the LHS is 2- while the total charge on the RHS is only 1-. There is a clear distinction between mass number and relative
The balanced ionic equation will be atomic mass: the mass number of an atom is the number of
protons and neutrons in the nucleus of the atom. It is ALWAYS
Cl2 + 2Br- - 2Cl- + Br2 a whole number. The relative atomic mass of an element is the
where the total charge on both sides of the equation is 2- average mass of its atoms compared to the mass of a Carbon-12
atom.

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 RELATIVE MOLECULAR MASS
If the calculations are correct, the total percentages of
all the elements present in a compound should add up to
Relative Molecular Mass is the
100%. Hence in example, % mass of carbon present is
Sum of the Ar for all atoms The symbol for relative calculated simply as 100 – 72.7 = 27.3%
present in the molecule molecular mass is Mr.

[EXAMPLE] [EXAMPLE 2]
Calculate the relative molecular mass of chloroform CHCl3 Calculate the percentage by mass of water in sodium carbonate crystals, Na2CO3 10H2O

Mr (CHCl3) = Ar (C) + Ar (H) + Ar (Cl) = 12 + 1 + 3×35.5 = 119.5 10  (2  1  16)

% H2O in NaCO3 10H2O   100%
2  23  12  3  16  10  (2  1  16)
[EXAMPLE 2]  62.9%
Calculate the relative molecular mass of copper (II) sulphate crystals, CuSO4 5H2O
The percentage by mass can also be used to calculate the mass of an element
Mr(CuSO4 5H2O) = Ar (Cu) + Ar (S) + 9×Ar (O) +10×Ar (H) in a given sample.
= 64 + 32 + 9×16 + 10×1
=250
Formula
The relative molecular mass is a ratio and therefore has no unit. Mass of element in sample  % of element in compound  sample mass
Mass of element in compound
  sample mass
Mr of comound
PERCENTAGE MASS OF AN ELEMENT IN A COMPOUND
The percentage by mass of an element present in a compound is fixed.
This percentage can be calculated using the formula [EXAMPLE]
Calculate the mass of copper in 32g of copper (II) sulphate.
Formula
mass of element in compound mass of Cu in CuSO4 64
% Mass of element in compound   100 Mass of Cu   32g   32g  12.8g
Mr of compound Mr of CuSO4 64  32  4  16

No. of atoms  Ar of element

  100 THE MOLE CONCEPT
Mr of comound  The “MOLE” is just a special number for particles (atoms, molecules, ions, etc).
ONE mole = 6 x 1023 particles.
The number is called the AVAGADRO CONSTANT
[Example] Calculate the percentage by mass of oxygen in carbon dioxide, CO2 The short form for
1 mol of carbon atom contains 6 x 1023 carbon atoms the mole is mol

2  16 1 mol of sodium ions contains 6 x 1023 sodium ions

% O in CO2   100%  72.7%
12  2  6 1 mol of water molecules contains 6 x 1023 water molecules

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 MOLAR MASS [EXAMPLE 3]
The molar mass of a substance is the mass of 1mol (6 x 1023) of the substance. How many moles does 66g of carbon dioxide contain?
66 1.5mol
No. of moles of Carbon dioxide   1.5
12  2  16
The mass of 1 mol of atoms equals to its relative atomic mass in grams.

Use a formula triangle to help remember formulae. Mass

[EXAMPLE] If the number of moles is required, put your palm
Ar (Al) = 27
The molar mass of aluminium is 27g over the term ‘No. of moles’ to obtain the equation. No. of Molar
Ar (Ar) = 40
The molar mass of argon is 40g moles mass
 The formula can be manipulated to give the mass of a substance, given the
The mass of 1 mol of molecules equals to its relative molecular mass in grams. number of moles present.
The formula is Mass = Number of moles×Molar mass
[EXAMPLE] The formula can be obtained from the formula by covering the term ‘mass’ in the
The molar mass of a sodium chloride (NaCl) is formula triangle.
23 + 35.5 = 58.5g Ar (Na) = 23
The molar mass of hydrogen (H) is Ar (Cl) = 35.5 [EXAMPLE]
2×1 = 2g What is the mass of 0.1 mol of carbon dioxide molecules?
Molar mass of CO2 = 12 + 2 x 16 = 44g
Mass = 0.1 x 44 = 4.4g
CALCULATING THE NUMBER OF MOLES [EXAMPLE 2]
The number of moles of a substance can be obtained using this What is the mass of 3 x 1023 atoms of copper (II) sulphate?

Numbler of particles in sample

No. of moles = 3  10  0.5mol
23
Formula Number of moles  
6  1023
mass of substance 6  1023
Number of moles   Molar mass of CuSO4 = 64 + 32 + 4 x 16 = 160 g
molar mass of substance
Mass = 0.5 x 160 = 80 g
[EXAMPLE]
How many moles of atoms are there in 10g of calcium atoms?
10  0.25mol
No. of moles   0.25
40

[EXAMPLE 2]
How many moles are present in 32g of sulphur dioxide, SO2?
32  0.5mol
No. of moles of sulphur dioxide   0.5
32  2  16

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EMPIRICAL FORMULAE [EXAMPLE 2]
30g of silicon oxide contains 14g of silicon. Find the formula of the compound.
An empirical formula is a chemical formula that shows
the simplest ratio between the atoms in a molecule. Mass of O in the compound is 30 – 14 = 16g
Si O
Mass (g) 14 16
IMPORTANT POINTS TO REMEMBER:
Molar mass (g) 28 16
14/28= 0.5 16/16= 1
The formula of a compound shows the ratio of each type of atom in the compound. No. of mol
-For example
Na2O means 2 atoms of Na combine with 1 0.5/0.5 = 1 1/0.5 = 2
Mol ratio
atom of O to form the compound. If we count ATOMS MOLES
using the mol instead, the formula will show  The empirical formula is SiO2
that 2 mol of Na atoms combine with 1 mol Ratio is the same
of O atom to form the compound.
In most cases, the mol ratio obtained in the last step contains
whole numbers or nearly whole numbers such as 1.96 which can be
Using this idea, you can form the empirical formula. easily rounded off to a whole number. However if the ratio
obtained contains numbers such as 2.5, do not round it off but
[EXAMPLE] use a multiple instead. For example, if the ratio is 1:2.5, then the
What is the empirical formula of a compound with composition of 80% copper and ratio 2:5 should be used in the empirical formula instead of 1:3.
20% sulphur?
Assume the mass
of the sample
Cu S compound is 100g The relationship between Molecular formula and Empirical formula
Step 1 Mass (g) 80 20
Step 2 Molar mass (g) 64 32
The molecular formula shows the actual number of each kind of atom in a
Step 3 No. of mol 80/64 = 1.25 20/32 = 0.625
compound. It is a multiple of the empirical formula.
Step 4 Mol ratio 1.25/0.625 = 2 0.625/ 0.625 = 1 No of mol
= mass /Molar
mass Formula
Divide each No. of mole so as to get simplest pair of Molecular formula = (Empirical formula)n where n = 1,2,3…
integers by the less one.

To find the value of n, we make use of the equation

From calculation, the simplest ratio of Cu to S is 2 to 1.

 The empirical formula is Cu2S

Formula
Relative molecular mass = n x relative empirical formula mass

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[EXAMPLE] CALCULATING THE MASS AND THE VOLUME OF GASES
The empirical formula of a compound is C2H4O. Its relative molecular mass is 88. One mole of any gas has a volume of 24dm3(24 000cm3) at room
find the molecular formula. temperature(25℃) and standard atmospheric pressure (1 atm).
This volume is called the molar volume of a gas.
1. Let the molecular formula be (C2H4O) n The volume occupied by 1mol of any
2. Find n: After finding the value of -The volume of gases is gas is constant, while the mass of
1 dm3 = 1000 cm3
n, do not leave the formula measured using either the 1mol of substance is variable.
Molecular mass = n x empirical formula mass 1 dm3 = 1 litre
88 = n x ( 2 x 12 + 4 x 1 + 16 ) as (C2H4O) 2 but expand to dm3 or cm3
1 cm3 = 1 ml
n = 88/44 = 2 give the formula C4H8O2
3. Put n:
The molecular formula is (C2H4O) 2 = C4H8O2 Here is the formula to calculate the number of moles of gas at room temperature and
pressure
[EXAMPLE 2] Formula
A compound of carbon and sulphur has a composition of 15.8% carbon and 84.2% sulphur. volume of gas in cm3 at r.t.p.
(a) Find the empirical formula Number of moles of gas 
24000cm 3
volume of gas in dm3 at r.t.p.
C S 
24dm 3
Mass (g) 15.8 84.2
Molar mass (g) 12 32
[EXAMPLE]
No. of mol 15.8/12 = 1.32 84.2/32 = 2.63
How many moles are there in 0.08dm3 of hydrogen gas at r.t.p.?
1.32/1.32 = 1 2.63/1.32 = 1.99
Mol ratio 0.08  1000
≈ 2 Number of moles of gas   0.003  0.003mol
24000

 The empirical formula is CS2 [EXAMPLE 2]

In an experiment, when hydrochloric acid was reacted with calcium carbonate at
(b) The relative molecular mass of the compound is 76. Find the molecular formula. room temperature and pressure, 48cm3 of carbon dioxide gas was produced.
Calculate the number of carbon dioxide molecules evolved.
1. Let the molecular formula be (CS2) n
2. Find n: It is possible for the value
48
Molecular mass = n x empirical formula mass of n to be 1. In this case, No. of moles of CO2 molecules =  0.002
24000
76 = n x ( 12 + 2 x 32 ) the empirical formula is
n = 76/76 = 1 also the molecular formula No. of CO2 molecules = 0.002 x 6 x 1023 = 1.20 x 1021
3. Put n: of the compound.
The molecular formula is (CS2) 1 = CS2

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 A formula triangle can also be used to remember this formula CALCULATION INVOLVING CHEMICAL EQUATION

The formula triangle can be arranged to calculate Most mole calculations can then be IMPORTANT MOLE FACTS
the volume of a gas, given the number of moles. calculated by using the idea of ratio 1 mole = 6 x 1023 particles
using the above facts and / or using 1 mole = Mr in g
Volume of gas ( in dm3) = Number of moles x 24 dm3 the formula n = m/Mr 1 mole of a GAS = 24 dm3 at r.t.p.

[EXAMPLE]
Calculate the volume of 0.016g of methane, CH4, under room conditions. Note: There are many other ways to solve mole calculations. If you are
confused, try looking up other ways in different text books to find a
0.016 method which suits you. You will only know by doing lots and lots of
No. of moles of CH4   0.001
12  4  1 PRACTICE!!!
Volume of CH4=0.001 x 24 = 0.024dm3
Calculations form chemical equations involves five main steps, although not all the
[EXAMPLE 2] steps are required in each set of calculations.
Calculate the mass of 1.5dm3 of carbon monoxide, CO, under room conditions.
[EXAMPLE]
Calculate the mass of water produced when 4.0g of methane is completely burnt in oxygen.
No. of mol of CO = 1.5 = 0.0625
24
Mass of CO = 0.0625 x (12 + 16) = 1.75g Step 1 Write a balanced chemical equation for the reaction and define the
problem to be solved by writing down all of the information that is given
Note: The volume of 1 mol of gas is 24 dm3 only at 25℃ (room in the question.
Make sure all the data
temperature). If the gas is cooled, the volume will decrease. CH4 + 2O2  CO2 + 2H2O for each compound
Then at 0℃ the volume of the gas is 22.4 dm3. 4.0g(given) ?(to be found) LINES UP with the
0℃ is called standard temperature. Step 2 Calculate the No. of mol of the substance formula in the chemical
whose mass is given. equation of the reaction.
[Example 3]
What is the volume of 66g of carbon dioxide at room temperature and pressure? mass of methane 4.0
  0.25
molar mass of methane 12  4  1
① Find no. of moles: Step 3 Write down the relevant mole ratio from the chemical equation.
m = 66 g Mr[CO2} = 12 + 2 x 16 = 44 n=? mol CH4 1

n = m / Mr = 66 / 44 = 1.5 moles mol H2O 2
② Find volume: Step 4 Calculate the No. of mol of the substance to be found.
1 mole = 24 dm3 2
mol CH4 1  mol H2O   mol CH4
1.5 moles = x dm3 { cross-multiply }  1
 x = 1.5 x 24 / 1 = 36 dm 3 mol H2O 2 2
mol H2O   0.25  0.5
The volume of 66g of CO2 is 36 dm3 at r.t.p. 1
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Step 5 Calculate the required mass (or volume) Step 1
2NaHCO3(s)  Na2CO3 (s) + CO2(g) + 2H2O(g)
No. of mol x molar mass of H2O Don’t forget the 16.8g(given) ?(to be found)
UNITS in your final
= 0.50 x (2 x 1 + 16)
answer!!! Step 2
= 9.0g
16.8 16.8
  0.2
23  1  12  3  16 84
 In a chemical equation, the balancing numbers IN-FRONT show how many
molecules or atoms take part in the chemical reaction. These numbers also show Step 3
HOW MANY MOLES take part in the reaction. mol Na2 CO3 1

mol NaHCO3 2
Step 4
You can interpret the chemical equation 2
mol Na2 CO3 1  mol Na2 CO3   mol NaHCO3
CH4 + 2O2  CO2 + 2H2O  1
As mol NaHCO3 2 2
mol Na2 CO3   0.2  0.1
1 molecule of CH4 reacts with 2 molecules of O2 to produce 1 molecule of 1
CO2 and 2 molecules of H2O Step 5
Even as No. of mol x molar mass of Na2CO3 = 0.1 x (2 x 23 + 12 + 3 x 16)
1 mol of CH4 reacts with 2 mol of O2 to produce 1 mol of CO2 and 2 mol of H2O = 0.1 x 106
= 10.6g

[EXAMPLE 3]
What mass of magnesium oxide is produced when 60g of magnesium burns in air?
It is important to identify between which and which the mole ratio is
needed to find the answer.
Step 1 2Mg + O2  2MgO
Then only the mol ratio between the substance given and the substance
60g (given) ?(to be found)
to be found is required.
By defending the problem in step 1, we will know which mol ratio is Step 2 60
needed to help solve the problem with as ease as possible.  2.5
24
Step 3 mol MgO 1
 1
[EXAMPLE 2] mol Mg 1
Calculate the mass of the solid product obtained when 16.8g of sodium hydrogen  mol MgO  1  mol Mg
Step 4 mol MgO
carbonate is heated strongly until there is no further change. The equation for the 1
mol Mg mol MgO  1  2.5  2.5
reaction is
2NaHCO3(s)  Na2CO3 (s) + CO2(g) + 2H2O(g) Step 5
No. of mol x molar mass of MgO = 2.5 x (24 + 16) = 2.5 x 40 = 100g

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[EXAMPLE 4] Step 1
Limestone decomposes when heated as shown in the equation. Which mass of Mg + 2HCl  Mg Cl2 + H2
limestone is needed to produce 84g of calcium oxide? 0.12g (given) ? (to be found)

Step 2
Step 1
0.12
CaCO3  CaO + CO2  0.005
24
? (to be found) 84kg (given)
Step 3
Step 2 mol H2 1
 1
84000 mol Mg 1
 15000
40  16 Step 4
Step 3 mol H2 1  mol H2  1  mol Mg
mol CaCO3 1  1
 1 mol Mg 1 mol H2  1  0.005  0.005
mol CaO 1 Step 5
Step 4
No. of mol x 24dm3(unit volume of gas) = 0.005 x 24
mol CaCO3  mol CaCO3  1  mol CaO
1 = 0.12dm3 (= 120cm3)
mol CaO mol CaCO3  1  1500  1500
Step 5
No. of mol x molar mass of CaCO3 = 1500 x (40 + 12 + 3 x 16) The calculation steps in this problem are similar to those in previous
= 150000g (= 150kg) examples until the last step where volume instead of mass was
required. The formula to use to calculate volume is
Volume = number of moles x 24 dm3
Be careful of the units used in the equation. The molar mass is in
grams. Hence, when calculating the number of moles of a
substance, the mass of the substance must also be in grams. [EXAMPLE 6]
What volume of hydrogen is required to react with 48 litres of oxygen to
produce steam?
2H2 (g) + O2(g)  2H2O(g)
 In order to calculate volume, you can take similar steps to those in previous
examples until last step. Step 1
2H2 (g) + O2(g)  2H2O(g)
[EXAMPLE 5] ?(to be found) 48dm3 (given)
Magnesium reacts with hydrochloric acid as shown. What it’s the volume of
hydrogen produced at r.t.p. when 0.12g of magnesium reacts? Step 2 1 litre = 1 dm3
48
2
Mg + 2HCl  Mg Cl2 + H2 24

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Step 3 (c) S is the excess reactant in this reaction. 0.1 mol of S will react Fe, leaving
mol H2 2 (0.2 – 0.1 = 0.1) mol of S behind.

mol O2 1
Step 4 Mass of excess S = 0.1 x 32 = 3.2g
2
mol H2 2  mol H2   mol O2
 1
mol O2 1 2 In part (b) of example, the mass of product was calculating using
mol H2   2  4
1 the number of moles of Fe, the limiting reactant. Remember that
Step 5 once the limiting reactant is used up, no more reaction takes
No. of mol x 24dm3(unit volume of gas) = 4 x 24 place. The amount of product formed will depend on how much
= 96dm3 (= 96liters) limiting reactant was present in the reaction.

CALCULATIONS INVOLVING CONCENTRATIONS

Sometimes the question requires the student to identify the limiting reactant
as part of the solution to the calculation. The limiting reactant is the reactant - The concentration of a solution
The concentration of a solution is the
that, once used up, will cause the reaction to stop. is given in either g/dm3 or mol/dm3
amount of a solute dissolved in a unit
volume of the solution
[EXAMPLE]
5.6g of iron burns in 6.4 g of sulphur to form iron(II) sulphide. To calculate concentration, the following formulae can be used
(a) Calculate and find out which reactant is the limiting reactant
(b) Calculate the mass of iron(II) sulphide formed after the reaction.
(c) Calculate the mass of the excess reactant that is left after the reaction.
Formula Mass
Mass of solute (g)
Concentration (g/dm ) 
3 （/Moles）
- WORKING OUT - Volume of solution (dm 3 )
(a) No. of moles of solute
Concentration (mol/dm3 ) 
5.6 Volume of solution (dm 3 ) Concentration Volume of
No. of mol of Fe   0.1
56 Fe + S  FeS (/Molarity) solution
6.4
No. of mol of S   0.2
32 [EXAMPLE]
From equation, 1 mol of Fe reacts with 1 mol of S. This means 0.1 mol of Fe will A solution of glucose contains 0.45g of glucose in 75cm3 of solution. What is the
react with 0.1 mol of S. Hence the limiting reactant is Fe. concentration of the glucose solution in g/dm3?

(b) 0.45g
Concentration   6.0 g/dm3
0.1 mol of Fe will produce 0.1 mol of FeS 75
dm 3
Mass of FeS = 0.1 x (56 + 32) = 8.8g 1000
 6.0g/dm3

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[EXAMPLE 2] MOLARITY OF A SOLUTION
What is the concentration in mol/dm3 of a solution of sodium hydroxide if it
contains 3.5g of sodium hydroxide in 100 cm3 of solution? The MOLARITY of a solution is - Concentration in mol / dm³ is
a common unit for concentration also called MOLARITY.
Mass 3.5
No. of moles of NaOH    0.0875
Molar mass 23  16  1
0.0875mol From the formula for the concentration, the formula for the Molarity is
Concentration   0.875 mol/dm3
100
dm 3

1000
 0.875mol/dm3
Formula
[EXAMPLE 3] Concentration (mol/dm3 ) 
The dilute sulphuric acid used in school laboratories usually has a concentration No. of moles of solute
Molarity 
of 2 mol / dm3. What is the mass of H2SO4 in 250cm3 of this acid? Symbol for Molarity is M Volume of solution (dm 3 )
(= mol / dm³)
No. of moles  Concentration (mol/dm3 ) x volume (dm3 )
2mol 250
  dm3  0.5
dm3 1000
Mass of H2SO4 = 0.5 x (2 x 1 + 32 + 4 x 16) = 0.5 x 98 = 49g [EXAMPLE]
What volume of water (in cm³) should be added to 4 g of sodium hydroxide to
[EXAMPLE 4] make a 0.2 M solution?
A beaker containing 500 ml of water is used to dissolve 117 g of common salt.
What is concentration in mol / dm3 of the solution? DATA: m = 4 g c = 0.2 M
１ Calculate number of moles: Mr[NaOH] = 23 + 16 + 1 = 40
DATA: V = 500 ml m = 117 g  n = m / Mr = 4 / 40 = 0.1 moles
WORKING OUT: { Common salt = sodium chloride } ２ Calculate volume in dm³: c = n / V and M = mole / dm³
１ Find no. of moles: m = 117 g Mr[NaCl] = 23 + 35.5 = 58.5  V = n / c = 0.1 / 0.2 = 0.5 dm³
n = m / Mr = 117 / 58.5 = 2 moles ３ Change dm³ to cm³:
２ Find volume in dm3: 1000 cm³ = 1 dm³
1 dm3 = 1000 ml { 1 dm3 = 1 L = 1000 ml} x cm³ = 0.5 dm³ { cross – multiply }
x dm 3
= 500 ml { cross-multiply }  x = 0.5 x 1000 / 1 = 500 cm³
x = 1 x 500 / 1000 = 0.5 dm3
３ Calculate concentration in mol / dm3: Answer: 500 cm³ of water is needed to dilute 4 g of NaOH to 0.2 M
c = n / V = 2 / 0.5 = 4 mol / dm3

Answer: The concentration of the salt solution is 4 mol / dm3.

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 There are 103 elements discovered at present.
Periodic Table can help you classify them and easy to know properties of each
element.
COTENT LEARNING OBJECTIVE (Pupils should be able to)
 Describe the Periodic Table as a method of classifying
elements and its use to predict properties of elements 5.1 Periodic trends
 Describe change form metallic to non-metallic character
Periodic Trends
across a period In Periodic Table Elements are arranged in order of
 Describe the relationship between Group number, number increasing proton number (atomic number)
of valence electrons and metallic or non-metallic character.
 Describe lithium, sodium and potassium in Group I as a
collection of relatively soft metals showing a trend in －A horizontal row in the Periodic Table is known as a Period
melting point and in reaction with water There are 7 Periods in the Periodic Table.
 Predict the properties of other elements in Group I, given －A vertical column in the Periodic Table is known as a Group
data, where appropriate. There are 8 groups in the Periodic Table.
 Describe chlorine, bromine and iodine in Group VII as a
Group properties collection of diatomic non-metals showing a trend in colour, The Periodic table are divided into sections as shown below
state, and in their displacement reactions GROUP I II III IV V VI VII VIII
 Predict the properties of other elements in Group VII, PERIOD /O
given data, where appropriate.
 Describe the noble gases as being inert. 1 H
 Describe the uses of the noble gases in providing an inert
atmosphere 2

3
NON - METALS
4

In the last page of this text book 5 METALS

PERIODIC TABLE is put as an appendix
6

Transition Semi - Metals

Reactive Metals Metals

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METALS AND NON-METALS  Form positively charged ions with variable charges
A ‘zig- zag’ diagonal line (staircase line) in the Periodic table divides metallic Examples: Copper forms either Cu+ or Cu2+
elements from non-metallic elements. Iron forms either Fe2+ or Fe3+

NON-METALS Some uses of transition metals

Non-metals are elements which do not have the properties of metal and always Many transition metals are often good catalysts in industry to speed up
form the negative ions when they react to form ionic compounds. reactions.
[Examples] The hydrogenation of oil to make margarine uses a nickel catalyst.
- Their states are often gases at room conditions (N, O, F, Cl, noble gases and The manufacture of ammonia uses an iron catalyst
etc) or low melting point solids (P, S, I and etc).
- They are Poor electrical and thermal conductors. Many transition metals are used to make alloys
[Example] Steel is made by mixing iron with a small amount of carbon
METALS - an alloy is a mixture of two or more elements
Metals are a class of chemical elements which always form positive ions when
they react to form compounds. Many transition metals are often useful engineering materials as strong and
hard metals
- They are often shiny solid
- They are good conductors of heat and electricity. SEMI-METALS
- They also form solid oxides that act as bases. Elements near the line (such as Boron and silicon) are called ‘semi-metals’.
Semi–metals have the characteristics of both metals and non-metals.
TRANSITION METALS They are often electrical semiconductors
Transition metals are found in the centre block of Periodic Table. whose physical properties resemble metals
-They are typical metal. but whose chemical properties resemble
- They are hard, strong metals with high melting and boiling points. PERIOD non-metals.
-They also have high density. A period is a horizontal row or elements
- The first 3 rows are called ‘short periods’.
They have partly filled inner electron shells - The next 4 rows, which include the transition metals, are called ‘long periods’.
which give them distinctive properties.
Elements in the same period have the same number of electron shells.
They also have the following properties:
 Form IONS in aqueous SOLUTION which are COLOURED Going across a period from left to right, the number of outermost electrons
Examples: Copper(II)  Blue increases by one every successive element
Iron (II)  pale green
Iron (III)  reddish brown (when solid)
yellow ( when in solution)

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 It becomes easier for an element to lose electrons going down a group. With an
=> Elements change from metallic to non-metallic character across a period. increase in the number of electron shells, the attraction between the positively
charged nucleus and the valence electrons are reduced. The elements become more
metallic in character.
METALS Semi NON-METALS
Reactive  Less reactive Metals Less reactive  Reactive
The element hydrogen is unique because a H atom can form
either H, by losing its one valence electron, or H, by gaining one
GROUP valence electron to complete its outer shell. Forming ions with
A group is a vertical column of elements 1+ charge is characteristic of Group I elements, while forming
Not the same total number of electrons
ions with 1- charge is typical of Group VII elements. This
Elements in the same group have the same number of outer shell electrons explains why hydrogen is placed by itself in the Periodic Table.
(valence electrons). This means that elements in the same group will have similar
chemical properties since they will form ions with the same charge.
SHEILDING
They will also form compounds with similar formulae.
When we talk about the reactivity of elements, sometimes we use this idea.

The group number is the same as the number of outer shell electrons Reactivity changes as you move down the Groups due to shielding.
 Valency of metals = Group number This is because each new e- shell is further out from the nucleus and the inner
Valency of non-metals = 8 – Group number e- shells shield the outer e-‘s from the positive nucleus.

As METAL atoms get bigger, the outer e- is more easily lost. This makes
Going down a group form top to bottom, the number of electron shells increases
METALS MORE REACTIVE as you go DOWN Groups I and II.
by one for every consecutive element

As NON-METAL atoms get bigger, the extra e- are harder to gain. This
=> Elements become more metallic in character, i.e. they lose valence
makes NON-METALS LESS REACTIVE as you go DOWN Groups VI and VII.
electrons more easily.

METALS NON-METALS
5.2 Group properties They all react with water to
As you go down a group As you go up a group form alkalis, hence their name.
Group I: ALKALI METALS
Elements in Group I are also known as alkali metals that are the elements in the
More reactive More reactive first group in the periodic table, which all have a single valence electron.

They are the most reactive metals in the Periodic Table.

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 PHYSICAL PROPERTIES CHEMICAL PROPERTIES REACTION OF GROUP I ELEMENTS WITH WATER
 Silvery / white in colour  All are VERY REACTIVE Group I elements become more reactive down the group.

 SOFT and EASY TO CUT  REACT vigorously with COLD WATER to These metals are stored under oil or in vacuum to prevent them
with a knife from reacting with water and/ or oxygen in the air.
 LOW DENSITIES FORM H GAS
[Example] Reaction of Group I elements with water
 relatively LOW MELTING  BURN in AIR with COLOURED FLAMES to
POINTS form OXIDES Li reacts violently
 Good conductors of heat  Alkali metals REACT with HALOGENS to Na reacts very violently, sometimes with an explosion
and electricity PRODUCE a NEUTRAL SALT which K reacts explosively
 All have 1 e- in outer shell dissolves to form a colourless solution

The compounds of Group I metals are all ionic. Group - Group I elements react with cold water to form metal hydroxides and
I metals always form ions with 1+ charge in their hydrogen gas.
compounds. Group I metal + water  metal hydroxide + Hydrogen gas
2Na(s) + 2H2O (l)  2NaOH (aq) + H2 (g)
- Metals in general are hard, dense with high melting and boiling points. Group I
metals are highly unusual because they are soft, easily cut and have low density
and low melting points.
The compounds of Group I elements
The first 3 elements in the
are usually colourless unless the
Group I: Alkali Metals – BEHAVIOUR TRENDS group can float on water
compound contain a transition element The metal hydroxide solutions formed
ATOMIC ALKALI Density Melting TRENDS: are all strong alkalis with pH values
NUMBER METALS 3
(g/cm ) point( Cº) more than 7.
3 Li Lithium 0.53 180 -Reactivity
11 Na Sodium 0.97 98 increase
19 K Potassium 0.86 64 -Densities Group VII: HALOGENS
37 Rb Rubidium 1.5 39 increase
1.9 29 -M.P. and B.P Elements in Group VII are also known as halogens that are the elements which
55 Cs Caesium
have seven valence electrons in their outermost shell.
87 Fr Francium - - decrease.
Their ions and compounds are called halides
Some physical properties of G I elements -Softer to cut

Francium is the most reactive They are very reactive non-metals. The name is derived form Greek
metal in the Periodic Table. and means “salt-markers.”

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 Each molecule is made up of 2 atoms DISPLACEMENT REACTIONS
joined together by a single covalent bond.
Group VII elements become less reactive down the group.
PHYSICAL PROPERTIES CHEMICAL PROPERTIES
 Non-metals which form
DIATOMIC MOLECULES with Elements can either be ionic or covalent. More reactive halogens will DISPLACE less reactive
COLOURED VAPOURS halogens from their aqueous salt solutions
 POISONOUS  REACT with METALS to form
(use fume cupboard) IONIC compounds (neutral SALTS)
 Poor conductors of heat and  REACT with another NON- [Example1]
electricity METALS to form MOELCULAR Cl2 (g) + 2NaI (aq)  2NaCl (aq) + I2 (l)
 All have one less e- in outer shell compounds
 form -1 ions Chlorine, being the more reactive halogen, will displace iodine from aqueous sodium
iodide. The reddish brown solution formed contains sodium chloride and aqueous
iodine.
Group VII: Halogens – BEHAVIOUR TRENDS
[Example2]
Fluorine and chlorine are gases; bromine
F2 (g) + 2NaCl (aq)  2NaF (aq) + Cl2 (l)
is a liquid while iodine is a solid.

ATOMIC Melting Boiling TRENDS: Fluorine, being more reactive than chlorine, will displace it from aqueous sodium
HALOGENS State Colour chloride. The yellowish solution formed contains sodium fluoride and chlorine.
NUMBER point(Cº) point(Cº)
9 F2 Fluorine -220 -188 Gas pale -Reactivity
yellow decreases [Example3]
17 Cl2 Chlorine -101 -35 Gas dense -M.P. and B.P Br2 (g) + 2KI (aq)  2KBr (aq) + I2 (s)
green increase.
35 Br2 Bromine -7 59 Liquid Reddish - State Bromine, being more reactive than iodine, will displace it form aqueous potassium
brown Gas to Solid iodide. The reddish brown solution formed contains potassium bromide and iodine.
53 I2 Iodine 114 184 Solid black - Colour
[Example4]
85 At2 Astatine - - Solid - darker
Physical properties of G VII elements
There is no reaction between iodine (I) and sodium chloride (NaCl). Iodine is less
reactive than chlorine and does not displace chlorine from sodium chloride.
Astatine is the least reactive The element Iodine has many colours,
element in Group VII. depending on what physical state it is in.
It is purple in gaseous state, black in
solid state, and forms a reddish brown
States are under Room Conditions solution when dissolved in water. Displacement reactions are redox reactions.

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 USE OF HALLOGENS
 Small amounts of fluorine is added to tap water and toothpaste to prevent Group O: Noble Gases – BEHAVIOUR TRENDS
tooth decay
 Chlorine is used to treat tap water and swimming pools to kill harmful ATOMIC Density Boiling TRENDS:
NOBLE GASES
germs and bacteria NUMBER (g/cm3) point( Cº)
 Iodine is used as an antiseptic; Small amounts of iodine are needed in our 2 He Helium 0.14 -269 -Reactivity
bodies to prevent goitre (swelling of thyroid gland) 10 Ne Neon 0.67 -246 inert
 Silver halides are used on black and white photographic film 18 Ar Argon 1.38 -186 -Densities
2AgBr + LIGHT  Br2 + 2Ag (silver metal deposit) 36 Kr Krypton 2.89 -157 increase
54 Xe Xenon 4.56 -108 - B.P
86 Rn Radon 7.70 -62 increase.
Group VIII/O: NOBLE GASES Some physical properties of G VIII elements
Group VIII elements are also known as noble gases or inert gases that are
extremely inert. Noble gases occupy 1% of the
atmosphere. Of all the noble gases,
Group O: Noble Gases – USES
ELECTRONIC STRUCTURE OF GROUP VIII ELEMENTS Argon is the most abundant in air.
Group VIII elements are the least reactive elements in the Periodic Table  Helium is used in balloons and airships because it is less dense than air (the
because all their outer shells are completely filled. second lightest gas and not flammable like hydrogen)
 Neon is used in advertising signs because it glows red when electricity is
X X
XX
× X
XX discharged through it.

XX X XX
X X  Argon is used to fill filament lamps (light bulbs). It prevents the filament
He Ne
X
X
X X X Ar X X inside the bulb form burning out.
× X X X  Krypton and Xenon are used in lamps in lighthouses, stroboscopic lamps, and
XX photographic flash units.
XX
4 20 40
Helium, 2He Neon, Ne
10 Argon, Ar
18  All these uses are because noble gases are CHEMICALLY INERT!!
The full electronic structure of the first 3 noble gases

NOBLE GASES all have the following PROPERTIES:

The terms ‘unreactive’ and ‘inert’ must be distinguished from each other.
All are colourless
A substance can be unreactive, but given the correct condition, it will undergo
All are gases that consist of single atoms.
reaction to form new substance. If a substance is said to be inert, then it is
All are monatomic gases with Low Melting and Boiling points.
stable and will not take part in reaction no matter what conditions are provided.
All are stable, hardly combine with other atoms (INERT)

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 6.1 Rate of Reaction

Rate of Reaction is the Speed at which the chemical reaction proceeds.

The speed of a chemical reaction refers to how fast reactants are used up or how
COTENT LEARNING OBJECTIVE (Pupils should be able to) fast products are formed in a reaction.
 Describe the effect of concentration, pressure, particle
size and temperature on the speeds of reactions and explain - Different chemical reactions have different speeds.
these effects in terms of collisions between reacting EXAMPLES
Rate of reactions ・ Reaction of potassium metal with water -> very fast
particles.
 Interpret data obtained from experiments concerned with ・ Resting of an iron nail in the presence of air and water -> slow, takes a few days.
speed of reaction. ・ Gold reacting with oxygen in the air -> no reaction, speed of reaction of gold
 Define oxidation and reduction (redox) in terms of oxygen with oxygen is zero.
Redox reaction
gain/loss
 Define exothermic and endothermic reactions MEASURING THE RATE OF REACTION
Energy changes  Describe bond breaking as an endothermic process and bond The speed of a reaction is defined as
forming as an exothermic process
change in amount of reactant or product
speed of reaction 
time

We can measure the speed of reaction by observing either how quickly the
reactants are used up or how quickly the products are forming.
Common methods are shown below.

1. Change in mass (usually gas given off)

 Any reaction that produces a gas can be carried out on a MASS BALANCE
and the mass disappearing is easily measured
2. Volume of gas produced
 Uses a GAS SYRINGE to measure the volume of gas produced
3. Precipitation
 Observe a MARKER through a solution which becomes CLOUDY as the
product precipitates and measures the time taken for the marker to
DISAPPEAR

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 ANALYSING GRAPHS FOR RATE OF REACTION (b) Which of the two reactions was the fastest? Suggest the reason for this
difference. (2 marks)
Amount of Product Made  Graph is steepest at START (c)(i) What was the mass lost for Reaction A after 30 seconds? (2 mark)
 REACTION is FASTEST at start (ii) After what time did Reaction B lose 25 g of mass? (1 mark)
A  Graph then starts to level out (d). Did Reaction A start with less calcium carbonate, more calcium carbonate or
 REACTIONS SLOWS as more the same amount of calcium carbonate as Reaction B? (1 mark)
B REACTANTS gets USED UP
 When graph is flat (level)
 REACTION has FINISHED
 Reaction A is faster than Reaction B ANSWER:
Time (a)
The greater the gradient is, Loss of Mass During Thermal Reaction A
the faster the reaction is. Reaction B
STEPS: How to Plot Graphs Decomposition of Calcium Carbonate
 Use the data range to choose a suitable SCALE for each axis.
 Note that the scale markings should be evenly spaced! 50
 PLOT each data point carefully

Loss of Mass (g)

40
 Draw a LINE OF BEST FIT B at
 LABEL each axis (include units) 30 25 g
 Give the graph a TITLE e.g. Graph of Loss of Mass with Time
20

[EXAMPLE] 10
Calcium carbonate when heated undergoes thermal decomposition to form
0
calcium oxide and carbon dioxide. The loss of mass during the reaction was
0 60 120 180 240 300 360 420 480
measured for two different reactions. A at 30 s
Time (sec) 0 60 120 180 240 300 360 420 480 Time (sec)
A Loss of Mass (g) 50 40 30 22 17 14 12 11 11
B Loss of Mass (g) 50 45 20 15 13 12 11 11 11 (b) Reaction B was faster (as it has the steeper curve). The reason for this is that
Reaction B was heated more strongly than Reaction A.
(a) Plot the results on the graph below. Label each reaction A or B (3 marks) (c) (i) Mass Lost = Original Mass – Mass at 30 s
= 50 g – 42 g = 8 g { 2 marks!! use graph & calculation}
(ii) 150 sec { use graph }
(d) Reaction A and Reaction B started with the same amount of calcium carbonate
because they both ended up with exactly the same amount of mass lost during the
reaction.

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 The reaction is complete once the gradient of the curve becomes zero. 6.2 Redox Reactions
In the above example, the reaction is completed in 420 seconds.
It is incorrect to say that since the reaction is completed in 420 Redox reactions are reactions that involve both oxidation and reduction
seconds, the reaction is half completed at 420/2 = 210 seconds. This is
because the rate of reaction changes with time – it is faster at the Oxidation is a chemical reaction involving the gain of oxygen.
beginning, becomes slower as the reaction proceeds and finally stops. To Reduction is a chemical reaction involving the loss of oxygen.
determine the time when the reaction is half completed, we need to look
at how long it takes for half the amount of reactant to be used. Oxidation and reduction reactions occur simultaneously.
If one reactant is oxidised, then the other reactant must be reduced.

FACTORS AFFECTING THE SPEED OF A REACTION LOSS or GAIN OF OXYGEN

A reaction is caused by the collision of particles of substances. When a substance gains oxygen during a chemical reaction, it is oxidised. If it loses
The factors below affect the collision and therefore affect speed of reaction. oxygen the substance is reduced.

1. Concentration of reactant – more concentrated reactants, faster reactions EXAMPLES

2. Pressure of reactant (gaseous reactions only) – higher pressure, faster reactions
3. Temperature – higher temperature, faster reactions H2 (g) + CuO(s)  Cu(s) + H2O(g)
4. Particle size of reactant – smaller particle reactants, faster reactions H2 is oxidised to H2O because it has gained oxygen.
5. Use of a Catalyst CuO is said to be reduced to Cu because it has lost oxygen.
I.e. INCREASE in SURFACE AREA

CATALYSTS Fe2O3 (s) + 3CO(g)  2Fe(s) + 3CO2(g)

CO is oxidised to CO2 because it has gained oxygen.
 Catalysts work best with a big surface area Fe2O3 is reduced to Fe because it has lost oxygen.
A CATALYST is a substance
[e.g.] powder, pellets or gauze
which INCREASES the speed
 Catalysts are specific to certain reactions
of reaction without being
 Enzymes are biological catalysts
changed or used up in the
 Catalysts are used to REDUCE COSTS in
reaction The definition using oxygen is the easiest to use. However, the
Industrial Reactions
 Catalysts LOWER the ACTIVATION ENERGY use is limited to reactions involving oxygen atoms.
Actually there are other definitions for redox reactions which
Activation energy is the minimum amount of energy required to start a
use gain or loss of hydrogen/electron, or oxidation number.
chemical reaction. This energy is used in the breaking of chemical bonds
EXAMPLES The most versatile definition by far is the one using oxidation
・IRON CATALYST is used to produce AMMONIA in the HABER PROCESS. numbers although we do not talk about those definitions here.
・A PLATINUM CATALYST is used in the production of nitric acid and in the
CATALYTIC CONVERTER in a car engine.
・The catalytic converter is used in the engine of the car to promote combustion
of the fuel to reduce pollution from unburnt exhaust gases.
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DEFINING OXIDISING and REDUCING AGENTS ENDOTHERMIC REACTIONS
Oxidising agents are substances that help oxidation take place. In the process, EXAMPLES
they become reduced. Similarly, reducing agents are substances that help  When ammonium chloride is dissolved in a
reduction take place. In the process, they become oxidised. An endothermic reaction is beaker of water, the temperature of the
a chemical reaction during solution drops from 28Cº to 22Cº
EXAMPLE CuS + 4H2O2  CuSO4 + 4H2O which heat is taken in,  Heat energy must be supplied during the
CuS is oxidised to CuSO4 as it has gained oxygen, and H2O2 is reduced causing a temperature drop thermal decomposition of calcium carbonate.
to H2O because it has lost oxygen. in the surroundings.  Light energy must be absorbed before
Since H2O2 causes CuS to become oxidised (by losing oxygen to it), it is photosynthesis by plants can take place.
the oxidising agent. On the other hand, CuS is the reducing agent since
it causes H2O2 to become reduced (by removing oxygen from it)
In a chemical reaction, bonds between reactants must be broken so
that the atoms can rearrange themselves to form products. Heat
6.3 Energy changes energy must be taken in by the reactants for bond breaking.
Hence, bond breaking must be endothermic.
When chemical reactions take place, energy is either taken in or given out from
the surroundings in the form of heat and/or light. We describe reactions as either
exothermic or endothermic, depending on whether energy is absorbed or given out. EXOTHERMIC REACTIONS ENDOTHERMIC REACTIONS

EXOTHERMIC REACTIONS
 HEAT is GIVEN OUT  HEAT is TAKEN IN

EXAMPLES
 Temperature RISES  Temperature DECREASES
 When sodium carbonate is dissolved in a
(or reaction requires heating)
An exothermic reaction is beaker of water, the temperature o the
 Energy released when forming
a chemical reaction during solution rises from 28Cº to 40Cº
bonds is GREATER than the energy  Energy absorbed when breaking
which heat is given out,  When methane is burnt, heat energy is
absorbed when breaking bonds bonds is GREATER than the energy
causing a temperature rise evolved and the temperature of the
released when forming bonds
in the surroundings. surroundings rises.
[e.g.] Combustion, Freezing and
 When acids react with alkalis, neutralisation
condensing, Neutralisation reactions, [e.g.] Decomposition of limestone,
takes place wit the evolution of heat. The
Haber process, Reduction of iron (III) Decomposition of halide crystals by
temperature of the solution formed rises.
in the blast furnace, Adding light, Melting and boiling,
concentrated H2SO4 to water, Adding Photosynthesis, Dissolving certain
After all the bonds in the reactants are broken, the atoms water to anhydrous CuSO4 salts (KCl or NH4NO3)
will form new bonds to give the products of the reaction; Heat Summary of exothermic and endothermic
energy will be released when these new bonds are formed.
Hence, bond forming is exothermic.

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 For human beings air and water are two of the commonest, indeed, the most
important chemical substances in the world. There are however, other classes of
chemical materials which are not only common but are also very important in our
everyday lives. These classes are the acids, bases and salts which are the subject
matter of this chapter.
COTENT LEARNING OBJECTIVE (Pupils should be able to)
 Describe the meaning of the terms acid and alkali in terms
of the ions they contain or produce in aqueous solution. Acids are chemical compounds which produce hydrated hydrogen
 Describe the characteristic properties of acids as in their ions H+ (aq) when in aqueous solution.
reactions
 Describe the characteristic properties of bases as in their Bases are chemical compounds that react with acids to form a salt
rections with acids and with ammonium salts and their and water
effects on indicator paper.
Acid, Base and
 Describe neutrality and relative acidity and alkalinity Alkalies are water-soluble bases which produce hydrated
Alkali
 Describe the formation of hydrogen and product of the hydroxide ions OH- (aq) when in aqueous solution.
reaction between; reactive metals and water/metals and
acids Salts are chemical compounds formed when the hydrogen of an
 Classify oxides as either acidic, basic, or amphoteric acid is partially or wholly replaced by a metal or other positive ion
related to metallic or non-metallic character. (E.g. ammonium ion).
 Describe and explain the importance of controlling acidity in
soil
 Describe the preparation, separation and purification of
Preparation of salts. 7.1 Acid, Base and Alkali
Salt  Suggest a methods of preparing a given salt from suitable
starting materials, given appropriate information PH SCALE
 Describe the use of aqueous sodium hydroxide and aqueous The pH scale shows the strength of an acid or alkali in an aqueous solution.
ammonia to identify the aqueous cations. It is a measure of the concentration of H+ ions present in the solution.
 Describe tests to identify the anions. pH is an abbreviation for “potential of hydrogen”
 Describe tests to identify the gases.
Identification
 Describe the identification of hydrogen using a lighted
test strong weak weak strong
splint ( water being formed)
 Describe the identification of oxygen using a glowing splint. 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
 Describe the identification of carbon dioxide using lime
water.
ACIDIC ALKALINE

NEUTRAL

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 The pH scale ranges from 0 to 14 ACIDS
 A pH value of less than 7 indicates that the solution is acidic. Acids are substance that will dissolve in water and undergo ionization to form
 A pH value of more than 7 indicates that the solution is alkaline. hydrogen ions. The table below shows some common acids found in the laboratory
 A pH value of 7 indicates that the solution is neutral. and the ions they contain.
It is neither acidic nor alkaline.
EXAMPLES Pure water, saltwater, and various organic liquids Name of acid Ions present Salt formed
Hydrochloric acid, HCl H+ , Cl- - chloride, -Cl
Sulphuric acid, H2SO4 H+, SO42- -sulphate, -SO4
INDICATOR Nitric acid, HNO3 H+, NO3- - nitrate, -NO3
We can check whether a solution is acidic or alkaline by indicators. Ethanoic acid, CH3COOH H+, CH3 COO - - ethanoate, -CH3 COO
An acid-base indicator changes colour, reversibly. Some indicators with change in Common Acids
colour are shown below.
Acids have acidic properties only Note that HCl in gaseous form is called
when they are dissolved in water. hydrogen chloride. If it is dissolved in water ,
Colour in: it will undergo ionisation to form a solution
INDICATOR
ACID ALKALI called hydrochloric acid.
Litmus paper Red Blue
Phenol phthalein Colourless Pink or red BASICITY of an acid = NUMBER OF H+ IONS produced
Methyl orange Red Yellow when aqueous
Bromothymol blue Yellow Blue [e.g.] Basicity of H2SO4 = 2 H + ions = 2

UNIVERSAL INDICATOR TYPES OF ACIDS

Universal indicator is a mixture of several indicators and turns a range of 1. MINERAL acids
colours corresponding to different pH values. -These are STRONG acids and they IONISE COMPLETELY
[e.g.] hydrochloric, sulphuric and nitric acids
← Acidic Neutral Alkaline 
pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 2. ORGANIC acids
Green- -These are WEAK acids and they only PARTIALLY IONISE
red orange yellow Green blue violet [e.g.] carbonic, acetic (vinegar) and citric acids
blue
Colour change in Universal indicator with pH
PHYSICAL PROPERTIES
 Acids taste sour.
The range of colours in different solutions of pH for the universal
[e.g.] vinegar and lemon (They contain ethanoic acid and citric acid respectively)
indicator approximates the rainbow colours – red, orange, yellow, green,
 Acids turn litmus paper red
blue, indigo and violet. Taking pH 7(neutral) to be green, colours to the left
 Acids have pH values less than 7
of green the rainbow indicate acidic solution, while colours to the right
indicate alkaline solutions.

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REACTION OF ACIDS REACTION OF ACIDS WITH CARBONATES
There are 3 common reactions of acid Acids react with carbonates to give salts, carbon dioxide gas and water.

[EXAMPLE]
1. Acid + base  salt + H2O
calcium carbonate + hydrochloric acid calcium chloride + carbon dioxide + water
2. Acid + metal  salt + H2
CaCO3 (s) + 2HCl (aq)  CaCl2 (aq) + CO2 (g) + H2O (l)
3. Acid + carbonate  salt + H2O + CO2

REACTION OF ACID WITH BASE

We can test for carbon dioxide gas Delivery tube
Acids react with bases to give salts and water only
using limewater (calcium hydroxide
This reaction is also called neutralisation solution). Carbon dioxide causes the
[EXAMPLE 1] Acid with Metal oxide
limewater to turn chalky.
Lime water
Hydrochloric acid + zinc oxide  zinc chloride + water Calcium carbonate
2HCl (aq) + ZnO (s)  ZnCl2 (aq) + H2O (aq) + dilute hydrochloric acid

[EXAMPLE 2] Acid with Metal Hydroxide

BASES AND ALKALIS
Hydrochloric acid + sodium hydroxide  sodium chloride + water A base is a substance that reacts with an acid to give a salt and water only.
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (aq) This reaction is called neutralisation.

REACTION OF ACID WITH METAL

Acid react with metal above hydrogen in the reactivity series to give salts and Neutralisation is the chemical reaction between a base and an
hydrogen gas. acid to form a salt and water

[EXAMPLE] Acid + Base  Salt + Water

Hydrochloric acid + magnesium  magnesium chloride + hydrogen gas. HCl (s) + NaOH (aq)  NaCl (aq) + H2O (l)
2HCl (aq) + Mg (s)  MgCl2 (aq) + H2 (g)

SOLUBLE BASES AND INSOLUBLE BASES

The metal reactivity series lists metals according
to their reactivity. SEE the topic ‘METAL’. .  Many bases are insoluble in water.
Bases that can dissolve in water form solutions called alkalis
Lighted
We can test for hydrogen gas wooden splint
using a lighted splint. The flame
Acidic solution
 Bases are usually metal oxides or metal hydroxides.
extinguishes with a ‘pop’ sound.

Metal 46
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Table below lists some common bases and alkalis. REACTION OF BASES
Insoluble bases Soluble bases There are 2 common reactions of bases
Name Formula Name Formula
Magnesium oxide MgO Sodium oxide Na2O 1. Base + acid  salt + H2O
Copper(II) oxide CuO Calcium hydroxide Ca(OH)2 2. Base(alkali) + ammonium salt  salt + ammonium gas + H2O
Lead(II) hydroxide Pb(OH)2 Ammonium hydroxide NH4OH
Common bases and alkalis
REACTION OF BASES WITH AMMONIUM SALTS
SIMPLE CLASSIFICATION ON BASES
Alkalis react with ammonium salts to produce salts, ammonia gas and water.
Bases
[EXAMPLE]
sodium hydroxide + ammonium chloride  sodium chloride + ammonia + water
NaOH (aq) + NH4Cl (aq)  NaCl (aq) + NH3 (g) + H2O (l)
Oxides Hydroxides

Soluble Insoluble Soluble Insoluble The ammonia gas evolved is pungent, Red litmus paper
 Alkalis Basic oxides Alkalis Basic hydroxide colourless and turns damp red litmus
paper blue Alkaline solution
It can be seen from this that all alkalis are bases,
not all bases are alkalis. Solubility depends on HEATED UP Ammonium salt
the combination of ions
involved. We will see that
later. SIGNIFICANCE OF pH MESUREMENTS
-
 Hydroxide ions, OH are produced when bases dissolve in water to form alkalis. Apart form enabling us to determine whether substances are acidic or alkaline, pH
values have very important significance and implications in industry, agriculture,
[EXAMPLE] Sodium hydroxide, NaOH pharmacy and medicine.
NaOH (aq)  Na+ (aq) + OH- (aq)
CONTROL OF pH IN AGRICULTURE
The ability of alkalis to neutralise acids is due to the
Most plants need a soil pH of 6.5 to 7.5 to grow well. If the ground is too acid,
presence of these hydroxide ions.
slaked lime (solid calcium hydroxide) can be added to neutralise the acid. This
PHYSICAL PROPERTIES OF ALKALIS process is called liming the soil.

 Alkalis feel slippery  Slaked lime is chosen because

 Edible alkalis have a bitter taste. 1. It is cheap and easily available.
 Alkalis turn litmus paper blue 2. Slaked lime is sparingly soluble in water. Once the aid is neutralised, the
 Alkalis have pH values greater than 7 excess base will remain as solid in the soil. It will not dissolve in water to
make the soil too alkaline.

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 BASIC OXIDES
Aqueous ammonia and aqueous sodium hydroxide are alkalis that can also neutralise;
however, slaked lime has an advantage over them.
Basic oxides are oxides of metals.
 The person spraying the solution (e.g. sodium hydroxide solution) will not know when -They react with acids to produce salt and water only.
enough alkali has been added to neutralise the acid if the products of neutralisation
appears as a colourless solution. Excess alkali will cause the ground to become alkaline. [EXAMPLE]
Copper (II) oxide + sulphuric acid  copper (II) sulphate + water
CuO + H2SO4  CuSO4 + H2O (aq)
OXIDES
Oxides are formed when substances burn in oxygen gas. Oxides have acidic, basic, Here, neutralisation takes place.
amphoteric or neutral character, depending on which type of oxide they belong to. Basic oxides that dissolve in water form solutions called alkalis

AMPHOTERIC OXIDES
Type of Oxide Examples
Some oxides of metals known as amphoteric oxides behave as acidic or
ACIDIC Carbon dioxide CO2, Sulphur dioxide SO2,
1. basic oxides.
(Non-metallic) Nitrogen dioxide NO2
- When they react with acids, they behave as basic oxides;
BASIC Magnesium oxide MgO, Calcium oxide CaO,
2. - When they react with alkalis, they behave as acidic oxides.
(Metallic) Sodium oxide Na2O
Zinc oxide ZnO, Aluminium oxide Al2O3,
3. AMPHOTERIC [EXAMPLE] Zinc oxide reacts with an acid and a base for neutralisation.
Lead(II) oxide PbO, Tin oxide and
Water H2O, Carbon monoxide CO,
4. NEUTRAL 1. Sulphuric acid + zinc oxide  zinc sulphate + water
Nitrogen monoxide NO
H2SO4 + ZnO  ZnSO4 + H2O
[acid] [base]
PROPERTIES OF DIFFERENT TYPES OF OXIDES
- In this case, zinc oxide is acting as a base

ACIDIC OXIDES
2. Zinc oxide + sodium hydroxide  sodium zincate + water
Acidic oxides are usually oxides of non-metals.
ZnO + 2NaOH  Na2ZnO2 + H2O
- They form acids (H + ions) when dissolved in water
[acid] [base]
- In this case zinc oxide is acting as an acid.
[EXAMPLE] Carbon dioxide + water  carbonic acid
CO2 + H2O  H2CO3
NEUTRAL OXIDES
Natural rain has a pH slightly lower 7. Carbon dioxide in the air will Neutral oxides do not dissolve in water to form acids nor do they react with
dissolve n rainwater to produce a weakly acidic solution of carbonic acid. bases to form salts. NEITHER acidic or basic properties

[EXAMPLES] Water, Carbon monoxide, Nitrogen monoxide, Sulphur monoxide, etc

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7.2 Salt Preparation The selection of salt preparation method is summarised below.
Salts are chemical
o Precipitation is carried out if on insoluble salt is required.
compounds formed
As you have seen, the reaction of an acid results o If a soluble salt is needed, it is prepared by Filter and crystallisation method
when the hydrogen of
in the products of a salt. or by Titration.
an acid is partially or
Now we shall see how to prepare the salt required.
wholly replaced by a
SELECTING THE CORRECT PREPARATION METHOD
Compounds in which the H+ions in an acid metal or other positive
have been replaced by ammonium ions, ion (E.g. ammonium ion). SALT
NH4+are called ammonium salts.

SELECTION OF METHOD
Insoluble Salt Soluble Salt If sodium, potassium
SOLUBILITY RULES and ammonium salt
The method chosen to prepare a salt depends on its solubility.
 The solubility depends on the combination of positive and negative ions. Starting with Starting with
an Insoluble base an Soluble base
SOLUBLE INSOLUBLE Precipitation
 ALL nitrates  ALL carbonates
[E.g.] Filter and
 ALL chlorides EXCEPT FOR: sodium carbonate, Titration
PbCl2, PbSO4 Crystallisation Method
EXCEPT FOR: silver chloride and lead potassium carbonate and ammonium
(II) chloride carbonate CaSO4, AgCl
[E.g.] [E.g.]
 ALL sulphates  ALL sulphides NaCl, K2SO4,
CuSO4, MgCl2, Ca(NO3) 2
EXCEPT FOR: calcium sulphate, barium EXCEPT FOR: sodium sulphide, NH4Cl
sulphate and lead (II) sulphate potassium sulphide and ammonium
sulphide
 ALL oxides
[EXAMPLES] Name the correct method to prepare
EXCEPT FOR: sodium oxide, potassium
(a) potassium chloride
Note that oxide (Group I) and ammonium oxide
KCl  soluble salt the base can contain K  soluble base
All sodium, potassium (even oALL hydroxides
other Group I elements)
 use TITRATION
EXCEPT FOR: sodium hydroxide,
Ammonium and Nitrate potassium hydroxide (Group I),
(b) zinc sulphate
compounds are soluble. ammonium hydroxide and calcium
ZnSO4 soluble salt the base can not contain Na, K or NH4 insoluble base
hydroxide
 use FILTER AND CRYSTALLISATION METHOD

Table: Solubility of compounds

(c) silver chloride
AgCl  insoluble salt  use PRECIPITATION

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PREPARETION OF INSOLUBLE SALT PROCUDURE OF PRECIPITATION
Insoluble salts are prepared by mixing solutions containing their positive and
negative ions using the method of PRECIPITATION.  Dissolve each reactant separately in water
 Mix chemically equivalent quantities of the reactant solutions
-The reactant solutions are chosen so that on exchanging ions the unwanted  Filter the solution and wash the precipitate in warm distilled water
product is still soluble but the given insoluble salt will form as a precipitate.  Dry the solid salt that was produced in an oven (105ºC)

Refer to the topic of

[E.g.] To prepare Lead(II) iodide, PbI2
SOLUBLE SALT + SOLUBLE SALT  SOLUBLE SALT + INSOLUBLE SALT ‘separation technique’
(precipitate)
Pb(NO3) 2 (aq) + 2NaI (aq)  PbI2 (s) + 2NaNO3 (aq)
Precipitate is an insoluble solid formed
when a chemical reaction occurs between The ionic equation is
To find the Reactant Solutions:
two dissolved ionic substances. Pb2+ (aq) + 2I- (aq)  PbI2 (s)
Choose 2 starting solutions.
 One must contain the positive ion of the insoluble salt required
 The other must contain h the negative ion of the insoluble salt required. The reactants involved in a
precipitation reaction must
[EXAMPLE] To make barium sulphate (BaSO4 ), be in solution form because
the ions must be able to
REACTANT IONS: AFTER MIXING: move and interact with one
Ba2+ NO3- NaNO3 (soluble) another when the reactants
+ Pb2+ I-
Na SO42- BaSO4 are mixed together - Na+
NO3
We can use barium nitrate and sodium sulphate. MIX
NO3- I- Na+
The easiest salt solution containing the positive ions is
the METAL NITRATE (as all nitrates are soluble). Lead nitrate Sodium iodide
When the ions in the insoluble solution solution
salt encounter each other, they
For a solution containing the negative ions, can use the
will attract each other to form
SODIUM SALT (as all sodium salts are soluble)
a solid that will sink to the
bottom of the container and be
Na+ NO3-
[EXAMPLE 2] To make silver chloride (AgCl), collected as the precipitate.
Precipitate NO3- Na+ Na+(aq) and
REACTANT IONS: AFTER MIXING: of lead iodide NO3-(aq) are
Ag+ NO3- NaNO3 (soluble) I- Pb2+ I- spectator ions
+
Na Cl- AgCl
We can use silver nitrate and sodium chloride Other salts prepared using
precipitations include silver
chloride, lead chloride and etc.
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PREPARETION OF SOLUBLE SALT Step1
As you have seen, we have two methods to prepare soluble salts depending on the An excess reactant ensures that
solubility of the base that would be a starting material. copper(II) excess all the acid has been used up
sulphate copper(II)
and water oxide
SOLUBLE SALTS are prepared using two methods:
Don’t boil. The acid is warmed to
1. Filter and crystallisation method increase the speed of reaction
Neutralising an ACID with EXCESS INSOLUBLE REACTANT between the reactants.
2. Titration
Neutralising an ACID with the EXACT AMOUNT of ALKALI Step2

Residue: excess
Mixture: copper(II) oxide
FILTER AND CRYSTALLIDATION METHOD copper(II) sulphate solution
This method is used for preparation of soluble salts when a suitable insoluble and excess copper (II) oxide
starting material can be found. Filtrate: copper(II)
The acid reacts with an EXCESS of insoluble reactant that can be: sulphate and water
1. METAL
Step3 and 4 Copper(II) evaporation
2. BASE (INSOLUBLE)
3. CARBONATE sulphate and
water crystallisation

Therefore to prepare a given salt, we need to choose the correct acid and a Don’t evaporate all the
suitable insoluble reactant (METAL, OXIDE, HYDROXIDE or CARBONATE). water. The filtrate is
heated until a thin crust
STEPS of crystals form on the Copper(II)
１ Neutralise the acid with and excess of the insoluble reactant surface of the liquid. sulphate
２ Filter off any unreacted reagent crystals
３ Evaporate the solution to the crystallisation point
Preparing a soluble salt by filter and crystallisation method.
４ Cool to produce crystals of the salt
５ Filter, wash and dry the crystals before collection.
If a metal carbonate is used to prepare a salt using this method,
[EXAMPLE] The preparation of copper (II) sulphate there will be bubbles of carbon dioxide gas as the metal carbonate is
added to the acid in step1. When there is no more bubble, all the acid
Starting materials; copper (II) oxide and dilute sulphuric acid. has been used up and we may proceed to next step.
CuO(s) + H2SO4 (aq)  CuSO4 (aq) + H2O (l)

This reaction is used below to illustrate the procedure.

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TITRATION Step 1
The soluble salts of ammonium and Group I metals (sodium, potassium and  Place the soluble dilute acid in a conical flask
indicator
lithium) are prepared using the TITRATION METHOD.  Add a few drops of indicator
(e.g. methyl orange) sodium
hydroxide
This is because all their compounds are soluble (including the metals
solution
themselves) and very reactive. The Group I metals are so reactive resulting
in too violent reaction that we CAN NOT USE EXCESS reactant. Using pipette, measure
25.0cm of sodium
hydroxide into conical flask
This method is used when it is not possible to find a suitable insoluble starting Step 2
burette
material like a metal, a metal oxide or a carbonate that can be easily filtered  From a burette, slowly add the alkali solution.
off at the end of the reaction. Ensure that the solution is mixed well.
V1
 When the indicator begins to change colour,
 TITRATION means using the EXACT quantities of reactants for the reaction. dilute the reaction should be slowed to a drip.
nitric acid
INDICATOR
In a titration, an indicator is needed At this point, just enough acid is added to
Acidic END POINT Alkaline sodium neutralise the alkali, all the alkali has reacted.
to show the endpoint of one reactant
solution (neutral) solution hydroxide
needed to exactly neutralise a given
RED GREY or GREEN
volume of the other reactant. COLOURLESS Step 3 The resulting mixture
A common indicator used in the Colour change in methyl orange  Once the colour change is complete, contains only sodium
laboratory is the screened methyl orange. nitrate and water
the reaction is complete (END POINT).
The burette should be turned off.
From the titration result, we can V２
know the exact volume of nitric
STEPS OF TITRATION
acid needed to react with 25.0cm sodium nitrate
To prepare a given salt, the most common procedure is to react the alkali of sodium hydroxide. Volume of and water
solution with the dilute acid using a burette. Indicator is used to determine nitric acid, Va= V２- V1
when the exact amount of reactant has been added.
Step 4
[EXAMPLE]  Evaporate the solution to crystallisation point
The preparation of sodium nitrate is used to illustrate the procedure.  Cool to produce crystals of the salt
 Filter, wash and dry the crystals
Starting materials: aqueous sodium hydroxide and dilute nitric acid
-In a strict titration, second titration should be carried out. The salt solution obtained
NaOH(aq) + HNO3(aq)  NaNO3 (aq) + H2O (l) in the first titration is thrown away because it is affected by the indicator. Second one
is done without it. The exact volume of acid to be added is obtained from the first.

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7.3 Identification Tests  The cations react with the hydroxide ions present in aqueous sodium hydroxide
or aqueous ammonia to form insoluble hydroxides. These insoluble hydroxides
IDENTIFY SALT SOLUTIONS appear as precipitates.
To identify any salt solutions, we can take the following steps; [EXAMPLE]
Fe2+ (aq) + 2OH-(aq)  Fe(OH)2
From NaOH or NH3 Green precipitate
1. IDENTIFY / TEST for the METAL cation present
-Some of these precipitate dissolve in excess aqueous sodium hydroxide to form
2. IDENTIFY / TEST for the SALT anion present soluble complex salts.
 SALT SOLUTION = {METAL} + {SALT}
Test1 Test2 These appear as colourless solution. This occurs for amphoteric
metal hydroxides (Al3+, Zn2+ and Pb2+) which react with the alkalis.
Cations are positively charged ions
Again in an excess of ammonium solution, Zn and Cu redissolve to
Anions are negatively charged ions
form soluble complex salts. These appear as colourless solution or
TEST 1:Identification of METAL CATIONS dark blue solution.

When testing for a cation using either aqueous sodium hydroxide or aqueous
ammonia, two observations will help identify the cation present: FLOW CHART
From the previous table
1. the colour of the precipitate formed on adding a few drops of chemical regent;  copper (II), iron(II) and iron(III) ions are easily identified by the
2. the solubility of the precipitate in excess chemical regent. characteristic colour of their precipitations.
 Aluminium, lead (II) and zinc ions all give the same observations when aqueous
- Table below summarises the test for cations. sodium hydroxide is used. However, only zinc ions will give a white precipitate
Effect of NaOH solution Effect of NH4OH solution soluble in excess aqueous ammonia; aluminium and lead ions do not.
Name of Cation
Colour of Colour of
METAL present IN EXCESS IN EXCESS
Precipitate Precipitate
Calcium Ca2+ white insoluble white insoluble  Solution containing Al3+ ,Pb2+or Zn2+
Magnesium Mg2+ white insoluble white insoluble ↓- Add a few drops of aqueous ammonia and shake
Iron (II) Fe2+ green insoluble green insoluble  White precipitate formed
Iron (III) Fe3+ brown insoluble brown insoluble - Add excess aqueous ammonia
Copper(II) Cau2+ blue insoluble blue dark blue soln
Zinc Zn2+ white colourless soln white colourless soln  In excess aqueous ammonia, precipitate is
Lead (II) Pb2+ white colourless soln white insoluble
Aluminium Al3+ white colourless soln white insoluble
INSOLUBE - Al3+or Pb2+ present SOLUBLE - Zn2+combined
Table: Test for Cations
‘sol’ means ‘SOLUTION’

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  To distinguish between aluminium and lead(II) ions, dilute hydrochloric acid [EXAMPLE]
or aqueous potassium iodide can be used: Which solution will form a brown precipitate if sodium hydroxide is added and
Al3+ (aq) + 3Cl-(aq)  AlCl3 a white precipitate if silver nitrate is added?
From hydrochloric acid Colourless solution
Pb2+ (aq) + 2Cl-(aq)  PbCl2 Test 1: NaOH  brown precipitate  Fe3+  iron (III) [cation]
From hydrochloric acid White precipitate Test 2: AgNO3  white precipitate  Cl-  chloride [anion]
-Similar results will be obtained if aqueous potassium iodide is used.
Aluminium ions will give a colourless solution of aluminium iodide while Salt solution = Iron (III) chloride
lead(II) ions will give a yellow precipitate of lead(II) iodide.

IDENTIFICATION OF GASES
TEST 2: Identification of SALT ANIONS
 Carbon dioxide, sulphur dioxide and chlorine are all acidic gases and will turn
- the table below summarises the tests for anions. moist blue litmus paper red. Hence, the blue litmus paper test is not a
ANION PRESENT Formula TEST and Result conclusive test; it only indicates the presence of an acidic gas. It is necessary
Carbonate CO32-  Add hydrochloric ACID to conduct confirmatory tests in order to conclude the presence of a
also  Carbon dioxide is produced particular gas.
Hydrogen Carbonate HCO3-  Turns limewater milky  Ammonia, chlorine and sulphur dioxide have characteristic smell and are thus
 Acidify by adding dilute nitric acid easily identified.
When testing for
Chloride Cl-  Add silver nitrate solution hydrogen gas, hold
 White precipitate forms (AgCl) Table below summarises the test for gases. the lighted splint at
 Acidify by adding dilute hydrochloric acid GAS FORMULA TEST and RESULT the mouth of the
Sulphate SO42-  Add barium chloride solution Hydrogen H2  Burns with a ‘POP’ sound test tube.
 White precipitate forms (BaSO4) Oxygen O2  Relights a glowing splint
 Acidify by adding dilute nitric acid Carbon Dioxide CO2  Turns limewater milky
When testing for
Iodine I-  Add lead(II) nitrate solution  Turns damp blue litmus red
oxygen gas, insert
 Yellow precipitate forms (PbI2) then bleaches litmus paper the glowing splint
Chlorine Cl2
Table Tests of Anions  Yellowish-green colour into the test tube
 Choking smell
 Turns damp red litmus BLUE
Ammonia NH3
When recording the observations after conducting tests for the  Pungent smell
carbonate and the nitrate ion, remember to include the smell and Hydrogen  Turns damp blue litmus RED
HCl
colour of the gas, the chemical test result for the gas as well as the Chloride  Choking smell
name of the gas. Simply copying from the data sheet provided as  Turns damp blue litmus RED
Sulphur Dioxide SO2
‘carbon dioxide produced’ or ‘ammonia produced’ is insufficient and  Choking smell
will lead to a loss of marks. Table: Test for gases

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 mix with a solution Sodium hydroxide mix with a solution
FLAME TESTS containing A2+ ions solution containing B2+ ions

 React a small quantity of compound with a couple of drops of concentrated

White precipitate C Blue precipitate D
hydrochloric acid
 Dip a clean piece of platinum wire into the mixture and put into the flame
of a Bunsen burner to precipitate add an to precipitate add an
 The flame colour in each reaction is shown below. excess of sodium hydroxide excess of dilute sulphuric
solution acid
METAL PRESENT CATION FLAME COLOUR
Sodium Na +
Orange-yellow Precipitate disappears Precipitate disappears
Potassium K+ Lilac-pink leaving a colourless solution leaving a blue solution E
Calcium Ca2+ Brick red
Barium Ba2+ Pale green (a) Suggest identities for the ions, A2+ and B2+, and substances C, D and E
Copper (II) Cu2+ Green ANSWER:
Lead (II) Pb2+ Blue A2+: add NaOH  white precip.  colourless soln in excess 
Table: Test for flames 2+ charge  zinc ion, Zn2+ (or Pb2+ ion but NOT Al3+ ion!!)
2+
B : add NaOH  blue precip.  copper ion, Cu2+
C: Zn(OH)2 (or Pb(OH)2) zinc hydroxide (or lead (ii) hydroxide}
REACTION SCHEMES D: Cu(OH)2 copper (ii) hydroxide
E: CuSO4 copper (ii) sulphate
When solving Reaction Schemes it is essential that you know all the reactions
typical for acids and alkalis and ammonium salts. (b) Write a chemical equation for any one of the reactions shown in the diagram
1. acid + metal  salt + H2 ANSWER:
2. acid + alkali / base  salt + H2O 2NaOH + Zn2+  Zn(OH)2 + 2Na+ or
3. acid + carbonate  salt + H2O + CO2 Zn(OH)2 + 2NaOH  Na2Zn(OH)4 or
4. ammonium solution + acid  ammonium salt + H2O 2NaOH + Cu2+  Cu(OH)2 + 2Na+ or
5. alkali + ammonium salt  salt + NH3 + H2O Cu(OH)2 + H2SO4  CuSO4 + 2H2O { easiest answer!! }
You may also need to know the identification tests for the various cations, anions
and gases. (c) Explain how the sodium hydroxide solution can be used to distinguish between a
solution containing an iron (II) compound and a solution containing an iron (III)
[EXAMPLE] compound.
The diagram below shows some properties and reactions of the ions, A2+ and B2+, ANSWER:
and the substances C, D and E. { 1997 Paper 3 Section B } When sodium hydroxide solution is added to a solution containing an iron (II)
compound, a green precipitate forms. When sodium hydroxide solution is added
to a solution containing an iron (III) compound, a red-brown precipitate forms.

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 In Periodic table we saw the most of the elements are metals. The non-metals
confined to the top right-hand corner of the periodic table. Of over 100 elements
which we know, only 21 are non-metals.
Now we shall investigate metals.

COTENT LEARNING OBJECTIVE (Pupils should be able to)

 Describe the general physical properties of metals 8.1 Properties of metal
Properties of  Explain why metals are often used in the form of alloys.
metals  Identify representations of metals and alloys from PHYSICAL PROPERTIES and USE
diagrams of structures In terms of appearance, Non-metals are different with each other while metals are
 Place in order of reativity calcium, copper, hydrogen , iron, all alike. Only copper and gold are coloured; all are shiny.
magnesium, potassium, sodium and zinc by reference to the Let’s see physical properties of typical metals comparing with those of non-metals.
reactions if any of the metals with water(or steam)
 Account for the apparent unreactivity of aluminium in terms METALS are… NON-METALS are…
Reactivity series  Conduct electricity and heat  Do not conduct electricity and heat
of the presence of an oxide layer which adheres to the
metal.  Shiny  Dull
 Deduce an order of reactivity from a given set of  Malleable and ductile  Brittle
experimental results.  High MP and BP  usually solid  Low MP and BP  usually liquid or gas
 Describe the ease in obtaining metals from their ores by  High densities  Low densities
Exraction of  Magnetic materials  Non-magnetic materials
relation the elements of the reactivity series.
metals  Strong and Tough
 Describe the effect of aluminium on human beings.
 Describe the essential reactions in the extraction of iron Except for Cu and Al Except for Mercury that is liquid at r.t.p.
from haematite and Gallium that melts below 30℃
Describe methods of rust prevention
Iron
 Describe the idea of changing the properties of iron by the
controlled use of additives to form alloys called steels
 State the uses of mild steel and stainless steel. Malleable means ‘be easily made into sheet without breaking.
 Describe the extraction and purification of copper from its Ductile means ‘be easily made into wire without breaking.
Copper ore.
 State the uses of copper related to its properties  Due to the properties above, metals are used in many ways
 State the uses of aluminium in air craft and food containers METAL USES USEFUL PROPERTY
Aluminium  State the uses of zinc for galvanising and for making brass Copper  Electric cables  Excellent conductor of electricity
(with copper) Aluminium  Soft drink cans  Does not corrode
 Coat ‘tin’ cans used
Tin  Non-poisonous
for food e.g. jam
Gold / Silver  Jewellery  Malleable and very unreactive

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ALLOYS 8.2 Reactivity Series
Pure metals are usually too soft and weak for most uses. To improve the strength and
hardness of pure metal, we do the following treatment.
The reactivity series is a lis of metals placed in
In pure metals, the atoms are arranged orderly in layers. order of their reactivity, as determined by
When a force is applied to the metal, the layers of metal atoms can their reaction with water and dilute acid.
slide one over another
PUSH

 To improve the strength

and hardness of pure metals, REACTION OF METALS WITH WATER
atoms of another element can Some metals react with water in the liquid state while others react only with it in the
Metal structure After being pushed
be added, usually in small amounts. gas state (steam).
These atoms prevent the atoms of the metal from sliding over on
another, making the metal stronger and harder and less likely to have - Table 1 lists the reaction of some metals with water
its shape distorted. The final product is an alloy of the metal Metal Observation/Equation
Reacts very violently. Enough heat is produced to ignite the
One metal hydrogen gas produced. The hydrogen burns with a blue flame.
An alloy is a mixture of two or Potassium (K)
E.g. copper
more elements that are usually 2K (s) + 2H2O(l)  2KOH(aq) + H2(g)
another metal metals except for carbon in steel.
E.g. zinc Reacts violently. The hydrogen gas produced may catch fire.
Sodium (Na)
2Na (s) + 2H2O(l)  2NaOH(aq) + H2(g)
ADVANTEGES OF ALLOYING
Reacts readily. Hydrogen gas and calcium hydroxide solution
1. Stronger and harder than the pure metals.
Calcium (Ca) are formed.
2. Improved metal appearance
2Ca (s) + 2H2O(l)  Ca(OH)2 (aq) + H2(g)
3. Increased resistance to corrosion
Reacts very slowly with cold water. A test tube of hydrogen
- Some examples of alloys are below
gas is produced only after a few days.
ALLOY MIXTURE OF USES USEFUL PROPERTY Magnesium (Mg)
Mild steel Iron and Carbon  car bodies  hard and strong 2Mg (s) + 2H2O(l)  Mg(OH)2 (aq) + H2(g)
Stainless Iron, Chromium  cutlery Zinc (Zn), Iron (Fe),
 corrosion resistant
steel and Nickel  surgical instrument Lead(Pb), Copper(Cu), Do not react with cold water
Brass Copper and Zinc  Screw  corrosion resistant Silver(Ag)
Bronze Copper and Tin  ornament  good appearance Table1 ; reaction of metals with water
Aluminium and  aircraft and
Duralumin  strong and lightweight
Magnesium bicycle frames
Solder Lead and Tin  welding metals  low MP
Pewter Tin and Lead  ornament  good appearance

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- Table 2 lists the reaction of some metals with steam REACTION OF METALS WITH DILUTE HYDROCHLORIC ACID
Metal Observation/Equation The reaction of metals with dilute acid is also taken into account for the reactivity
The hot magnesium reacts violently with steam to form series (The dilute acid is hydrochloric acid here).
magnesium oxide (a white powder) and hydrogen gas. A bright
Magnesium (K) white glow is produced during the reaction. - Table 3 lists the reaction of some metals with dilute hydrochloric acid.
Metal Observation/Equation
Mg (s) + H2O(l)  MgO(s) + H2(g) Potassium(K), Explosive reaction. Reaction is not usually carried out because
Hot zinc reacts with steam to produce zinc oxide and hydrogen Sodium(Na) it is too dangerous to do in a laboratory.
gas. Zinc oxide is yellow when hot and white when cold. Reacts vigorously to give hydrogen gas and calcium chloride.
Zinc (Zn)
Calcium (Ca)
Zn (s) + H2O(l)  ZnO(s) + H2(g) Ca (s) + 2HCl(aq)  CaCl2(aq) + H2(g)
Red hot iron reacts slowly with steam to form hydrogen gas Reacts rapidly to give hydrogen gas and magnesium chloride.
Iron (Fe)
and tri-iron tetraoxide. Magnesium (Mg)
Lead(Pb), Copper(Cu), Mg (s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Do not react with steam
Silver(Ag) Reacts moderately fast to give hydrogen gas and zinc chloride.
Table2 ; reaction of metals with water Zinc (Zn)
Zn (s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Reacts slowly to give hydrogen gas and iron(II) chloride.
 FROM THE OBSERVATIONS of the reactions of metals with water Iron (Fe)
Fe (s) + 2HCl(aq)  FeCl2(aq) + H2(g)
1. When metals react with water or steam, metal hydroxides or metal oxide and Copper(Cu),
Do not react with cold water
hydrogen gas are formed. Silver(Ag)
Table3 ; reaction of metals with dilute hydrochloric acid
Metal + water  Metal hydroxide + Hydrogen gas
Metal + steam  Metal oxide + Hydrogen gas
If a piece of aluminium foil is reacted with hot dilute hydrochloric
acid, the initial rate of reaction will be very slow as the acid reacts
Note that magnesium reacts with both water and steam.
When it reacts with water, the product is magnesium hydroxide; with the layer of aluminium oxide on the surface of the foil.
when it reacts with steam. The product is magnesium oxide. Once the oxide layer is removed, the reaction will speed up as
aluminium is a reactive metal.
2. The more vigorous the reaction, the more reactive the metal.

Potassium, sodium, calcium are reactive metals. We can draw the reactivity series of metals from the reactions of metals with
Magnesium, zinc and iron are fairly reactive metals. water and dilute acid as shown below.
Lead, copper and silver are unreactive metals.

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 THE METAL REACTIVITY SERIES DISPLACEMENT REACTIONS
Displacement reactions from solutions can be predicted using the reactivity series.
Potassium K The more reactive metal has the
VERY higher tendency to lose valence More reactive metals will displace a less reactive metal from its compound
Sodium Na
REACTIVE electrons and form positive ions.
or solution (a colour change of the solution is often observed)
Calcium Ca
Magnesium Mg
FAIRLY A metal higher in the series will displace a metal lower in the series.
Aluminium Al
REACTIVE
[EXAMPLE 1] Iron + copper (II) nitrate solution
Zinc Zn
Hydrogen is included in the series although
NOT VERY Iron Fe it is a non-metal. It serves as a reference Mg (s) + 2HCl(aq)  MgCl2(aq) + H2(g)
REACTIVE Tin Sn point in the series
Lead Pb
A brown metallic deposit of coper metal will form as the solution turns from blue to
HYDROGEN H2  Metals above hydrogen will react with
pale green due to the formation of iron (II) ions.
Copper Cu dilute acids to give hydrogen gas.
NOT AT ALL Silver Ag  Metals below hydrogen will not react with
dilute acids to produce hydrogen gas. [EXAMPLE 2] Iron + zinc (II) sulphate solution
REACTIVE Platinum Pt
Gold Au
Iron is lower than zinc in the reactivity series. Since it is less reactive than zinc,
no displacement reaction will take place.

Important Note: Aluminium is placed higher in the reactivity series although it [EXAMPLE 3] Copper + silver nitrate solution
shows no observable reaction with dilute hydrochloric acid. It appears less Since copper is above silver in the reactivity series, copper will displace silver from
reactive due to the protective layer of aluminium oxide (Al2O3) that keeps silver nitrate solution.
the metal inside.
Cu (s) + 2AgNO3 (aq)  Cu(NO3) 2 (aq) + 2Ag(s)

A layer of silver will form on the copper metal. The solution will also turn from
The metal reactivity series may differ from book to book, colourless to blue due to the formation of copper(II) ions.
depending on how many metals are included in it.

Generally
Group I metals will be located at the top of the series, since
they are the most reactive metals in the Periodic Table. Displacement reactions also take place for
Group II metals, Group III metals and finally, the transition Group VII elements. The more reactive
metals will follow them. halogen will displace the less reactive halogen
from a solution containing its ions.

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8.3 Extraction of metals When carbon is used to extract a metal from its metal oxide, a redox reaction takes
place. Carbon is said to be the reducing agent as it reduces the metal oxide to the metal
The method of extraction of a metal from its compounds is determined by its by removing oxygen from it
position in the metal reactivity series. The more reactive the metal, the harder it is
to extract the metal from its compounds EXAMPLE: Copper (II) oxide and carbon 2CuO + C  CO 2 + 2Cu
Black reddish brown

-There are 2 methods for extracting metals from their ores:

1. Reduction of the molten metal compound
SCRAP METALS AND RECYCLING
2. Electrolysis of the molten metal compound
Metal ores are finite and limited and expensive to mine. It is essential that we
recycle those scrap metals that are still useful. Iron, steel, copper and aluminium
are the most easily recycled metals.
Extraction of metals and the reactivity series
ADVANTAGES of RECYCLING
Most reactive Metals more reactive are  Saves energy and reduces greenhouse gas emissions (CO2)
Potassium K
extracted by electrolysis of their  Preserves non-renewable material
Sodium Na
molten salts or molten ores  Reduces land degradation, air and water pollution through mining
Calcium Ca
 Reduces the amount of land fill required for disposal of scrap metal
Magnesium Mg
Aluminium Al

Reactivity Zinc Zn
decreases Iron Fe
Tin Sn Metals less reactive are Recycling is sometimes not feasible because of the costs
Lead Pb extracted by reducing the ore involved. Transportation, sorting through waste and cleaning
Copper Cu with carbon or carbon monoxide the scrap metal, etc. may cost more than extracting the metal
Silver Ag from its ores. This is true for some cheaper metals.
Platinum Pt
The least reactive metals
Gold Au
Least reactive e.g. silver or gold even occur
native i.e. unreactive

Electrolysis involves the use of large amounts of electricity

and is a very expensive process compared to reduction using
carbon. It is only used to extract very reactive metals because
their compounds are too stable to be reduced using carbon.

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8.4 Iron PROCESS
Let us see the process in terms of the reactions that take place in the furnace.
EXTRACTION OF IRON
Iron is extracted from its ore haematite, Fe2O3, by reduction using carbon COKE (carbon)
LIMESTONE (CaCO3)
in a BLAST FURNACE Carbon burns in air to form carbon
Limestone decomposes to form
dioxide:
calcium oxide and carbon dioxide:
RAW MATERIALS C(s) + O2 (g)  CO2 (g)
 Haematite (Fe2O3 containing iron(III))
 Coke (carbon) CaCO3(s)  CaO (s) + CO2 (g)
Carbon dioxide combines with more
 Limestone (CaCO3) coke to form carbon monoxide:
BLAST FURNACE CO2 (g) + C (s)  2CO (g)
-The diagram below shows the blast furnace

Equal amount of
Iron ore and Coke IRON ORE (HAEMATITE)
with Lime stone
The waste hot gases are used to
-Iron ore (Haematite) contains IRON (III) OXIDE (Fe2O3)
heat the incoming hot air.
Waste gases and IMPURITIES (e.g. sand SiO2)
Waste gases (The toxic gas carbon monoxide
out out recombines with oxygen to form
the greenhouse gas carbon dioxide.) REDUCING IRON ORE TO IRON

800 °C Carbon monoxide gas reacts with iron (III)

oxide to form molten iron:
Preheated 1500 °C Preheated The molten slag is less dense 3CO (g) + Fe2O3(s)  3CO2 (s) + 2Fe(l)
air air than the molten iron, so that it
2000 °C Carbon monoxide acts as the reducing agent in
can floats on iron and be gained
from the top tap. the reaction. The liquid iron formed flows to
tap the base of the blast furnace.
SLAG
tap The molten slag can be used for REMOVING IMPURITIES
IRON Molten road building or to make fertiliser The basic calcium oxide is used to remove acidic
Molten slag
impurities (e.g. sand, SiO2):
iron
CaO(s) + SiO2 (s)  CaSiO3 (l)

The raw iron obtained in this process is known as cast iron or pig iron.
The liquid slag flows to the base of the blast
It can be purified further by bubbling oxygen gas through it to burn away
furnace and floats on top of the molten iron
impurities. The purified iron is then used to make alloys such as steel.

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STEEL RUSTING
The iron that is formed in a blast furnace is not hard enough to be used industrially as Rusting is the corrosion of iron or steel to form hydrated iron (III) oxide Fe2O3 . nH2O
the metal is too soft, so that the iron is alloyed into steel.
FORMATION OF RUST Rusting is a redox reaction
There are many types of steel depending on the
Steel is an alloy made
type and amount of additives to it.
by mixing iron with For rusting to occur, both AIR (oxygen) and WATER must be present.
There are 2 kinds of steel; carbon steels and
carbon or other metals.
alloy steels. The chart below shows that.
[EXPERIMENT]
Iron nails are put in various test tubes. Let’s see formation of rusting in each of them.

Mild steel is softer and more QUICKEST SLOW NO RUSTING

easily shaped. It is used to make RUSTING RUSTING
car bodies and machinery.
nail nail nail nail
Carbon steels contain
tap tap boiled
mainly iron and carbon
water water water drying
High carbon steel is strong but is filled agent
brittle. It is used to make knives,
e.g. CaCl2
hammers, chisels, saws and other
STEEL cutting and boring tools. (1) (2) (3) (4)
RUST FORMS NO RUST FORMS

Alloy steels contain iron and

carbon and a transition From the conditions for rusting (THE PRESENCE OF AIR AND WATER),
element such as manganese, you can tell the results above.
Stainless steel is an example
nickel, chromium, tungsten of an alloy steel.
or vanadium. It does not rust, is extremely  In the test tube (1), enough amounts of air and water are there.
durable and resistant to  In the test tube (2), it looks like that only water is there. But small amount
corrosion even upon heating. It of air can exist in water, so that rusting can occur.
is used to build chemical plants,  In the test tube (3), unlike the tube (2), there is no air although water is
in making of cutlery and filled. It is because air that used to exist in water has been removed by
surgical instruments.
boiling.
 In the test tube (4), there is no air with during agent.

 In the test tubes that satisfy the condition for rusting, rusting occur.

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 8.5 Copper
 Unlike aluminium which reacts with oxygen in the air to form a
protective layer on the metal surface, rust is brittle and flaky. EXTRACTION OF COPPER
The irons underneath will eventually rust and flake away Copper is an unreactive metal so it can be extracted from its ore, by heating
with carbon
The overall reaction that takes place in rusting is given by
the equation COPPER ORES
4Fe(s) + 3O2(g) + 2nH2O(l)  2Fe2O3 . nH2O(s)
 CUPRITE, Cu2O (by heating with carbon)
- This is an oxidation reaction that takes place slowly. In  MALACHITE, CuCO3-Cu(OH)2 (by decomposing on heating)
this process, iron is first oxidised to iron(II) ions before
the iron are further oxidised to iron(III) ions.
PROCESSING COPPER ORES INDUSTRIALLY
The two main ways to process copper ores industrially are:
 FLOTATION, roasting and SMELTING
PREVENTING THE FORMATION OF RUST  LEACHING with dilute sulphuric acid (more commonly used in Zambia)
There are 2 main ways of preventing rusting of iron or steel. then using ELECTROLYSIS or adding SCRAP IRON
BARRIER PROTECTION
1. Coat the iron/steel object with a layer of substance that will stop oxygen PURIFYING EXTRACTED COPPER INDUSTRIALLY
in the air and/or water form reaching the metal.  Very pure copper is needed for electrical conductors
[EXAMPLES] painting, oil or greasing  ELECTROLYSIS is used to produce VERY PURE COPPER
2. coat the iron/steel object with a less reactive metal or with plastic
[EXAMPLE] steel food cans coated with tin (tin-plating)
ELECTROLYSIS to PURIFY COPPER
SACRIFICIAL PROTECTION - +
CATHODE ANODE
Coat the iron/steel object with a more - ve + ve
Galvanizing is a method of
reactive metal. The more reactive metal Starts as
protecting a metal (e.g. iron or
will corrode in place of iron. pure copper Impure copper which
steel) from corrosion by covering
and more pure dissolves away
it with a thin layer of ZINC - - + +
[EXAMPLE] galvanizing copper adds to it Cu2+
through dipping or electroplating. - - Cu2+ + +
- - Cu2+ + +
Copper (II) sulphate - - Cu2+ + +
In this method, a reactive metal is used for a coating metal, but not all
reactive metals are suitable. For example, magnesium is not used as a coating on solution containing - -- + + + + Sludge formed
Cu2+ ions (electrolyte) - from impurities
an iron or steel object because it will react with the oxygen in the air to form
magnesium oxide. Magnesium oxide flakes easily and will come off the surface,
exposing more magnesium for reaction. In this way, a magnesium coating will Pure copper is deposited on Copper dissolves from the
wear out very quickly. Hence magnesium is not suitable for the coating metal. the pure cathode (-) impure anode (+)
Cu2+ (aq) + 2e-  Cu (s) Cu (s)  Cu2+ (aq) + 2e-
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 STEPS:  A molten state is needed for electrolysis. This can be very expensive. Al 2O3
 The ANODE (positive electrode) is made from impure copper has a very high melting point over 2000°C. Instead the Al2O3 is dissolved in
At this electrode, the copper atoms give up e- to form Cu2+ ions which molten cryolite (a less common ore of aluminium). This only requires a
dissolve in the solution temperature of about 900°C, which is much cheaper.
 These Cu2+ ions are then ATTRACTED to the negative electrode
 The CATHODE (negative electrode) starts as a thin piece of very pure USES OF ALUMINIUM AND ALLOYS
copper  Overhead electrical cables are made of aluminium as it is lightweight and a
At this electrode, Cu2+ ions gain e- to form Cu atoms which deposit on the good conductor of electricity
cathode which increases in size  Cooking utensils and food containers are made of aluminium as it does not
 The impurities in the anode fall to the bottom as a sludge as the anode corrode (due to its protective oxide layer) and is a good conductor of heat
dissolves away  Aircraft and bicycle frames are made from aluminium alloy(duralumin) as they
are strong and lightweight
USES OF PURE COPPER
Copper is used to make electrical wiring and heat exchangers because it is an
excellent conductor of electricity and heat
SUMMARY Reactivity Series and Reactions of Metals
COPPER ALLOYS
 Brass is an alloy of copper and zinc and is used to make musical instruments REACTION ACTION OF

METAL
REACTION WITH REACTION
and bimetallic strips WITH HEAT ON
DILUTE ACID WITH AIR
 Bronze is an alloy of copper and tin and is used to make trophies WATER CARBONATE
 Both of these alloys are non-corroding K Burns very
React with Violent reaction with dilute
Na easily with No reaction
cold water acids
Ca a bright
8.6 Aluminium Mg flame
React with
Al Burn slowly to
steam
Aluminium and its alloys have the following properties. Zn React fairly well with dilute form oxide
Electrolysis is Reacts acids with decreasing ease
 It has low density Decompose to
 It has good electrical and heat conductivity. expensive as it uses Fe reversibly form oxide
lots of electricity with steam React slowly
 It is resistant to corrosion. and CO2
 It is a relatively strong metal. Pb May react slowly if warmed with air when
No reaction with dilute heated
Cu No reaction acids, may react with
EXTRACTION OF ALUMINIUM
with water or concentrated acids
Aluminium is a reactive metal so must be extracted from its ore by electrolysis steam Decomposes
Ag No reaction No reaction to form Ag,
ALUMINIUM ORE O2 and CO2
 bauxite, Al2O3

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 9.1 WATER

Water is the most abundant liquid on earth – it covers 70% of the earth’s surface.
COTENT LEARNING OBJECTIVE (Pupils should be able to)
Water is used at home for drinking, cooking, cleaning and washing. Now let us see
 Explain the effects of water pollution
this important liquid.
 Suggest ways of reducing water pollution
Water  Describe in outline the purification of water supply in terms
WATER PURIFICATION
of filtration and chlorination.
 State uses of water in industry and the home.
WATER TREATMENT
 Describe the volume composition of clean air
Treatment of drinking water is carried out at the waterworks.
Air  Name common pollutants
- Three main stages Sedimentation, Filtration and Chlorination are involved
 State the sources of each of the following pollutants
 Explain the use of hydrogen in a manufacture of ammonia 3 STAGES IN WATER TREATMENT
and of margarine and as fuel in rockets
 Name the uses, oxygen tents in hospitals, and with
acetylene (a hydrocarbon) in welding. RAW WATER
 Describe the need for nitrogen, phosphorus and potassium Raw water is first screened to remove large solid impurities.
Common
compounds in plant life.
Non-metals  Describe the essential conditions for the manufacture of - Alum (a coagulating agent) is used to make solid
ammonia by the haber process
particles stick together.
 Name the uses of ammonia in the manufacture of fertilisers
such as ammonium sulphate and nitrate. 1. SEDIMENTATION - The solid clumps sink to the bottom in the
 Discuss the effect of chemical fertilizers on the soil. sedimentation tank and are removed.

Lime (calcium oxide) is added to reduce acidity.

Activated carbon is added to remove foul odour and taste

WATER WORKS Water tower 2. FILTRATION - The water is filtered to remove any remaining
Raw solid particles
Water

- Chlorine is added to kill germs and bacteria
3. CHLORINATION present in the water.
Screening
Gravel
Tank Chlorination
Filter
Sedimentation plant The water supplied to our homes does not have to be
Tank Water pure but safe to drink. It is called potable (drinkable).
CLEAR WATER
Supply

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 Important Note: Chlorine is a highly poisonous substance. It is important to use INDUSTRIAL USES OF WATER
the right quantity even when using ‘Chlorine’ for chlorinating water at home. Water is used in many different ways by industries.

DESALINATION INDUSTRIAL USES EXAMPLES

Desalination is the process of removing dissolved salts from seawater. The sea  Beer making
As an essential ingredient for a product
thus provides a ready source of drinking water.  Whisky production
 Producing electricity in a coal
Coolant
-Two methods of desalination are commonly used. or oil fired power station
1. Distillation; seawater is evaporated and the pure water vapour formed is Source of energy  Hydroelectricity
condensed, e.g. solar distillation Raw material in manufacturing process  Paper manufacture
2. Reverse osmosis; Pure water is extracted from seawater using a semi- Polar solvent  Dissolving ionic compounds
permeable membrane under high pressure.

WATER POLLUTANTS
Water from rivers and lakes contains dissolved mineral salts, organic matter as
9.2 AIR
well as some pollutants.
We human beings can not survive without air. Earth is surrounded by the
atmosphere that contains air. Let’s see this important gas.
SOURCE OF
POLLUTANT HARMFUL EFFECTS
POLLUTANT
COMPOSITION OF AIR
Aquatic life cannot survive in low pH
Clean, dry air is a mixture of gases.
ACIDS water. Low pH water also causes Acid rain Air also contains water vapour in
poor growth of vegetation. variable amounts, depending on the
Causes eutrophication – excessive % COMPOSITION OF AIR humidity of the surroundings
NITRATES growth of vegetation which uses up
Excess fertilisers
AND dissolved O2. This causes the fish to
washed off from crops
PHOSPHATES die. After the vegetation decays,
NOBLE
the water becomes stagnant OXYGEN GAS
Waste from industries 21% mainly
HEAVY
Poisonous to mankind involved in mining and ARGON
METALS
processing metals NITROGEN
Untreated household 78% 0.94%
Health problems such as infections.
SEWAGE waste and excretion
Can also cause eutrophication.
from animals
Kills aquatic life as oxygen can non CARBON
DIOXIDE 0.04%
OIL longer pass through and dissolve in Ships with oil spills
water
The percentages of the gases that make up air will vary
slightly form place to place, depending on local conditions.
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AIR POLLUTION METHANE CH4
Air is said to be polluted when it contains chemicals in high enough concentrations SOURCE:
to harm living things or damage non-living things.  Bacterial decay of vegetable matter, animal dung and rubbish buried in landfills

COMMON AIR POLLUTANTS EFFECTS:

Common air pollutants include  It can combine with oxides of nitrogen in the presence of sunlight to form
o Sulphur dioxide SO2 photochemical smog.
o Nitrogen oxides NO and NO2  It is also a green house gas that can cause global warming.
o Carbon monoxide CO
o Methane CH4
o Lead compounds Global warming is the gradual change in world climate
caused by the greenhouse effect.
SOURCES OF AIR POLLUTANTS It is thought that it may cause changes in weather
Now let’s look at these sources of air pollutants respectively. patterns, causing droughts and storms, and even melting
the polar ice caps to bring about severe flooding
SULPHUR DIOXIDE SO2 CARBON MONOXIDE CO
SOURCE: SOURCE:
 Burning of fossil fuels containing sulphur and sulphur compounds  Carbon monoxide is produced during the INCOMPLETE combustion of
e.g. coal, natural gas or petroleum carbon containing compounds i.e. insufficient O2
- power stations - areas with high concentration of vehicles due to car exhaust
- car exhaust - faulty gas appliances
- areas where combustion takes place with poor ventilation
EFFECTS: e.g. using a brazier inside
 Sulphur dioxide irritates the eyes and causes breathing difficulties MINIMISING MEASURES:
 Produces acid rain EFFECTS:  Use more air during combustion
 Carbon monoxide poisoning  Cars use a catalytic converter
MINIMISING MEASURES: causes suffocation
 treat exhaust gases with wet calcium hydroxide to remove SO2 because the haemoglobin in our blood reacts more readily with CO than O2

NITROGEN OXIDES NO and NO2 LEAD COMPOUNDS

SOURCE: SOURCE:
 At high temperatures, the N2 and O2 in air combine to form nitrogen oxides  Lead compounds are added to some fuels to make the car engines run properly
- exhaust from car engines EFFECTS:
- power stations and factories  Lead can cause brain damage and is especially harmful to young children
- naturally occur from bush fires and lightning
EFFECTS:
MINIMISING MEASURES:
 Produces acid rain MINIMISING MEASURES:  Use lead-free petrol
 Cars use a catalytic converter 67
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ACID RAIN 9.3 Common non-metallic elements
.It is a long time since the effect of acid rain was known commonly. The acid rain
causes serious damages to the environment over wide areas in the world. It is
about time for us to face this global problem earnestly. HYDROGEN
Hydrogen is the first and lightest element in the periodic table. It is the most
SOURCE: abundant element in the universe (it is present in water and in all organic
 Sulphur dioxide in the air reacts with oxygen and water to form sulphuric compounds) and is the main constituent of stars.
acid which dissolves in rain clouds to form acid rain with a pH of 4 even down to 2
 Nitrogen oxides also form nitric acid with air to form acid rain PREPARATION
 In industry: Steam reforming method which is the reaction of methane and
EFFECTS: steam is applied.
 Corrodes metal structures e.g. bridges and cars CH4(g) + H2O(g)  CO2 (g) + 4H2 (g)
 Corrodes limestone buildings as cement contains carbonate that readily
 In a laboratory: There are 3 simple ways Check the topic of
react with the acids
 Endangers aquatic life as fish and plants can not survive in acidic water  Reaction of reactive metal and water ‘acid, base and salt’
 Causes the soil to become acidic, causing plants to die more readily  Reaction of metal with acids
 Electrolysis acidified water
MINIMISING MEASURES: 2H+ + 2e-  H2 (collects at anode)
 Remove the acidic gases, NO2 and SO2, at the source (see above points)
 Reduce water and soil acidity using slaked lime, Ca(OH)2 IDENTIFICATION
When a lighted splint is held at the mouth of a test tube containing hydrogen
gas, the gas burns explosively, making a “pop” sound.

You know, there are many global problems. Apart form acid rain, another USES
common problem on the Earth is global warming caused by the green hous effect  Hydrogen is used in the Haber Process to produce ammonia
 Hydrogenation is used to change vegetable oils into margarine.
 The green house effect is the trapping
-Vegetable oil is unsaturated and the hydrogen breaks the double carbon
of heat energy in the atmosphere
bonds to form saturated margarine.
because of the effects of greenhouse
 Rockets burn liquid hydrogen as a fuel with liquid oxygen to form water. This
gases. The infrared radiation (heat
is a very lightweight fuel.
energy) is given off from the earth’s
surface as it is warmed up by the Sun

Hydrogen is so flammable that there would be a

Green houses gases are gases in the atmosphere which absorb infra-red radiation,
risk of explosion and it would be hard to liquefy for
causing an increase in air temperature.
storage in modern fuel. But it may have a use as a
The most important is carbon dioxide, which is increased by burning fossil fuels and by
deforestation, which reduces the amount of carbon dioxide removed by photosynthesis. common non-polluting fuel for road vehicles in the
Another is methane, a by-product of rice farming and cattle-rearing. near future.

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OXYGEN
Oxygen is the most important gas in the air. It is a colourless, odourless gas.
3. RUSTING occurs when iron comes into contact with water and
oxygen to form rust. Rust is hydrated iron (iii) oxide, Fe2O3.nH2O.
PREPARATION
 In industry: Fractional distillation of liquid air is applied
Combustion, respiration and rusting are all processes using up oxygen.

 In a laboratory: Thermal decomposition with catalyst of manganese dioxide.

 Combustion of fuels and respiration produce carbon dioxide. However
 Heating Hydrogen peroxide onto manganese dioxide powder
the approximate composition of gases in air remain unchanged overall
2H2O2  2H2O + O2 (MnO2 as catalyst)
because PHOTOSYTHESIS by green plants converts carbon dioxide
 Heating Potassium chlorate with manganese dioxide
back into oxygen and sugar using sunlight.
2KClO3  2KCl + 3O2 (MnO2 as catalyst)
6CO2 + 6H2O + energy  C6H12O6 + 6O2

IDENTIFICATION
When a glowing splint is held at the mouth of a test tube containing oxygen
gas, the splint relights.
NITROGEN
USES Nitrogen is the first element in Group V of the periodic table. It is a colourless,
 Oxygen cylinders in hospitals help people with breathing problems odourless gas which makes up 78% of the air. It is an unreactive gas but does have
 To burn acetylene gas in an oxyacetylene torch when welding steel some uses.
 Rockets burn liquid oxygen as a fuel with liquid hydrogen to form water.
This is a very lightweight fuel. PREPARATION
 Oxygen masks are used in an aircraft if there is an air leak / low pressure  In industry: Fractional distillation of liquid air is applied
 To kill bacteria in the treatment of sewerage
 Used in the production of steel to oxidise any impurities in iron before IDENTIFICATION
producing the type of steel required It does not support burning of other substances.

THREE CHEMICAL PROCESSES INVOLVING OXYGEN: AMMONIA

1. COMBUSTION takes place when any substance reacts with oxygen Ammonia is a colourless, pungent gas, NH3, that is less dense than air.
to produce heat. If flames are produced it is called burning. It is the most soluble of all gases and dissolves in water to form an
e.g. C + O2  CO2 2H2 + O2  2H2O 4Na + O2  2Na2O alkali called aqueous ammonia NH3 (aq).
S + O2  SO2 2Mg + O2  2MgO 2Fe + 3O2  Fe2O3 It is only common alkaline gas and makes most red litmus paper blue.

2. RESPIRATION is the oxidation of sugars in our body to produce energy. Commercially ammonia is very important and prepared by HABER PROCESS
C6H12O6 + 6O2  6CO2 + 6H2O + energy
(This is the reverse of photosynthesis). In a laboratory, ammonia is prepared by heating an
ammonium salt with a base.
[e.g] ammonium nitrate + sodium hydroxide
3. RUSTING occurs when iron comes into contact with water and
69 NH4NO3 + NaOH  NaNO3 +H2O + NH3
oxygen to form rust. Rust is hydrated iron (iii) oxide, Fe2O3.nH2O.
K7, 000
Combustion, respiration and rusting are all processes using up oxygen.
Combustion of fuels and respiration produce carbon dioxide. However the
 FERTILISERS
HABER PROCESS Plants need three essential elements: nitrogen, phosphorous and potassium.
This is the process for the manufacture of ammonia gas from direct Because of a growing demand for food to feed an increasing population, farmers need
combination of NITROGEN and HYDROGEN gases. to rely on fertilisers to provide essential elements needed for crops.
Ammonium nitrate, NH4NO3, is an especially good fertiliser as is contain nitrogen
The Haber Process is a REVERSIBLE reaction from two sources (NH4 and NO3).However EXCESS nitrate fertiliser washed into
streams and rivers can cause EUTROPHICATION.
N2 (g) + 3H2 (g) 2NH3 (g) + heat
Eutrophication is when the excess fertilisers cause the plant life to grow too
much, they then die and bacteria then takes over processing the decaying matter,
FLOW DIAGRAM OF HABER PEOCESS this uses up the oxygen and causes the animal life to also die. The water then
The nitrogen and becomes stagnant.
hydrogen which DO NOT
REACT are REUSED
CARBON
PRESSURE of Carbon is the lightest non-metallic element in GroupIV of the periodic table.
350 atm. is observed TEMPERATURE of It forms the basis of life chemistry. It forms allotropes.
450°C is used
Text box
ALLOTROPES
Allotropes are solid forms of an element with different molecular structures.
IRON is put DIAMOND and GRAPHITE occur naturally as allotropes of carbon.
INCOMING hydrogen as CATALYST
and nitrogen are  DIAMOND is suitable for
mixed together in a  Cutting and Grinding tools because it is
3:1 RATIO
the hardest naturally occurring substance

 GRAPHITE is suitable for

 Lubricant because it is soft and flaky
The ammonia which
due to a layered structure which is held
forms is passed
Hydrogen is obtained through a COOLER to
by a week interaction.
GRAPHITE DIAMOND
from water or natural LIQUEFY and removed LIME
gas (methane)
One of common compounds that carbon takes part in is LIME STONE.
Nitrogen is obtained
There are some kinds of lime. Are you clear which is which?
easily from air (which is
78% nitrogen)
LIME STONE calcium carbonate CaCO3  used for manufacture of iron, making cement.
USES OF AMMONIA:
LIME or QUICK LIME calcium oxide CaO  used for NEUTRALISATION of acidic soil
 to make fertilisers e.g. ammonium nitrate, ammonium sulphate
SLAKED LIME calcium hydroxide Ca(OH)2  used for LIME WATER
 to make nitric acid

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 10.1 Introduction of Organic Chemistry

Organic Chemistry is the branch of chemistry concerned with the

compounds of carbon (except carbonates and oxides of carbon).
COTENT LEARNING OBJECTIVE (Pupils should be able to) “Organic” relates to living “organisms,” and all organic compounds
 Describe a homorogous series as a group of compounds with a
are or have been associated with living material.
general formula, similar chemical properties and showing a
Introduction of
gradation in physical properties.
Organic
 Describe the general characteristics of any homologous series
Chemistry  Define the functional group. Food, fibres, fuels, tyres, plastics and most medicines are all carbon containing
 Name, and draw the structure of the organic compounds compounds known as organic compounds. Although they are made of a few elements
 Describe the properties of alkans such as carbon, hydrogen and oxygen, there are such a large variety of compounds.
 Describe the properties of alkenes
It is because of the ability of carbon atoms to form strong covalent bonds.
 Distinguish between saturated and unsaturated hydrocarbons:
Here we learn general properties of organic compounds and try to classify them.
 Describe the manufacture of alkenes and of hydrogen by cracking
Hydrocarbons
hydrocarbons.
 Name natural gas and petroleum as sources of fuels
 Describe the separation of petroleum by fractional distillation. GENERAL PROPERTIES of ORGANIC COMPOUNDS:
 Name the uses of petroleum fractions
 Describe the properties and use of alcohols  They do not conduct electricity
 Describe formation of ethanol by fermentation  They have low melting point
Alcohols and  Describe the formation of ethanoic acid by the oxidation of  They are flammable and volatile (evaporate so easily)
Acids ethanol  They are insoluble in water but soluble in organic solvents like
 Describe the reaction of ethanoic acid with ethanol to form the
ethanol, acetone, etc.
ester, ethyl ethanoate
 Most of them burn forming carbon dioxide and water
 Describe the structure of the polymer product from a given
[E.g.] Methane : CH4 + 2O2  CO2 + 2H2O
monomer and vice versa.
 Describe the pollution problems caused by non-biodegradable
plastics.
 Identify carbohydrates, proteins and fats as natural polymers
 Describe the formation of addition polymers
Polymer HOMOLOGOUS SERIES
 Describe the formation of condensation polymers
 Describe some natural polymers as possessing the same linkages as Homologous Series is a group of
some synthetic polymers The chemical and physical properties of
compounds with increasing
 Describe the hydrolysis of carbohydorates gives simple sugars an organic compound are determined by number of carbon atoms where
 Describe the hydrolysis of proteins to amino acids its FUNCTIONAL GROUP. Organic
each member differs from the
 Describe soap as a product of hydrolysis of fats compounds with the same functional next consecutive member by
group are grouped into a family called a
another –CH2 unit.
HOMOLOGOUS SERIES.

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 Organic compounds are named according to how many carbon atoms they contain
PREFIX (start with) + SUFFIX (end with)
and which functional group they possess. Table below gives the prefixes and the
1 C atom  “METH” suffixes assigned.
Alkane  “ane”
2 C atoms  “ ETH” [EXAMPLE 1]
Alkene  “ene”
3 C atoms  “ PROP” An organic molecule belongs to the alcohol series and contains 4 carbon atoms.
Alcohol  “ol”
4 C atoms  “BUT” Since the names of alcohols end with ‘-ol’, the molecule will be called butanol.
Carboxylic acid  “oic acid”
5 C atoms  “PENT”
[EXAMPLE 2]
The name of the molecule with formula C2H5 COOH is propanoic acid, since it
There are many homologous series and each series is given a name.
contains 3 carbon atoms and belongs to the carboxylic acid series.
All homologous series have the following characteristics:

1 They have the same general formula.

DRAWING STRUCTURAL FORMULAE OF ORGANIC COMPOUNDＳ
2 They have similar chemical properties because they have the same functional
group, i.e. they undergo the same type of reactions. To describe organic compounds, we often use STRUCTURAL FORMULAE as well
3 They show a trend in physical properties as the molecular mass increases. as molecular formulae. You can find the same number of lines as bond
- As the number of carbon atoms INCREASES: around the atom
A functional group is an
 Melting points and boiling points INCREASE
atom or group of atoms RULES: * Each C has FOUR (4) bonds
 Flammability DECREASES (don’t catch fire so easily)
that give an organic * Each H, OH or COOH has ONE (1) bond
 Viscosity INCREASES (don’t flow so easily) molecule its typical
 Volatility DECREASES (don’t evaporate so easily) chemical properties. STEPS: [E.g.] Ethane CH4, Ethene CH6 and Ethanol CH3OH
1. Write the correct No. of C atoms for compound
Table lists some homologous series of organic compounds.
HOMOLOGOUS SERIES GENERAL FORMULA FUNCTIONAL GROUP C C C C C C
Ethane Ethene Ethanol
ALKANES CnH2n+2 Nil
You could also say
ALKENES CnH2n -C=C- Double bonds -C-C- Single Bonds 2.Draw the bonds connecting these atoms
ALCOHOL CnH2n+1OH -OH Hydroxyl group  REMEMBER DOUBLE BOND for ALKENES
-C=O
CARBOXYLIC ACIDS CnH2n+1COOH | Carboxyl group C–C C=C C-C
O-H 3.Add in any functional groups (OH or COOH) (usually to last C atom)
ESTER CnH2n+1COOCmH2m+1 -COO- Ester functional group
Table : Common homologous series C–C C=C C – C -OH
‘n’ stands for NO. of Carbon atom 4. Lastly fill in the correct no. of H atoms (so that each C atom has 4 bonds)
A FEW OF RULES in a molecule of the compound
NAMING ORGANIC COMPOUNS

=
72
Ethane Ethene Ethanol K7, 000
 In ALCOHOL series In ETHER series
DOUBLE-CHECK;
ALL bonds are drawn
Isomers of DIFFERENT HOMOLOGOUS SERIES have
(no missing bonds)
ISOMERS DIFFERENT CHEMICAL PROPERTIES

Isomers are different compounds which have

the SAME MOLECULAR formula but
DIFFERENT STRUCTURAL formula

 The more carbon atoms the more isomers which are possible

[EXAMPLE 1] C4H10
H
|
H H H H
H H—C—H H
| | | |
| | |
H— C— C — C —C —H
H — C — C — C — H
| | | |
| | |
H H H H
H H H
Butane Isobutane
B.P. = -0.5°C B.P. = - 12°C

Isomers of the SAME HOMOLOGOUS SERIES have

SIMILAR CHEMICAL PROPERTIES
but DIFFERENT PHYSICAL PROPERTIES (like B.P, M.P.)

[EXAMPLE 2] C2H6O
H H H H
| | | |
H— C— C — OH H— C— O —C —H
| | | |
H H H H

Ethanol Diethyl ether

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 10.2 Hydrocarbons Table below shows the first 4 members of the alkane series.
Oil is an essential item to us. Crude oil is mainly composed of some hydrocarbons.
No. of Physical State
It can be separated into Petrol, Kerosene, Diesel oil and so on. CHEMICAL
C ALKANE FORMULA
STRUCTURAL FORMULA at Room
We are going to see the hydrocarbons that are basic organic compounds. Temperature
atoms
H
Hydrocarbons are organic All hydrocarbons have covalent molecules.
|
They are found naturally in PETROLEUM
compounds that contain only 1 Methane CH4 H—C—H GAS
carbon and hydrogen atoms and NATURAL GAS. |
H
H H
ALKANES and ALKENES are HYDROCARBONS (contain only C and H atoms) | |
Alcohols and Carboxylic Acids are not hydrocarbons (contain O atoms as well) 2 Ethane C2H6 H— C—C —H GAS
| |
H H

H H H
SATURATED Hydrocarbons (e.g. ALKANES) | | |
We can classify hydrocarbons 3 Propane C3H8 H— C— C —C —H GAS
 Contain ONLY SINGLE C-C bonds | | |
in terms of bonding.
We will see this issue later. H H H
UNSATURATED Hydrocarbons (e.g. ALKENES)
H H H H
 Contain DOUBLE C=C bonds | | | |
4 Butane C4H10 H— C— C —C —C —H GAS
| | | |
Where ‘n’ is the number of H H H H
ALKANES General formula: CnH2n+2 carbon atoms in one molecule. Table; Properties of the first 4 members of the alkane series.
The alkanes are a family of hydrocarbons, i.e. they contain hydrogen and carbon
atoms only.
Every name ends
They are the main hydrocarbons found in petroleum and natural gas.
with ‘-ane’.

PROPERTIES OF ALKANES Alkanes are covalent compounds with weak intermolecular forces
between the molecules. As the number of carbon increase, the melting
 Alkanes have ALL C-C SINGLE BONDS point and boiling point increase; the first four members are gases, the
 Alkanes are insoluble in water next thirteen members are liquids and the rest are solids.
 Alkanes become more viscous, i.e. more difficult to pour out as the number
of carbon atoms increase.

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 The carbon atoms in alkanes are held together only by –C-C- SINGLE ALKENES General formula: CnH2n
COVALENT BONDS. Thus alkanes are said to be SATURATED. The alkenes also form a family of hydrocarbons-they contain only carbon atoms
and hydrogen atoms.
An organic molecule is said to be saturated if it contains only single
carbon-carbon covalent bonds. In all organic compounds, each carbon atom
They are formed when petroleum fractions undergo cracking.
will form 4 covalent bonds, while H will form 1 covalent bond. If oxygen
atoms are present, each oxygen atom will form 2 covalent bonds. PROPERTIES OF ALKANES
Table below shows the first 3 members of the alkene series.
No. of Physical State
CHEMICAL REACTION OF ALKANES CHEMICAL
C ALKENE FORMULA STRUCTURAL FORMULA at Room
Alkanes are fairly unreactive molecules as their single bonds are strong. atoms Temperature
They are used mainly as fuels to provide heat energy.
H H
Thus they don’t form polymers
COMBUSTION OF ALKANES 2 Ethene C2H4 | | GAS
H— C = C —H
Alkanes burn in air (oxygen) to form carbon dioxide and water.
H H H
[EXAMPLE 1] Methane + oxygen  carbon dioxide + water vapour | | |
CH4 + 2O2  CO2 + 2H2O 3 Propene C3H6 H— C—C = C —H GAS
|
Alkanes can be used as fuels
H
[EXAMPLE 2]
When there is not enough air, burning is incomplete. In this case, soot and H H H H
carbon monoxide are also produced. | | | |
4 Butene C4H8 H— C—C —C = C —H GAS
| |
Ethane + insufficient oxygen
H H
 carbon + carbon monoxide + carbon dioxide + water vapour
3C2H6 + 6O22  4C + CO + CO2 + 9H2O Table: Properties of the first 3 members of alkene family
.
Every name
SUBSTITUTION REACTION
ends with ‘-ene’.
In presence of SUNLIGHT, it undergoes a Substitution Reaction with chlorine Note that alkene family starts with ethene where n=2.
to form chloroalkanes. (i.e. H atoms replaced by Cl atoms) Methene, where n=1 to give the formula CH2, does not exist

[EXAMPLE]
The formulae of each
Methane + Chlorine  Chloromethane + hydrogen chloride member differs from
CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g) The alkenes contain carbon – carbon double bonds (- C = C -).
the previous one by an
This reaction does not take place in This carbon double bond is known as the functional group of
extra
the dark. Sunlight is needed to the alkene family. All alkenes must have this functional group. -CH2- group
+ Cl-Cl  Cl + H-Cl provide energy to break the Cl-Cl
bond to produce chlorine atoms which
then react with the alkane molecule.
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 Any organic compounds with a CARBON = CARBON DOUBLE BOND is said to ADDITION OF HYDROGEN
be UNSATURATED. If a molecules has more than one set of carbon – carbon Alkenes undergo addition reaction with hydrogen gas in the presence of a
double bonds, it is said to be polyunsaturated. nickel catalyst to form alkanes.

CHEMICAL REACTION OF ALKENES [EXAMPLE]

+ H-H  C2H4 + H2  C2H6
Alkenes are more reactive than alkanes because of the carbon = carbon double
bond. The reaction of alkenes takes place at the carbon = carbon double bond. ethene + H2  ethane
During a reaction, the carbon = carbon double bond opens up, allowing the
addition of other molecule onto the alkenes:  This process is known as HYDROGENATION.
Thus they can form polymers Hydrogenation is used in MARGARINE manufacture to change UNSATURATED
VEGETABLE OILS into a solid product.
 The unsaturation in the alkene molecule
is destroyed. A saturated product is
-C=C- + X-Y  formed in which the double bond is COMBUSTION OF ALKENES
replaced by single bonds. An addition Alkenes burn in plenty of air (oxygen) to form carbon dioxide and water.
X Y
reaction is said to have taken place.
[EXAMPLE] ethane + oxygen -> carbon dioxide + water vapour
C2H4 (g) + 3O2(g)  2CO2(g) + 2H2O(g)
An addition reaction is a reaction in
which one molecule adds to another to
Alkenes will produce soot and carbon monoxide when
form a single molecule product.
there is insufficient oxygen for complete combustion.

In addition reactions, molecules are always added across a carbon = carbon TEST FOR UNSATURATION
double bond, i.e. the addition is across adjacent carbon atoms. Hence the We can use the addition reaction as a test to find out if a hydrocarbon is an
final structure of the product will always take the appearance above. alkane or alkene. Fig below shows the testing process.
Liquid alken Gaseous
alkenes
ADDITION OF STEAM
Alkenes react with water (steam) in the presence of phosphoric (V) acid
(H3PO4) catalyst at high temperature and pressure to form alcohols.

[EXAMPLE]
Bromine Bromine
C2H4 + H2O  C2H5OH Bromine becomes Bromine becomes
ethene + steam  ethanol solution SHAKE ! solution
+ H-O-H  colourless colourless

O-H The reddish-brown colour of the bromine solution is decolourised as the bromine
is used in the reaction
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 - This is an addition reaction of bromine [EXAMPLE 1] Big alkane  Smaller alkane + Alkene
C20H42  C12H26 + C8H16
ADDITION OF AQUEOUS BROMINE  Cracking is also used to make hydrogen gas
Alkenes undergo addition reaction with the aqueous bromine. [EXAMPLE 2] Big alkane  Alkene + Alkene + Hydrogen
C18H38  C8H16 + C10H20 + H2
[EXAMPLE] In a laboratory, in order to form alkenes from paraffin oil (big alkane C10H22 ),
C2H4 + Br2  C2H4Br2 Cracking takes place in the way (CATALYTIC CRACKING) shown below.
Ethene +bromine  1,2-dibromoethane + Br-Br 
Br Br
The bromine molecule adds onto the Broken pot
double bond of the ethene molecule. (Catalyst)

FROM THE OBSERVATIONS (Alkene)

 The reddish brown colour of bromine is

ALKENE
quickly decolourised, i.e. the colour of the
 mixture in the test tube changes from
What is shaken with a reddish brown to colourless.
solution of bromine is ….

 There is no reaction. Alkanes do not On heating over a suitable catalyst, it will break down into
ALKANE undergo addition reactions because they smaller molecules. These molecules may include alkenes.
are saturated.
 Cracking is essential to match the demand for fractions containing smaller molecules
PREPARATION OF ALKENES from the refinery process. Some of these smaller molecules are used as chemical
Alkenes are formed when petroleum fractions undergo CRACKING, while Alkanes feedstock, while others are used to produce high grade petrol for motor vehicles.
are the main hydrocarbons found in petroleum and natural gas.

CRACKING SOURCDES OF ENERGY

Big hydrocarbon molecules can be broken up into smaller molecules by a Most of our energy comes from the burning of CRUDE OIL and NATURAL GAS.
process called cracking. The big molecules are passed over a solid CATALYST In some countries, solid COAL is used as fuel.
Crude oil is also known
(aluminium oxide or silicon (V) oxide) at a high temperature (about 600℃),
as PETROLEUM
where they break up to give smaller molecules. FOSSIL FUELS
Fossil fuels are found in the form of crude oil, natural gas and coal
The products of cracking CAN NOT BE PREDICTED accurately. They are formed as dead plant and animal material are subjected to intense
What we know is that at the end of the process, smaller hydrocarbon pressure and heat over millions of years.
molecules (either alkanes or alkenes) and/or hydrogen may be formed. Consequently, these fuels are NON－RENEWABLE energy sources.

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NATURAL GAS 10.3 Alcohols and Carboxylic acids
Natural gas is mostly methane gas (CH4). It burns cleanly in air to form carbon Most of organic compounds in living things contain oxygen. What we have seen
dioxide gas and water: CH4(g) + 2O2 (g)  CO2 (g) + + 2H2 O(g) are hydrocarbons which have carbon and hydrogen only. Now we are going to see
This reaction is highly EXOTHERMIC. Alcohols and Carboxylic acids which contain oxygen.
COAL
Coal is mainly carbon, with small amounts of hydrogen, oxygen, nitrogen and sulphur.
ALCOHOLS General formula: CnH2n+1OH
When it burns in air, the main products are carbon dioxide and water:
Coal + Oxygen in air  Carbon dioxide + Water
Alcohols are - colourless, flammable liquids Functional group of the alcohols
At the same time, small amounts of soot, oxides of sulphur - good solvent and fuel.
and nitrogen and ash (a solid residue) are formed. - soluble in water. -OH (hydroxyl group)
- Coal is not a clean fuel. The sulphur dioxide and nitrogen dioxide gases present in
the waste gases of a coal burning power station are removed by passing them PROPERTIES OF ALCOHOLS
through wet limestone before the waste gases are emitted into the atmosphere. Table below shows the first 4 members of the alcohol series.
No. of Physical State
CRUDE OIL (petroleum) CHEMICAL
C ALCOHOL STRUCTURAL FORMULA at Room
Crude oil (petroleum) is a mixture of hydrocarbons with different carbon chain length. FORMULA Temperature
atoms
Petroleum is quite useless as a mixture; H
 It is usually refined by fractional distillation to separate out its different |
LIQUID
compounds to make useful fuels and petrochemicals. 1 Methanol CH3OH H — C — OH
| (B.P.=64℃)
- Crude oil is separated into 7 fractions H
Boiling H H
FRACTION USE No. of C atoms | |
Point LIQUID
Below 40°C Petroleum Gas Gas fuel 1~3
2 Ethanol C2H5OH H — C — C — OH
| | (B.P.=78℃)
40 -75 °C Petrol / Gasoline Car fuel 4~8 H H
75-150°C Naphtha Chemical feed stock 7 ~ 14
H H H
160-250°C Paraffin / Kerosene Stove fuel / Jet fuel 11~ 15
| | |
250-300°C Diesel Diesel fuel 16 ~ 20 LIQUID
3 Propanol C3H7OH H — C — C — C — OH
300-350°C Lubricant oil Lubricating oil, waxes 20 ~ 35 | | | (B.P.=97℃)
and polishes H H H
Over350°C Bitumen Making roads More than50 H H H H
| | | |
The process is unable to give pure fractions because the boiling Molecules with a short LIQUID
points of the hydrocarbons found in crude oil are too close for 4 Butanol C4H9OH H — C — C — C — C — OH
carbon chains have | | | | (B.P.=117℃)
efficient separation. low boiling points H H H H
Hence, in table, the data quoted shows a range of boiling points while those with long
instead of a single boiling point. carbon chains have Every name Table: Properties of the first 4 members of alcohol family
high boiling points. ends with ‘-ol’.
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 [EXAMPLE]
Ethanol can be oxidized into an organic acid called ethanoic acid
C2H5OH + 2[O]  CH3COOH + H2O
ethanol + from  ethanoic acid + water
oxidising agent
They are NEUTRAL
Alcohols do not have the same empirical formula.
For example, METHANOL is CH3OH, or CH4 O. liquids
ETHANOL is C2H5OH or C2H6O. Both formulae
+ 2[O]  + H2O
cannot be reduced to any simpler form.

As the number of carbon atoms in the alcohol increases,

1 the boiling point increases Oxidation of alcohols can also take place they are left exposed to
2 the solubility in water decreases oxygen in the air for a few days.
For example, if ethanol is left exposed in the air it turns “SOUR.” This
is the common fate of wines and beers which are opened but no drunk.
CHEMICAL REACTION OF ALKENES
he reason is that the ethanol has been oxidised to ethanoic acid

COMBUSTION OF ALCOHOLS
C2H5OH + O2  CH3COOH + H2O
Alcohols burn in plenty of air (oxygen) to give carbon dioxide and water vapour.
The product is a dilute solution of ethanoic acid called vinegar.
This reaction takes place in the presence of bacteria in the air.
[EXAMPLE] Ethanol + oxygen  carbon dioxide + water vapour
2C2H5OH + 6O2  4CO2 + 6H2O
ETHANOL C2H5OH
The reaction gives out lots of heat energy and is exothermic. This is the commonest alcohol, and is a colourless, water-soluble liquid.
In some countries such as Brazil, ethanol is sometimes used USE of ethanol is:
as a fuel in cars in place of petrol. 1 as fuel for vehicles It is sometimes more economical to
2 as solvent for paints and varnishes produce ethanol from ethene gas
3 in alcoholic drinks such as beer and wine obtained by cracking petroleum
fractions.
 Ethanol is Produced by fermentation or As for Process of the addition
OXIDATION OF ALCOHOLS
reaction, you can refer to ‘ ALKENE’
Alcohols can be oxidized to carboxylic acids. the addition reaction of an ethene with steam
This reaction takes place in the presence of an oxidizing agent such as acidified
potassium manganate (VII) or acidified potassium dichromate (VI) FERMENTATION
This is one of methods to prepare ethanol. Fermentation is the conversion of
This process is used all over the world sugars into ethanol and carbon dioxide
for baking, wine-making, and brewing gas by the action of micro-organism
beer. such as yeast, in the absence of air.
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 A solution containing glucose ( a sugar ) is mixed with water and yeast and allowed PHYSICAL PROPERTIES OF CARBOXYLIC ACIDS
to react for a few days in the absence of air.
Table below shows the physical properties of the first 4 members of the series
YEAST If air is present in the mixture,
Glucose Ethanol + Carbon dioxide No. of Physical State
oxidation of ethanol by the CARBOXYLIC CHEMICAL STRUCTURAL
37°C
bacteria in the air will take place C at Room
ACID FORMULA FORMULA Temperature
C6H12O6  2C2H5OH + 2CO2 and the end products will be atoms
water and ethanoic acid.
H— C =O
CONDITIONS FOR FERMENTATION Methanoic LIQUID
1 HCOOH |
 This process must take place in the absence of air (oxygen). Acid OH (B.P.=101℃)
H
 This process takes place at an optimum temperature of 37 ℃. |
Ethanoic LIQUID
2 CH3COOH H— C—C =O
If the temperature goes above 40 ℃, the enzymes in Acid | | (B.P.=118℃)
yeast which catalyse the reaction become denatured H OH
so that they can no longer act as catalysts.
H H
| |
Propanoic LIQUID
3 C2H5COOH H —C — C — C = O
Acid | | | (B.P.=141℃)
H H OH

H H H
| | |
Butanoic LIQUID
4 C3H7COOH H —C —C— C — C = O
CARBOXYLIC ACIDS General formula: CnH2n+1COOH Acid | | | | (B.P.=164℃)
H H H OH
The carboxylic acids form a homologous series. Functional group of
Table: Properties of the first 4 members of carboxylic acid family
They are generally WEAK ACIDS. the Carboxylic acids
So they exhibit normal acidic properties.
-COOH or
The general formula for the carboxylic acids is CnH2n+1COOH,
They exist mainly as molecules and do not form where ‘n’ STARTS WITH 0 for the first member of the series.
hydrogen ions as easily as mineral acids. That is why
the acidity of compounds in this series is weak
The first 4 members are all liquids at room temperature.
As the number of carbon atoms in the molecule increase, the boiling point increase
DO YOU REMEMBER?
Since solution of carboxylic acids are acidic, they will undergo typical The most important carboxylic acid is ETHANOIC ACID. It is used for flavourings
reactions of acids – they will react with metals above hydrogen in the and as a preservative
reactivity series to form hydrogen, with metal carbonates to form salt, carbon
dioxide and water, and with bases to form salt and water. 80
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 PREPARATION OF CARBOXYLIC ACIDS ESTERS General formula: CnH2n+1COOCmH2m+1
Carboxylic acids are prepared by oxidation of alcohols. Functional group
Esters are organic compounds formed of the esters
You can check the previous page! by the reaction of a carboxylic acid and alcohol.
OXIDATION IN AIR
When a solution of a carboxylic acids, for example, ethanol is exposed to air, Esters are volatile fragrant substance -COO-
the oxygen present slowly oxidises ethanol into ethanoic acid in the presence of or
USE OF ESTERS
bacteria. Vinegar, which is a solution of ethanoic acid in water, is made this way
 Flavourings in food
 Ingredients in Perfume (sweet smelling)
OXIDATION OF ETHANOL USING OXIDISING AGENT
The orange acidified potassium dichromate solution turns green in this reaction.
DETERMINATION of NEMES and CHENICAL FORMULAE
The names and formulae of the esters formed follows;
CHEMICAL REACTIO OF CARBOXYLIC ACIDS
Carboxylic acids react with alcohols to form water in a reaction called
esterification. Concentrated sulphuric acid (H2SO4) is used as a catalyst.
Reaction: ORGANIC ACID + ALCOHOL  ESTER + WATER
Name: [ ] anoic acid + { } anol  { } yl [ ] anoate + water
[EXAMPLE] ethanoic acid + ethanol ethyl ethanoate + water
Formula: [ ] COOH + { } OH  [ ] COO { } + H 2O

[EXAMPLES]
+ + H2O I) ethanoic acid + ethanol  ethyl ethanoate + water
CH3COOH + C2H5OH  CH3COOC2H5 + H2O
II) propanoic acid + butanol  butyl propanoate + water
C2H5COOH + C4H9OH  C2H5COOC4H9 + H2O
ETHANOIC ACID loses the –OH group
III) propanoic acid + methanol  methyl propanoate + water
while ETHANOL loses the –H group
C2H5COOH + CH3OH  C2H5COOCH3 + H 2O
to form WATER.

The organic compound formed in esterification is called ESTER

Esters also form a homologous series. SOME ORGANIC REACTIONS

alkane + O2  CO2 + H2O

alkene + O2  CO2 + H2O
Esterification is not the same as neutralization even though water
alkene + H2  alkane
is produced in both reactions. In neutralization, the hydrogen ion
alkene + H2O  alcohol
reacts with the hydroxide ion to form water. In esterification, an
alcohol reacts with a carboxylic acid to form water. alcohol + carboxylic acid  ester + H2O
COMBUSTION: alcohol + O2  CO2 + H2O
OXIDATION: alcohol + 2[O]  carboxylic acid + H2O
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10.4 Polymers SYNTHETIC POLYMERS
Most organic compounds have at most a few tens of atoms. Some have as many as Synthetic polymers are substances such as plastics, and man-made fibers such as
one million atoms. Large molecules can natural materials like proteins, DNA, nylon and terylene. They are fomed by either addition or condensation polymerisation.
cellulose or starch, or Man-made materials like plastics. Such large molecules have
properties which depends on the functional groups they contain and the overall USES OF SYNTHETIC POLYMERS
shape of the molecule itself.
 Poly(ethene): plastic bags, mineral water bottles, cling film
POLYMERS are very large molecules that are formed when thousands of smaller  Nylon: can be made into fibres to make strong ropes(e.g. fishing lines) or
units of identical molecules called MONOMERS are joined together. woven into cloth to make sleeping bags, parachutes, etc.
 Terylene: can be made into fibres and woven into cloth

M -M-M-M-M-M-M-M- Polymers are also called

POLLUTION PROBLEMS OF SYNTHETIC POLYMERS
Monomer molecule polymer molecule MACROMOLECULES

 Plastics burn easily and may produce poisonous gases on combustion.

 The process of joining monomers Bonding that takes place
They need to be coated with fire
to form a polymer is called POLYMERISATION. between monomers are
retardants to reduce the risk of fire If you look at it as a bright side,
covalent bonds
plastics are resistant to corrosion
There are 2 types of polymerisations;
１ADDITION polymerisation – the addition reaction between monomers  They are also non-biodegradable, i.e. they are not decomposed by bacteria in
２CONDENSATION polymerisation – the successive linking together of monomers the ground.
Disposal of plastics is difficult and gives rise to environmental
 Polymers occur NATURALLY or may be SYNTHETIC (man-made) . pollution when they are incinerated or buried in landfills.
Hence there are 2 groups of polymers.
１SYNTHETIC polymers
２NATURAL polymers It is often called PLASTIC NATURAL POLYMERS
Proteins, carbohydrates and gats are natural polymers found in plants and animals
as well as in our food.
All natural polymers are biodegradable,
Addition polymers; unlike most synthetic polymers
PROTEINS
[e.g.] Poly, P.V.C. Teflon
Proteins are needed by both plants and animals mainly for growth, and also to
Synthetic polymers provide enzymes an some energy. Proteins are made by polymerizing amino acids.
Condensation polymers;
POLYMERS [e.g.] Terylene, Nylon
Part of the structure of a protein molecule

Natural polymers Condensation polymers;

[e.g.] Proteins, Fats
Monomers: Amino acids

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FATS HYDROLYSIS OF NATURAL POLYMERS
Fats are naturally occurring esters of a fatty acid and glycerol. The condensation polymerisation reaction in natural polymers is easily reversible
by the process of hydrolysis
They are solid at body temperature and are widely
Hydrolysis is the chemical
used by plants and animals as a means of storing For example, if starch hydrolysis takes
food which can be used as a fuel(energy). reaction of a compound with
place, starch (polymer) will break down
water which causes it to break
into smaller molecules, eventually into
down.
sugars (monomers).
the structure of fats
Monomers: Glycerol Fatty acids
HYDROLYSIS
C3H5(OH)3
Polymer to Monomer
SOAP is formed by boiling fats
STARCH SUGARS
with aqueous sodium hydroxide.
This is because the fatty acids
PROTEINS AMINO ACIDS
react with sodium hydroxide to form
their sodium salts.
FATS FATTY ACIDS  These salts have a WATER-LOVING
END and a WATER-HATING END.
Fatty acids are general names for hydrocarbons which The water-hating end attaches itself
contain the –COOH. Carboxylic acids belong to this. to the grease and the water-loving
end causes the grease to detach i.e.
STARCH HYDROLYSIS
dissolve therefore ‘cleaning’
CARBOHYDRATES STARCH is broken down to form SUGAR by:
Carbohydrates are important nutrients for energy to plants and animals.  ACID HYDROLYSIS (heated with dilute acid)
Examples of carbohydrates are starch, sugar, glycogen and cellulose Acid hydrolysis is slow but eventually the starch is It takes place in the
Carbohydrates are made from small sugar molecules joined together. broken down into glucose, which is the monomer and stomach of mammals.
will not undergo further hydrolysis.

Part of the structure of a Starch molecule

 ENZYME HYDROLYSIS ( by enzyme amylase)
Enzyme breaks down the starch down into the Amylase is found in
disaccharide maltose which contains two glucose saliva in the mouth.
units minus a water molecule.
Monomers: Sugar
[EXERCISE]
Chemical formula of sugar is The diagram below gives a summery of the breakdown of starch to maltose and glucose and
C6H12O6 then to ethanol. (From 2003 national exam.)
General formula of carbohydrates is (a) Name the processes represented by the letters A and B.
Cx(H2O)y (b) What is the purpose of the yeast in process A?

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 YEAST A
(C6H10O5)n Step 3 join the molecules together. The structure can be represented simply as
(C6H12O6)n
SALIVA ACID YEAST shown on the right.
Starch MALT Maltose Glucose Ethanol
Where ‘n’ stands for the
C12H22O11
+ number of monomers in
CO2 the structure
ACID B
ANSWER
(a) A: Fermentation B: Hydrolysis The group –CH2CH2- in the simplified structure is
(b) For catalyst / For the reaction to speed up called the REPEATING UNIT of the polymer.

- Table shows monomers, polymers and their uses.

FORMATION OF POLYMERS MONOMER POLYMER USES
As we have seen, we can classify polymers into two groups in terms of how to form.

 Plastic bags
ADDITION POLYMERS
 Cling film
Addition polymers are synthetic polymers made from unsaturated monomers
 Waterproof sheets
(e.g. alkenes) through an addition reaction. In addition polymerisation,
 Plastic plates
monomers add onto one another to form a single polymer
.
[EXAMPLE] Formation of poly (ethene)
Polyethene is made from ethene molecules. The molecules contain a carbon-
carbon double bond ( -C=C-) that can add onto one another.  Water pipes
 Waterproof sheets
The steps below show how to draw the structure of poly(ethene).  Electrical insulators
Step 1 Draw some ethene molecules side by side:
This is an ADDITION
REACTION. To join monomers
togeter they must havfe either
C = C double bonds or reactive
 Coating for non-stick
functional groups that will link
Step 2 open the double bonds in the molecules: cooking utensils
them together on the left as
 Sealing and Bearings
well as on the right to form a
chain strucure.
Hence monomers should be
UNSATURATED
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 [EXERCISE] Step1 draw the monomers alternately

The structure of a polymer is shown.

From what hydrocarbon is the polymer made?
Draw its structure Step2 take away one water molecule from each pair of monomer molecules.

ANSWER The molecules are said to

condense together to give
To determine the structure of the monomer water.
first identify the repeating unit in the
=
polymer. The monomer is obtained by Step3 join the remaining parts of the monomers together.
‘closing the ends’ of the repeating unit to The repeating unit of terylene
The polymer is an addition polymer. obtain the double bond.
The monomer from which it is made must
contain a carbon – carbon double bond. The repeating unit is 

Hence, the structure of

terylene can be represented
CONDENSATION POLYMERS The units in terylene are joined by the
as
Condensation polymers are made from monomers containing alcohol, carboxylic group of atoms.
acid or amino functional groups which link together. Condensation polymers can We say that terylene has an ESTER LINKAGE.
be natural or synthetic  Polymers containing ester linkages are n
also known as POLYESTERS
In condensation polymerization, monomers join together to form a polymer with
the elimination of small molecules such as water or ammonia. Natural polymers can be formed by ester linkages as well as synthetic polymers.
-Table below shows example polymers which have an ester linkage
[EXAMPLE 1] Formation of terylene POLYMER
SYNTHETIC POLYMERS NATURAL POLYMERS
LINK
Two different monomers join together to form terylene. POLYESTERS
The diol and the dicarboxylic acid can react as monomers to form an ester linkage. ESTER

A condensation polymer
may contain two kinds
of monomer.

Diol is an alcohol containing Dicarboxylic acid is an carboxylic Monomers; Alcohol +Acid

2 -OH groups in its acid containing 2 -COOH groups in Monomers; Glycerol + Fatty acid
molecules its molecules 85
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[EXAMPLE 2] Formation of Nylon Natural polymers can be formed by amide linkages as well as synthetic polymers.
Two different monomers join together to form nylon. -Table below shows example polymers which have an amide linkage
The acid and the amine ends on the monomers can react to form an amide likage. POLYMER
SYNTHETIC POLYMERS NATURAL POLYMERS
LINK
POLYAMIDES
AMIDE
AMINE form a homologous group
which has –NH2 as its functional
group.
 It is derived from ammonia by
Diamine is an amine Dicarboxylic acid
replacement of one or more H
containing 2 -NH2 atoms by hydrocarbons.
groups in its molecule [E.g.] Aniline C6H5NH2
Monomers; Amine +Acid Monomers; Amino acid
Step1 Draw the monomers alternately

[EXAMPLE 3] CARBOHYDRATES
As you have seen, carbohydrates (starch, sugars and cellulose) are natural
Step2 Take away one water molecule from each pair of monomer molecules condensation polymers made up of smaller sugar molecules joined together
in order for molecules to condense together to give water.

Monomer  Repeating unit

Hence, the structure of

terylene can be represented
Step3 join the remaining parts of the monomers together.
as
The repeating unit of terylene
The polymer of Carbohydrates
is like n

Hence, the structure of It has single monomers.

terylene can be represented
The units in nylon are joined by the
as
group of atoms.
We say that terylene has an AMIDE LINKAGE. The units in carbohydrates are joined by the
 Polymers containing amide linkages are n
also known as POLYAMIDES  We say that carbohydrates have CELLULOSE LINKAGES.

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[EXERCISE 1]
Which of the following structures represents Terylene?
(From 2004 National Exam.)

 READ the QUESTION carefully

look out for KEY WORDS for CHEMISTRY
collect any DATA relevant for the question

 THINK how you can use what you KNOW to ANSWER the question
Use the PERIODIC TABLE if necessary

 ANSWER what the question ASKS FOR

Look out for KEY WORDS to answer the EXAM QUESTION
e.g. state, define, describe, calculate, draw
Use the number of MARKS as a guide to HOW DETAILED your answer
should be:
ANSWER: A EACH MARK should be ANOTHER POINT for your answer!!
Terylene is a condensation polymer. ALWAYS SHOW any WORKING OUT
Therefore, the key to work out is to identify the linkages of the polymers.
Terylene has ESRTER LINKAGES -COO.
As you can find between monomers, Figure A has ester linkages.
Common Exam Terminology
“State”: one or two words
“Define”: one sentence
[EXERCISE 2]
“Describe” / “Discuss” / “Explain”: a few sentences
Which pair of substances both contain the linkage shown?
“Compare“ / “Contrast”: differences / similarities between
“Draw”: diagram
A. nylon and terylene
“Calculate”: calculations with working out
B. sugars and protein
C. nylon and protein
D. terylene and poly(ethene)

ANSWER: C
This is AMIDE LINKAGE. Hence polymers to be answered should be polyamides
Nylon is a synthetic polyamide, while protein is a natural polyamide.

KILONA LONA has finished. But your learning still continues.

As long as you’re studying, you’re making a progress.
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[EXAMPLE 1]
(a) State the name of your Chemistry teacher [1 mark]
ANSWER: MR. T. IGUCHI
(b) Define the role of a Chemistry teacher [1 mark]
ANSWER: To teach chemistry to students
(c) Describe your Chemistry teacher in terms of physical appearance
THINK: 5 marks = 5 points!! [5 marks]
 Average height NOT (as need physical appearance):
 Short hair  Punctual
 Wears glasses  Friendly
 Mkuwa (Mzungu)  Strict
 Male  Too lenient
Now write these 5 points into short sentences 
ANSWER:
My chemistry teacher is a male of average height. He has black short hair and
sometimes wears glasses. He comes from Japan.

[EXAMPLE 2]
(a) State the S.I. unit of matter [1 mark]
ANSWER: kilogram

(b) Define diffusion [1 mark]

ANSWER: Diffusion is the movement of particles from an area of high
concentration to an area of low concentration

(c) Describe in terms of Kinetic Theory the process of melting [3 marks]

THINK: melting = solid  liquid {heating}
Kinetic Theory = MOVING particles
 Answer should include all these details and cover at least 3 points (3 marks!!)
ANSWER:
 As a solid is heated, the particles begin to vibrate more and more
 Eventually the particles vibrate so much that they overcome the attractive
forces and begin to move about like particles in a liquid
i.e. the solid has melted to become liquid

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