KAWAMBWA TECHNICAL HIGH SCHOOL
DEPARTMENT OF NATURAL SCIENCES
CHEMISTRY NOTES Grade 10 -12
INTRODUCTION
CHEMISTRY is a branch of science which studies matter and it‟s applications. It deals with both
organic and inorganic matter. Chemistry finds its applications in several industrial processes such as
1. Fractional distillation of petroleum
2. Beer manufacturing processes
3. The chemical processes in the mining industry etc
THE NATURE OF MATTER
MATTER is anything that occupies space and has weight. The three states of matter are
1. Solid
2. Liquid
3. Gas
SOLID
1. Definition: a solid is a substance which has both a fixed shape and a fixed volume
2. Arrangement of particles: in a solid the particles are very close to each other and are tightly
packed.
3. Movement of particles: particles in solids are not free to move about but can vibrate about their
fixed positions.
4. Inter-particle distance: the inter-particle distance in solids is very small
5. Inter-particle forces of attraction: the inter-particle forces of attraction in solids are very strong
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HEATING A SOLID
When the temperature of a solid s increased to such an extent that that it melts; the following changes
occur.
1.
2.
3.
4.
Arrangement of particles: the particles are set apart and become loosely packed.
Movement of particles: the particles tend to have more freedom to move about
Inter-particle distance: the inter-particle distance increases
The inter-particle forces of attraction: the inter-particle forces of attraction become weaker than
before.
LIQUID
1. Definition: a liquid is a substance which has no fixed shape but has a fixed volume
2. Arrangement of molecules: in a liquid the particles are slightly close to each other and are
loosely packed. The molecules in liquids slid on each other.
3. Movement of particles: molecules in liquids are free to move about and slide on each other.
4. Intermolecular distance: the intermolecular distance in liquids is slightly larger than that in
solids.
5. Intermolecular forces of attraction: the intermolecular forces of attraction in liquids are
weaker than that in solids
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COOLING A LIQUID
When the temperature of a liquid is reduced to such an extent that that it solidifies; the following changes
occur.
1. Arrangement of molecules: the molecules are brought closer and become tightly packed.
2. Movement of molecules: the movement of molecules reduces and the particles become
tightly packed.
3. Intermolecular distance: the inter-particle distance reduces.
4. The intermolecular forces of attraction: the inter-particle forces of attraction become
stronger.
GASES
1.
2.
3.
4.
5.
Definition: a gas is a substance which neither has neither a fixed shape nor a fixed volume.
Arrangement of molecules: in a gas the molecules very far away from each other
Movement of particles: molecules in gases move continuously and randomly.
Intermolecular distance: the intermolecular distance in gases very large.
Intermolecular forces of attraction: the intermolecular forces of attraction in gases are very
weak.
THE CHANGE OF STATE
Melting: is the change of state from solid to liquid. E.g. Changing of ice to water.
Melting Point: is the temperature at a substance changes from solid to liquid.
Evaporation/Boiling: is the change of state from liquid to gas. E.g. changing of water to water vapor
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Boiling Point: is the temperature at which a substance changes from liquid to gas.
Condensation: is the change of state from gas to liquid. Egg. The changing of water vapor to water.
Freezing: is the change of state from liquid to solid. E.g. the changing of water to ice.
Freezing Point: is the temperature at which a substance changes from liquid to solid state.
Sublimation: is the change of state from solid directly to gas or vice versa without passing through
the liquid state.
The following substances can sublime:
1. Iodine
2. Ammonium chloride
3. Naphthalene
SUMMARY OF PHASE CHANGE:
PROCESS:
A: Melting.
B: Freezing.
C: Sublimation.
D: Evaporation.
E: Condensation.
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THE PHASE DIAGRAM
Exercise
1. When a substance is heated it changes its physical state according to the model shown.
a) What term is used to describe this change of state : …….Ans: Boiling.
b) Describe what happens to the arrangement and movement of particles when A changes to B
Ans: Arrangement: the particles become scattered and move further away from each other
hence increasing the interparticle distance.
Movement: the particles acquire more freedom to move about and begin to move
randomly and continuously.
c) What happens to the temperature of a pure substance when it melts?
Ans: the temperature will remain constant for some time until a substance melts completely.
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2. The diagram below represents models of physical changes between the three states of matter
a) Name the process
(i)
A: ans: melting
(ii)
B: ans: Freezing
(iii)
C: ans: Sublimation
b) Substance C melts at 232 °C and boils at 2623°C. The diagram below shows how the particles
of substance C are arranged at 200 °C.
Draw similar diagrams in the boxes provided to show how the particles would be arranged at
:
3. Study the following information given in the table
R
T
S
P
Melting point -101°C -73°C -77°C -137°C
Boiling point -35°C -10°C -33°C -0.5°C
a) (i) which substance (s) is/are solid(s) at -75°C
(ii) below what maximum temperature do substance R and P exist as solids
b) Substance B melts from 10°C to 18°C. is substance B a pure substance? State a reason for
your answer.
Ans:
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(a) (i) Substance T (i) below -101 °C (b) Ans: it is not a pure substance Reason: A pure
substance must have a constant melting point.
DIFFUSION
Diffusion is the movement of particles from the region of higher concentration to the region of
lower concentration of the same substancedown the concentration gradient.
Diffusion is said to occur only in
(a) Liquids
(b) Gas
EXAMPLES OF DIFFUSION
1. Evidence of diffusion of a solid: when a crystal of potassium manganate(VII) is dropped
into water, its purple colour slowly spreads throughout the liquid until a uniform purple
colour is observed.
2.
Evidence of diffusion of a liquid: When perfume is sprayed in one corner of a room,
the particles spread until the scent is distributed in all parts of the room.
3. Evidence of diffusion of a gas: when a few drops of bromine is put into a gas jar as
shown below, the bromine will vapourise to fill up the gas jar. A gas jar full of air is then
placed on top of the gas jar full of bromine vapour and the jar lids are removed. The
reddish brown vapour spreads throughout the gas jar over a period of tme, even though
the bromine vapour is denser than air.
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Explanation: the particles of bromine moved around randomly throughout the gas jar. Some bromine
particles moved up from the lower jar into spaces between the air particles in the upper jar. Likewise, the
air particles would have moved downwards into the bottom jar until the two different kinds of particles
get evenly mixed up. Hence the uniformity in colour of the particles is observed in both jars overtime.
RATE OF DIFFUSION
The rate of diffusion is dependent on two factors:
1. Temperature
2. Molecular mass of the particles.
Temperature:
The higher the temperature, the faster the rate of diffusion. This is due to the fact that particles at a
higher temperature have more kinetic energy and hence are able diffuse at greater speeds.
Molecular mass:
At a given temperature, lighter particles diffuse faster than heavier particles.
Evidence of the effect of molecular mass:
1. The reaction between hydrogen chloride gas and ammonia gas can be used to show that
molecular mass affects the rate of diffusion.
A piece of cotton wool is soaked in concentrated hydrochloric acid which gives off hydrogen chloride
gas with molecular mass of 36.5. Another piece of cotton wool is soaked in concentrated ammonia
solution which gives off ammonia gas with molecular mass of 17. Ammonia gas reacts with hydrogen
chloride to produce ammonium chloride which forms a white ring near the hydrochloric acid side.
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Explanation: Since the ammonia gas molecules are lighter, they will move faster than the hydrogen
chloride molecules which are heavier. So the two kinds of molecules will meet and react nearer to the
hydrochloric acid side.
2. A beaker of nitrogen gas is inverted over a porous pot containing carbon monoxide as shown
below.
Nitrogen gas with molecular mass of 28 and carbon monoxide with molecular mass 28 both have
the same density and hence their rates of diffusion are the same. Since the number of molecules
that remain in the porous pot are the same, resulting in no change of pressure, hence no change in
the water level.
3. A beaker of oxygen gas is inverted over a porous pot containing carbon monoxide as shown
below.
Oxygen gas has molecular mass of 32 while carbon monoxide a molecular mass of 28. Clearly
oxygen gas has a greater density than carbon monoxide. Therefore the diffusion rate of carbon
monoxide is greater than that of oxygen. This means that the rate at which carbon monoxide
diffuses out of the pot will be greater than the rate at which oxygen gas diffuses into the pot. This
lowers the pressure in the pot and the water level moves upwards towards the porous pot.
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THE LABORATORY GLASSWARE AND APPARATUS
B: PIPETTE: is filled by suction and is used to accurately measure fixed volumes such as 10 cm3, 20
cm3, 25 cm3 e.t.c.
A: BURETTE: is used to measure accurately a range of volumes with an accuracy of 0.1 cm3 and it has a
tap used to deliver controlled volumes in other containers.
C: MEASURING CYLINDER: is used to measure approximate volumes of liquids.
D: CONICAL FLASK: is used for estimating volumes of liquids, storage of liquids and it is used for
carrying out chemical reactions.
E: FLAT BOTTOMED FLASK: is used for estimating volumes of liquids, storage of liquids and it is
used for carrying out chemical reactions.
J: ROUND BOTTOMED FLASK: is used for estimating volumes of liquids, storage of liquids and it is
used for carrying out chemical reactions. This flask is better suited for carrying out chemical reactions
involving heating.
F: BEAKER: is used only for estimating volumes and storage of chemicals.
G: TEST TUBE: is mainly used for storage of chemical samples and liquids.
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H: TRIPOD STAND AND WIRE GAUZE: are used to provide support when heating reagents.
I: BUNSEN BURNER: is used as a source of heat during heating of reagents.
Yellow flame: is produced when the air hole is closed. This flame produces pollutant gas such as carbon
monoxide.
Blue flame: is produced when the air hole is half open and it is most generally used.
Blue – Green Flame: is produced when the air hole is completely open. It is used for strong heating.
K: ELECTRONIC BALANCE: is used to measure the mass of chemical samples.
L: TRIPLE BEAM BALANCE: is used to measure the mass of substances in the laboratory.
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COLLECTION OF GASES
The method used to collect a gas depends entirely on its properties.There are therefore three
methods used to collect a gas
1. Upward Delivery
2. Downward Delivery
3. Downward dispalcement of water.
UPWARD DELIVERY
Gases which are less dense than air are best collected using the upward delivery method
which is sometimes called downward displacement of air.
Examples of gases collected by this method are : hydrogen gas and ammonia gas
DOWNWARD DELIVERY
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Gases which are denser than air are best collected using the downward delivery method
which is also called upward displacement of air.
Examples of gases collected by this method are : hydrogen chloride and carbon dioxide
DOWNWARD DISPLACEMENT OF WATER.
This method is only suitable for gases whch are insoluble in water.
Examples of gases collected by this method are : hydrogen gas and methane gas
SEPARATION TECHNIQUES
A mixture: is a combination of two or more substances which are physically combned and can only be
separated by physical means.
Examples: i) A mixture of salt and water
ii) a mixture of sugar and wate
iii) crude oil which is a mixture of fuels such as petrol, kerosine, diesel etc.
A solvent : is a liquid substance in which a solute dissolves. E.g water.
A solute: is a solid substance which dissolves in a solvent. E.g sugar, salt.
A solution: is a uniform mixture of a solvent and a solute. E.g salt solution (mixture of salt and water)
COMMON SEPARATION TECHNIQUES
1. Filtration
2. Crystallisaton
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3.
4.
5.
6.
7.
8.
9.
10.
Simple dstillation
Fractional distillation
Chromatography
Separating funnel
Decantation/ sedimantation
Maganetism
Centrifugation
Sublimation
FILTRATION
Filtration is used to separate an insoluble solid from a liquid.
The insoluble solid which remains on the filter paper is called Residue while the liquid whch passes
through a filter paper is called filtrate.
INDUSTRIAL APPLICATION OF FILTRATION
1. In domestic water treatment to remove suspended particles from water.
2. Separation of pencilin from yeast.
CRYSTALLISATION
Crystallisation is used used to separate pure solids in form of crystals from impurities suspended in
solution.
Example: mixture of salt, water and impurities
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The main difference between crystallisaton and evaporation is that in evaporation the entire liquid is
evapotated while in crystallisation only part of the liquid is evaporated in order to saturate the
solution for crystals to grow. A saturated solution is the one with the maximum number of solutes
which can dissolve in a solvent.
INDUSTRIAL APPLICATION OF CRYSTALLISATION
1. In the prepartion of salts such as copper(II) sulphate
2. Obtaining pure sugar
3. Purification of antibiotics
DISTILLATION
Distillation: is a separation technique used to separate pure liquids from liquids- liquid mixtures
by employing evapoartion and condansation.
THE TWO TYPES OF DISTILLATION
1. Simple distillation
2. Fractional distillation
SIMPLE DISTILLATION
Simple distillation: is a separation technique used to obtain a pure liquid from a solution of a
solid.
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INDUSTRIAL APPLICATION OF SIMPLE DISTILLATION
1. To obtain pure water from sea water.
FRACTIONAL DSTILLATION
Fractional distillation : is the separation technique used to separate miscible liquids by using their
boiling points. E.g ethanol and water.
Miscible liquids: are liquids which do not show any layer of separation when they are mixed.
1. Ethanol-water mixture
2. Petrol-kerosine mixture
3. Crude oil.
The liquid with the lowest boiling point will always be distilled first. In the mixture of ethanol
with boiling point (78°C) and water with boiling point (100°C), ethanol will be distilled first.
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INDUSTRIAL APPLICATION OF FRACTIONAL DISTILLATION
1. Separation of liquid air into oxygen, nitrogen and other useful gases
2. Separation of crude oil into petrol, kerosine, diesel and other useful products.
3. Separation of fermented liquor into ethanol and water.
CHROMATOGRAPHY
Chromatography: is a separation technique used separate to
1. Identify a substance
2. Determine the purity of a substance
3. Separate two or more substances with different solubilites in the same solvent.
Its is used mainly to separate dissolved solids such as dyes and pigments by using their solubility
in a solvent.
The main principals of chromatography are
1. Different substances have different solubilities in the same solvent.
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2. The more soluble substance will get carried further by the solvent ahead the less soluble
ones.
There are two types of paper chromatography. And these are ascending and descending order.
FACTORS AFFECTING CHROMATOGRAPHY
1. The separating or stationary media used
2. The surface tension of the solvent
3. The viscosity of the solvent.
Steps for carring out chromatography:
1. Use a pencil and not ink to draw the start line on the paper because the ink from a pen
contains dyes which can also elute or separate and complicate the chromatogram
2. Place the filter paper in the solvent with a start line and spot of sample slightly above the
solvent level. If the startline is below the solvent level, the spot of mixture may dissolve into
the solvent instead of travelling up the chromatography paper.
3. The beaker must be covered when the chromatography is being carried out in order to reduce
evaporation of the solvent from the beaker and from the paper.
RESULTS OF CHROMATOGRAPHY
The piece of paper used in chromatograph which shows the results of separation is called
chromatogram.
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The chromatogram can be summarised as follows:
1. Ink X has 4 dyes whereas Y has only 3.
2. Ink X and Y are made up of 2 common dyes.
The chromatogram below shows one unknown sample and 4 pure substances (A,B,C,D)
1.
2.
3.
4.
How many components are in the unknown sample
Ans: 3 components
Which pure substances are contained in the known sample.
Ans: A,C
which pure substances are not present are not present in the unknown sample:
Ans: B,D
suggest a reason why D did not produce any spot . Ans: because D is insoluble in the solvent.
ATTEMPTS
1. The diagram below shows a chromatogram obtained using solutions of three single dyes (blue,
green and red) and four other solutions (A,B,C and D).
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(a) Which of the solutions A,B,C or D contains
(i)
One dye only ………………………………………………………………………....[1]
(ii)
Three of the dyes ……………………………………………………………………..[1]
(iii)
Green and red only ……………………………………………………………………[1]
(b) In preparing the chromatogram, the following instructions were given. Suggest a reason for each
instruction
(i)
The start line should be drawn with a pencil rather than ink
………………………………………………………………………………………..[1]
(ii)
At the end of the experiment, the solvent front should be near the top of the paper
………………………………………………………………………………………...[1]
(iii)
The spots of solutions and dyes on the starting line should be small
………………………………………………………………………………………….[1]
Solutions
(c)
(iv)
(v)
(vi)
Solution B………………………………………………………………………....[1]
Solution A……………………………………………………………………..[1]
Solution D ……………………………………………………………………[1]
(iv)
(v)
(vi)
Ink contains dyes which may contaminate the solvent and give inaccurate results…...[1]
This is to make sure that the separation of dyes is complete …………….…………...[1]
To prevent spreading of the dyes sideways and thereby getting mixed up with other spots
next to them.[1]
(d)
RF VALUES
1. Calculate the RF values for the components A,B,C and D
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INDUSTRIAL APPLICATION OF CHROMATOGRAPHY
1.
2.
3.
4.
Separating amino acids from proteins
Separating antibiotic drugs from their growing media
Separating pigments from plants.
Indentifying the flavouring components in foodstuff.
SUBLIMATION
Sublimation is the separation technique used to separate substances which can sublime.
e.g. 1. Mixture of iodine and sand
2. mixture of ammonium chloride and sodium chloride
MAGNETIC SEPARATION
Magnetic separation: is the separation technique used to separate magnetic materials from non magnetic
materials.
e.g mixture of iron fillings and sulphur powder.
DECANTATION/ SEDIMENTATION
This is the separation technique used to separate insoluble solids from a liquid mainly by allowing the
mixture to settle so that the solids settle to the bottom while the liquid is poured off or decanted. E.g.
mixture of mealie meal and water.
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SEPARATING FUNNEL.
A separating funnel is a separating technique used to separate immiscible liquids. The immiscible liquids
are liquids which show layers of separation between them because they don‟t mix. E.g cooking oil and
water.
CENTRIFUGATION
Centrifugation is the separating technique used to separate small suspended solids from a mixture with a
liquid which cannot be effectively separated by filtration. During centrifugation, the mixture is put in a
test tube which is mounted on a rotor of a centrifuge. The mixture is then span or rotated at high speed to
allow the solids to settle down while the liquid remains on top.
e.g separation of blood into the liquid part and the solid part
SUMMARY
S/N
1
2
3
TECHNIQUE
Filtration
Crystallization
Fractional distillation
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SUBSTANCES SEPARATED BY THIS
TECHNIQUE
Insoluble solid and liquid
Pure soluble solids from its solution
Miscible liquids with different boiling
EXAMPLES
Sand and water
Salt solution
Ethanol
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points
Pure solvent from its solution
4
Simple distillation
5
Insoluble solids from solvents.
6
7
Decantation/
sedimentation
Separating funnel
Floatation
8
Magnetic separation
9
Paper chromatograph
Magnetic materials from non magnetic
materials.
Dissolved solid pigments and dyes
Immiscible liquids
Less dense solids and liquids
Crude oil
Salt solution e.g sea
water
Mealie meal and water.
Oil and water
Charcoal and salt in
water
Iron fillings and sulphur
powder
Dyes and pigments.
THE CRITERIA OF PURITY
The criteria of purity: is the method used to identify a pure substance.
1.
2.
3.
4.
By measuring the boiling point of a substance
By measuring the melting point of a substance
By measuring the freezing point
By determining the density of a substance.
The melting point, boiling point, freezing point and density of a pure substance are always constant.
For instance the density of water is 1g/cm3 and its boiling point is 100 °C. this shows that water is a
pure substance.
If the boiling point of water varies from 100 °C to 105 °C then the water must contain impurities.
THE BROWNIAN MOTION
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During the Brownian motion experiment, the smoke is put in a glass tube and a source of light is
projected from one side of the glass tube. The microscope is then used to observe the motion of
smoke particles.
1. The air particles cannot be seen with our naked eyes. So their motion cannot easily be observed.
2. Therefore, smoke is used to clearly observe the motion of air particles. The air particles
continuously collide with the smoke particles and this exerts some force on the smoke particles
setting them in motion.
3. The smoke particles will be moving in a zig-zag motion as shown below.
4. The Brownian motion experiment shows that air particles move freely, continuously and
randomly.
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





The periodic table is a table of elements arranged in strict order of their atomic numbers
The horizontal rows are called periods while the vertical columns are called groups. There are 7
periods and 8 groups in the periodic table.
In each group the elements exhibit similar chemical and physical properties because they have
similar electronic structures. Elements in the same group form ions with same formula and have
the same number of outer electrons.
There is a change in each period from metallic to non metallic character as one moves from
the left to the right of the periodic table. This means that all the elements on the left of the
periodic table are metals while those on the right are non-metals.
The elements in the periodic table are arranged based on the following guidelines.
1. The number of electrons in the outer most shell of an atom of any element determines the
group.
2. The number of shells determines the periods
For example: calcium has electronic structure 2,8,8,2. Since there are 2 electrons in its outermost
shell, calcium must be put in GROUP 2 and since it has 4 SHELLS it must be in PERIOD 4.
NB : Hydrogen is placed between the group 1 and group 7 of the periodic table mainly because hydrogen
can form a single positive charge by losing one electron like the elements in group 1 and can form a
single negative charge by gaining one electron like the elements in group 7.
GROUP 1:
The alkali metals
The elements in group 1 are called alkali metals because they form oxides which when dissolved in water
form solutions which are highly alkaline or caustic. These elements include lithium, sodium and
potassium.
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Physical Properties of the group 1 metals.
1.
2.
3.
4.
5.
They all have one electron in the outermost shell of their atoms e.g Li(2,1) Na(2,8,1) and
K(2,8,8,1)
They are very soft and silvery metals which can easily be cut with a razor blade
They have low melting and boiling points
They have low densities and can even float on water
They are good thermal and electrical conductors
Chemical properties
1. They are the most reactive metals in the periodic table
2. They are stored under oil because they react vigorously with water and air [their reaction with
water is highly exothermic and can even cause a fire ]
3. Their reactivity increases with increase in the atomic numbers down the group.
4. They burn in oxygen to produce oxides
Metal + oxygen → metal oxide
(i)
4Li(s) + O2(g) → 2Li2O(s)
(ii)
4Na(s) + O2(g) → 2Na2O(s)
(iii)
4K(s) + O2(g) → 2K2O(s)
5. They also react with water to produce metal Hydroxide and hydrogen gas.
Metal + water →metal hydroxide + hydrogen gas
(i)
2Li(s) + 2H2O(l) → 2LiOH (aq) + H2(g)
(ii)
2Na(s) + 2H2O(l) → 2NaOH (aq) + H2(g)
(iii)
2K(s) + 2H2O(l) → 2KOH (aq) + H2(g)
6. Group 1 metals burn in chlorine with a bright flame to produce chlorides
2Na (s) + Cl2(g) → 2NaCl(s)
GROUP 2 Alkaline Earth Metals
The group 2 elements are called alkaline earth metals because they form oxides which are slightly soluble
in water and form alkaline solutions. They are also found on the earth‟s surface. These metals include
beryllium, magnesium and calcium
Physical properties
1. They all have 2 electrons in the outermost shell of their atoms e.g Be(2,2), Mg(2,8,2) and
Ca(2,8,8,2)
2. They have slightly higher melting and boiling points than the elements in group 2
3. They high densities
4. They are good thermal and electrical conductors
Chemical properties
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7. They are also reactive metals but not as reactive as the elements in the sodium family
8. Their reactivity increases with increase in the atomic numbers down the group.
9. They burn in oxygen to produce oxides
Metal + oxygen → metal oxide
(iv)
2Be(s) + O2(g) → 2BeO(s)
(v)
2Mg(s) + O2(g) → 2MgO(s)
(vi)
2Ca(s) + O2(g) → 2CaO(s)
10. They also react with water to produce metal Hydroxide and hydrogen gas. The only exception
which only reacts with hot water (steam)
Metal + water →metal hydroxide + hydrogen gas
(i)
Ca(s) + 2H2O(l) → Ca(OH)2 (aq) + H2(g)
(ii)
Mg(s) + H2O(g) → MgO(s) + H2(g)
GROUP 7 : The halogens
The non-metallic elements in group 7 of the periodic table are called the halogens. The halogens are also
sometimes called the „salt markers‟. These include fluorine, chlorine, bromine, iodine and astatine.
Physical properties
1.
2.
3.
4.
They have 7 electrons in the outermost shell of their atoms e.g F(2,7) and Cl(2,8,7)
They are the most reactive non-reactive metals which are never found in Free State in nature.
They exist as diatomic molecules e.g F2, Cl2, Br2, I2 and At2
The first two are gases, the third is a liquid while the rest are solids at room temperature
5. They are non metals with very low melting and boiling points
6. They are poor conductors of both heat and electricity
7. Their colours increase in intensity down the group
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8. Their densities, melting and boiling points increase with increase in atomic numbers down the
group.
Chemical properties
1. Their reactivity decreases with increase in atomic numbers down the group.
2. They react with group one metals to form salts
2Na (s) + Cl2(g) →2NaCl(s)
3. They undergo displacement reactions with other halide ions. The most reactive halogen can
displace the less reactive halogen from its compound
Cl2(g) + 2KBr(aq) → 2KCl(aq) + Br2(aq)
Cl2(g) + 2KI(aq) → 2KCl(aq) + I2(aq)
F2(g) + 2KBr(aq) → 2KF(aq) + Br2(aq)
Br2(g) + 2KI(aq) → 2KBr(aq) + I2(aq)
Use of halogens
1. Fluorine is used in toothpaste to help prevent tooth decay
2. Chlorine is put in tap water to kill germs and bacteria
3. Iodine is put in table salt to help prevent a condition called goitre.
GROUP 8 : The noble gases
The group O elements are called the noble gases because they are generally uncreative- they are
chemically stable. They are inert as they have a full outermost shell. These include Helium, neon and
argon
General properties
1.
2.
3.
4.
5.
They all have 8 electrons in the outermost shell except Helium e.g He(2), Ne(2,8) and Ar(2,8,8)
They have very low melting and boiling points
They are all colourless gases
They are poor thermal and electrical conductors
They exist as monatomic molecules. e.g He, Ne, Ar
Use of noble gases.
1. Helium is used to fill hot air balloons due to its low density.
2. Argon is widely used in light bulbs and street lamps to create an inert atmosphere
3. Neon is used in advertising signs as it glows red when electricity is discharged through it.
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TRANSITION ELEMENTS


The transition elements form part of the periodic table between Group IIA and group IIIA of the
periodic table.
They are also called Heavy metals. These include Zinc, Copper, Manganese, iron, Lead e.t.c
General Properties
1.
2.
3.
4.
They have extremely high densities
They have high melting and boiling points
They are strong and hard metals
They form coloured compounds
s/n Coloured compound
1
Copper(II ) sulphate solution
2
Iron (II) sulphate solution
3
Potassium Manganate(VII) solution
Colour
blue
Green
Purple
5. Most of them are multivalent
s/n Transition element Valency
1
Copper (I)
1
2
Copper (II)
2
3
Iron (II)
2
4
Iron (III)
3
6. They are not as reactive as the group 1 and II elements
7. They are good catalysts in many chemical reactions
8. They often form strong alloys among each other.
Use of transition elements
1.
2.
3.
4.
Due to high melting and boiling point, tungsten is used to make filaments in electric bulbs
They are used for making alloys e.g steel
They are used as catalysts e.g iron in the manufacture of ammonia
They are used for making machine parts due to their hardness
SEMI METALS (MATALLOIDS )


Semi metals are elements in the periodic table which exhibit both metallic and non-metallic
character
They are widely used for making computer chips
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A chemical bond is formed when atoms of elements combine together. Atoms combine differently to
form bonds. The three most important types of bonding are Electrovalent, covalent and metallic
bonding.
Electrovalent or ionic bonding

Ionic bonding is the type of bonding which involves the transfer of electrons from a metal to a
non metal. It is only possible for this type of bond to occur between Metals and non metals
because naturally metals are electron donors while non metals are electron accepters. When a
metal donates electrons it acquires a positively charged ion while non metals accept the donated
electrons and form negatively charged ions. These oppositely charged ions attract each other
through strong electrostatic van der Waal‟s forces of attraction and form ionic bonds.
Dot and cross diagrams
1. Sodium Chloride ; NaCl
Sodium(2,8,1) donates one electron to chlorine(2,8,7) ,forming Na+ and Cl – ions respectively.
The oppositely charged ions attract each other by electrostatic forces of attraction forming an
ionic bond.
2. Magnesium oxide, MgO
Magnesium (2,8,2) donates two electron to oxygen (2,6) ,forming Mg2+ and O2 – ions
respectively. The oppositely charged ions attract each other by electrostatic forces of attraction
forming an ionic bond.
3. Magnesium chloride, MgCl2
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Magnesium (2,8,2) donates one electron to each of the two atoms of chlorine (2,8,7) ,forming
Mg2+ and Cl– ions respectively. The oppositely charged ions attract each other by electrostatic
forces of attraction forming an ionic bond.
Crystalline lattice of ionic compounds
Sodium chloride, magnesium oxide and calcium chloride are all ionic compounds. They are
formed purely between metals and non metals. In these compounds the oppositely charged ions
attract each other by electrostatic van der Waal‟s forces of attraction and result in strong
electrostatic ionic lattice. In the crystal lattice of sodium chloride, sodium ion (Na+) is centrally
positioned and is bonded to six chloride ions (Cl–).
1. Sodium chloride crystal lattice
2. Magnesium oxide crystal lattice
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CHARACTERISTICS OF IONIC COMPOUNDS
1. Ionic compounds are formed between metals and non metals and involve the transfer of
electrons from one atom to another
2. All ionic compounds conduct electricity in molten or aqueous state
3. Ionic compounds are soluble in water
4. They have high melting and boiling point
5. They are non volatile
COVALENT BONDING
Covalent bonding is the type of bonding formed by sharing electrons between two non metallic atoms.
Each bonding atom contributes an equal number of electrons to the shared pairs. There are various types
of covalent bonding, but the most common and obvious ones are:
1. Single covalent bonding:
Single covalent bonding occurs when each bonding atom contributes one electron to the shared
pair.
(a) Hydrogen gas ; H2
(b) Water, H2O
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(c) Methane gas, CH4
2. Double covalent bonding
Double covalent bonding occurs when each bonding atom contributes two electrons to the shared
pairs.
(a) Carbon dioxide, CO2
3. Triple covalent bonding
Triple covalent bonding occurs when each bonding atom contributes three electrons to the shared
pairs.
(a) Nitrogen gas, N2
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CHARACTERISTICS OF COVALENT COMPOUNDS
1. Covalent compounds are strictly formed between atoms of non metals and involve the sharing of
electrons
2. All covalent compounds are cannot conduct electricity in any physical state
3. They have very low melting and boiling points because the intermolecular forces of attraction
between molecules are very weak
4. They are insoluble in water but soluble in organic solvents
5. They are highly volatile
METALLIC BONDING
Metals are electropositive with the ability to give away their outermost electrons to achieve the noble gas
structure. Metallic bonding is the type of bonding in which metallic positively charged ions are held
together by delocalized sea of electrons. Metals are good conductors of electricity in either solid or
molten state due to the presence of these free mobile electrons. These mobile electrons are free to move
from one atom to another within the metallic crystalline lattice.
Relative atomic mass(Ar): is the average mass of one atom of an element compared with 1/12th of the
mass of one atom of carbon 12 isotope.
The relative atomic mass has no units.
s/n Atom/element Relative atomic mass
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1
2
3
4
5
6
7
8
9
10
11
Hydrogen
Carbon
Calcium
Chlorine
Magnesium
Oxygen
Potassium
Zinc
Copper
Sulphur
Sodium
1
12
40
35.5
24
16
39
65
64
32
23
Relative molecular mass(Mr): is the total mass of one mole of a compound compared with 1/12th of the
mass of one atom of carbon -12 isotope.
The symbol of relative molecular mass is Mr and it has no units.
Example
 Calculate the relative molecular masses of the following compounds
1. Carbon dioxide: CO2
2. Water ; H2O
carbonate; CaCO3
4. Ammonium carbonate ;( NH4)2CO3
3. Calcium
5. Hydrated copper(II) sulphate; CuSO4.5H2O
THE MOLE CONCEPT



The masses of the basic units of matter i.e atoms, molecules and ions are too small to be
measured individually. It is therefore more convenient to measure them in a large number.
Avogadro’s constant; 6.0 x 1023 is the number of particles contained in one mole of a substance.
A mole: is the amount of substance containing an Avogadro‟s number of particles.
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𝑚𝑎𝑠𝑠 𝑜𝑓 𝑎 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒
𝑚𝑎𝑠𝑠
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =𝑚𝑜𝑙𝑎𝑟
𝑚
𝑛 =𝐴𝑟
or 𝑛 =𝑚
𝑀𝑟
Example
1. Calculate the number of moles in
(a) 80g of calcium
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =80
40 = 2 moles
(b) 19.5g of potassium
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =19.5
39 = 0.5 moles
(c) 4g of hydrogen gas, H2
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =42 = 2 moles
(d) 71g of chlorine gas, Cl2
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =71
35.5 = 2 moles
(e) 4g of oxygen gas,O2
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =432 = 0.125 moles
(f) 25g of hydrated copper sulphate, CuSO4.5H2O
We first calculate the relative molecular mass of this compound.
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =25
250 = 0.1 moles
(g) 9g of water, H2O
9
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 =18
= 0.5 moles
2. Calculate the mass of
(a) 0.25 moles of copper.
Mass = n x Ar = 0.25 x 64 = 16g
(b) 0.01 moles of zinc
Mass = n x Ar = 0.01 x 65 = 0.65g
(c) 0.02 mole of ammonium sulphate, (NH4)2CO3
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Mass = n x Mr = 0.02 x 96 = 1.92g
(d) 0.03 moles of urea, (NH2)2CO
Mass = n x Mr = 0.03x 60 = 1.8g
The number of particles
Examples:
1. How many particles in the following
(a) 0.03 moles of sodium
Number of atoms = n X 6.0 x 1023 = 0.03 mol X 6.0 x 1023 = 1.8 x 1022 atoms of sodium
(b) 48g of maganesium
𝑚𝑎𝑠𝑠 48𝑔
n = 𝐴𝑟 = 24 = 2 mol
Number of atoms = n X 6.0 x 1023 = 2 mol X 6.0 x 1023 = 1.2 x 1023 atoms of magnesium
(c) 0.5 mol of oxygen gas
Number of molecules = n X 6.0 x 1023 = 0.5 mol X 6.0 x 1023 = 3 x 1023 molecules of
oxygen gas
(d) 22g of carbon dioxide
𝑚𝑎𝑠𝑠 22𝑔
n = 𝑀𝑟 = 44 = 0.5 mol
Number of molecules = n X 6.0 x 1023 = 0.5 mol X 6.0 x 1023 = 3 x 1023 molecules of
carbon dioxide
2. How many moles are in 8 x 1020 atoms of copper
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Percentage Mass of an element in a compound
1. Calculate the percentage by mass of
(a) Carbon in carbon dioxide , CO2
The relative molecular mass of CO2 must first be determined. Clearly from this calculation
the mass of Carbon in CO2 is 12.
𝟏𝟐
% C = x 100% = 27.27%
𝟒𝟒
(b) Nitrogen in the ammonium sulphate fertilizer, (NH4)2SO4
%N=
𝟐𝟖
𝟏𝟑𝟐
x 100% = 21.21%
2. The newly manufactured ammonium sulphate fertilizer weighs 396 tonnes. What is the mass of
nitrogen (in tonnes) in this fertilizer.
Determining The Empirical Formula
Empirical formula: is the simplest formula of a compound formed from its mass compositions.
Rules.
1. Identify the elements present in a given compound
2. Identify the mass or percentage composition for each element.
3. Divide each individual mass or percentage composition by the relative atomic mass of each
element
4. Identify the smallest ratio and divide it throughout
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5. The resulting simplest ratios should be taken to represent the number of atoms of each element in
the compound
6. Note that if decimal numbers are obtained the simplest even number can be multiplied
through.
Examples.
1. A sample of iron sulphide contains 5.373g of iron and 4.627g of sulphur. What is the empirical
formula of the compound?
2. A compound has the composition of by mass of 29.4% calcium , 23.5% sulphur and 47.1%
oxygen. Find its empirical formula.
3. What is the empirical formula of a hydrocarbon containing 85.7% carbon?
4. Sodium carbonate has the molecular formula of NayCO3. The relative molecular formula of this
compound is 106. What is the value of y and hence what is the molecular formula of the
compound?
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5. When iron is heated in a stream of dry chlorine, it produces a chloride which contains 34.5% by
mass of iron.
(a) Calculate the empirical formula of this chloride.
(b) If the relative molecular mass of this chloride is 325, what is its molecular formula.
6. What is the empirical formula of glucose; C6H12O6?
7. Vitamin C has the following structural formula.
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What is the empirical formula of vitamin C ?
Ans: its molecular formula is C6H8O6, but its empirical formula is C3H4O3
VOLUME OF GASES:




The volume of one mole of any gas at room temperature and pressure(r,t.p) (temp. 25°C
and pressure of 1atm) is 24 dm3
The volume of one mole of any gas at standard temperature and pressure(s,t.p) (temp.
0°C and pressure of 1atm) is 22.4 dm3
Volume = number of moles X 24dm3
Volume = n X 24dm3
1 litre = 1000cm3= 1000ml = 1dm3
The volume of any gas is directly proportional to the number of moles.
Examples
1. What is the volume at r.t.p of
(a) 0.05 mol of carbon dioxide
Volume = 0.05 mol X 24dm3 = 1.2dm3
(b) 4mol of hydrogen gas
Volume = 4 mol X 24dm3 = 96dm3
(c) 14g of carbon monoxide
𝑀𝑎𝑠𝑠
14
n = 𝑀𝑟 = 28 = 0.5 mol Volume = 0.5 mol X 24dm3 = 12dm3
(d) 4g of oxygen gas
𝑀𝑎𝑠𝑠
4
n = 𝑀𝑟 = 32 = 0.125 mol
Volume = 0.125 mol X 24dm3 = 3dm3
2. How many moles are in
(a) 250cm3 of carbon dioxide at r.t.p
We first convert 250 cm3 to dm3
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(b) 1.2 dm3 of nitrogen gas.
Concentrations Of Solutions
Concentration or molarity is the amount of solute in 1 dm3 of a solution
𝒎𝒂𝒔𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆
Concentration = 𝒗𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏
Concentration = molarity =
in g/dm3
𝒏𝒖𝒎𝒃𝒆𝒓 𝒐𝒇 𝒎𝒐𝒍𝒆𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆
𝒗𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏
in mol/ dm3 or molar
Concentration can only be called molarity if it is expressed in mol/dm3. In both expressions for
concentration the volume must always be in dm3.
2mol/dm3 = 2 molar.
Examples
1. A solution contains 5.0g of HCl in 1.0dm3. calculate its concentration in both g/dm3 and mol/dm3
𝒎𝒂𝒔𝒔
𝟓𝒈
(a) C = 𝒗𝒐𝒍𝒖𝒎𝒆 = 𝟏dm 3 = 5g/dm3
𝒎𝒂𝒔𝒔
𝟓
𝒎𝒐𝒍𝒆𝒔
𝟎.𝟏𝟒 𝒎𝒐𝒍
(b) n= 𝑴𝒓 = 𝟑𝟔.𝟓 0.14 mol
M= 𝒗𝒐𝒍𝒖𝒎𝒆 = 𝟏 𝒅𝒎𝟑 = 0.14mol/dm3
2. A 20 cm3 solution contains 5.0g of H2SO4. Calculate its molarity.
n=
𝒎𝒂𝒔𝒔
𝑴𝒓
=
𝟓
𝟗𝟖
𝟐𝟎
𝒎𝒐𝒍𝒆𝒔
𝟎.𝟎𝟓 𝒎𝒐𝒍
= 0.05 mol volume =𝟏𝟎𝟎𝟎 = 0.02dm3 M= 𝒗𝒐𝒍𝒖𝒎𝒆 = 𝟎.𝟎𝟐 𝒅𝒎𝟑 = 2.5 mol/dm3
calculations involving chemical equations
The following equations can be read stoichiometrically as
1. CaCO3 → CaO + CO2
1 mol
1mol 1mol
One mole of calcium carbonate decomposes to one mole of calcium oxide and one mol of
carbon dioxide
2. N2 + 3H2 → 2NH3
1mol 3mol
2mol
One mole of nitrogen gas reacts with 3 moles of hydrogen and produces 2 moles of ammonia
gas
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3. Fe2O3 + 3CO →2Fe + 3CO2
1mol
3mol
2mol 3mol
One mole of iron(III) oxide reacts with 3 moles of carbon monoxide and produces 2 moles of
iron and 3 moles of carbon dioxide.
Examples
1. Zinc carbonate reacts with dilute hydrochloric acid as follows
ZnCO3(s) + 2HCl(aq) → ZnCl2(aq) + H2O(l) + CO2(g)
What mass of zinc chloride is produced from the reaction of 36.5g of HCl ?


If the answer is required in form of mass and what has been provided is form of mass,
then all calculations should be done in form of mass.
According to the question we shall only relate HCL to ZnCl2 to find the solution.
2. A total volume of ammonia gas produced during the Haber process is 600cm3. The equation
below shows the reaction involved in this process.
N2(g) + 3H2(g) → 2NH3(g)
What volume of hydrogen gas is required to yield the said volume of ammonia?
Mole – mass calculations
1. A solution Q contains 276g of potassium carbonate (K2CO3) in 1 dm3
(a) Calculate the relative molecular mass of potassium carbonate
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(b) What is the concentration in, mol/dm3 of solution Q
 We first change the mass of K2CO3 into moles
(c) Potassium carbonate reacts with dilute sulphuric acid as follows
K2CO3(s) + H2SO4 (aq) → K2SO4 (aq) + H2O(l) + CO2(g)
(1) How many moles of sulphuric acid are reacted?
 We know that the moles of K2CO3 supplied are 2 moles calculated from 276g
 So we use stoichiometric mole ratios to find the moles of H2SO4 reacted.
(2) How many moles of carbon dioxide are produced
(3) What volume of carbon dioxide is liberated when measured at r.t.p.?
LIMITING REACTANT, REACTANT IN EXCESS , THEORETICAL YIELD , ACTUAL
YIELD , PERCENTAGE YIELD AND PERCENTAGE PURITY


In most chemical reactions , there are normally two or more reactants involved in the reaction
The two reactants do not normally finish at the same time, one finishes faster than the other.
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









The limiting reactant: is the reactant which gets depleted (finishes) faster than the other in a
chemical reaction.
In short the reactant that is completely used up during a chemical reaction is called the limiting
reactant. It is limiting in the sense that when it finishes the reaction stops.
The reactant in excess : is the reactant which is not completely used up at the end of a chemical
reaction
The limiting reactant determines the amount of product to be produced.
For instance, if 20g of sodium can be burnt completely in 30cm3 of dry chlorine, but 40cm3 of
chlorine is supplied, despite this excess supply of chlorine, only 30cm3 of chlorine will react with
20g of sodium. Therefore, 10cm3 of chlorine will remain unreacted, while the 20g of sodium has
completely reacted. We can tell that sodium is the limiting reactant while chlorine is the
reactant in excess.
The amount of product is always calculated using the limiting reactant. Because once the limiting
reactant runs out no more products can be produced.
Theoretical yield: is the maximum amount of product that can be formed when the limiting
reactant reacts completely.
Remember that both theoretical and actual yield refer to THE REACTION PRODUCTS
ONLY and not the reactants.
Actual yield: is the amount of product that is actually formed from the laboratory experiment.
Percentage yield: is the ratio of the actual yield to the theoretical yield expressed as a
percentage.

Percentage purity: is the ratio of the mass of a pure substance to the mass of the
sample expressed as a percentage.
Example
1. Sodium chloride can be prepared by the reaction of sodium metal with chlorine gas.
2Na(s) + Cl2(g)→ 2NaCl(s)
Suppose that 6.70 mol of Na reacts with 3.20 mol Cl2.
(a) What is the limiting reagent (reactant)?
 According to the reaction we need 2 moles of sodium to react with 1 mole of chlorine.
 So we need to know how many moles of chlorine can react with 6.70 mol of Na
 From this analysis we shall be able to tell which reactant is in excess and which one is
limiting.
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
From the calculation above 6.70 mol of sodium needs 3.35 mol of Cl2, but only 3.20 mol
of Cl2 is available. This means that chlorine is in short supply and all the available 3.20
moles of chlorine will react while some sodium metal will remain in excess.
(b) What is the mass of NaCl produced during the reaction?
 The calculation of the product, sodium chloride, must be done using the mount of
limiting reactant used up.
 In this case we shall be using 3.20mol of chlorine. It is always important to remember
that the moles of chlorine that reacted are 3.20 mol and not 3.35 mol calculated. This
value of 3.35mol refers to the number of moles of chlorine that would have reacted with
exactly 6.70 mole of sodium.
2. Calcium carbonate is decomposed by heating, as shown in the following equation.
CaCO3(s) → CaO(s) + CO2(g)
a. What is the theoretical yield of CaO if 24.8 g of CaCO3 is heated?
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b. What is the percent yield if 13.1g CaO is produced?
 Actual yield = 13.1 g from the laboratory
 Theoretical yield = 13.8 g estimated stoichiometrically from the equation
3. Manganese dioxide (MnO2) reacts with concentrated hydrochloric acid according to the following
reaction. MnO2(s) + 4HCl(aq) → MnCl2(aq) +2H2O(l) + Cl2(g)
A 4.35g sample of Manganese dioxide was added to 1.0 mol/dm3 of hydrochloric acid. 48 cm3 of
HCl was needed to react with Manganese dioxide in the given sample. Calculate the percentage
purity of Manganese dioxide.
 Our first assumption is that the sample of Manganese dioxide contains impurities. It is
not possible that, the entire 4.35g provided reacted with the acid.
 In order to be sure how much Manganese dioxide actually reacted with the acid, we must
find the number of moles of the acid in 48 cm3 of 1.0mol/dm3 HCl
86

There is now need to use the equation to find out how many moles of MnO2 in the sample
actually reacted


So the amount of pure MnO2 in the sample is 1.044g.
We can now calculate the percentage purity
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ACIDS:
An acid; is a chemical substance which when dissolved in water gives hydrogen as the only
positively charged ion.
The presence of an acid can only be determined by the presence of H+ ions
Classification of acids
Acids are classified into two classes and these are mineral and organic acids
Inorganic or Mineral acids: are artificial acids made from the combination of mineral
substances and are mainly found in the laboratory.
Examples of mineral acids
1. Hydrochloric acid ………HCl
2. Sulphuric acid ……………H2SO4
3. Nitric acid ………………….HNO3
Organic acids: are acids obtained from natural sources such as food staffs
Examples of organic acids
1. Ethanoic acid ………….CH3COOH in vinegar and tomato juice
2. Citric acid – found in citrus foods like lemons, oranges and grapefruit
3. Lactic acid – found in sour milk and yoghurt, and in muscle respiration
4. Tartaric acid – found in grapes
5. Formic acid – found in bee stings
Strength of an acid






The strength of an acid depends on the concentration H+ ions it produces in solution
The more hydrogen ions an acid produces in solution, the stronger the acid is.
The instrument used to measure the potential of hydrogen ions in solution is called pH meter.
pH is the measure of how acidic or alkaline the solution is.
The pH meter runs from 1 to 14. pH 7 indicates neutral media. All pH values below 7 indicate
acidic media while those above 7 indicate alkaline media.
Note that the higher the concentration of H+ ions in solution, the lower the pH value and the
more acidic the solution is.
TYPES OF ACIDS
There are two types of acids and these are strong and weak acids
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Strong acids:
A strong acid is one that is completely ionized in water or aqueous solution
e.g 1. Sulphuric acid ……………………………………….H2SO4 (aq) → 2H+ (aq) + SO4 2– (aq)
2. Hydrochloric acid ……………………………….. HCl (aq) → H+ (aq) + Cl– (aq)
3. Nitric acid ………………………………………………… HNO3 (aq) → H+ (aq) + NO3 – (aq)
A weak acid is one that is only partially ionized in water or aqueous solution. The ionization of most
weak acids is reversible. A reversible reaction: is a reaction which can either proceed to the product or
reactant side depending on the equilibrium conditions available.
Dilute acid: solution containing small amount of acid dissolved in water
concentrated acids: solution containing large amount of acid dissolved in water
Physical properties acids
1. Acids have a sour taste
2. Acids have pH below 7
3. Acids turn blue litmus paper red
Indicators and pH values
The universal indicator solution shows a rainbow of colours in neutral, acidic or alkaline media.
These colours include ROYGBV i.e Red, orange, yellow, green, blue and violet. The green
colour indicates neutral media (pH 7), all colours on the left (red, orange and yellow) indicate
acidic media while those on the left (blue, violet) indicate alkaline media.
Other indicators
s/n Indicator
1
Blue Litmus paper
2
Methyl orange
For the N.S. Dept © 2015
Colour in acidic media
Red
Red
Page 49
3
4
Bromothymol blue Yellow
Phenolphthalein
colourless
Chemical properties of acids.
There are three characteristic chemical reactions of acids and these are:
1. Acids react with fairly reactive metals to produce a salt and hydrogen gas. Hydrogen gas
burns with a „pop‟ sound.
Metal + acid → salt + hydrogen
Ca(s) + H2SO4 (aq) → CaSO4 (aq) + H2(g)
Mg(s) + 2HCl(s) → MgCl2 (aq) + H2(g)
(No reaction with metals below hydrogen in the electrochemical series)
2. Acids undergo neutralization reactions with bases (oxides and hydroxides) to produce a salt
and water only.
Base(hydroxide) + acid → salt + water
1. 2NaOH (aq) + H2SO4 (aq) → Na2SO4 (aq) + 2H2O(l)
2. NaOH (aq) + HCl (aq) → NaCl (aq) + H2O(l)
Base(oxide) + acid → salt + water
1. CuO(s) + 2HNO3(aq) → Cu(NO3)2(aq) + H2O(l)
2. CaO(s) + H2SO4 (aq) → CaSO4(aq) + H2O(l)
3. Acids react with carbonates to produce a salt, water and carbon dioxide.
Carbonate + acid → salt + water + carbon dioxide
1. CaCO3(s) + H2SO4 (aq) → CaSO4(aq) + H2O(l) + CO2(g)
2. ZnCO3(s) + 2HCl (aq) → Zn(Cl)2 (aq) + H2O(l) + CO2(g)
Ionic equations
Rules
1. Complete the given equation and balance it carefully
2. Indicate state symbols i.e (g) for a gas, (s) for a solid, (l) for a liquid and (aq) for aqueous
substances or substances which are soluble in water.
3. Ionize the terms in the equation by using the following guidelines:
(a) The solid products shall not be ionized
(b) All free or uncombined elements shall not be ionized
(c) All molecules be it liquids or gases shall not be ionized.
(d) The solid compound on the reactant side must be ionized.
4. Get rid of spectator ions or ions existing on both sides of the equation.
Example
1. Acid + metal → salt + hydrogen gas
Ca(s) + H2SO4 (aq) → CaSO4 (aq) + H2(g)
Ca(s) + 2H+(aq) + SO42– (aq) → Ca2+(aq) + SO42– (aq) + H2(g)
Ca(s) + 2H+(aq) → Ca2+(aq) + H2(g)
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2. Base (hydroxide) + acid →salt + water
2NaOH (aq) + H2SO4 (aq) → Na2SO4 (aq) + 2H2O(l)
2Na+(aq) + 2OH –(aq) + 2H+(aq) + SO42– (aq) → 2Na+(aq) + SO42– (aq + 2H2O(l)
2OH –(aq) + 2H+(aq) → 2H2O(l)
2H+(aq) + 2OH –(aq) → 2H2O(I)
3. Base (oxide) + acid → salt + water
CaO(s) + H2SO4 (aq) → CaSO4(aq) + H2O(l)
Ca2+(aq) + O2– (aq) + 2H+(aq) + SO42– (aq) → Ca2+(aq) + SO42– (aq) + H2O(l)
2H+(aq) + O2– (aq) → H2O(l)
4. Acid + carbonate → salt + water + carbon dioxide
ZnCO3(s) + 2HCl (aq) → Zn(Cl)2 (aq) + H2O(l) + CO2(g)
Zn2+(aq) + CO32– (aq) + 2H+(aq)+2Cl –(aq) → Zn2+(aq) + 2Cl –(aq) + H2O(l) + CO2(g)
2H+(aq) + CO32– (aq)→ H2O(l) + CO2(g)
Bases
A base is a chemical substance which reacts with an acid to produce a salt and water only. Bases are
usually metal oxides or metal hydroxides
Examples of bases
1.
2.
3.
4.
5.
6.
7.




Sodium hydroxide ….NaOH
Sodium oxide …………Na2O
Potassium hydroxide…KOH
Potassium oxide ………K2O
Calcium hydroxide ….Ca(OH)2
Ammonia……………NH3
Ammonium hydroxide ….NH4OH
An alkali is a soluble base. All soluble bases are alkalis.
Not all bases are soluble. Most bases of group 1 metals are alkalis while those of group two
elements are either slightly soluble or insoluble.
Bases containing transition metals or group three metals are usually insoluble in water.
The AMASOPO bases are alkalis i.e Ammonium hydroxide, sodium and potassium (hydroxide,
oxide) are soluble.
s/n Base
Formula Solubility in water
1
Ammonium hydroxide NH4OH
Soluble
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2
3
3
4
5
6
Sodium hydroxide
Sodium oxide
Potassium hydroxide
Calcium hydroxide
Calcium oxide
Copper (II) hydroxide
NaOH
Na2O
K2O
Ca(OH)2
CaO
Cu(OH)2
Soluble
Soluble
Soluble
Slightly soluble
Insoluble
Insoluble
Physical Properties of bases(alkalis)
1. Bases have a bitter taste and feel soapy to a touch
2. Bases have a pH above 7 or between 8 and 14
3. Bases turn red litmus paper blue
Other indicators
s/n Indicator
Colour in Alkaline media
1
Red Litmus paper Blue
2
Methyl orange
yellow
3
Bromothymol blue blue
4
Phenolphthalein
Red
Chemical properties of bases
There are two characteristic reactions of bases. These are
1. Bases undergo neutralization reactions with acids to produce salt and water only. A
neutralization reaction is the reaction between a base and an acid to produce a salt and water.
Base(hydroxide) + acid → salt + water
NaOH (aq) + HNO3 (aq) → NaNO3 (aq) + H2O(l)
Mg(OH)2 (aq) + 2HCl (aq) → MgCl2 (aq) + 2H2O(l)
Base(oxide) + acid → salt + water
ZnO(s) + 2HNO3(aq) → Zn(NO3)2(aq) + H2O(l)
CaO(s) + H2SO4 (aq) → CaSO4(aq) + H2O(l)
2NH3 (aq) + H2SO4 (aq) → (NH4) 2SO4 (aq)
2. Bases react with ammonium salts to produce a salt, water and ammonia gas
Base(alkalis) + ammonium salt → salt + water + ammonia gas
NaOH (aq) + NH4NO3 (aq) → NaNO3 (aq) + H2O (l) + NH3 (g)
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Classification of oxides
There are four kinds of oxides and these include
1.
2.
3.
4.
Basic oxides
Acidic oxides
Neutral oxides
Amphoteric oxides
Basic oxides
Basic oxides are oxides of metals which react with acids to produce a salt and water only
Examples:
1. Calcium oxide
2. Copper(II) oxide
CaO(s) + H2SO4 (aq) → CaSO4(aq) + H2O(l)
CuO(s) + 2HNO3(aq) → Cu(NO3)2(aq) + H2O(l)
Acidic oxides
Acidic oxides are oxides of non metals which when dissolved in water produce acidic solutions.
Examples
1. Sulphur dioxide
2. Carbon dioxide
3. Nitrogen dioxide
SO2(g) + H2O(l) → H2SO3(aq)
CO2(g) + H2O(l) → H2CO3(aq)
NO2(g) + H2O(l) → HNO3(aq)
Neutral oxides
Neutral oxides: are oxides of non metals with a single oxygen atom in their molecules. They do not
exhibit either acidic or alkaline properties.
Examples
1. Carbon monoxide ….CO
2. Nitrogen monoxide …NO
3. Water ……….H2O
Amphoteric oxides
Amphoteric oxides: are oxides of metals which react with both acids and bases to produce a salt and
water.
Examples.
1. Zinc oxide
2. Lead (II) Oxide
For the N.S. Dept © 2015
ZnO(s) + 2HNO3(aq) → Zn(NO3)2(aq) + H2O(l)
ZnO(s) + 2NaOH(aq) → Na2ZnO2(aq) + H2O(l)
Sodium zincate
PbO(s) + 2HNO3(aq) → Pb(NO3)2(aq) + H2O(l)
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PbO(s) + 2NaOH(aq) → Na2PbO2(aq) + H2O(l)
Sodium plumbate
Al2O3(s) + 6HNO3(aq) → 2Al(NO3)3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH(aq) → 2NaAlO2(aq) + H2O(l)
Sodium aluminate
3. Aluminium oxide
Soil acidity
Crops grow well in a nearly neutral soil with the pH between 6.5 and 8.5. If the soil is acidic a
base is added to neutralize the acid in the soil. Most farmers use solid calcium hydroxide (slaked
lime) to neutralize acidic soil. This is because slaked lime is only slightly soluble in water and
thus does not make the soil any more alkaline.
Salts
A salt is a substance formed when hydrogen in an acid is wholly or partially replaced by a metal. Most
salts are named after the acids from which they are formed.
s/n
1
2
3
4
Acid
Sulphuric acid
Hydrochloric acid
Nitric acid
Carbonic acid
Formula
H2SO4
HCl
HNO3
H2CO3
Salt
Sulphates
Chlorides
Nitrates
Carbonates
Classification of salts
Salts are classified as Normal and acidic salts.
A normal salt: is a salt formed when all the hydrogen atoms in an acid are replaced by a metal.
e.g 1. Sodium chloride NaCl
2. copper(II) sulphate CuSO4
An acidic salt is a salt formed when only part of the hydrogen atoms in an acid are replaced by a metal.
An acidic salt still contains hydrogen atoms in its molecule.
e.g. 1. Calcium hydrogen sulphate
2. sodium hydrogen carbonate
Ca(HSO4)2
NaHCO3
Preparation of salts
The method used to prepare a salt depends entirely on its solubility in water.
SOLUBILITY RULES:
1. All AMASOPO (Ammonium, sodium and potassium) salts are soluble
2. All nitrates are soluble
3. All sulphates are soluble except BACALES i.e Barium, calcium, lead and silver sulphates
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4. All chlorides are soluble except LES i.e Lead and silver chlorides
5. All carbonates are insoluble except AMASOPO i.e ammonium, sodium and Potassium
carbonates
Preparation of insoluble salts





Insoluble salts are prepared by the process known as precipitation
Precipitation is the method used to prepare an insoluble salt from two soluble salts
The solid product produced during precipitation is called precipitate(ppt).
A double decomposition reaction occurs during precipitation. This means that ions are
exchanged between salts.
Some of the insoluble salts prepared by precipitation include
1. Barium sulphate BaSO4
2. Calcium sulphate CaSO4
3. Lead (II) sulphate PbSO4
4. silver sulphate Ag2SO4
5. Lead (II) Chloride PbCl2
6. silver Chloride AgCl
7. Lead (II) iodide
PbI2
Mechanism of preparing an insoluble salt




During precipitation the two starting materials are SOLUBLE salts.
One of the two salts contains the Cation (positive ion) of the salt to be prepared while the
other contains the Anion (negative ion) of the same salt.
The choice of soluble salts is made simple by the use of AMASOPO (ammonium,
sodium and potassium ) salts and nitrates because all these salts are always soluble.
For instance, if we are asked to prepare barium sulphate, BaSO4. We need to separate the
two ions i.e Barium and sulphate. To barium we add a nitrate to make a soluble Barium
Nitrate while to sulphate we can add sodium to make soluble sodium sulphate
Ba(NO3)2 Na2SO4 BaSO4(s) NaNO3
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Examples
1. Preparation of lead (II) iodide, PbI2
Starting materials:
(a) Lead(II) nitrate; Pb(NO3)2
(b) Potassium iodide; KI
Equation
Pb(NO3)2(aq) + 2KI (aq) → PbI2(s) + 2KNO3 (aq)
Preparation method:

The solution of lead(II) nitrate is added t the solution of potassium iodide. The precipitate
(Lead(II) iodide) is filtered and removed. The precipitate is washed with distilled water and
dried between pieces of filter paper.
Preparation of soluble salts

There are two methods used to prepare insoluble salts and these are
(a) Crystallization
(b) Titration
Crystallization


Crystallization is the method used to prepare all soluble salts except (AMASOPO) salts.
This method involves the following three reactions
(a) Action of acids on a fairly reactive metal
(b) Action of acid on an insoluble base
(c) Action of acid on an insoluble carbonate
 One important fact to remember is that during the preparation of a soluble salt using this
method is that one of the starting materials must be an acid.
Example:
Preparation of copper(II) sulphate, CuSO4
Starting materials:
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(a) Copper(II) Oxide, CuO
(b) Sulphuric acid, H2SO4
Equation
CuO (s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
Method of preparation
 The copper(II) oxide powder is continuously added to dilute sulphuric acid whilst stirring
until the oxide is in excess. The excess oxide is then filtered off. The salt solution is then
heated(evaporated) until a saturated solution is obtained. The saturated solution is cooled to
grow salt crystals. The crystals are then filtered, washed and dried between two pieces of
filter paper.
Precautions
1. The water should not be evaporated completely and cooling must be done shortly after saturation.
2. Cooling reduces the solubility of the soluble salt to saturation point until salt crystals are formed
3. The salt crystals must not be heated directly to the source of heat, but must be dried between two
pieces of filter paper. This is to prevent the crystals from losing the water of crystallization.
Titration:
Titration is the method used to prepare a soluble salt from the reaction between an acid and a an alkalis

One important condition for this method is that one of the starting materials is a dilute acid
while the other is a soluble base.



Only AMASOPO (Ammonium, sodium and potassium) salts are prepared by titration
This method works on the basis of neutralization.
Since no solid is involved in this method, the end point of the reaction can only be
determined by the use of indicator solution.
Example
Preparation of sodium chloride, NaCl
Starting materials
(a) Sodium hydroxide, NaOH
(b) Dilute Hydrochloric acid, HCl
Equation
NaOH(s) +2HCl(aq) → NaCl(aq) + H2O(l)
Preparation method

About 25 cm3 of sodium hydroxide of known concentration is pipetted into the conical flask .Two
drops of phenolphthalein are then added to the flask . The burette is filled with dilute
hydrochloric acid. The acid is later delivered from the burette bit by bit into the flask until the
For the N.S. Dept © 2015
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pink colour of the indicator changes to colourless. The volume of the acid used is noted. The
solution formed is a neutral solution of sodium chloride. This solution is then heated to evaporate
the water to saturation point. The solution is cooled and crystals are formed. The pure crystals of
sodium chloride are filtered, washed and dried between two filter papers.
Identification of ions
Test of Anions :
Anion
Carbonate (CO2–
)
Chloride (Cl – )
Iodide (I – )
Nitrate (NO3 –)
Sulphate (SO42–
)
Test
Add any dilute acid
Test result
Effervescence, carbon dioxide
produced
White precipitate
Acidify with dilute nitric acid, then add
aqueous silver nitrate
Acidify with dilute nitric acid, then add
aqueous lead (II) nitrate
Add aqueous sodium hydroxide then
aluminium foil; warm carefully
Acidify with dilute nitric acid, then add
aqueous barium nitrate
Yellow precipitate
Ammonia produced
White precipitate
Equations
1.
2.
3.
4.
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
2KI(aq) + Pb(NO3)2(aq) → PbI2(s) + 2KNO3(aq)
(NH4)2SO4(aq) + Ba(NO3)2 → BaSO4(s) + 2NH4NO3(aq)
Test for aqueous ammonia
Cation
Aluminium (Al3+)
Effect of aqueous sodium hydroxide
White ppt, soluble in excess giving a colourless
solution
Effect of aqueous ammonia
White ppt, insoluble in excess
Ammonium
(NH4+)
Calcium (Ca2+)
Copper(II), (Cu2+)
Ammonia produced on warming
-
White ppt, insoluble in excess
Light blue ppt, insoluble in excess
Iron(II), (Fe2+)
Iron(III), (Fe3+)
Green ppt, insoluble in excess
Red-brown ppt, insoluble in excess
Zinc (Zn2+)
White ppt, soluble in excess giving a colourless
solution
Light blue ppt, soluble in excess
giving a dark blue solution
Green ppt, insoluble in excess
Red-brown ppt, insoluble in
excess
White ppt, soluble in excess
giving a colourless solution
Equations
1. Al3+(aq) + 3NaOH(aq) → Al(OH)3(s) +3 Na+(aq)
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2.
3.
4.
5.
6.
7.
NH4+(aq) + NaOH(aq) → H2O(l) + NH3(g) + Na+(aq)
Ca2+ (aq) + 2NaOH(aq) → Ca(OH)2(s) + 2Na+(aq)
Cu2+ (aq) + 2NaOH(aq) → Cu(OH)2(s) + 2Na+(aq)
Fe2+ (aq) + 2NaOH(aq) → Fe(OH)2(s) + 2Na+(aq)
Fe3+ (aq) + 3NaOH(aq) → Fe(OH)3(s) + 3Na+(aq)
Zn2+ (aq) + 2NaOH(aq) → Zn(OH)2(s) + 2Na+(aq)
Test for gases
Gas
Test and test result
Ammonia (NH3)
Turns damp red litmus paper blue
Carbon dioxide (CO2)
Turns lime water milky
Chlorine (Cl2)
Bleaches damp litmus paper
Hydrogen (H2)
“pops” with a lighted splint
Oxygen gas (O2)
Relights a glowing splint
Sulphur dioxide (SO2)
Turns aqueous potassium dichromate(VI) green
Equations
1. NH3(g) + H2O(l) → NH4OH(aq)
2. CO2(g) + Ca(OH)2(aq)→CaCO3(s) + H2O(g)
3. H2(g) + O2(g) → H2O(l)
GENERAL PROPERTIES OF METALS
1. metals have High density, melting and boiling points
2. Malleable and ductile
malleable: can be bent and beaten into different shapes ductile: can be stretched to form wires
3. metals are excellent conductors of electricity and heat
4. metals are electropositive – they for positive ions by losing electrons
Metals always from positive ions
5. metals are sonorous : ability to produce sound when struck
6. they are lustrous: they have a shine appearance when polished
7. they are solids at room temperature and pressure except mercury.
REACTIVITY SERIES OF METALS
‘POSOCAMAZILHCOMES’
1. Potassium
2. Sodium
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3. Calcium
4. Magnesium
5. Aluminium
6. Zinc
7. Iron
8. Lead
9. Hydrogen
10. Copper
11. Mercury
12. Silver
DISPLACEMENT REACTIONS
1. The more reactive metals high in the reactivity series can displace weak metals lower in the
reactivity series from their compounds. For instance, zinc can displace copper from all its
compounds
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
This explains why copper (II) sulphate solution cannot safely be stored in a zinc container. The
copper (II) sulphate solution can be contaminated with Zinc sulphate produced from the reaction
between copper (II) sulphate solution and atoms of zinc in a zinc container.
2. All metals above hydrogen in the reactivity series can displace hydrogen from acids and water.
While those metals below hydrogen cannot displace hydrogen from any compound.
Zn (s) + H2 SO4(aq)→ ZnSO4(aq) + H2(g)
3. Aluminium though high in the reactivity series does not react with cold water because it forms an
oxide layer which prevents further reactions
REDUCTION OF METAL OXIDES BY HYDROGEN GAS AND CARBON MONOXIDE
1. The removal of oxygen from a substance is called reduction.
2. Both hydrogen and carbon monoxide can reduce METAL OXIDES to free METALS and
WATER OR CARBON DIOXIDE.
3. Note that the oxides of potassium and sodium can NEVER be reduced by either hydrogen gas
or carbon monoxide gas.
4. Other OXIDES of metals can be reduced to free metals.
e.g CaO(s) + H2(g) →Ca(s) + H2O(l)
CuO(s) + CO(g) →Cu(s) + CO2(g)
Fe2O3(s) + 3CO(g) →2Fe(s) + 3CO2(g)
THERMAL DECOMPOSITION OF METAL CARBONATES
1. Most metal carbonates decompose into metal oxides and carbon dioxide
2. Note that potassium carbonate and sodium carbonate can NEVER BE decomposed
3. Carbonates of metals between calcium and copper can decompose into metal oxides and
carbon dioxide
(a) CaCO3 (s) → CaO(s) + CO2 (g)
(b) MgCO3 (s) → MgO(s) + CO2 (g)
(c) CuCO3 (s) → CuO(s) + CO2 (g)
4. Less reactive metals such as silver form carbonates which decompose completely into a metal,
oxygen gas and carbon dioxide.
(d) Ag2CO3 → Ag(s) + O2(g)+ CO2 (g)
METAL EXTRACTION
The method used to extract a metal depends entirely on its reactivity. The more reactive the metal
is, the harder it is to extract from its compounds. All reactive metals are always extracted by
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electrolysis.
s/n Element
1
Potassium
2
Sodium
3
Calcium
4
Magnesium
5
Aluminium
6
Zinc
7
Iron
8
Lead
9
Copper
10
Silver
11
Gold
Method of extraction
These metals can only be extracted by electrolysis of their molten
salts e.g. Na from molten NaCl
These metals can be extracted by heating with a chemical reducing
agent such as hydrogen, carbon and carbon monoxide
These metals occur naturally and remotely in the ground
Metals are extracted from their ores. A METAL ORE is a chemical substance from which a
precious mineral is obtained. The metals ores are normally in form of OXIDES, SULPHIDES
OR CARBONATES. The following are the CHIEF ORES for the listed metals.
s/n Metal
1
Iron
2
Copper
3
4
Ore
Haematite
Cuprite
Malachite
Aluminium Bauxite
Zinc
Zinc blende
Chemical present
Iron (III) oxide , Fe2O3
Cu2O
CuCO3.Cu(OH)2
Aluminium oxide, Al2O3
Zinc Sulphide, Zn2S
ALLOYING METALS
Pure metals are usually too soft and weak for most uses. To improve the strength and hardness, metals are
usually alloyed. An alloy is a uniform mixture of a metal with another element. This other element can be
a metal or a non metal.
Pure metal
alloy
Advantages of alloying metals
1. Alloying improves the hardness of metals
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2. It also improves the resistance of a metal to corrosion
3. Alloying improves the appearance of the metal
SOME COMMON ALLOYS
S/N ALLOY
1
Mild steel
2
Stainless
steel
3
4
5
brass
bronze
solder
ELEMENTS
PRESENT
IN THE ALLOY
Iron, carbon
Iron, chromium, nickel
Copper , zinc
Copper , tin
Lead , tin
PROPERTIES
USES
It is hard
Making car bodies
It is strong
Its resistant to
corrosion
Making machinery
Making cutlery
It is hard
Resistance to corrosion
Good appearance
Low melting point
Making surgical
instruments
Making screws
Making ornaments
Welding metals
RECYCLING METALS
Recycling metals refers to the conversion of scrap metal into useful products again.
ADVANTAGES OF RECYCLING METALS
1.
2.
3.
4.
5.
6.
Recycling saves energy
Recycling preserves scarce and non renewable materials
Recycling creates employment
It is cheaper to recycle some metals like aluminium than it is to extract them from their ores.
Reduces air and water pollution due to the mining activities.
Reduces the amount of land required for the disposal of scrap metal
DISADVANTAGES OF RECYCLING METALS
1. Recycling is a time consuming process and requires massive human resources and efforts.
2. It is expensive to transport scrap metal and sort through wastes for scrap metal.
IRON
Uses of iron
1. It is used for making alloys such as steel.
2. It is in construction work
Rusting of iron
Rusting is the corrosion of iron or steel to form hydrated iron (III) oxide ; Fe2O3 . nH2O. The only two
conditions required for rusting to occur are : WATER and OXYGEN. Only when the two conditions are
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present that is when rusting can occur.
Experiment
METHODS USED FOR PREVENTING RUSTING
1.
2.
3.
4.
Painting
Oiling or greasing
Metal plating – coating iron with a less reactive metal such as tin (tin – plating) prevents rusting
Galvanizing – coating iron with a more reactive metal such as zinc. This is also called sacrificial
protection. The more reactive metal corrodes instead of the iron.
IRON EXTRACTION
Iron is extracted from its CHIEF ORE Haematite, iron (III) Oxide; Fe2O3. One other important ore of
iron is magnetite: Fe3O4, though it is not used for extraction.
RAW MATERIALS:
3. Coke; C
1. Haematite; iron (III) oxide; Fe2O3 2. Limestone; calcium carbonate (CaCO3)
THE BLAST FURNACE
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Process of extraction:










The three raw materials; haematite, limestone and coke are collectively fed from the top of the
furnace as charge.
The hot air (oxygen) is blasted from the bottom of the furnace at a temperature of about 1500 °C.
Coke burns in hot air to produce carbon dioxide. This reaction is highly exothermic and raises the
temperature in the furnace to about 2000°C.
C (s) + O2 (g) → CO2(g)
Carbon dioxide(CO2) is reduced to Carbon monoxide(CO) by more coke
CO2 (g) + C (s) → CO(g)
CO is a reducing agent which reduces haematite to iron metal. This is a REDUCTION reaction.
Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2 (g)
The molten iron obtained is still filled with sand particles called silica (SiO2). Limestone is used
to remove these impurities.
The calcium carbonate is first decomposed to Calcium Oxide(CaO) and carbon dioxide(CO2)
CaCO3(s) → CaO(s) + CO2(g)
The calcium oxide(quick lime ) reacts with sand (silica) to produce silicon Dioxide (CaSiO2)
CaO(s) + SiO2(s) → CaSiO3(s)
Slag, being less dense than iron, floats on iron and is tapped off from the top tap
Molten iron is removed from the bottom as it is much denser than slag.
NB: slag is used in road surfacing and it is sometimes used as fertilizer.
ALUMINIUM
Properties
For the N.S. Dept © 2015
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1.
2.
3.
4.
It is a light weight metal (it has low density)
It is a very good conductor of heat and electricity
It is highly resistant to corrosion (due to its protective oxide layer.)
It is a relatively strong metal
s/n Use
1
Making overhead electrical cables
2
3
4
Making kitchen utensils and food
containers
It is used for making aircraft bodies
It is used for making bicycle and
window frames
Useful Property
It has low density and it is a good conductor of
electricity
It is a good conductor of heat and does not
corrode easily
It is strong and has low density
It is strong and has low density
EXTRACTION OF ALUMINIUM
Aluminium is extracted from its CHIEF ORE Bauxite; Al2O3. Aluminium is a fairly reactive metal
which is extracted by electrolysis.
Electrolysis is the process by which an electrolyte is split up or decomposed by an electric current
passing through it.
An electrolyte: is an aqueous ionic compound, an aqueous acid or alkalis which allows an electric
current to pass through and get decomposed by it.
An Electrode: is a conductor or electric plate which allows electric current to enter or leave an
electrolyte.
The two main types of electrodes are anode and cathode
Anode: is an electrode connected to the positive terminal of the battery and this is where negative ions
are attracted.
Cathode: is an electrode connected to the negative terminal of the battery and this is where positive ions
are attracted.
PROCESS OF EXTRACTING ALUMINIUM





Aluminium is extracted from bauxite (Al2O3) by the process of electrolysis.
The solid bauxite is first heated in order to make a molten electrolyte
This process is however very expensive and uneconomical because the melting point of bauxite
is very high about 2045 °C . This means that huge energy is required to melt bauxite.
In order to lower the working temperature and make the process economical, bauxite (Al2O3) is
dissolved in cryolite (Na3AlF6) which reduces the melting point of bauxite from 2045 °C to
900°C.
The mixture of bauxite and cryolite becomes the electrolyte.
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





Electrolyte: molten mixture of cryolite(Na3AlF6) and bauxite (Al2O3)
Electrodes : (i) Anode : Graphite
(ii) Cathode: Graphite
Ions present : Al3+ and O2 –
Reaction at the anode:
The O2 – ions migrate to the anode and get discharged by losing electrons to form oxygen gas
2O2 –(l) → O2(g) + 4e–
At a working temperature of 900 °C, oxygen gas produced at the anode reacts with carbon in the
graphite anode to produce carbon dioxide. This explains why the graphite anode is continuously
replaced at the anode as it is continuously depleted by this oxidation with oxygen.
Reaction at the cathode:
The Al3+ ions migrate to the cathode and get discharged by gaining electrons to form aluminium
metal. The aluminium metal is deposited at the cathode and settles at the bottom.
Al3+(l) + 3e–→ Al (l)
The molten aluminium metal is tapped off at intervals through an outlet at the bottom.
Overall reaction:
Al2O3(l)→Al(l) + O2(g)
COPPER
USES OF COPPER




Copper is widely used for making electrical wires because it is an excellent conductor of
electricity
It is used for making boilers and water pipes mainly because it is resistant to corrosion and
conducts heat,
Copper is also used for making bullets due to its high density and poisonous nature.
Copper is also used for making alloys.
COPPER EXTRACTION
Copper is extracted mainly from the following ores;
1. Cuprite : copper(I) oxide , Cu2O
2. Malachite: Copper (II) carbonate – copper(II) hydroxide, CuCO3 . Cu(OH)2
3. Chalcocite : copper(I) sulphide; Cu2S
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4. Copper pyrites: copper –iron sulphide; CuFeS2.
Copper is an uncreative metal which is extracted by heating with carbon.
COPPER PROCESSING


the copper ore is first crushed before being taken for flotation, roasting and smelting
the copper ore is sometimes leached in sulphuric acid before being taken for electrolysis
CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
ELECTROLYSIS


very pure copper is needed for electrical conductors
electrolysis is therefore used to produce very pure copper


Electrolyte: copper (II) sulphate solution
Electrodes: (i) Anode: impure copper
(ii) cathode: pure copper
Ions present: Cu2+ , SO42+ , H+ and OH–
Reaction at the anode:
Copper dissolves in solution from the impure anode by losing electrons
Cu(s)→ Cu2+(aq) + 2e–
The impurities at the anode falls away as sludge and settles at the bottom.
This explains why the anode reduces in both size and weight.
Reaction at the cathode:
The Cu2+ ions migrate to the cathode and get discharged by gaining electrons forming copper
atoms.
The yellowish green copper metal is deposited on the cathode
Cu2+(aq) + 2e– → Cu(s)
This explains why the cathode increases in both size and mass
Overall reaction: there is no overall reaction as the concentration of the reactants remains
the same.




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AIR
Air is a mixture of different gasses. The earth is surrounded by the atmosphere which contains air. In air
there is oxygen needed by animals for respiration and carbon dioxide essential for photosynthesis in
plants.
General composition of air.
1.
2.
3.
4.
5.
6.
Nitrogen gas
Oxygen gas
Carbon dioxide
Noble gases
Water vapour
Dust particles.
The composition of dry and clean air is :
1.
2.
3.
4.
Nitrogen gas
Oxygen gas
Carbon dioxide
Noble gases mainly argon
79%
20%
0.04%
0.96%
The percentage of oxygen can be calculated using the apparatus below.
For the N.S. Dept © 2015
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Oxygen % =
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑜𝑥𝑦𝑔𝑒𝑛 𝑖𝑛 𝑎𝑖𝑟
𝑖𝑛𝑖𝑡𝑖𝑎𝑙 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑎𝑖𝑟
𝑋 100%
Example:







The only gas in air that supports combustion is oxygen
If for instance the initial volume of air in the two syringes is 80cm3.
After burning copper powder the volume of air remaining is 64cm3.
Volume of air used up during combustion = 80cm3 - 64cm3= 16cm3.
The obvious assumption is that 16cm3 of air used up during combustion is definitely oxygen gas
as it is the only gas that supports combustion.
16cm 3.
Therefore: oxygen % = 80cm 3. 𝑥100% = 20%
Hence the percentage composition of oxygen in air is calculated as 20%
AIR POLLUTION
Air pollution: is the presence of a substance in air to such an extent that it becomes harmful to both
living things and damages the environment.
An air pollutant: is a substance which accumulates in air to such an extent that it becomes harmful to
both the living things and damages the environment.
The major common air pollutants:
1.
2.
3.
4.
5.
6.
Sulphur dioxide SO2
Carbon monoxide CO
Nitrogen oxides NOX
Lead compounds
Methane gas CH4
Chlorofluorocarbons CFCs
SULPHUR DIOXIDE SO2
Source
1.
2.
3.
4.
Burning of fossil fuels containing sulphur e.g coal and petroleum products
In Car exhaust gases
From the copper processing industries
From the sulphuric acid manufacturing plants
Harmful effects
1. It irritates the eyes and causes breathing difficulties
2. It causes acid rain which damages buildings made of limestone, destroys vegetation and affects
aquatic life
Preventive measure:
1. Exhaust gases from factories in industries are sprayed with calcium hydroxide or limestone
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which absorb the sulphur dioxide
CARBON MONOXIDE CO
Source
1. From incomplete combustion of carbon containing compounds
2. In Car exhaust gases
Harmful effects
1. Carbon monoxide poisoning causes suffocation as it reduces the oxygen carrying capacity of the
blood.
Preventive measure:
1. Use more air during combustion
2. Use a catalytic converter to convert carbon monoxide to carbon dioxide in vehicles.
NITROGEN OXIDES NOX
Source
1. At high temperatures especially during bush fires and lightning, the nitrogen in the air combines
with oxygen to form nitrogen oxides.
2. From car exhaust gases
3. From nitric acid manufacturing plants
Harmful effects
1. Nitrogen oxides form acid rain which affects vegetations, buildings and aquatic life
2. Nitrogen oxides can damage lungs
Preventive measure:
1. The cars should use a catalytic converter to convert nitrogen oxides to harmless nitrogen gas
2. Lower the temperature of the burning fuel
METHANE GAS, CH4
Source
1. Bacterial decay of vegetable matter, animal dung and rubbish.
Harmful effects
1. It is a green house gas which may cause global warming
2. It can combine with nitrogen oxides to produce a photochemical smog
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Preventive measure:
1. Use methane gas a fuel for cooking.
LEAD COMPOUNDS
Source
1. From Car exhaust gases as lead compounds are added to some fuels in order for the car engeines
to run properly
Harmful effects
1. Lead poisoning can cause brain damage and is especially harmful to young children
Preventive measure:
1. Use unleaded fuels
CHLOROFLUOROCARBONS CFCs
Source
1. They are used as refrigerants in air conditioners and refrigerators
2. They are released into the air when aerosols are used.
Harmful effects
1. The chlorine produced when CFCs are decomposed by ultraviolet rays from the sun destroys the
ozone layer.
Preventive measure:
1. Use CFC –free aerosols
WATER


Water is the most abundant liquid on earth
Liquid water covers about 70% of the earth‟s surface
Domestic use of water
Water for domestic use does not have to be pure but safe to drink
1. For cooking
2. For washing
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3. For drinking
Industrial use of water
1.
2.
3.
4.
Water is an essential ingredient in beer brewing
it is used as a coolant in oil fired and electricity powered machines
water is an important source of hydroelectric power
water is a universal solvent and dissolves several chemicals in industries
water pollution
Water pollution: is the presence of a substance in water to such a high concentration that it becomes
harmful to both aquatic life and damages the environment.
The major water pollutants:
1.
2.
3.
4.
5.
Acids
Nitrates and phosphates
Heavy metals
Sewage
Oil
ACIDS
Source of pollutant
1. Acid rain
Harmful effects
1. Kills aquatic creatures. Marine life cannot survive in low pH
2. Damages vegetation in the water.
3. Makes the water unfit for domestic use
NITRATES AND PHOSPHATES
Source of pollutant
1. Excess fertilizers washed off from farmers crops
Harmful effects
1. Causes eutrophication which is the excessive growth of vegetation in the water which removes
dissolved oxygen from the water and suffocates marine life
HEAVY METALS
Source of pollutant
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1. Wastes from the mining industry
Harmful effects
1. Poisonous to mankind
SEWAGE
Source of pollutant
1. Untreated household wastes
2. Excretion from animals
Harmful effects
1. Serious health infections to humans
2. Causes eutrophication in the water sources.
OIL
Source of pollutant
1. Oil spills from water vessels e.g ships
Harmful effects
1. Suffocates and kills marine life
WATER TREATMENT
Water is purified for both industrial and domestic use.
The main stages involved in water treatment are:
1. Sedimentation
2. Filtration
3. Chlorination
Screening:
The raw water obtained from the rivers is first screened to remove large solid impurities.
Sedimentation:
At this stage the coagulating agent called alum (aluminium sulphate) is added to the water to make the
suspended particles stick together into bigger sizes. This is called coagulation.
After coagulation, sedimentation takes place in which the solid particles settle to the bottom of the
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sedimentation tank so that they can easily be removed.
Filtration:
The water is then filtered to remove all the remaining solid particle. It is at this stage that activated
carbon is added to remove foul odour from the water. Lime (calcium oxide) is also added to water to
remove acidity.
Chlorination:
Chlorine is added to the water in order to kill the germs and bacteria.
Water Distribution:
At this stage the water can be distributed in water pipes to consumers as it is clean and safe to drink but it
is not pure. This means that the water contains mineral salts needed by the human body.
AMMONIA PRODUCTION IN THE BLAST FURNACE
Properties of ammonia
1. It is the only alkaline gas which turns dump red litmus blue
2. It has a pungent smell
3. It has the relative molecular mass of 17.
Uses of ammonia
1. Ammonia is used in the manufacture of nitric acid
2. It is used in the manufacture of fertilizers such as ammonium sulphate and ammonium nitrate.
Plants need NPK (nitrogen, phosphorus and potassium) for proper growth. The two fertilizers
mentioned provide nitrogen essential for vegetative growth and making proteins in plants.
3. It is used in the manufacture of explosives
4. It is used for making nylon
Chemical properties of ammonia:
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1. Aqueous ammonia (NH3) reacts with acids to produce ammonium salts
2NH3(aq) + H2SO4(aq) → (NH4)2SO4(aq)
NH3(aq) + HNO3(aq) → NH4NO3(aq)
2. Ammonia gas in the laboratory is produced from the reaction between an ammonium salt and an
alkalis(soluble base)
NH4NO3(aq) + NaOH(aq) → NaNO3 (aq) + H2O(l) + NH3(g)
THE HABER PROCESS
The raw materials for the reaction.
1. Hydrogen gas(H2) from the reaction between methane and steam [CH4(g) + H2O(g) →CO(g) +
3H2(g)]
2. Nitrogen gas (N2) from fractional distillation of liquid air.
Reaction conditions
1. Iron as catalyst
2. Temperature of about 450°C
3. Pressure of about 200 atm – 350 atm
Chemical equation
The haber process




The two gases nitrogen and hydrogen are mixed together in the ratio 1 : 3 and the mixture is
passed over heated finely divided iron catalyst.
At a temperature of 450°C and pressure of 200 atm to 350 atm the two gases combine to form
ammonia gas.
The reaction is reversible and highly exothermic. Since the reaction is exothermic relatively low
temperature favours a shift of equilibrium to the right and more ammonia is produced
The unreacted nitrogen and hydrogen gases are recycled back to the reaction chamber and passed
over heated iron catalyst.
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REACTION FLOW CHART
HYDROGEN GAS
Properties
1. Hydrogen is the first and the lightest element ever known
2. It is also the most abundant element in the universe
3. it puts out a flame with a „pop‟ sound
uses of hydrogen
1. It is used as fuel in rocket engines. [ it is the best fuel as its product during combustion is water
which is not a pollutant]
2. It is used in the manufacture of margarine.
3. It is used in the manufacture of ammonia in the Haber process
Industrial preparation of hydrogen gas
Hydrogen gas is prepared by the process known as steam reforming during which methane gas reacts
with steam to produce carbon monoxide and hydrogen gas. This occurs at a temperature of about 800°C
to 1000°C and pressure of 10 – 50atm.
CH4(g) + H2O(g) → CO(g) + H2(g)
Electrolysis is another method used industry to prepare hydrogen gas.
laboratory preparation of hydrogen gas
Hydrogen in the laboratory is prepared by two processes
1. The action of a fairly reactive metal on water
2. The action of a fairly reactive metal on an acid
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The action of a metal on water
When calcium metal reacts with water, calcium hydroxide and hydrogen gas are produced
Ca(s) + H2O(g) → Ca(OH)2(aq) + H2(g)
The action of a metal on dilute acid
When zinc metal reacts with dilute hydrochloric acid, hydrogen gas is produced together with zinc
Chloride salt
Zn(s) + 2HCl(aq) → Zn(Cl)2(aq) + H2(g)
OXYGEN GAS
Properties of oxygen gas
1. Oxygen makes up about 20% of the atmospheric air
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2. The two allotropes of oxygen are oxygen gas (O2) and ozone(O3)
3. Oxygen also has two isotopes namely oxygen-16 (168O) and oxygen-17 (178O)
4. Oxygen rekindles or relights a glowing splint.
Uses of oxygen
1.
2.
3.
4.
5.
Oxygen is used to produce an oxyacetylene flame for welding
It is used in steel making
It used in hospitals to aid patients who have difficulties in breathing
It is used by mountain climbers and sea divers
It is used by animals and humans for respiration
Industrial preparation of oxygen gas
Oxygen gas is prepared industrially by the following processes
1. Electrolysis of acidified water [2H2O (l) → O2(g) + 2H2(g)]
2. Liquefaction and fractional distillation of liquid air
Liquefaction and fractional distillation of liquid air




The air is first cooled to remove dust particles, carbon dioxide and water vapour
The air is then compressed to a pressure of 100 atmospheres. Compressing the air raises its
temperature.
The compressed is then cooled and allowed to expand as result of which it turns into a liquid
(liquefaction)
The liquid air is then fractionally distilled and nitrogen gas is obtained at –196°C , nobles gases
at –186°C while oxygen gas is collected at –183°C
Laboratory preparation of oxygen gas
Oxygen gas is prepared in the laboratory from three reactions
1. Thermal decomposition of sodium nitrate
2. Thermal decomposition of potassium manganate(VII) or potassium permanganate
3. Catalytic decomposition of hydrogen peroxide
Thermal decomposition of sodium nitrate or potassium manganate(VII)
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2NaNO3(s) → NaNO2(s) + O2(g)
[sodium nitrate →sodium nitrite + Oxygen gas]
2KMnO4(s) →K2MnO4(s) + MnO2(s) + O2(g)
[potassium manganate(VII) →potassium manganate(VI) + manganese dioxide + oxygen gas ]
Catalytic decomposition of hydrogen peroxide


Dilute hydrogen peroxide decomposes smoothly to water and oxygen gas when
dropped in a catalysts of powdered manganese dioxide (MnO2)
2H2O2(l) →2H2O(l) + O2(g)
CARBON AND ITS ALLOTROPES
Properties

Carbon is one of the most abundant elements on earth. It is a major constituent in food stuffs and
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


many other organic compounds. Carbon forms the basis of life chemistry.
It has two main isotopes and these are carbon – 12(126C) and Carbon – 14(146C).
Allotropes: are the different crystalline forms of the same element existing in the same physical
state but with different physical properties. The two most important crystalline allotropes of
carbon are
(1) Diamond and
(2) Graphite
One other allotrope of carbon, though not very common, is buckminsterfullerene which consists
of 60 (C60) carbon atoms arranged in a circular fashion.
Diamond
Diamond is the hardest substance ever known. In its molecule each carbon atom is bonded to four other
carbon atoms forming a tetrahedral crystal. There are no free electrons in the molecule of diamond. This
explains why diamond does not conduct electricity.
Properties of diamond
1.
2.
3.
4.
It has very high melting and boiling point
It is a very poor conductor of electricity [no free electrons]
It is the hardest substance ever known
It has a density of 3.5g/cm3
Uses of diamond
1. It is used for making jewels due to its brilliant lustre
2. Diamond is used for making drill bits due to its hardness
3. It is also used for making cutting and grinding tools mainly because of its hardness
Graphite
Graphite has carbon atoms arranged in layers. There are weak van der waal‟s forces of attraction between
layers, allowing the layers to slide over each other. This gives graphite a soft, smooth and greasy feeling.
In its molecule each carbon atom is bonded to three other carbon atoms while the forth electron is
delocalized.
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Properties of graphite
1.
2.
3.
4.
It is a soft, shiny and black substance
It is a good conductor of electricity
it has a density of 2.25g/cm3
it writes well on paper
uses of graphite
1. It is used as a lubricant in moving parts of machines because it is soft and layered.
2. it is used for making electrodes because it conducts electricity
3. it is used for making leads in pencils because it writes well on paper.
SULPHUR AND SULPHURIC ACID
Properties of sulphur.



Sulphur is a crystalline yellowish solid
The most important isotopes of sulphur are sulphur – 32 (3216S)and sulphur – 34(3416S)
Sulphur occurs as a free element. It is also found in the following compounds
(1) Galena – lead(II) sulphide PbS
(2) Zinc blende – zinc sulphide ZnS
(3) Gypsum – calcium sulphate CaSO4
The allotropy of sulphur
At ordinary temperature Sulphur exists as molecules containing eight sulphur atoms (S8).
The two most important allotropes of sulphur are
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(i)
(ii)
Rhombic sulphur
Monoclinic sulphur
Uses of sulphur
1.
2.
3.
4.
It used in vulcanization of rubber
Sulphur is used in the production of sulphur dioxide
It is used as a fungicide
It is used in the preparation of gun powder
Uses of sulphur dioxide
1.
2.
3.
4.
it is used as a bleaching agent
it is used as a food preservative
it is used in the manufacture of sulphuric acid
it is used in the vulcanization of rubber.
SULPHURIC ACID
Sulphuric acid is manufactured by the Contact process
Raw materials
1. sulphur
2. air
3. water
Reaction Conditions
1. Temperature of 450° C
2. Pressure of 1 atm
3. Vanadium (V) oxide(vanadium pentoxide); V2O5 as catalyst
The contact process
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
Sulphur is burned in the air to produce sulphur dioxide
S(s) + O2(g) → SO2(g)

The sulphur dioxide is cooled and further reacted with oxygen to produce sulphur trioxide
2SO2(g) + O2(g) → 2SO3(g)
In the presence of a catalyst vanadium(V) oxide , at a temperature of 450° C and pressure of
1atm sulphur trioxide is dissolved in concentrated sulphuric acid to produce fuming
sulphuric acid called oleum (H2S2O7)
H2SO4(l) + SO3(g) → H2S2O7(l)
When oleum is finally dissolved in water, sulphuric acid is formed
H2S2O7(l) + H2O(l)→ 2H2SO4(aq)


Uses of sulphuric acid
1.
2.
3.
4.
It used to make detergents
It is used for making fertilizers
It is used in making dyes
It is used as an electrolyte in car batteries.
NITRIC ACID
Properties of nitric acid
1.
2.
3.
4.
Pure nitric acid is colourless
It is a fuming and highly corrosive liquid
It is a strong acid
It is highly volatile
Laboratory reparation of nitric acid
Nitric acid is prepared in the laboratory from the reaction between sodium nitrate and sulphuric acid
NaNO3(s) + H2SO4(aq) → NaHSO4(aq) + HNO3(aq)
Industrial preparation of nitric acid

Nitric acid is prepared by the process known as the OSTWALD process
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Raw materials
1. Ammonia gas
2. Air
3. Water
Reaction conditions
1. Temperature of 900°C
2. Platinum – Rhodium alloy(Pt-Rh) as catalyst
3. Pressure of about 8 atm
The Ostwald process





Ammonia is oxidized to Nitrogen monoxide and water at a temperature of 900°C,
pressure of 8 atm and over a catalyst of platinum-rhodium alloy
4NH3(g) + 5O2(g) → 4NO (g) + 6H2O(l)
Nitrogen monoxide produced further reacts with air at 25 °C to produce Nitrogen Dioxide
2NO(g) + O2(g) → 2NO2 (g)
The nitrogen dioxide is then dissolved in water to produce Nitric acid and nitrogen monoxide
3NO2(g) + H2O(g) → 2HNO3(aq) + NO(g)
Nitrogen monoxide produced as a side product is recycled back to the reactor.
If very concentrated nitric acid is needed, the dilute nitric acid is fractionally distilled with
sulphuric acid.
Uses of nitric acid
1. it is used in the manufacture of explosives e.g (TNT) trinitrotoluene
2. it is used in the manufacture of fertilizers
3. it is also used in the manufacture of dyes
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A redox reaction: is the one where both oxidation and reduction occur simultaneously.
Redox in terms of oxygen transfer




oxidation : is the gain in oxygen
NB: Any substance which gains oxygen is oxidized
Reduction: is the loss in oxygen
NB: Any substance which loses oxygen is said to be reduced
Oxidizing agent: is a substance which adds oxygen to another substance
NB: An oxidizing agent is always reduced.
Reducing agent: is a substance which removes oxygen from another substance
NB: A reducing agent is always oxidized.
Examples:
1. CuO (s) + H2(g) → Cu(s) + H2O(l)
Oxidized: H2 (gained oxygen)
Reduced: CuO (lost oxygen)
Oxidizing agent: CuO (added oxygen to H2)
Reducing agent: H2 (removed oxygen from CuO)
2. Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
Oxidized: CO (gained oxygen)
Reduced: Fe2O3 (lost oxygen)
Oxidizing agent: Fe2O3 (added oxygen to CO)
Reducing agent: CO (removed oxygen from Fe2O3)
Redox in terms of Hydrogen transfer (OIL RIG )

oxidation : is the loss in hydrogen
NB: Any substance which loses hydrogen is oxidized
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


Reduction: is the gain in hydrogen
NB: Any substance which gains hydrogen is said to be reduced
Oxidizing agent: is a substance which gains hydrogen.
NB: An oxidizing agent is always reduced.
Reducing agent: is a substance which loses hydrogen
NB: A reducing agent is always oxidized.
Examples
1. H2S(g) + Cl2(g) → 2HCl(g)
+ S(s)
Oxidized: H2S (lost hydrogen)
Reduced: Cl2 (gained hydrogen)
Oxidizing agent: Cl2 (gained hydrogen)
Reducing agent: H2S (lost hydrogen)
Redox in terms of electron transfer (OIL RIG )




oxidation : is the loss of electrons
NB: Any substance which loses electrons is oxidized
Reduction: is the gain in electrons
NB: Any substance which gains electrons is said to be reduced
Oxidizing agent: is a substance which gains elctrons.
NB: An oxidizing agent is always reduced.
Reducing agent: is a substance which loses electrons
NB: A reducing agent is always oxidized.
Examples:
1. Zn(s) + CuSO4(s) → ZnSO4(aq) + Cu(s)


Zinc is oxidized as it loses electrons to form Zn2+ (Zn to Zn2+ ) [it is a reducing agent]
Copper is reduced as it gains electrons to form Cu(s) (Cu2+ to Cu ) [ it is an oxidizing
agent]
Redox in terms of oxidation numbers


oxidation : is the increase in oxidation number
NB: Any substance whose oxidation state increases is said to be oxidized
Reduction: is the decrease in oxidation number.
NB: Any substance whose oxidation state decreases is said to be reduced.
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Oxidation state guidelines
1. The oxidation number of any element in its free state or uncombined state is zero(0)
s/n Element Oxidation state
1
H2(g)
0
2
Mg(s)
0
3
Cl2(g)
0
2. For a single monatomic ion its oxidation number is equal to its charge.
s/n Element Oxidation state
1
Mg2+
+2
2
O2–
-2
3
Cl–
-1
3. The total sum of oxidation numbers of elements in a compound is equal to its overall charge.
4. The oxidation number of oxygen in all of its compounds is –2.
5. The oxidation number of hydrogen in almost all its compounds is + 1 except in hydrides where it
shows – 1 oxidation state.
6. The oxidation number of almost all metals is equal to their ionic charge
s/n Element Oxidation state
1
Mg2+
+2
2
Na+
+1
3+
3
Al
+3
Examples
1. What is the oxidation state of
(a) Sulphur in sulphur dioxide, SO2
 We know that the overall charge for SO2 is zero , 0
 we also know the oxidation state of oxygen as –2.

(b) Manganese in manganate(VI), MnO42The overall charge is now –2
For the N.S. Dept © 2015
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Some common oxidizing agents
1. Oxygen gas, O2
2. Chlorine gas, Cl2
3. Ozone, O3
4. Nitric acid, HNO3
5. Potassium dichromate, K2Cr2O7
Some common reducing agents
1. Hydrogen gas, H2
2. Carbon monoxide, CO
3. Carbon, C
4. Sulphur dioxide, SO2
EXPERIMENT TO TEST FOR AN OXIDIZING AGENT
1. When aqueous chlorine is added to aqueous potassium iodide containing starch, a dark blue
colourization occurs indicating the presence of iodine (I2) in solution.
Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(s)
2. When aqueous copper (II) sulphate is added to aqueous potassium iodide, a dark blue
colourisation is seen indicating the presence of iodine(I2) in solution.
2CuSO4(aq) + 4KI(aq) → 2K2SO4(aq) + 2CuI(aq) + I2(s)
EXPERIMENT TO TEST FOR A REDUCING AGENT
1. When acidified potassium manganate(VII), KMnO4, is added to any suspected reducing
agent, the purple colour of the permanganate is decolourised to a pink colour.
2. When a suspected reducing agent is passed through a solution of potassium dichromate
,K2Cr2O7, the orange colour of the dichromate is discharged to a green colour.
Attempts
1. For each of the reactions below give the formula of the substance which is oxidized.
(a)
Zn (s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
Oxidized: Zn (its oxidation number increases from 0 to +2)
(b)
Cl2(g) + KI(aq) → 2KCl(aq) + I2(s)
Oxidized: KI (Iodine(I) in KI has its oxidation number increased from -1 to 0)
(c)
Fe 2+ (aq) + NO3– (aq) + 2H+(aq) →Fe3+(aq) + NO2(g) + H2O(l)
Oxidized: Fe 2+ (its oxidation number increases from +2 to +3)
(d)
Cl2(g) + Br –(aq) →Br2(g) + 2Cl –(aq)
Oxidized: Br –(its oxidation number increases from - 1 to 0)
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(e)
(f)

Fe 2+ (aq) + Cl2(g) → Fe3+(aq) + 2Cl –(aq)
Oxidized: Fe 2+ (its oxidation number increases from +2 to +3)
Fe 2+ (aq) + 8H+(aq) + MnO4 – (aq) → Fe3+(aq) + 4H2O(l) + Mn2+
Oxidized: Fe 2+ (its oxidation number increases from +2 to +3)
The speed of a chemical reaction is very important in a chemical industry.
Collision theory






A chemical reaction can only occur when the particles of the reactants collide with each other.
However, not all collisions will result in the formation of products
Therefore, effective collisions are required in a chemical reaction. Effective collisions only occur
when the reactant particles have enough energy to overcome the activation energy of the reaction
and when particles collide in the correct orientation.
Activation energy is the minimum amount of energy required to make the reaction take place.
The speed of any chemical reaction depends on the number of effective collisions between reactants.
The greater the number of effective collisions, the higher the rate of reaction
Factors that affect the rate of chemical reaction
1.
2.
3.
4.
5.
Particle size
Pressure
Temperature
Catalyst
Concentration of the reactants.
Effect of Particle size.

Decreasing the size of reactant particles will increase the speed of reaction. This is because by
breaking the reactant particles, the surface area increases, and this in turn results in more particles
being able to collide more frequently.
Effect of Pressure.

If one of the reactants is a gas, increase in pressure will increase the speed of reaction. Increasing
the pressure brings the reactant particles closer to each other in a smaller volume, resulting in
increased effective collisions.
Effect of Temperature .

Increasing the temperature of the reactants will result in increased speed of reaction. This is because
at high temperatures, the reactant particles will have greater kinetic energy, resulting in more
frequent effective collisions.
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Effect of Concentration

An increase in the concentration of one or more of the reactants will increase the speed of reaction.
This is due to the fact that increasing the concentration introduces more particles in a given volume
and this results in frequent effective collisions.
Effect of Catalyst


A catalyst is a chemical substance which alters the rate of chemical reaction without itself
undergoing any chemical change.
Addition of a catalyst to a chemical reaction will increase the speed of the chemical reaction because
the catalyst lowers the activation energy of the reaction. Apart from colliding with each other, the
reactant particles can also collide effectively with the catalyst. This will in turn increase the rate of
reaction.
Measurement of the rate of reaction
The speed of any reaction can be determined by the following three methods.
1. Measuring the time taken for the reaction to be completed.
2. Measuring the volume of gaseous product over a fixed time interval
3. Measuring the amount of reactant left over a fixed time interval.
Measuring the time taken for the reaction to be completed.



Two pieces of magnesium ribbon 2 cm in length each are put in two separate beakers
A sample of hydrochloric acid is added to one beaker while sulphuric acid is added to the other.
Time taken for the magnesium ribbon to react in each beaker is recorded.
For the N.S. Dept © 2015
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


With HCl, the relative rate of reaction = 1/100= 0.01 s –1
With H2SO4, the relative rate of reaction = 1/50 = 0.02 s –1
There the speed of reaction between magnesium and sulphuric acid is twice as fast as that with
hydrochloric acid.
Measuring the volume of gaseous product over a fixed time interval


Calcium carbonate is reacted with dilute hydrochloric acid as shown by the equation
CaCO3(s) + 2HCl (aq) →CaCl2(aq) + H2O(l) + CO2(g)
Every half a minute the volume of carbon dioxide produced is measured and
recorded using the apparatus below.
 The volume is tabulated as follows
Time(s)
0.5 1 1.5 2 2.5 3 3.5 4
Volume of CO2(cm3) 20 35 40 40 40 40 40 40

The graph is then drawn to determine the rate of reaction
For the N.S. Dept © 2015
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
The graph is steepest at the start of the experiment showing that the reaction is
fastest at the start.
 When the graph is flat, the reaction is complete.
 Some important questions
(a) Calculate the average rate of reaction during the first 0.5 seconds.
 A tangent must be drawn on the graph where t= 0.5 seconds and its gradient is equal
to the rate of reaction
(b) Use your graph to determine the volume of carbon dioxide obtained after 0.8 seconds
Volume of carbon dioxide at 0.8 seconds = 30 cm3
(c) How long did it take for the reaction to end?
Time taken for the reaction to come to an end= 1.5 seconds
Measuring the amount of reactant left over a fixed time interval.




Calcium carbonate is first added to sulphuric acid in a beaker.
During this reaction carbon dioxide is produced and lost to the atmosphere
The decrease in the total mass of the reactants can be monitored over a fixed time
The beaker and its contents are put on the electronic balance.
For the N.S. Dept © 2015
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
The loss of mass is measured and recorded as follows.
Time(s)
Loss in
mass
0
60
120
180
240
50
40
30
22
17
300 360
14

The graph is plotted as follows

Some important questions
(a) What was the mass lost after 30 seconds?
Mass lost = 50g- 44g = 6g
(b) After what time did the reaction lose 25g
Mass remaining = 50g – 25g = 25g
For the N.S. Dept © 2015
12
420 480
11
11
Page 93
From the graph time = 150 sec.
Comparing the rates of reaction

Two experiments for the reaction between Calcium carbonate, CaCO3 and hydrochloric acid,
HCl are carried out using the same concentration of HCl but different sizes of calcium
carbonate.

The volume of carbon dioxide produced in each case are recorded as follows
Time(s)
0
0.5 1 1.5 2 2.5 3 3.5 4
0
20 35 40 40 40 40 40 40
Experiment 1.
Volume of
CO2(cm3)
0
10 20 30 40 40 40 40 40
Experiment 2
Volume of
CO2(cm3)

Some important questions
(a) What is the total volume of carbon dioxide collected in each experiment?
Experiment 1, volume = 40 cm3
Experiment 2, volume = 40 cm3
(b) How long did it take for each reaction to reach completion?
Experiment 1, time for the reaction to stop = 1.5seconds
Experiment 2, time for the reaction to stop = 2 seconds
(c) What conclusion can be drawn from the graphs drawn?
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The graph levels off much earlier and it is steeper in experiment 1 with small pieces of
calcium carbonate. Therefore, experiment 1 with smaller sizes of calcium carbonate is
faster than experiment 2.

When chemical reactions occur, energy is either taken in or given out to the surroundings in form
of heat or light energy. Reactions are described as either exothermic or endothermic, depending
on whether the energy is absorbed or given out to the surroundings.
Exothermic reactions




Exothermic reaction: is the chemical reaction which involves the loss of heat energy to the
surrounding resulting in an increase in the surrounding temperature.
In an exothermic reaction the total energy of the products is lower than that of the reactants due
to the loss of energy to the surrounding.
The change in heat energy ∆𝐻 is negative (∆𝑯 = – ve ) in an exothermic reaction. The negative
sign indicates energy loss
When two atoms are joined together to form a chemical bond, heat energy is given out. Bond
making is therefore an exothermic reaction.
Examples of exothermic reactions
1. Combustion of methane gas is highly exothermic
CH4(g) + O2(g) →CO2(g) + H2O(g) ∆𝐻 = –882 KJ
2. The production of ammonia in the Haber process is also exothermic
N2(g) + H2(g) → 2NH3(g)
∆𝐻 = –184KJ
3. Neutralization reactions between acids and alkalis are exothermic.
4. Addition of sulphuric acid to water is also exothermic
5. All freezing and condensing reactions are exothermic
Endothermic reactions
Endothermic reaction: is the chemical reaction in which heat energy is absorbed from the surrounding
resulting in a temperature drop in the surrounding.



The temperature of the surrounding decreases and the container becomes colder.
In an endothermic reaction the total energy of the products is greater than that of the reactants
The change in heat energy ∆𝐻 is positive (∆𝑯 = + ve ) in an endothermic reaction. The positive
sign indicates gain in energy.
 When bonds are being broken, heat is absorbed. Bond breaking is therefore an endothermic
reaction.
Examples of endothermic reactions
1. The decomposition of limestone
CaCO3(s) →CaO(s) + CO2(g)
∆𝐻 = +222 KJ
2. Photosynthesis
6CO2(g) + 6H2O(g) → C6H12O6(g) + 6O2(g)
∆𝐻 = + 2816KJ
3. Decomposition of silver halide crystals by light
For the N.S. Dept © 2015
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AgCl(s) → Cl2 + Ag(s) ∆𝐻 = + 254KJ
4. Dissolution of ammonium nitrate in water.
Energy diagrams
Exothermic reactions
1. Combustion of methane gas is highly exothermic
CH4(g) + O2(g) →CO2(g) + H2O(g) ∆𝑯 = –882 KJ
2. The production of ammonia in the Haber process is also exothermic
N2(g) + 3H2(g) → 2NH3(g)
∆𝑯 = –184KJ
Endothermic reactions
For the N.S. Dept © 2015
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1. The decomposition of limestone
CaCO3(s) →CaO(s) + CO2(g)
∆𝑯 = +222 KJ
2. Photosynthesis
6CO2(g) + 6H2O(g) → C6H12O6(g) + 6O2(g)
∆𝑯 = + 2816KJ
Bond energies

The bond energy is the energy absorbed in breaking a covalent bond.
Covalent bond Bond energy,KJ
H–H
436
Cl – Cl
242
C–C
348
C–H
412
O–H
463
Cl – H
431
N–H
388
O=0
496
C=O
743
For the N.S. Dept © 2015
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945
C=C
838
EXAMPLES.
1. The decomposition of ammonia is represented by the equation below.
2NH3(g) → N2(g) + 3H2(g)
(a) Calculate the heat of reaction for this decomposition reaction.
∆𝑯 = 6(N – H) – [
+3H–H]
= 6 X 388 - [ 945 + 3 X 436]
= 2328 - 2253
= + 75 KJ
(b) What type of reaction is this decomposition
It is an endothermic reaction
(c) What will be the heat of reaction if 4 moles of ammonia decomposed
(d) Draw the energy diagram for the decomposition of ammonia.
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Electrolysis: is the process by which an electrolyte is decomposed by the passage of electricity through it.
An electrolyte: is an ionic compound which conducts an electric current when molten or dissolved in
water and is decomposed by electric current.
Electrolytes are usually molten or aqueous solutions of ionic compounds(salts) and aqueous solutions
of acids or alkalis
NB: Ionic compounds (salts) conduct electricity by the movement of ions. All ionic compounds do not
conduct electricity in solid state because their ions are held in fixed positions by strong electrostatic
forces. In order for the ions to move, the solid must be in molten state or dissolved in water thereby
making the ionic compound conduct electricity.
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TYPES OF ELECTROLYTE
NON- ELECTROLYTES
WEAK ELECTROLYTES
STRONG ELCTROLYTES
All organic liquids e.g
(a) Ethanol
(b) Tetrachloromethane
(c) Pure water
(d) Sugar
(e) Molten sulphur
(f) All covalent compounds
Weak acids e.g
(a) Ethanoic acid
(b) Carbonic acid
(c) Sulphurous acid
Weak alkalis e.g
(a) Limewater (Ca(OH)2)
Strong acids e.g
(a) Sulphuric acid
(b) Hydrochloric acid
Strong alkalis e.g
(a) Sodium hydroxide
Salts
(a) Sodium chloride
(b) Copper(II) sulphate
(c) Lead(II) Bromide
ELECTRODES
Electrodes: are metallic conductors or terminals through which electrons enter or leave the electrolyte.
There are two types of electrodes and these are (a) anode and (b) cathode
ANODE:
An anode: is an electrode connected to the positive terminal of the battery





Anions (negative ions) are attracted to the anode.
Oxidation is the loss of electrons
Oxidation occurs at the anode where anions discharge by losing electrons
The anion which loses electrons becomes a neutral atom.
The lost electrons at the anode are pumped to the cathode by the battery leaving the anode
positively charged.
CATHODE:
Cathode: is an electrode connected to the negative terminal of the battery.



Cations (positive ions) are attracted to the cathode.
Reduction is the gain of electrons
Reduction occurs at the cathode where cations discharge by gaining electrons.
ELECTROLYSIS OF MOLTEN ELCTROLYTES


Electrolysis of molten electrolytes involves simple binary ionic electrolytes containing only
two ions i.e – metallic ion and non-metallic ion.
The most common molten electrolytes are:
(i)
Lead(II) Bromide (PbBr2)
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

(ii)
Sodium chloride (NaCl)
The electrodes used during the electrolysis of the molten electrolyte are inert electrodes. E.g
carbon and platinum
During this type of electrolysis the electrolyte is usually in solid state, it is only made molten
by direct heating.
ELECTROLYSIS OF MOLTEN LEAD(II) BROMIDE – PbBr2
Electrolyte: molten lead(II) bromide(PbBr2)
Electrode: carbon anode and carbon cathode
Ions present: lead(II) ions Pb2+ , Bromide ions Br –
Reaction at the anode:


The bromide ions(Br –) drift or migrate to the anode and get discharged by losing electrons
to form bromine atoms thereby getting oxidized.
The bromine atoms combine to form bromine gas. The brownish bromine gas is seen at the
anode.
2Br – (l) → Br2 (g) + 2e –
Reaction at the cathode:


The lead (II) ions (Pb2+) migrate to the cathode and get discharged by gaining electrons to
form lead atoms thereby getting reduced.
A silvery deposit of lead metal is seen at the cathode.
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Pb 2+ (l) + 2e – → Pb(l)
Overall reaction:
ELECTROLYSIS OF MOLTEN SODIUM CHLORIDE – NaCl
For the N.S. Dept © 2015
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Electrolyte: molten sodium Chloride (NaCl)
Electrode: carbon anode and carbon cathode
Ions present: sodium ions Na+ , chloride ions Cl –
Reaction at the anode:


The chloride ions(Cl –) drift or migrate to the anode and get discharged by losing electrons
to form chlorine atoms thereby getting oxidized.
The chlorine atoms combine to form chlorine gas. The yellowish- green chlorine gas is seen
at the anode.
2Cl – (l) → Cl2 (g) + 2e –
Reaction at the cathode:


The sodium ions (Na+) migrate to the cathode and get discharged by gaining electrons to
form sodium atoms thereby getting reduced.
A shiny, silvery deposit of sodium metal is seen at the cathode.
Na+ (l) + e – → Na(l)
Overall reaction:
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ELECTROLYSIS OF MOLTEN MAGNESIUM OXIDE – MgO
Electrolyte: molten magnesium oxide (MgO)
Electrode: carbon anode and carbon cathode
Ions present: magnesium ions Mg2+, oxide ions O 2–
Reaction at the anode:


The oxide ions (O 2– ) drift or migrate to the anode and get discharged by losing electrons to
form oxygen atoms thereby getting oxidized.
The oxygen atoms combine to form oxygen gas. The bubbles of the colourless and odourless
For the N.S. Dept © 2015
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oxygen gas are liberated at the anode.
O 2– (l) → O2 (g) + 2e –
Reaction at the cathode:


The magnesium ions (Mg2+) migrate to the cathode and get discharged by gaining electrons
to form magnesium atoms thereby getting reduced.
A shiny, silvery deposit of magnesium metal is seen at the cathode.
Mg 2+ (l) + 2e – → Mg(l)
Overall reaction:
ELECTROLYSIS OF AQUEOUS ELECTROLYTES
The factors that affect the electrolysis of aqueous electrolytes
1. Concentration: the level of concentration of the electrolyte will determine which ions will
discharge.
2. Types of electrodes: the electrodes used can either be inert or active electrodes depending on
required outcome of electrolysis.
NB: Every aqueous solution must contain water which ionizes into Hydrogen ion and hydroxide ion.
Therefore, an aqueous electrolyte must contain atleast 4 or 3 ions
For example1: aqueous sodium chloride produces: H+ , OH – , Na+ and Cl – in solution.
For example2: aqueous hydrochloric acid produces: H+ , OH – and Cl – in solution.
The question of which ion out of the four will discharge is based on the selective discharge rule or
potential discharge rule.
THE SELECTIVE DISCHARGE RULE
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The selective discharge rule states that the less reactive metal (or less electropositive metal) will be
easier to discharge than the more reactive metal ( or more electropositive metal).
i.e POSOCAMAZINILHCOMES
ACRONYM ION NAME
REACTIVITY
+
PO
K
Potassium
Most reactive
SO
Na+
Sodium
CA
Ca+2 Calcium
M
Mg+2 Magnesium
A
Al3+ Aluminium
Z
Zn2+ Zinc
I
Fe2+ Iron
NI
Ni2+ Nickel
L
Pb2+ Lead
H
H+
Hydrogen
2+
CO
Cu
Copper
ME
Hg2+ Mercury
S
Ag+ Silver
GO
Au+ Gold
PLA
Pt2+
Platinum
Least reactive
Similarly, the less reactive (or less electronegative) non- metals are easier to discharge than the most
electronegative non-metal.
The actual selective discharge rule:
CATHODE REACTIVITY
K+
Difficulty to discharge
+
Na
Ca+2
Mg+2
Al3+
Zn2+
Fe2+
Ni2+
Pb2+
H+
Cu2+
Hg2+
Ag+
Au+
Pt2+
Easier to discharge
ANODE
SO4 2+
NO3 –
Cl –
Br –
I–
OH –
REACTIVITY
Difficulty to discharge
Easier to discharge
IONS THAT WILL NEVER BE DISCHARGED
1. Ions of very reactive metals such as potassium, sodium, calcium, magnesium and aluminium
can never be discharged whether in dilute or concentrated solutions.
2. Sulphate and nitrate ions will never discharge
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CONCENTRATED ELECTROLYTES OR SOLUTIONS
ANODE:
 All the negative ions in aqueous solution will drift to the positively charged anode.
 In concentrated aqueous solutions the halide ions (Cl – , Br – or I –) are preferentially
discharged by virtue of being highly concentrated compared to the Hydroxide ions(OH –)
CATHODE :


All the positively charged ions migrate to the cathode
In a concentrated aqueous solution containing H+ and either Ni2+ or Pb2+ ions, the Ni2+
or Pb2+ will preferentially be discharged by virtue of their concentration instead of the
H+ ions.
DILUTE AQUEOUS ELCTROLYTES
 For all dilute aqueous electrolytes the discharge rules apply at both the anode and
cathode.
THE ELECTROLYSIS OF CONCENTRATED AQUEOUS SODIUM CHLORIDE - NaCl
Electrolyte: concentrated sodium Chloride (NaCl)
Electrode: carbon anode and carbon cathode
Ions present: Na+ , Cl – , H+, OH –
Reaction at the anode:
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

Both hydroxide ions ( OH – ) and chloride ions(Cl –) will migrate to the anode but only the
Cl – ions will preferentially be discharged or oxidized because they are in a higher
concentration.
A yellowish green chlorine gas Cl2 is observed at the anode.
2Cl – (aq) → Cl2 (g) + 2e –
Reaction at the cathode:


The sodium ions (Na+) and the hydrogen ions (H+) will both migrate to the cathode but only
H+ ions will preferentially be discharged or reduced.
Bubbles of hydrogen gas are liberated at the cathode.
2H+ (aq) + 2e – → H2(g)
Overall reaction:
THE ELECTROLYSIS OF DILUTE AQUEOUS SODIUM HYDROXIDE – NaOH
Electrolyte: Dilute sodium hydroxide (NaOH)
Electrode: carbon anode and carbon cathode
Ions present: Na+ , H+, OH –
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Reaction at the anode:


Hydroxide ions ( OH – ) will migrate to the anode and get discharged or oxidized by
losing electrons to form water and oxygen gas.
Bubbles of the colourless and odourless oxygen gas are seen at the anode .
4OH – (aq) → 2H2O(l) + O2(g) + 4e –
Reaction at the cathode:


The sodium ions (Na+) and the hydrogen ions (H+) will both migrate to the cathode but only
H+ ions will preferentially be discharged or reduced because H+ ion is lower in the
electrochemical series.
Bubbles of hydrogen gas are liberated at the cathode.
2H+ (aq) + 2e – → H2(g)
Overall reaction:
THE ELECTROLYSIS OF DILUTE SULPHURIC ACID OR THE ELECTROLYSIS OF
WATER – H2SO4
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Electrolyte: Dilute sulphuric acid (H2SO4)
Electrode: carbon anode and carbon cathode
Ions present:
H+, OH –, SO42 –
Reaction at the anode:


Both the Hydroxide ions ( OH – ) and sulphate ions (SO42 –) will migrate to the anode but
only the Hydroxide ions ( OH – ) will preferentially be discharged or oxidized by losing
electrons to form water and oxygen gas.
Bubbles of the colourless and odourless oxygen gas are seen at the anode .
4OH – (aq) → 2H2O(l) + O2(g) + 4e –
Reaction at the cathode:


The hydrogen ions (H+) will migrate to the cathode and get discharged or reduced forming
hydrogen atoms which letter combine to form hydrogen gas.
Bubbles of hydrogen gas are liberated at the cathode.
2H+ (aq) + 2e – → H2(g)
Overall reaction:
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NB: Water is a non- electrolyte but with the addition of sulphuric acid, water can be electrolyzed to
hydrogen and oxygen in the ratio 2: 1.
THE ELECTROLYSIS OF CONCENTRATED HYDROCHLORIC ACID. - HCl
Electrolyte: concentrated aqueous hydrochloric acid (HCl)
Electrode: carbon anode and carbon cathode
Ions present: H+, Cl –, OH –
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Reaction at the anode:


Both hydroxide ions (OH – ) and chloride ions(Cl –) will migrate to the anode but only the
Cl – ions will preferentially be discharged or oxidized because they are in a higher
concentration.
A yellowish green chlorine gas Cl2 is observed at the anode.
2Cl – (aq) → Cl2 (g) + 2e –
Reaction at the cathode:


The hydrogen ions (H+) will drift to the cathode and get discharged or reduced forming
hydrogen atoms which letter combine to form hydrogen gas.
Bubbles of hydrogen gas are liberated at the cathode.
2H+ (aq) + 2e – → H2(g)
Overall reaction:
ELECTROLYSIS BY THE TYPE OF ELCTRODES
There are two types of electrodes and these are
1. Inert electrodes
2. Active electrodes
Inert electrodes: are electrodes which do not react with the electrolyte or the products of
electrolysis. E.g carbon (graphite) and platinum
Active electrode: are electrodes which react with the electrolyte or the products of electrolysis.
E.g metals such as : copper, silver, aluminium e.t.c.
THE ELECTROLYSIS OF AQUEOUS COPPER (II) SULPHATE – CuSO4 USING INERT
ELECTRODES.
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Electrolyte: Dilute aqueous copper (II) sulphate (CuSO4)
Electrode: carbon anode and carbon cathode
Ions present: Cu2+, SO42- , H +, OH –
Reaction at the anode:


Both the Hydroxide ions ( OH – ) and sulphate ions (SO42 –) will migrate to the anode but
only the Hydroxide ions ( OH – ) will preferentially be discharged or oxidized by losing
electrons to form water and oxygen gas.
Bubbles of the colourless and odourless oxygen gas are seen at the anode .
4OH – (aq) → 2H2O(l) + O2(g) + 4e –
Reaction at the cathode:


Both the hydrogen ions (H+) and copper(II) ions (Cu2+) will drift to the cathode but only the
copper(II) ions (Cu2+) will preferentially be discharged by gaining electrons to form copper
atoms.
The pink copper metal will be deposited at the cathode.
Cu+2 (aq) + 2e – → Cu(s)
Overall reaction:
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THE ELECTROLYSIS OF AQUEOUS COPPER (II) SULPHATE – CuSO4 USING ACTIVE
(COPPER ) ELECTRODES
Electrolyte: Dilute aqueous copper (II) sulphate (CuSO4)
Electrode: Anode: Impure copper anode
cathode: pure copper cathode
Ions present: Cu2+, SO42- , H +, OH –
Reaction at the anode:

Both the Hydroxide ions ( OH – ) and sulphate ions (SO42 –) will migrate to the anode but
neither of the two ions will discharge. Instead the copper atoms in the anode dissolve in
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
solution by losing electrons to form copper(II) ions (Cu2+ ) . This is because it is easier for
copper to lose electrons than it is for OH – ions.
The anode will reduce in mass and size as the copper atoms dissolve away.
Cu(s)) → Cu 2+(aq) + 2e –
Reaction at the cathode:



Both the hydrogen ions (H+) and copper(II) ions (Cu2+) will drift to the cathode but only the
copper(II) ions (Cu2+) will preferentially be discharged by gaining electrons to form copper
atoms.
The pink deposit of copper metal will be seen at the cathode.
The mass of the cathode will significantly be increased due to the copper deposits.
Cu+2 (aq) + 2e – → Cu(s)
Overall reaction:
There is no overall reaction because the concentration of the electrolyte remains contant. The copper(II)
sulphate solution remains blue because the copper ions discharging at the cathode are replenished by the
copper dissolving from the anode.
ELECTROLYTIC CELL Vs
ELECTROCHEMICAL CELL.
1. Electrolytic cell
 An electrolytic cell is the type of cell which converts electrical energy to chemical
energy.
 It uses a battery as a source of power or electricity to effect the chemical reactions in the
cell.
 An electrolytic cell mainly involves inert electrodes or some electrodes of the same
kind.( e.g. copper anode and copper cathode)
 The two electrodes are always put in the same container.
2. Voltaic cell or electrochemical cell or simple cell or galvanic cell
 Electrochemical cell or voltaic cell is the type of cell which converts chemical energy
to electrical energy.
 The electrodes in a voltaic cell are metals of two different metals with different
reactivity.
 In a voltaic cell, a resistor or a voltmeter is connected in place of a battery.
 The electrodes are usually dipped in an aqueous solution of an ionic salt or an acid.
 The two electrodes may both be dipped in one cell with a common electrolyte or they
may be put in separate containers such that the electrolyte in each cell must be a salt of
the metal being used as an electrode.
 If two separate containers are used they are normally connected by a salt bridge which
completes the circuit by allowing the flow of ions from one to another.
 The more reactive metal (more electropositive metal) has a higher tendency of losing
electrons. Therefore, in a voltaic cell a more reactive metal supplies the electrons
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




through the external circuit to the less reactive metal.
The more reactive metal with a higher tendency of losing electrons is called a negative
terminal while the less reactive metal is called the positive terminal.
Oxidation therefore occurs at the negative terminal while reduction occurs at the
positive terminal.
The anode is the negative terminal which a more reactive metal
The cathode is the positive terminal which is the less reactive metal.
The wider the difference in reactivity between the two metals used as electrodes, the
larger the voltage supplied into the external circuit.
NB: The selective discharge does not apply at the anode but only applies at the cathode.
EXAMPLE OF VOLTAIC CELL
Electrolyte: Cell One : ZnSO4
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Cell Two : CuSO4
Electrode:
Anode : Zinc
cathode: Copper.
Ions present: Cell One : Zn 2+ , SO42+ , H+, OH –
Cell Two : Cu2+ , SO42+ , H+, OH –
Reaction at the anode:

The atoms in zinc will dissolve in solution by losing electrons through the external circuit to
the copper electrode thereby getting oxidized to Zn2+ ions.
Zn(s) → Zn2+ (aq) + 2e –
Reaction at the cathode:

Both the copper ions (Cu2+ ) and hydrogen ions (H+) will drift to the cathode but only Cu2+
ions will be discharged by gaining electrons thereby getting reduced to pink copper metal
which is deposited at the cathode.
Cu+2 (aq) + 2e – → Cu(s)
Overall reaction:
EXAMINATION QUESTIONS
1. A silver rod and a copper rod were connected together dipped in dilute sulphuric acid as shown in
the diagram. An electric bulb connected between the metals glowed.
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(a) At which rod does oxidation occur?
ANS; iron rod
(b) What term is used to describe the device above?
ANS: electrochemical cell or voltaic cell or simple cell
(c) Which of the rods is the ..
(i)
Anode?
Ans: iron rod
(ii)
Cathode?
Ans: silver rod
(iii)
The positive pole?
Ans: slver rod
(d) State atleast three differences between the above cell and the electrolytic cell.
Ans: (i) the above cell converts chemical energy into electrical energy while the electrolytic
cell converts electrical energy to chemical energy.
(iii)
the above cell uses two metals of different reactivivity as electrodes while an
electrolytic cell mainly uses inert electrodes.
(iv)
In the above cell oxidation occurs at the negative terminal while in an ectrolytic
cell oxidation occurs at the positive electrode.
(e) What energy changes occur in the device?
ANS: chemical energy → electrical energy.
(f) How would the replacement of the iron rod with zinc alter the brightness of the bulb?
ANS: the bulb will grow brighter than before because of the increase in voltage produced as
zinc and silver metals are further apart in terms reactivity.
APPLICATION OF ELECTROLYSIS
Electrolysis is widely used in the following processes
1. Extraction of copper
2. Extraction of aluminium
3. Electroplating process
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4. Manufacture of sodium hydroxide
EXTRACTION OF COPPER
Properties of copper
1.
2.
3.
4.
It is generally unreactive
It does not corrode
It is an excellent conductor of electricity and heat.
It is hard and poisonous
COPPER
USES OF COPPER




Copper is widely used for making electrical wires because it is an excellent conductor of
electricity
It is used for making boilers and water pipes mainly because it is resistant to corrosion and
conducts heat,
Copper is also used for making bullets due to its high density and poisonous nature.
Copper is also used for making alloys.
COPPER EXTRACTION
Copper is extracted mainly from the following ores;
5.
6.
7.
8.
Cuprite : copper(I) oxide , Cu2O
Malachite: Copper (II) carbonate – copper(II) hydroxide, CuCO3 . Cu(OH)2
Chalcocite : copper(I) sulphide; Cu2S
Copper pyrites: copper –iron sulphide; CuFeS2.
Copper is an uncreative metal which is extracted by heating with carbon.
COPPER PROCESSING


the copper ore is first crushed before being taken for flotation, roasting and smelting
the copper ore is sometimes leached in sulphuric acid before being taken for electrolysis
CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
ELECTROLYSIS


very pure copper is needed for electrical conductors
electrolysis is therefore used to produce very pure copper
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




Electrolyte: copper (II) sulphate solution
Electrodes: (i) Anode: impure copper
(ii) cathode: pure copper
Ions present: Cu2+ , SO42+ , H+ and OH–
Reaction at the anode:
Copper dissolves in solution from the impure anode by losing electrons
Cu(s)→ Cu2+(aq) + 2e–
The impurities at the anode falls away as sludge and settles at the bottom.
This explains why the anode reduces in both size and weight.
Reaction at the cathode:
The Cu2+ ions migrate to the cathode and get discharged by gaining electrons forming copper
atoms.
The yellowish green copper metal is deposited on the cathode
Cu2+(aq) + 2e– → Cu(s)
This explains why the cathode increases in both size and mass
Overall reaction: there is no overall reaction as the concentration of the reactants remains the
same.
EXTRACTION OF ALUMINIUM
Aluminium is extracted from its CHIEF ORE Bauxite; Al2O3. Aluminium is a fairly reactive metal
which is extracted by electrolysis.
Uses of aluminium
1. It is used for making overhead electrical cables due to its low density and good electrical
conductivity
2. It is used for making kitchen utensils because it is a good conductor of heat and it has low
density
3. It is used for making aircraft bodies because of its density and resistance to corrosion
4. It is used for making bicycle frames because of it resistance to corrosion and being a light weight
metal.
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PROCESS OF EXTRACTING ALUMINIUM











Aluminium is extracted from bauxite (Al2O3) by the process of electrolysis.
The solid bauxite is first heated in order to make a molten electrolyte
This process is however very expensive and uneconomical because the melting point of bauxite
is very high about 2045 °C . This means that huge energy is required to melt bauxite.
In order to lower the working temperature and make the process economical, bauxite (Al2O3) is
dissolved in cryolite (Na3AlF6) which reduces the melting point of bauxite from 2045 °C to
900°C.
The mixture of bauxite and cryolite becomes the electrolyte.
Electrolyte: molten mixture of cryolite(Na3AlF6) and bauxite (Al2O3)
Electrodes : (i) Anode : Graphite
(ii) Cathode: Graphite
Ions present : Al3+ and O2 –
Reaction at the anode:
The O2 – ions migrate to the anode and get discharged by losing electrons to form oxygen gas
2O2 –(l) → O2(g) + 4e–
At a working temperature of 900 °C, oxygen gas produced at the anode reacts with carbon in the
graphite anode to produce carbon dioxide. This explains why the graphite anode is continuously
replaced at the anode as it is continuously depleted by this oxidation with oxygen.
Reaction at the cathode:
The Al3+ ions migrate to the cathode and get discharged by gaining electrons to form aluminium
metal. The aluminium metal is deposited at the cathode and settles at the bottom.
Al3+(l) + 3e–→ Al (l)
The molten aluminium metal is tapped off at intervals through an outlet at the bottom.
Overall reaction:
Al2O3(l)→Al(l) + O2(g)
ELECTROPLATING
Electroplating : is the coating of metallic substance or plastic material with a non-corrosive metallic
substance.
The two main advantages of electroplating are:
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1. It improves the resistance to corrosion
2. It also improves the appearance of a substance.
 The most common metals used for electroplating are gold and silver because of their brilliant
lustre and resistance to corrosion. The metallic substance to be plated must be connected to the
cathode.
 The anode is usually a metal which is used for plating .
 The electrolyte should be a salt of a metal which is plated (anode)
ELECTROPLATING AN ALUMINIUM SPOON USING SILVER METAL.
Electrolyte: Aqueous Ag2SO4
Electrode:
Anode : Silver
cathode: Aluminium spoon.
Ions present: Ag + , SO42+ , H+, OH –
Reaction at the anode:

Both the sulphate (SO42+) ions and hydroxide ( OH – ) ions will migrate to the anode but
neither of them will discharge because silver loses electrons more readily than the two ions.
Therefore, atoms in silver will dissolve in solution by losing electrons through the external
circuit thereby getting oxidized to Ag + ions.
Ag(s) → Ag+ (aq) + e –
Reaction at the cathode:
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

Both the silver ions (Ag + ) and hydrogen ions (H+) will drift to the cathode but only Ag +
ions will preferentially be discharged by gaining electrons thereby getting reduced to silver
metal
Ag+ (aq) + e – → Ag(s)
The silver metal is deposited and coated on the aluminium spoon.
Overall reaction:
There is no overall reaction because the concentration of the electrolyte remains unchanged.
MANUFACTURE OF SODIUM HYDROXIDE BY THE ELECTROLYSIS OF BRINE (SODIUM
CHLORIDE)
The source of brine(NaCl)
From the electrolysis of concentrated sea water.
The use of sodium hydroxide (NaOH)
1. It is used for making soap
2. It is used for making bleaching agents
There are two electrolysis processes used to manufacture sodium hydroxide from brine. These include:
1. The diaphragm cell
2. The mercury cell
THE DIAPHRAGM CELL
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The porous membrane: Allows for the exchange of ions between two electrolytes.
Electrolyte: Concentrated brine or concentrated aqueous sodium chloride (NaCl)
Electrode:
Anode : Titanium
cathode: Nickel
Ions present: Na + , Cl – , H+, OH –
Reaction at the anode:
Both hydroxide ions ( OH – ) and chloride ions(Cl –) will migrate to the anode but only the
Cl – ions will preferentially be discharged or oxidized because they are in a higher
concentration.
 A yellowish green chlorine gas Cl2 is observed at the anode.
2Cl – (aq) → Cl2 (g) + 2e –
 The sodium ion (Na+ ) will diffuse through the porous membrane towards the cathode

Reaction at the cathode:

The sodium ions (Na+) and the hydrogen ions (H+) will both migrate to the cathode but only
H+ ions will preferentially be discharged or reduced.
 Bubbles of hydrogen gas are liberated at the cathode.
2H+ (aq) + 2e – → H2(g)
 The removal of hydrogen ions leaves a high concentration of hydroxide ions (OH – ) around
the cathode
 Therefore, the sodium ions (Na+) are drawn from the anode through the porous membrane
to the cathode.
 The reaction occurs between the two ions forming sodium hydroxide, NaOH
Na+ (aq) + OH – (aq) → NaOH(aq)
Reaction at the cathode:
NaCl(aq) + 2H2O (l) →2NaOH(aq) + H2(g) + Cl2(g)
THE MERCURY CELL
Disadvantage of manufacturing NaOH using a mercury cell.
1. Mercury is very expensive.
2. Mercury is porous
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Electrolyte: Concentrated brine or concentrated aqueous sodium chloride (NaCl)
Electrode:
Anode : Graphite (carbon)
Cathode : Mercury
Ions present: Na +, Cl –, H+, OH –
Reaction at the anode:


Both hydroxide ions ( OH – ) and chloride ions(Cl –) will migrate to the anode but only the
Cl – ions will preferentially be discharged or oxidized because they are in a higher
concentration.
A yellowish green chlorine gas Cl2 is observed at the anode.
2Cl – (aq) → Cl2 (g) + 2e –
Reaction at the cathode:


The sodium ions (Na+) migrate to the cathode and get reduced by gaining one electron.
In the process the sodium mercury amalgam is produced.
Na+ (aq) + e – + Hg(l) → Na/Hg(l) … sodium mercury amalgam
 The sodium mercury amalgam reacts with cold water to produce sodium hydroxide and
hydrogen gas.
 Na/Hg(l) + 2H2O(l) → 2NaOH(aq) + H2(g) + 2Hg(l)
Overall reaction :
2 NaCl(aq) + 2H2O (l) →2NaOH(aq) + H2(g) + Cl2(g)
FARADAY’S LAW
FACTORS THAT DETERMINE THE MASS OF PRODUCTS DURING ELECTROLYSIS
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1. The magnitude of electric current flowing through the cell
2. The time of flow of electric current.
3. The charge on the element being deposited.
Faraday’s law: States that the mass of the product deposited or liberated during electrolysis is
directly proportional to the electric current passing through the cell.
Charge(Q) = Current(A) X time(s)
Where Q= Charge in coulombs
I = current(A)
t = time(s)
It was experimentally discovered that 96500C of charge needed to liberate one mole of the singly charged
ion during electrolysis.
Therefore, 96500C is called the FARADAY‟s CONSTANT
Q= nF
Where Q= Charge in coulombs
n = number of moles
F =faraday‟s constant (96500C)
Summary of the needed expressions:
Q = it
,
Q = nF
ATTEMPTS
1. Given that 2A of electric current flows through a cell in 20 minutes. What is the volume of
oxygen gas at r.t.p produced at the anode during he electrolysis of dilute aqueous sodium
chloride.
2. Calculate the mass of copper deposited when a current of 0.9A passed through a solution of
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copper (II) sulphate for 1hour 20 min [Mr = 64]
3. When moltn lead (II) bromide is electrolyzed by a current of 0.8A for 100A. Calculate the mass
of lead metal deposited at the cathode.
4. Calculate the quantity of charge required to deposit 10g of silver metal during electrolysis.
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NB: End of electrolysis.
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Organic chemistry: is the branch of science which studies carbon and its compounds except carbonates
and oxides.


The food we eat, plastics, rubber, drugs and detergents are all organic comounds.
Carbon has a tendency of forming covalent bonds with other carbon atoms. Apart from bonding
with its own kind, carbon also forms covalent bonds with other elements such as hydrogen,
nitrogen, oxygen and sulphur. Organic chemistry, however, is centered on compounds of carbon
and hydrogen.
Hydrocarbons
Hydrocarbons : are organic compounds containing carbon and hydrogen only.
These include:
1. Alkanes
2. Alkenes
3. Alkynes
Homologous series
A homologous series: is a group of organic compounds having the same general formula,
functional group and chemical properties.
Examples of Homologous series
1.
2.
3.
4.
5.
6.
Alkanes
Alkenes
Alkynes
Alcohols
Carboxylic acids
Esters
Characteristics of a homologous series
1.
2.
3.
4.
5.
All the members conform to the same general formula.
Each member of the homologous series differs from the next by - CH2
Members of the homologous series have the same chemical properties
They have the same functional group.
Their physical properties such as melting and boiling points increase with increase in
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the number of carbon atoms.
Functional Groups


A functional group: Is the special group of atoms available in homologous series compounds
which is responsible for the chemical properties of the compound.
All compounds in homologous series have functional groups except alkanes.
s/n
1
2
3
4
5
6
Homologous series
Alkanes
alkenes
Alkynes
Alcohols
Carboxylic acid
Esters
Functional group name
Carbon - carbon single bond
Carbon - carbon double bond
Carbon - carbon triple bond
Hydroxyl group
Carboxylic group
Carbonyl group
Functional group formula
C–C
C=C
C=C
-OH
-COOH
-COO
NOMENCLATURE
Nomenclature: refers to system of naming organic compounds.
s/n
1
2
3
4
5
6
7
8
9
10
Number of carbon atoms
One
Two
Three
Four
Five
Six
Seven
Eight
Nine
Ten
Name
MethEth PropBut PentHexHept OctNonDec-
ALKANES.





Alkanes are saturated hydrocarbons with a general formula CnH2n+2.
n represents the number of carbon atoms.
Saturated hydrocarbons: are hydrocarbons with single bonds between carbon atoms
alkanes have no functional group though sometimes the C – C is considered as a
functional group for alkanes
The
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