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 SS 3
FIRST TERM NOTES ON
CHEMISTRY
 TABLE OF CONTENTS

FIRST TERM
WEEK 1: ALKANOIC ACIDS

WEEK 2: ALKANOATES

WEEK 3: FATS AND OILS

WEEK 4: AMINO ACIDS

WEEK 5: NATURAL AND SYNTHETIC POLYMERS

WEEK 6: CARBOHYDRATES

WEEK 7: METALS AND THEIR COMPOUNDS

WEEK 8: CALCIUM AND ALUMINIUM

WEEK 9: TIN AND COPPER

WEEK 10: IRON

 WEEK: 1
Chemistry SS 3 First Term
Topic: Alkanoic Acids

Alkanoic acids are also known as carboxylic acids. A carboxylic acid can he identified
from the carboxyl functional group and the ‘-oic’ name ending.
General formula of carboxylic (alkanoic acids) : C nH2n+1COOH or R-COOH
Examples of Carboxylic (Alkanoic) Acids
longest C-C functional
Formula Name occurrence
carbon chain single bonds group
-COOH methanoic acid
HCOOH C1 : meth -an- ants
(oic acid) (formic acid)
-COOH ethanoic acid
CH3COOH C2 : eth -an- vinegar
(oic acid) (acetic acid)
-COOH propanoic acid dairy
C2H5COOH C3 : prop -an-
(oic acid) (propionic acid) products
-COOH butanoic acid rancid
C3H7COOH C4 : but -an-
(oic acid) (butyric acid) butter
-COOH pentanoic acid valerian
C4H9COOH C5 : pent -an-
(oic acid) (valeric acid) root
 Structural Examples

of
carboxylic acids

Nomenclature (Naming Alkanoic Acids)

Alkanoic acids are named as follows:
·The number of carbon atoms in the longest carbon atom chain is noted
·The corresponding alkane is then named
·Finally, the name of this corresponding alkane is modified by removing the ane and
replacing it with alkanoic acid
For example:
thylpentanoic acid ,
2-ethylpentanoic acid ,
benzoic acid

Preparation of ethanoic acid

The Complete Oxidation of Ethanol to Ethanoic Acid by acidified sodium
heptaoxochromate(VI) solution. Ethanol undergoes oxidation first to ethanal and
then to ethanoic acid.
The reaction of sodium dichromate(VI) solution with ethanol gives a carboxylic
acid, ethanoic acid, a dilute solution of which is sold as vinegar.
Ethanoic acid can also be prepared by
 distilling anhydrous sodium ethanoate, CH 3COONa with concentrated H2SO4
 or boiling methyl cyanide CH3CN with an acid
CH3COONa(aq)+ H2SO4(aq) → CH3COOH(g) + NaHSO4 (aq)
CH3CN(aq) + HCl(aq) → CH3COOH(g) + NH4Cl(aq)

Physical Properties
1. Ethanoic acid is usually a colourless liquid with a characteristics sharp and pungent
smell
2. It is soluble in water and a dilute solution has the usual sour taste of acid
3. It has a boiling point of 1180C
4. Pure anhydrous ethanoic acid freezes into ice-like crystals at temperature below
170C
5. It turns blue litmus paper red
6. Dilute solution has a sour taste

Chemical Properties of Carboxylic (Alkanoic) Acids

1. Carboxylic (alkanoic) acids are weak acids, the acid dissociation constant, K a, is
small.
2. Soluble carboxylic (alkanoic) acids dissociate slightly in water.
Neutralization Reactions
Neutralization: acid + base → salt + water
Carboxylic (alkanoic) acid + base → salt (metal alkanoate) + water
RCOOH + MOH → RCOO–M+ + H2O
e.g. CH3COOH + NaOH → CH3COO–Na+ + H2O
Ethanoic acid + sodium hydroxide → sodium ethanoate + water
Soluble salts of long-chain (fatty) acids are soaps
e.g. C17H35COOH + NaOH → C17H35COO–Na+ + H2O
Stearic acid + sodium hydroxide → sodium stearate + water
Reaction with Carbonates
acid + carbonate → salt + carbon dioxide gas + water
Carboxylic (alkanoic) acid + metal carbonate → metal alkanoate + carbon dioxide +
water
e.g. 2CH3COOH + Na2CO3 → 2CH3COO–Na+ + CO2 + H2O
 Ethanoic acid + sodium carbonate → sodium ethanoate + carbon dioxide + water
e.g. CH3COOH + NaHCO3 → CH3COO–Na+ + CO2 + H2O
Ethanoic acid + sodium bicarbonate → sodium ethanoate + carbon dioxide + water
Reaction with Active Metals
Acid + metal → salt + hydrogen gas
Carboxylic (alkanoic) acid + metal → metal alkanoate + hydrogen
e.g. 2CH3COOH + 2Na(s) → 2CH3COO–Na+ + H2(g)
Ethanoic acid + sodium → sodium ethanoate + hydrogen
Esterification Reactions
Esters are produced in a condensation reaction between a carboxylic (alkanoic)
acid and an alkanol (alcohol).
This is known as an esterification reaction.
carboxylic (alkanoic) acid + alkanol (alcohol) ester + water
e.g. 2CH3COOH + CH3OH CH3COOCH3 + H2O
Ethanoic acid + methanol methyl ethanoate + water

Uses
1. It can be used as a solvent
2. It is used in food industries as vinegar for preserving and flavouring food
3. It is used In the manufacture of cellulose ethanoate which is used for making
synthetic fibres, such as rayon
4. It is used in making compounds like ethyl ethanoate, ethanoic anhydride (used in
aspirin), cellulose ethanoate (used for packing), propanone etc

Assessment

1. Ethanoic acid is usually a colourless liquid with a characteristics sharp and

……. smell
a. pungent
b. annoying
c. foul
d. dirty
2. Esters are produced in a condensation reaction between a carboxylic (alkanoic)
acid and an alkanol (alcohol) known as …….
a. Polymerization
b. Esterification
c. Dehydration
d. Hydrolysis
3. Ethanoic acid reacts with active metal to give off ……….. gas
a. Oxygen
b. Hydrogen
c. Carbon
d. OH
4. Boiling point of Ethanoic acid is …….
a. 1170C
b. 1190C
c. 1180C
d. 1200C
5. It is used in food industries as vinegar for preserving and flavouring food.
True/False

Answers

1. A
2. B
3. B
4. C
5. True
 WEEK 2
Chemistry SS3 First Term
Topic: Alkanoates
Alkanoates or Esters
Ethyl Ethanoates
All esters are similar chemically although they may vary in degree of reactivity. Ethyl
ethanoate is one of the simple esters. Its molecular formula is CH 3COOC2H2.
Properties
Ethyl ethanoate is prepared by the esterification between ethanol and glacial
ethanoic acid at 150oC in the presence of concentrated tetraoxosulphate (vi) acid.
C2H5OH(aq) + CH3COOH(l) = CH3COOC2H5(l) + H2O(l)

Preparation of Alkanoates

1. Acylation. Esters are prepared by the acylation of alcohols or phenols. The acylating
agents can be any of the following:
Carboxylic acid / H 2SO 4 or Acyl chloride or Acid anhydride
(i) Condensation of alcohols with carboxylic acid This reaction involves
esterification of alkanols by alkanoic acid

(ii) Esterification through acid derivatives

2. By Reaction of Acids with Diazomethane. Acids on being treated with ethereal

solution of diazomethane yield methyl esters.
RCOOH + CH2N2 à RCOOCH3 + N2
Acid Diazo Ester
methane

3. By Tischenko Reaction. When aldehydes containing a-hydrogen atoms are treated

with aluminium ethoxide. They undergo condensation to produce esters.
Physical Properties of Alkanoates
1. Physical state. Esters are colourless, volatile and oily liquids with a characteristic
fruity smell. The smell of the most of the flowers and fruits is due to esters present
in them. The characteristic tastes and smells of different esters find applications in
the manufacture of artificial flavouring and perfuming agents. The flavours of some
of the esters are given below:
Ester Flavour Ester Flavour
n-Pentyl ethanoate Banana Amyl but)’rate Apricot
Octyl ethanoate Orange Isobutyl Raspberry
Ethyl butanoate Pineapple Benzyl ethanoate Jasmine

2. Solubility. Esters are sparingly soluble in water but are quite miscible in organic
solvents like alcohols and ethers. In fact, most of the esters are themselves very
good solvents for plastics and nitrocellulose.
3. Boiling points. The boiling points of esters are always less than the corresponding
carboxylic acids because esters do not form hydrogen bonds.

Chemical Properties

1. Hydrolysis.Esters are hydrolysed slowly by water at boiling temperature. The

reaction is catalysed by small amount of acid or base. The basic hydrolysis is also
known as saponification. It is because of the fact that the esters with high molecular
mass acids (C12-C17) give soap on hydrolysis with a base. Soaps are sodium or
potassium salts of carboxylic acids with high molecular mass (C 12-C17).
Ethyl ethanoate can be hydrolyzed by water into its component acid and alkanol
again. The reaction is catalyzed by hydrogen or hydroxide ions i.e dilute acid or
alkaliCH3COOC2H5(l) + H2O(l) → C2H5OH(aq) + CH3COOH(l)
2. Reduction. Esters are reduced to alcohols by the reducing agents like
(sodium/ethanol) or (lithium aluminium hydride).
3. Reaction with Ammonia. Esters on treatment with alcoholic ammonia yield acid
amides. This reaction is known as ammonolysis of esters.
4. Reaction with Phosphorus Pentachloride. Esters are converted into acid chlorides
and alkyl halides by heating with phosphorus pentachloride.
RCOOR’ + PC15 → RCOCl + R’Cl + POC13
 C6H5COOC2H5 + PCl → C6H5 COCl + C2H2OH
Ethyl benzoate Benzoyl chloride

5. Alcoholysis. An ester on refluxing with a large excess of an alcohol in the presence

of a little acid or alkali, undergoes exchange of alcohol residues, i.e., alkoxy parts
as shown below:
This reaction is known as alcoholysis or trans-esterification.

Uses of Esters
They are mainly used as solvents for cellulose nitrate and quick drying substances
like paints, nail varnishes, lacquer and adhesives. The commonly known thinner
water is a mixture of esters. Esters are used in perfumes and cosmetics and artificial
flavouring for foods. Certain volatile esters are used as solvents for lacquers,
paints, and varnishes; for this purpose, large quantities of ethyl acetate and butyl
acetate are commercially produced.
1. Esters that are have fragrant odours are used as a constituent of perfumes,
essential oils, food flavourings, cosmetics, etc
2. Esters are used as an organic solvent
3. Natural esters are found in pheromones
4. Naturally occurring fats and oils are fatty acid esters of glycerol
5. Phospoesters form the backbone of DNA molecules
6. Nitrate esters, such as nitroglycerin, are known for their explosive properties
7. Polyesters are used to make plastics
8. Esters are used to make surfactants E.g. soap, detergents

ASSESSMENT

1. How is Ethyl ethanoate prepared?

2. What are some of the uses of Esters?
3. What are the chemical properties of Esters?
 WEEK: 3
Chemistry SS3 First Term
Topic: Fats and Oils
Fats and oils are naturally-occurring esters of three long carboxylic acids called
fatty acid with a special type of alcohol called glycerol ( propane- 1,2,3-triol).
Glycerol has three carbon atoms, each with an –OH group on. These fats and oils
belong to a group of compounds called lipids.
Fats are solid at room temperature, whereas oils are liquids. Animals and plants
produce oils and fats as an energy store. Fats and oils are the most abundant lipids
in nature. They provide energy for living organisms, insulate body organs, and
transport fat-soluble vitamins through the blood.

Structures of Fats and Oils

Fats and oils are called triglycerides because they are esters composed of three
fatty acid units joined to glycerol, a trihydroxyl alcohol:

Saturated and Unsaturated Fats and oils

Fatty acids can be saturated or unsaturated
The saturated fatty acids have single bond in their hydrocarbons while the
unsaturated ones have double bonds. Esters produced from saturated fatty acids
are usually solids at room temperature and they are called fats while the esters
produced from unsaturated fatty acids are liquids at room temperature and are
called oils.

Sources of Fats and Oils

 Animal sources generally provide fats, for example, dripping from beef, lard from
pork and tallow from lamb. These fats are solid at room temperature and only
become liquid when heated.
 Vegetable sources generally provide oils. Ester oils can be obtained from olives and
seeds such as corn seed, sunflower seed, peanuts and soya beans, liquids at room
temperature.
 Marine sources can provide both fats and oils, Sea mammals providing fats and oils
being obtained from fish.

Properties of Fats and Oils

1. Both fats and oils are insoluble in water and they decompose at temperature
above 300oC.
2. Since they are esters their main reactions are saponification and hydrolysis.

Uses of Fats and Oil

1. Fats and oils are used as essential ingredients of food
2. Sometimes they are used in natural form such as groundnut oil, palm oil for such
industrial products as margarine
3. They are used as raw materials for making soap e.g. palm oil and coconut oil
4. They are used in making paint, candles and varnishers

Hydrolysis of Fats and Oils

Fats and oils contain ester links formed when glycerol has reacted with fatty acids.
These ester links can be broken when they react with water, splitting the molecule
back into an acid and an alcohol.
This process is known as hydrolysis of an ester. It can be described using the
following word equation:
ester + water → acid + alcohol

Hydrogenation of Oils
The hydrogenation of oils is also known as hardening of oil. This is done by passing
hydrogen into unsaturated oil at about 180 oC and 5 atmosphere of pressure and in
the presence of finely divided nickel as catalyst. In the process, the unsaturated
part of the oil is saturated and the oil becomes hardening into fat.
Margarine is made by mixing such hardening oils with vitamins, salt, skimmed milk
and colouring materials.
Formation of Soap (Saponification)
Saponification is a process by which triglycerides are reacted with sodium or
potassium hydroxide to produce glycerol and a fatty acid salt, called ‘soap’. In other
word, Saponification is the alkaline hydrolysis of the fatty acid esters.
Natural soaps are sodium or potassium salts of fatty acids, originally made by
boiling lard or other animal fat together with lye or potash (potassium hydroxide).
Hydrolysis of the fats and oils occurs, yielding glycerol and crude soap.
In the industrial manufacture of soap, tallow (fat from animals such as cattle and
sheep) or vegetable fat is heated with sodium hydroxide. Once the saponification
reaction is complete, sodium chloride is added to precipitate the soap. The water
layer is drawn off the top of the mixture and the glycerol is recovered using vacuum
distillation. Dyes, perfumes and disinfectants are added as required before the
soap is passes into bars.

Detergents
Detergents are substances which act with water to make things clean. The can be
conveniently classified into two main types – soapy and soapless detergents
Soapy Detergents: are detergents made from soap. They are made by heating
vegetable oils like palm oil with a strong alkali like sodium hydroxide. Soap is also
biodegradable, i.e. it can easily be decomposed by bacteria into simple inorganic
substances.
Soapless Detergents: Most of the soapless detergents are made from petroleum
fractions. Usually the hydrocarbon is reacted with sulphuric acid, and the product
neutralized with sodium hydroxide to obtain the soapless detergent. They are non-
biodegradable and so create water pollution problems when their forms clog up
waterways.
Detergent helps to remove dirt in different ways:
1.) They help water to spread out and completely soak a surface. They are good
wetting agents.
2.) They form an emulsion with the dirt. They are emulsifying agents.
3.) As a wetting agent, detergent help to reduce the surface tension of water,
which tends to pull water molecules together and making water itself to be a bad
wetting agent.
 As an Emulsifying Agent
Detergent has two main parts;
1. A long hydrocarbon chain which is soluble in the grease or oil of the dirt. This
part is not soluble in water and described as hydrophobic
2. An ionic or polar part which in not soluble in the dirt but is soluble in water. This
part is described as hydrophilic, COO– Na+ or COO– K+
Mode of Action
When detergent is added to grease, the grease soluble hydrophobic tail dissolve in
the grease and the hydrophilic parts that bring along with a negative charge are
insoluble in the grease and so they remain outside. The surface of the grease
becomes negatively charge, and if stirring is applied, very large of colloidal particles
of grease and detergent are formed. They are held in aqueous solution by the
attraction of the hydrophilic parts for the water molecules, but are prevented from
coming together because similarly charged on surface of the particles cause
repulsion between each other. Thus they form an emulsion of dirty water.

Differences between Soaps and Soapless Detergent

1) Soap from scum with hard water (which contains dissolve calcium or magnesium
ions), while soapless detergent is not affected.
2) Soaps are made from edible animal or vegetable oils while soapless detergents
are made from petroleum, which is cheaper and edible.

Assessment

1. ……… is a process by which triglycerides are reacted with sodium or potassium

hydroxide to produce glycerol and a fatty acid salt, called ‘soap’.
2. Fats are ……. at room temperature.
3. ……… are substances which act with water to make things clean
4. Mention 4 sources of fats and oil
 Answers

1. Saponification
2. Solid
3. Detergents
 WEEK: 4
Chemistry SS3 First Term
Topic: Amino Acids

Amino acids are the basic structural units of proteins. Amino acids are organic
compounds which contain both an amino group and a carboxyl group, that is, any
of a group of organic molecules that consist of a basic amino group (−NH 2), an acidic
carboxyl group (−COOH). Each molecule contains a central carbon (C) atom,
termed the α-carbon, to which both an amino and a carboxyl group are attached.
The remaining two bonds of the α-carbon atom are generally satisfied by
a hydrogen (H) atom and the R group. The formula of a general amino acid is:

The amino acids differ from each other in the particular chemical structure of
their R group.
There are 20 naturally occurring amino acids of biological importance. The human
body can synthesize all of the amino acids necessary to build proteins except for
the ten called the “essential amino acids”. They can be supplied by a combination
of cereal grains (wheat, corn, rice, etc.) and legumes (beans, peanuts, etc.). The 10
amino acids that we can produce are alanine, asparagine, aspartic acid, cysteine,
glutamic acid, glutamine, glycine, proline, serine and tyrosine. Tyrosine is produced
from phenylalanine, so if the diet is deficient in phenylalanine, tyrosine will be
required as well. The essential amino acids (that we cannot produce internally) are
arginine (required for the young, but not for adults), histidine, isoleucine, leucine,
lysine, methionine, phenylalanine, threonine, tryptophan, and valine. These amino
acids are required in the diet. Plants, of course, must be able to make all the amino
acids. Humans, on the other hand, do not have all the enzymes required for the
biosynthesis of all of the amino acids.”
Peptides
Peptides are amides formed by the interaction between amino groups and carboxyl
groups. The bond joining the α – amino group of one amino acid and the α –
carboxyl group of another amino acid is called peptide bond. Two amino acids react
to form a dipeptide, three to form a tripeptide etc. If a numbe of amino acids are
linked by peptide bonnds, a polypeptide is formed. A polypeptide chain has an
amino acid end and a carboxyl end. It consists of a regular repeating part or main
chain and a variable part. or side chain.

Proteins
Proteins are highly complex substance that is present in all living organisms.
Proteins are of great nutritional value and are directly involved in the chemical
processes essential for life. A protein molecule is very large compared with
molecules of sugar or salt and consists of many amino acids joined together to form
long chains, much as beads are arranged on a string. There are about 20 different
amino acids that occur naturally in proteins. Proteins of similar function have
similar amino acid composition and sequence. When proteins are eaten, they are
hydrolysed in the stomach to give a large variety of amino acids, which are then
carried by the blood stream to all parts of the body. Some are used to build cell
proteins in the growing body or repair damaged cells. Proteins are made up of
polypeptide chains. The amino acid sequence in the polypeptide chain of a protein
is specified by genes.
Glycine and alanine can combine together with the elimination of a molecule of
water to produce a dipeptide.

In each case, the linkage shown in blue in the structure of the dipeptide is known
as a peptide link.
If you joined three amino acids together, you would get a tripeptide. If you joined
lots and lots together (as in a protein chain), you get a polypeptide.
Occurrence
 Proteins are found in all living systems as structural components and as
biologically important substances such as hormones, enzymes and pigments.
Proteins in our food can be divided into first class and second class proteins. First
class proteins contain essential amino acids and they are mainly of animal origin.
Examples are lean meat, fish, eggs, milk and cheese. Second class proteins are
mainly vegetable proteins such as beans and peas.
The relationship between amino acid side chains and protein conformation
The defining feature of an amino acid is its side chain (at top, blue circle; below, all
colored circles). When connected together by a series of peptide bonds, amino
acids form a polypeptide, another word for protein. The polypeptide will then fold
into a specific conformation depending on the interactions (dashed lines) between
its amino acid side chains.
Some important proteins are:
1. Insulin (a hormone)
2. Haemoglobin (Oxygen-carrying pigment in the blood).
3. Ribonuclease (an enzyme)
4. Collagen (a muscle protein)
The hydrolysis is catalysed by an acid or base. The amino acids obtained on
hydrolysis can be separated and identified by using paper chromatography.

Test for Proteins

1. Biuret Test: If copper (II) tetraoxosulphate (VI) is added to a solution of a protein
and the resulting solution made alkaline with sodium hydroxide solution, a violet
colour develops.
2. Million’s Reagent Test: Add a drop or two of million’s reagent to some egg – white
solution in a test tube. The formation of a white precipitate which turns brick red
on heating indicates the presence of protein.
3. Trioxonitrate (V) acid Test: Add three or four drops of concentrated trioxonitrate
(V) acid to 2cm3 of egg-white solution. The formation of an intense yellow colour
indicates the presence of proteins.
 Properties
Denaturation – Proteins usually form colloidal solutions. When such solutions are
heated, the proteins precipitate or coagulate. This is due to the irreversible changes
in the molecular shapes of the proteins, and the proteins are said to have been
denatured. Proteins are easily denatured by

 temperatures above 40C

 certain organic solvents and chemical reagents
 variations in pH
Hydrolysis
Proteins can be hydrolyzed to give amino acids by boiling them with solutions of
hydrochloric acid. Hydrolysis is carried out using suitable enzymes.

Assessment
Explain the denaturation of protein.
Mention 3 examples of proteins
 WEEK: 5
Chemistry SS3 First Term
Topic: Polymers

Polymers are giant molecules which are formed by joining together a large number
of much smaller molecules thus, forming a long chain. The smaller molecule, which
is the starting material, is known as monomer (meaning one unit). The molecular
size of a given polymer is not fixed.
There are two types of polymers: synthetic and natural.
Synthetic polymers: are derived from petroleum oil, and made by scientists and
engineers. Examples of synthetic polymers include nylon, polyethylene, polyester,
Teflon, and epoxy.
Natural polymers: occur in nature and can be extracted. They are often water-
based. Examples of naturally occurring polymers are silk, wool, DNA, cellulose and
proteins. Rubber and many resins are also natural polymers found in
plants.Polymerization is any process in which relatively small molecules,
called monomers, combine chemically to produce a very large chain-like or
network molecule, called a polymer. The monomer molecules may be all alike, or
they may represent two, three, or more different compounds. Usually at least 100
monomer molecules must be combined to make a product that has certain unique
physical properties—such as elasticity, high tensile strength, or the ability to form
fibres. The formation of stable covalent chemical bonds between the monomers
sets polymerization apart from other processes, such as crystallization, in which
large numbers of molecules aggregate under the influence of weak intermolecular
forces.
Two classes of polymerization usually are condensation and addition
polymerization.
Condensation polymerization is the process whereby two or more monomers link
together to form the polymer with the elimination of a small molecule.
In condensation polymerization, each step of the process is accompanied by
formation of a molecule of some simple compound, often water. Two most
important condensation polymers are nylon and terylene.
Addition polymerization is the process whereby two or more of the same
monomers link together to form the polymer without elimination of any small
molecules. In addition polymerization, monomers react to form a polymer without
the formation of by-products. Addition polymerizations usually are carried out in
the presence of catalysts, which in certain cases exert control over structural details
that have important effects on the properties of the polymer. Addition polymers
include poly(ethene), poly(chloroethene), perspex, etc.
Other Terms

Linear polymers: which are composed of chain-like molecules, may be viscous

liquids or solids with varying degrees of crystallinity. A number of them can be
dissolved in certain liquids, and they soften or melt upon heating.

Cross-linked polymers: in which the molecular structure is a network, are

thermosetting resins (i.e., they form under the influence of heat but, once formed,
do not melt or soften upon reheating) that do not dissolve in solvents.

Note: Both linear and cross-linked polymers can be made by either addition or
condensation polymerization.

Plastics
Plastics are natural/synthetic materials. They are produced by chemically modifying
natural substances or are synthesized from inorganic and organic raw
materials. They are giant molecules which are products of polymerization of simple
unsaturated compounds like ethene, propene or substituted unsaturated
compounds like styrene (phenylethene) and vinyl chloride chloroethene. Plastics
also include synthetic fibres like nylon and terylene. Plastics are either:
Thermoplastic: This can soften upon heating and return to their original form.
They are easily molded and extruded into films, fibers and packaging. Examples
include polyethylene (PE), polypropylene (PP) and polyvinyl chloride (PVC).
Thermoset or thermosetting plastics: Once cooled and hardened, these plastics
retain their shapes and cannot return to their original form. They are hard and
durable. Thermosets can be used for auto parts, aircraft parts and tires. Examples
include polyurethanes, polyesters, epoxy resins and phenolic resins.
Polyethylene, LDPE and HDPE: The most common polymer in plastics is
polyethylene, which is made from ethylene monomers (CH2=CH2). The first
polyethylene was made in 1934. Today, we call it low-density polyethylene (LDPE)
because it will float in a mixture of alcohol and water. In LDPE, the polymer strands
are entangled and loosely organized, so it’s soft and flexible. It was first used to
insulate electrical wires, but today it’s used in films, wraps, bottles, disposable
gloves and garbage bags.
In the 1950s, Karl Ziegler polymerized ethylene in the presence of various metals.
The resulting polyethylene polymer was composed of mostly linear polymers. This
linear form produced tighter, denser, more organized structures and is now called
high-density polyethylene (HDPE). HDPE is a harder plastic with a higher melting
point than LDPE, and it sinks in an alcohol-water mixture. HDPE was first introduced
in the hula hoop, but today it’s mostly used in containers.
Polyvinyl Chloride (PVC): PVC is a thermoplastic that is formed when vinyl chloride
(CH2=CH-Cl) polymerizes. When made, it’s brittle, so manufacturers add a
plasticizer liquid to make it soft and moldable. PVC is commonly used for pipes and
plumbing because it’s durable, can’t be corroded and is cheaper than metal pipes.
Over long periods of time, however, the plasticizer may leach out of it, rendering it
brittle and breakable.
Polystyrene (Styrofoam): Polystyrene is formed by styrene molecules. The double
bond between the CH2 and CH parts of the molecule rearranges to form a bond
with adjacent styrene molecules, thereby producing polystyrene. It can form a hard
impact-resistant plastic for furniture, cabinets (for computer monitors and TVs),
glasses and utensils. When polystyrene is heated and air blown through the
mixture, it forms Styrofoam. Styrofoam is lightweight, moldable and an excellent
insulator.
Polypropylene (PP): In 1953, Karl Ziegler and Giulio Natta, working independently,
prepared polypropylene from propylene monomers (CH2=CHCH3) and received
the Nobel Prize in Chemistry in 1963. The various forms of polypropylene have
different melting points and hardnesses. Polypropylene is used in car trim, battery
cases, bottles, tubes, filaments and bags.
Nylon: Nylon fibre is obtained by heating hexane dioc acid (adipic acid) with hexane
1,6 diamine. Nylon was originally developed as a textile but is available in many
forms with vastly different properties. Engineering nylon grades are easy to
machine with good resistance to biological attack. Unfortunately nylons can absorb
moisture from the atmosphere and can degrade in strong sunlight (they are
unstable in ultraviolet light) unless a stabilizing chemical is added at the initial
manufacture of the plastic. Nylons are easy to mould. Nylons also have a natural
‘oily’ surface that can act as a natural lubricant. Nylons are used for everything from
clothes through to gears and bearings. It is also slippery and can be used to make
washers, spacers and bushes.
Epoxy Resin: is a thermosetting polymer formed from reaction of an epoxide resin
with polyamine hardener. The resin consists of monomers or short chain polymers
with an epoxide group at either end. Most common epoxy resins are produced from
a reaction between epichlorohydrin and bisphenol A. Epoxy resin has various uses
such as; the resins that are cured through exposure with ultra violet light are
normally used in optoelectronics, fibre optic and dentistry. Industrial tooling
applications normally use resin to make laminates, fixtures, castings and moulds.
In the electronic world, epoxy resin is used to make transformer, insulator, switch
gear and generators.
Terylene: This is a polyester that is also known as Dacron in the U.S.A. It is formed
by the condensation of benzene- 1,4- dicarboxylic acid (terephthalic acid) and
ethane- 1,2- diol (ethylene glycol), using an acid catalyst.

Terylene is another synthetic fibre which is mainly used in the manufacture of

synthetic textiles. Terylene is mostly used for clothing, ropes, sheets, sails and many
others.

Natural Rubber
Rubber is obtained from the rubber tree, Hevea brasiliensis. When the bark of the
tree is cut, a thick white liquid called latex oozes out. if the latex is collected and
heated, it changes into an elastic solid called rubber. This rubber is of little use
 because it is soft and sticky. Chemically, it consists of 2-methyl buta-1,3-diene
monomers known previously as isoprene. When the monomers polymerize, they
form a long polymeric chains which have only a limited number of cross-links
between them. By adding sulphur and heating, the soft rubber becomes hard.
Vulcanization is chemical process by which the physical properties of natural
or synthetic rubber are improved. In its simplest form, vulcanization is brought
about by heating rubber with sulfur. Finished rubber has higher tensile strength
and resistance to swelling and abrasion, and is elastic over a greater range of
temperatures.

Synthetic Rubber
1,3-Butadiene is an important industrial chemical used as a monomer in the
production of synthetic rubber. Buta-1,3-diene co-polymerize with phenylethene
to form the tyre of synthetic rubber known as bunas or styrene butadiene rubber
(SBR).

Assessment

1. ……… are giant molecules which are formed by joining together a large number of
much smaller molecules thus, forming a long chain.
2. Explain Epoxyresin.
3. What do you understand by Polystyrene?

Answers

1. Polymers
 WEEK: 6
Chemistry SS 3 First Term
Topic: Carbohydrates

Carbohydrates (saccharides) – Molecules consist of carbon, hydrogen and oxygen

atoms. It is a major food source and a key form of energy for most organisms. When
combined together to form polymers, carbohydrates can function as long term
food storage molecules, as protective membranes for organisms and cells, and as
the main structural support for plants and constituents of many cells and their
contents. Carbohydrates are naturally occurring organic compounds. The general
molecular formular is Cx(H2O) y.

Types of Carbohydrates
Carbohydrates can be classified into two i.e. simple and complex sugars.
1. Simple carbohydrates: These are also called simple sugars. They are crystalline,
soluble in water and have a sweet taste. Structurally, they can be further divided
into monosaccharides e.g. glucose, and disaccharides e.g. sucrose
2. Complex carbohydrates: These are also called polysaccharides. They are non-
crystalline, insoluble and tasteless substances, e.g. starch and cellulose. Starches
include grain products, such as bread, crackers, pasta, and rice.

Monosaccharides
A monosaccharide or simple sugar has a formula that is some multiple of CH 2O. For
instance, glucose (the most common monosaccharide) has a formula of C 6H12O6.
This is the smallest possible sugar unit. Examples include glucose, galactose or
fructose. Monosaccharides cannot be split into smaller units by the action of dilute
acids. Monosaccharides are classified according to the number of carbon atoms
they possess: trioses have three carbon atoms; tetroses, four; pentoses, five;
hexoses, six; etc. Each of these is further divided into aldoses and ketoses,
depending on whether the molecule contains an aldehyde group (–CHO) or a
ketone group (–CO–). For example glucose, having six carbon atoms and an
aldehyde group, is an aldohexose whereas fructose is a ketohexose. These
aldehyde and ketone groups confer reducing properties on monosaccharides: they
can be oxidized to yield sugar acids.

Glucose
Glucose commonly known as grape sugar or dextrose is present in fruits such as
grapes, honey and also sap of plants. It is the main source of energy for mineral
tissues and its present in the blood of animals.
Preparation
Glucose can be prepared from the hydrolysis of starch with dilute acid.
(C6H10O5)n + nH2O → nC6H12O6
Properties
If Glucose is heated with concentrated teteraoxosulphate(vi) acid, it will be
dehydrated to form a black residue of carbon.

Disaccharides
A disaccharide is a derived or condensed from two molecules of monosaccharides
by the elimination of one molecule of water.
monosaccharide + monosaccharide ↔ dissacharide + water

Sucrose
Sucrose or cane sugar is the common granulated sugar which we use to sweeten
food. It occurs naturally in many plants and fruits. e.g. pineapple, carrots, sorghum
and sap of sugar maple tree.
Preparation
Sucrose is prepared from juices of sugar cane and sugar beet. The cane or beet is
shredded and crushed between rollers and the juice is extracted with water
warmed to about 800C. The solution is then purified by treatment with slaked lime
and carbon(iv) oxide. The purified solution is concentrated by distillation under
reduced pressure. On cooling, the concentrated solution, brown crystals of sugar
separate out. The remaining liquid called molasses still contains a reasonable
 amount of sugar and is used in ethanol production. The brown sugar obtained is
impure. It is refined by treatment with slaked lime and carbon (iv) oxide, and
decolorized with animal charcoal.
Properties
Sucrose is a colourless crystalline solid. It has a very sweet taste and dissolves
readily in water but not alkanol.
Sucrose chars on strong heating or warming with concentrated tetraoxosulphate
(vi) acid.
If sucrose is heated to a temperature of about 210 celcius which is above melting
point but below its charring temperature, a yellowish – brown substance known
as caramel. Caramel is used for flavouring and in confectionery.
sucrose + water ↔ glucose + fructose
Sucrose is used to sweeten food and beverages, and for preserving food. It is also
used to produce ethanol by fermentation.

Polysaccharides
Preparation of Starch
The raw material to be used is peeled cassava tubers which should be washed and
ground into pulp. Water is then added to the pulp to extract starch. It forms
suspension and this can stay for sometime before the water above is decanted and
starch residue is allowed to dry.
Physical Properties of Starch
1. Starch is a white odourless, tasteless powder with the formula (C 6H10O5)n
2. It is insoluble in cold water but soluble in hot water forming a viscous solution
which sets into a jelly on cooling
Chemical Properties of Starch
1. Starch gives the familiar characteristics deep blue colour with iodine solution
2. Hot dilute acids hydrolyse starch into maltose and glucose
3. It does not reduce Fehling’s solution
4. It decomposes on heating in the presence of the enzyme diastase to form maltose
sugar.
Test for Starch
 Add a few drops of iodine to some boiled starch, a dark blue colouration which
disappears on cooling results.
Uses of Starch
1. It is used for stiffening linen
2. It is used to produce ethanol and glucose
3. It is used mainly as food

Cellulose
Cellulose is the highest of the polysaccharides. It is the main component of plant
cell walls and plant fibres. The principal industrial sources are cotton and wood
each of which contains about 50% of cellulose. Other sources of cellulose for textile
purposes are floxi china grass, hemp and jute.

Physical Properties of Cellulose

1. It forms transparent fibres when it is pure
2. It is insoluble in water and in most organic solvents
Chemical Properties of Cellulose
1. Cellulose can be completely hydrolysed to glucose by hot acids
2. Hydrolysis of cellulose can also be carried out readily by the enzyme cellulose which
is produced by micro organisms present in the digestive system of termites and
herbivorous animals.
Uses of Cellulose
1. It is used in the manufacture of explosives, surface coatings, paper, textiles and
ropes
2. In the manufacture of gum, cotton and explosives

Assessment

1. Chemical formular for starch is ………

a. (C5H10O5)n
b. (C6H12O5)n
 c. (C6H10O5)n
d. (C6H11O5)n
2. A …………… is derived or condensed from two molecules of single sugar by the
elimination of one molecule of water.
a. Disaccharide
b. Monosaccharide
c. Polysaccharide
d. Starch
3. Carbohydrates consists of ……
a. C, H, N
b. C, H, O
c. C and H
d. C and O
4. sucrose + water ↔ glucose + ………
a. glucose
b. galactose
c. fructose
d. maltose
5. Test for starch -Add a few drops of iodine to some boiled starch, a ………..
colouration which disappears on cooling results
a. Blue
b. Bluish-green
c. Dark green
d. Dark blue

Answers

1. C
2. A
3. B
4. C
5. D
 WEEK 7
Chemistry SS 3 First Term
Topic: Metals and Their Compounds
Metals
Physical Properties of Metals
The physical properties of elements are dependent on

 the arrangement of their atoms or molecules in crystal lattices when in solid state
 the bonds that bind the atoms or molecules in the solid, liquid or gaseous state.
Most metals are solid at room temperature and exist as crystal lattices in which
their atoms are held together by strong metallic bonds. metals have the following
physical properties

1. High melting and boiling points

2. Characteristic lustre
3. Malleability – can be hammered into sheets
4. Ductility – can be drawn into a thin wire
5. Sonorousity – give off a note when hit
6. Hard but not brittle with great tensile strength
7. Relatively high densities
8. Good conductors of heat and electricity
Some metals do not exhibit all the above properties e.g
Mercury is a liquid with a melting point of -39oC. Sodium and potassium are light,
soft metals with low melting points of 97oC and 63oC respectively.

Chemical Properties

1. ionization behaviour – metallic ions have few valence electrons and so have a great
tendency to form positive ions by losing electrons. i.e. they are electropositive
2. reducing and oxidizing agents – metals are reducing agents because they donate
electrons readily during chemical reactions.
3. reaction with acids – a metal is more electropositive than hydrogen readily
displaces the hydrogen ion from an acid. This is a redox reaction with the metallic
 ions donating electrons to form metallic ions and the hydrogen ions accepting
electrons to form gaseous hydrogen.
4. nature of oxides – most metals react with oxygen to form basic oxides which are
mainly ionic compounds.Soluble basic oxides form alkalis. Some metals like
aluminium and zinc form amphoteric oxides.

Occurrence of Metals
Element which have low chemical reactivity generally occur native or free or
metallic state. Eg. Au, Pt, noble gas etc. element which are chemically reactive,
generally occur in the combined state. Eg. Halogen, chalcogens etc. the natural
materials in which the metals occur in the earth are called minerals. The minerals
from which the metals is conveniently and economically extracted is called an ore.
All the ores are minerals but all ores cannot be ores. Ores may be divided into four
groups.

 Metallic core (siderophile) of the earth crust contains (Mn, Fe, Co, Ni, Cu, Ru, Rb,
Pd, Ag, Re, Os, Ir, Pt, Au). Entire composition of metals in the earth crust may be
given as, Au(8.3%); Ca(3.6%); Na(2.8%); K(2.6%); Mg(2.1%); Ti(0.4%); Mn(0.1%);
Fe(5.1%) other metals (0.1%).
 Native ores: These ores contains metals in free state, e.g. silver, gold, platinum,
mercury, copper, etc. These are found usually associated with rock or alluvial
materials like clay, sand, etc. sometime lumps of pure metals are also found. These
are termed nuggets. Irons is found in free state are meteorites which also have 20
to 30% nickel.
 Sulphurised and arsenical ores: these ores consist of sulphides and arsenides in
simple and complex forms of metals. Some or ores are:

Metal Name of the ore Compositions

Pb Galena Pbs
Zn Zinc blender Zns
Ag Cinnabar Hgs
Fe iron pyrites Fes2
Ni Kufer nickel NiAs
Cu Copper pyrites Cu2s

III. Oxidized ores: In these ores, metals are present as their oxides or oxysalts such
as carbonates, nitrates, sulphates, phosphates, silicates, etc.
Important ores of this groups are listed below,

Oxides
Haemalite Fe2O3
Magnetite Fe2O4
Limonite Fe2O3.3H2O
Bauxite Al2O3.2H2O
Corundum Al2O3
Diaspore Al2O3.H2O
Chromite FeO.Cr2O3
Chromeochre FeO.Cr2O3
Tinstone (Cassiterite) Cr2O3
Chrysoberyl BeO.Al2O3
Cuprite (Rubby copper) Cu2O
Pyrolusite MnO2
Zincite Zno
Rutile TiO2
Llmenite FeO.TIO2

Carbonates
Magnesite MgcO3
Lime stone CaCO3
Dolomite CaCO3.MgCO3
Calamine ZnCO3
Malachite CuCO3.Cu(OH)2
Azurite Cu(OH)2.2CuCO3
Cerussite PbCO3
Siderite FeCO3
 Nitrates
Chile saltpeter NaNO3
Salt petre KNO3
Sulphates
Epsom salt MgSO4.7H2O
Barites BasO4
Gypsum CaSO4.2H2O
Glauber’s salt Na2SO4.10H2O
Anglesite PbSO4
Schonite K2SO4.MgSO4.6H2O
Polyhalite K2SO4.MgSO4.CaSO$.2H2O
Phosphates and Silicates
Lepidolite (Li, Na, K)2Al2(SiO3)(F,OH)2
Petalite liAl(Si2O5)2
Triphylite (Li, Na)3PO4,(Fe, Mn)3(PO4)2

 Halide ores : Metallic halides are few in nature, chlorides are most common . for
example,
Comman salt NaCl.

Extraction of Metals
General Principles
Metals found in combined forms exist as positive ions. During extraction, the
metallic ions can be reduced to their corresponding metal atoms. This can be done
electrolytically or by chemical and thermal methods. The method chosen depends
on the stability of the ore which in turn depends on the position of the metal in the
activity series.
1. Mining of ore containing rock
The composition of rock around the world varies greatly and locations with metal
bearing ore have been sought ever since man was able to extract metals. Nowadays
the search is still going on for important deposits of rock with high percentages of
the mineral in question. This search is now taking place under the sea and in other
inhospitable environments.
Recently, for example, rock containing an appreciable percentage of rare earth
elements has been discovered under the pacific ocean. This is a particularly
important discovery as virtually 99% of known working deposits are in china and
rare earths are essential in the manufacture of the strong neodymium magnets
needed for the computer industry
2. Separation, purification or preparation of useful ore
Very few metal ores occur in a pure enough form to be used directly in the
extraction process. The first stage is to separate the useful ore from the rock. This
may not be necessary in some cases, for example, the extraction of iron, but
essential in the extraction of aluminium.
This separation may be physical, such as floatation, or chemical such as digestion
of the required compound in a strong base or acid followed by re-precipitation and
filtration.
Most ores are either oxides or sulfides. The sulfides are usually converted to oxides
by roasting in air. This tends to release sulfur in the form of sulfur(IV) oxide, a
pollutant and acidic gas. However, it is also a useful gas in that it is used for the
manufacture of sulfuric acid by the contact process.
3. Extraction of metal from ore
Metals are all electropositive and need to be reduced to become metallic elements.
Hence, all extraction processes use reduction. For the less reactive metals chemical
reduction suffices, but for the more reactive metals electrochemical reduction is
needed.
4. Purification of metal
Metals that are extracted by reductive processes usually need to be further
processed to make them industrially useful.

Uses of Metals
Metals are very useful to people. They are used to make tools because they can be
strong and easy to shape. Iron and steel have been used to make bridges, buildings,
or ships.
Some metals are used to make items like coins because they are hard and will not
wear away quickly. For example copper (which is shiny and red in color), aluminium
 (which is shiny and white), gold (which is yellow and shiny), and silver and nickel
(also white and shiny).
Some metals, like steel, can be made sharp and stay sharp, so they can be used to
make knives, axes or razors.
Rare metals with high value, like gold, silver and platinum are often used to make
jewellery. Metals are also used to make fasteners and screws. Pots used for cooking
can be made from copper, aluminium, steel or iron. Lead is very heavy and dense
and can be used as ballast in boats to stop them from turning over, or to protect
people from ionizing radiation.

Assessment

1. Most metals exist in nature as

a. crusts
b. alloys
c. ores
d. felspar
2. Most metals are malleable with high densities and have high boiling points except
a. Zn
b. K
c. Sn
d. Ca
3. Method adopted in extracting a particular metal from its ore depends on
a. the fragile nature of the metal
b. the location of the ore in the earth’s crust
c. the stability of the ore which depends on the position of the ore in E.C.S
d. the availability of power in the country
4. Metals found in combined forms exist as ……… ions
a. positive
b. negative
c. neutral
d. none of the above
 Answers

1. C
2. B
3. C
4. A
 WEEK 8
Chemistry SS 3 First Term
Topic: Calcium and Aluminium
Calcium

Calcium is too reactive to occur as a free metal in nature. It occurs abundantly in

the combined state as calcium trioxocarbonate (iv) in limestone, marble, chalk,
aragonite, calcite, coral, dolomite, calcium fluoride. etc In Nigeria, limestone is
found at Nkalagu in Ebonyi, Ewekoro at Abeokuta and Ukpilla in Delta state.

Extraction
Since calcium are very stable, metallic calcium is commonly extracted
electrolytically from fused calcium chloride a byproduct of solvay process. Some
calcium chloride is usually added to the fused calcium chloride to lower the melting
point from 850oC to about 650oC. The mixture is placed in a large crucible lined on
the inside with graphite which serves as the anode of the cell. The cathode consists
of iron rod which just touches the surface of the electrolyte. As electrolysis
proceeds metallic calcium collects on the cathode which is gradually raised so that
an irregular stick of calcium is formed on it. Chlorine is liberated at the cathode.

Chemistry of the Reaction

At the cathode – the calcium ions receive two electrons each to become reduced
to the metal.
At the anode – two chloride ions give up an electron each to become atomic
chlorine. The two atoms then combine to become liberated as a gaseous molecule.
Cl– → Cl– + e–
Cl + Cl → Cl2
Overall electrolytic reaction
Ca2(l) + 2Cl–(l) → Ca(s) + Cl2(g)
 Physical Properties of Calcium

1. Appearance – Silvery grey solid

2. Relative density is 1.55
3. Calcium is malleable and ductile
4. It has relatively low tensile strength
5. Melting point is 850oC
6. Calcium is a good conductor of heat and electricity

Chemical Properties of Calcium

Reaction with air – Calcium is a very electropositive and reactive metal. On
exposure to air, it rapidly tarnishes ad loses its metallic lustre due to the formation
of white film of calcium oxide or quick lime on the surface of the metal. When
calcium is heated in air, it burns with a brick red flame to form calcium oxide
2Ca(s) + O2(g) → 2CaO(s)
CaO(s) + H2O(l)→ Ca(OH)2(g)
Reaction with non-metals – on heating, calcium combines directly with nitrogen,
chlorine, Sulphur and hydrogen
3Ca(s) + N2(g) → Ca3N2(s)
Ca(s) + Cl2(g) → CaCl2(s)
Reaction with water – Calcium reacts slowly with cold water and rapidly with warm
water to form calcium hydroxide and hydrogen
Ca(s) + 2H2O(g) → Ca(OH)2(aq) + H2(g)
Reaction with Ammonia – if ammonia is passed over heated calcium, it reacts as
follows
3Ca(s) + 2NH3(g) → Ca3N2(s) + 3H2(g)
Test for calcium ions
Flame test – calcium compounds give an orange – red colour to a non-luminous
flame. Moisten the unknown compound with a few drops of concentrated
hydrochloric acid. Dip the tip of a clean platinum wire into the mixture and hold it
in a non-luminous Bunsen flame. If a bright brick red flame through a blue glass is
produced, the unknown ions of the compound are calcium ions.
 With sodium hydroxide – Add a few drops of NaOH solution to an unknown salt.
The formation of white precipitate which is insoluble in excess sodium hydroxide
indicate the presence of calcium ions

Uses of Calcium

1. Calcium is used as a deoxidant in steel castings and copper alloys.

2. It is also used in the manufacture of calcium fluoride and calcium hydride.
3. It is used in the extraction of uranium
4. It is needed in the diet of young children for development of strong bones and teeth
5. Calcium metal is used as a reducing agent in preparing other metals such as thorium
and uranium. It is also used as an alloying agent for aluminium, beryllium, copper,
lead and magnesium alloys.
6. Calcium compounds are widely used. There are vast deposits of limestone (calcium
carbonate) used directly as a building stone and indirectly for cement.
7. Gypsum (calcium sulfate) is used by builders as a plaster and by nurses for setting
bones, as ‘plaster of Paris’.

Aluminium
Aluminium is the most common metal in the Earth’s crust, making up 7.5% by mass.
Its main ore is bauxite-a clay mineral which you can think of as impure aluminium
oxide. It is the most important element in group III.

Extraction of Aluminium
Aluminium is obtained largely from the ore bauxite (Al2O3.2H2O). Its production is
a two-step process: the purification of bauxite and extraction by electrolysis.

Purifying the bauxite (aluminium oxide) – the Bayer Process

Crushed bauxite is treated with moderately concentrated sodium hydroxide
solution. The concentration, temperature and pressure used depend on the source
of the bauxite and exactly what form of aluminium oxide it contains. Temperatures
are typically from 140°C to 240°C; pressures can be up to about 35 atmospheres.
High pressures are necessary to keep the water in the sodium hydroxide solution
liquid at temperatures above 100°C. The higher the temperature, the higher the
pressure needed.
With hot concentrated sodium hydroxide solution, aluminium oxide reacts to give
a solution of sodium aluminate (III) (NaAl(OH) 4).

Extraction of Aluminium by Electrolysis

After purification, aluminium oxide is mixed with cryolite (sodium aluminium
fluoride) Na3AlF6 to lower the melting point from 2000º to 1000º, which saves
money. This mixture is heated and the molten liquid used as the electrolyte. Both
electrodes are made of graphite (carbon). The anode (+ve) is graphite and the
cathode (-ve) is a graphite lining to a steel case.
The anode disintegrates. The hot oxygen produced here reacts with the hot carbon
anode to give carbon dioxide. Hence it must be replaced regularly.
 Aluminium ions are attracted to the cathode (the negative electrode) and are
reduced to aluminium by gaining electrons.
Al3+ (l) + 3e– → Al (l)
The molten aluminium produced sinks to the bottom of the cell.
The oxide ions are attracted to the anode and lose electrons to form oxygen gas.
2O2- (l) → O2 (g) + 4e–
Note: The extraction of aluminium is an expensive process because the large
amount of electricity needed to keep the electrolytes molten is expensive. Hence
using cryolite saves energy and money, as it acts as a solvent for the aluminium
oxide and melts at a much lower temperature.

Physical Properties of Aluminium

1. Aluminium is a silvery white metal which is comparatively soft

2. It is a strong, malleable metal element.
3. It has a low density.
4. It is resistant to corrosion
5. It is a good conductor of heat and electricity.
6. It can be polished to give a highly reflective surface.
Chemical Properties of Aluminium

 Action with air: Aluminium burns in air at high temperature to form the oxide
and the nitride
4Al (s) + 3O2 (g) ——> 2Al2O3 (s)

 Reaction with non-metals: On heating, aluminium combines directly with non-

metals like the halogens, sulphur, nitrogen, phosphorus and carbon with
evolution of heat
2Al (s) + 3Cl2 ——> 2AlCl3 (s)

 Action with acids: Aluminium reacts more rapidly with the concentrated
hydrochloric acid to displace hydrogen but more slowly with dilute one
2Al (s) + 6HCl (aq) ——> 2AlCl3 (aq) + 3H2 (g)

 Reaction with alkalis: Aluminium reacts with both sodium and potassium
hydrogen solutions giving hydrogen gas and soluble tetrahydroxoaluminate (III)
2Al (s) + 2NaOH (aq) + 6H2O (l) ——> 2NaAl(OH)4 (aq) + 3H2 (g)

 Reaction with Iron (III) oxide: Aluminium reduces iron (III) oxide to molten
iron. The reaction is used in thermit process and it gives out a great deal of
energy
2Al (s) + Fe2O3 (s) ——> Al2O3 (s) + 2Fe (s)

Uses

1. Low density and strength make aluminium ideal for construction of aircraft,
lightweight vehicles, and ladders.
2. An alloy of aluminium called duralumin is often used instead of pure
aluminium because of its improved properties.
3. Easy shaping and corrosion resistance make aluminium a good material for
drink cans and roofing materials.
4. Corrosion resistance and low density leads to its use for greenhouses and
window frames.
 5. Good conduction of heat leads to its use for boilers, cookers and cookware.
6. Good conduction of electricity leads to its use for overhead power cables
hung from pylons (low density gives it an advantage over copper).
7. High reflectivity makes aluminium ideal for mirrors, reflectors and heat
resistant clothing for fire fighting.

Assessment
Explain the method of extraction of Aluminium in few lines and bullet points.
Mention 3 uses of Calcium
 WEEK 9
Chemistry SS3 First term
Topic: Tin and Copper
Tin
Tin does not occur naturally as a free element. The main source is the mineral
cassiterite or tin stone, SnO2. In Nigeria was mined in 1930 at Jos Plateau.

Extraction from its Ores

Tin ore is crushed and washed with water. This is called the concentration of the
ore. The tin ore is roasted in air to remove impurities, such as arsenic and
antimony as volatile oxides. The product is mixed with powdered charcoal and
heated to 1300oC. to reduce the oxide. Molten tin is tapped off. Iron compounds,
which might be present as impurities are removed by electromagnetic separation.
The metal is extracted from its ore by carbon reduction. The concentrated ore is
mixed with coke and heated in a furnace.
SnO2(s) + 2C(s) → Sn(l) + 2CO(g)
The tin obtained is purified. It is separated from copper, iron and any other element
present as impurities by either thermal -heating beyond it melting point of 232 K,
and running off the molten tin, leaving behind any less fusible impurities – or by
electrolytic means.
Much purer tin is obtained by the electrolysis of aqueous solution of tin(II) chloride,
SnCl2 – the impure tin is made anode, while the cathode is pure tin.

Test for Tin ions

Hydrogen Sulphide – Pass hydrogen sulphide into a solution of the unknown salt
acidified with dilute hydrochloric acid. Tin(ii) ions are present if a brown precipitate
which dissolves in yellow ammonium sulphide and in hot concentrated
hydrochloric acid is obtained
Mercury (ii) chloride- The formation of a white precipitate of mercury(i) chloride
would indicate the presence of tin(ii) ions.
Physical Properties of Tin

1. Tin is solid, with silvery white appearance with lustre.

2. It’s melting point is 232oC.
3. It is malleable and soft (enough to be cut with a knife).
4. It is a good conductor of heat and electricity.
5. It exists in three different forms. These are grey tin of density 5.76 g/cm3; white
tin of density 7.28 g/cm3; and rhombic tin of density 6.6 g/cm3. These allotropes
(the different forms) can be converted from one to another by changes in
temperature.
grey white rhombic
Grey 13.2oC White 161oC Rhombic

6. Tin is not ductile enough to be drawn into wires.

Chemical properties of Tin

1. Reaction with Oxygen – it does not react with oxygen, except at temperature
above 1300oC.
Sn(s) + O2(g) → SnO2(s) at temperatures greater than 1300oC
Hence, It does not corrode or explode in air.

2. With nitrogen and carbon – no reaction occurs.

3. With non metals, example, chlorine – it reacts when heated with chlorine to
form tin(IV) chloride.
Sn(s) + 2Cl2(g) → SnCl4(s)

4. With acids – it reacts with acids to different degrees at different concentrations

and temperatures. With dilute HCl – no reaction occurs.
With Concentrated HCl the reaction is rapid, producing tin(II) chloride.
Sn(s) + 2HCl(aq) → SnCl2(s) + H2(g)
With dilute H2SO4 – no reaction occurs. With hot concentrated H2SO4 a reaction
occurs to release SO2.
Sn(s) + 2H2SO2 → SnSO4(aq) + SO2(g) + 2H2O(g)
With dilute HNO3 – tin reacts with dilute HNO3, forming Sn(NO3)2 and hydrogen.
With conc. HNO3 – tin reacts with conc. HNO3 , forming SnO2.

5. With alkalis – it forms trioxostannate(IV) salts and hydrogen with conc. solutions
of alkalis.
Sn(s) + 2NaOH(aq) + H2O(l) → Na2SnO3(aq) + 2H2(g)

Uses of Tin
There are a number of ways tin can be used. These include:

1. For coating steel – the coating is done by electrolytic method. Tin prevents
corrosion in steel, such may be used to can food and drinks.
2. For making alloys together with lead, antimony and copper. Alloys of tin are
important, such as soft solder, pewter, bronze and phosphor bronze. A niobium-
tin alloy is used for superconducting magnets.
3. Due to its resistance to atmospheric corrosion and low melting point, it can be
used to make sheet glass.
4. The most important tin salt used is tin(II) chloride, which is used as a reducing
agent and as a mordant for dyeing calico and silk. Tin(IV) oxide is used for
ceramics and gas sensors. Zinc stannate (Zn2SnO4) is a fire-retardant used in
plastics.
5. Some tin compounds have been used as anti-fouling paint for ships and boats,
to prevent barnacles. However, even at low levels these compounds are deadly
to marine life, especially oysters. Its use has now been banned in most
countries.
6. A very important application of tin is tin-plating. Tin-plating is the process by
which a thin coat of tin is placed on the surface of steel, iron, or another metal.
Tin is not affected by air, oxygen, water, acids, and bases to the extent that steel,
iron, and other metals are. So the tin coating acts as a protective layer.
7. Another tin alloy is Babbitt metal. Babbitt metal is a soft alloy made of any
number of metals, including arsenic, cadmium, lead, or tin. Babbitt metal is
used to make ball bearings for large industrial machinery. The Babbitt metal is
laid down as a thin coating on heavier metal, such as iron or steel. The Babbitt
metal retains a thin layer of lubricating oil more efficiently than iron or steel.
Copper
Copper was one of the first metals discovered and used by man. It is a stable metal
readily obtained from its compounds. Copper ores are widely found around the
world. The main ores are copper pyrites (CuFeS2), malachite (CuCO3.Cu (OH)2),
chalcocite (CuS2) and cuprite (CuO).

Extracting copper from its ores

The method used to extract copper from its ores depends on the nature of the ore.
Sulphide ores such as chalcopyrite (copper pyrites) are converted to copper by a
different method from silicate, carbonate or sulphate ores.
The process:
The concentrated ore is heated strongly with silicon dioxide (silica) and air or
oxygen in a furnace or series of furnaces.

 The copper(II) ions in the chalcopyrite are reduced to copper(I) sulphide (which
is reduced further to copper metal in the final stage).
 The iron in the chalcopyrite ends up converted into an iron(II) silicate slag which
is removed.
 Most of the sulphur in the chalcopyrite turns into sulphur dioxide gas. This is
used to make sulphuric acid via the Contact Process.
An overall equation for this series of steps is:

The end product of this is called blister copper – a porous brittle form of copper,
about 98 – 99.5% pure.
Purification of copper
When copper is made from sulphide ores by the first method above, it is impure.
The blister copper is first treated to remove any remaining sulphur (trapped as
bubbles of sulphur dioxide in the copper – hence “blister copper”) and then cast
into anodes for refining using electrolysis.
Electrolytic refining
The purification uses an electrolyte of copper(II) sulphate solution, impure copper
anodes, and strips of high purity copper for the cathodes.
The diagram shows a very simplified view of a cell.

For every copper ion that is deposited at the cathode, in principle another one goes
into solution at the anode. The concentration of the solution should stay the same.
All that happens is that there is a transfer of copper from the anode to the cathode.
The cathode gets bigger as more and more pure copper is deposited; the anode
gradually disappears.
In practice, it isn’t quite as simple as that because of the impurities involved.

Physical properties of copper

1. Copper is a heavy, reddish-brown metal

2. It is very malleable and ductile
3. It has a density of 8.95 g cm-1
4. It is a good conductor of heat and electricity
5. It has a high melting point of 1083oC
6. It has a boiling point of 2300oC
7. It also forms alloys very readily
Chemical properties of Copper

 Reaction with Air: It is resistant to pure dry air, but over a long period of time in
a moist, impure atmosphere, it becomes coated with green, basic copper(II)
tetraoxosulphate (VI) CuSO4.3Cu(OH)2 and trioxocarbonate (IV).
On heating in air or oxygen, copper is readily oxidized to give black copper (II)oxide
2Cu(s) + O2(g) ——–> 2CuO(s)

 Effect of Acids: Copper is lower than hydrogen in the electrochemical series,

hence, it is not capable of displacing hydrogen from dilute acids. It is however
attacked by oxidizing acids like trioxonitrate (V) acid and tetraoxosulphate (VI)
acids
3Cu(s) + 8HNO3(aq) ——> 3Cu(NO3)2(s) + 4H2O(l) 2NO(g)

 Displacement Reaction: Because of its low position in the activity and

electrochemical series, copper is easily displaced from its compound
CuSO4(aq) + Fe(s) ——-> FeSO4(aq) + Cu(s)

 Hydrogen gas reduces copper oxides to the metal

CuO(s) + H2(g) ——> Cu(s) + H2O(l)

Test For Copper(II) Ions

With sodium hydroxide: Add a few drops of sodium hydroxide solution to a solution
of copper salt. The formation of a blue precipitate which is insoluble in excess
sodium hydroxide confirms the presence of copper(II) ions
Cu2+(aq) + 2NaOH(aq) ——-> Cu(OH)2(s) + Na+(aq)

Uses of Copper
Amongst other things copper is used for:

1. Electrical wiring. It is a very good conductor of electricity and is easily drawn out
into wires.
2. Domestic plumbing. It doesn’t react with water, and is easily bent into shape.
3. Boilers and heat exchangers. It is a good conductor of heat and doesn’t react
with water.
4. Baking brass. Brass is a copper-zinc alloy. Alloying produces a metal harder than
either copper or zinc individually. Bronze is another copper alloy – this time with
tin.
5. Coinage. In the UK, as well as the more obvious copper-coloured coins, “silver”
coins are also copper alloys – this time with nickel. These are known as
cupronickel alloys. UK pound coins and the gold-coloured bits of euro coins are
copper-zinc-nickel alloys.

Assessment

1. Tin exists in 3 different forms. Mention them?

2. How can Tin be extracted?
3. The method of extraction of copper depends on what?
4. Mention 4 uses of copper
 WEEK 10
Chemistry SS3 First Term
Topic: Iron

Iron
Iron is the most important element in the industry. It is the second most abundant
element on the earth crust after aluminium, but often occurs as a free metal.
The common ores are haematite found in united states, Australia and USSR. It can
also occur as impure iron (III) oxide (Fe2O3), Magnetite or magnetic iron ore (Fe3O4)
is found in Sweden and in North America.
Siderite or spathic iron ore, (FeCO3), found in Great Britain. Iron also occurs as iron
pyrites (FeS2) and limonite (Fe2O3.3H2O).
Iron is widely present as trioxosilicate (IV) in clay soils. Iron ore is available in Itakpe,
Ajaokuta, Jebba and Lokoja all in Kwara State (Nigeria).

Extraction of Iron
The extraction of iron from iron ore (haematite), using coke, limestone and air in a
blast furnace
Haematite is basically iron oxide, and the oxygen must be removed to leave the
iron behind. Reactions in which oxygen is removed are called reduction
reactions. Since carbon is more reactive than iron, it can displace the iron from its
oxide. Hence the method for extraction of iron is called ‘reduction by carbon’.
Coke is impure carbon, and it burns in the hot air blast to form carbon dioxide. This
is a strongly exothermic reaction which makes it an important reaction, as it helps
heat up the blast furnace. The iron ore, coke and limestone enter the blast furnace
at the top. The hot waste gases at the top of the furnace are piped away and used
to heat the air blast at the bottom.
C (s) + O2 (g) ——> CO2 (g)
At high temperatures in the furnace, the carbon dioxide is reduced by more carbon
to give carbon monoxide.
CO2 (g) + C (s) ——> 2CO (g)
It is the carbon monoxide which is the main reducing agent in the furnace-especially
in the cooler parts.
Assuming that the iron ore is haematite, Fe 2O3:
Fe2O3 (s) + 3CO (g) ——> 2Fe (l) + 3CO2 (g)
Due to the high temperatures, the iron produced melts and flows to the bottom of
the furnace, where it can be tapped off.
In the hotter parts of the furnace, some of the iron oxide is also reduced by carbon
itself.
Fe2O3 (s) + 3C (s) ——> 2Fe (l) + 3CO (g)
Notice that carbon monoxide is formed, rather than carbon dioxide, at these
temperatures.
However some use this equation instead:
iron oxide + carbon → iron + carbon dioxide
2Fe2O3 + 3C → 4Fe + 3CO2
The limestone is added to the furnace to remove impurities in the ore which would
otherwise clog the furnace with solid material.
The furnace is hot enough for the limestone (calcium carbonate) to undergo
thermal decomposition. It splits up into calcium oxide and carbon dioxide. This is
an endothermic reaction (it absorbs heat) and it is important not to add too much
limestone to avoid cooling the furnace.
CaCO3 (s) ——> CaO (s) + CO2 (g)
Calcium oxide is a basic oxide, and its function is to react with acidic oxides such as
silicon dioxide, SiO2. Silicon dioxide is the main constituent of sand, and is typical of
the sort of impurities that need to be removed from the furnace.
CaO (s) + SiO2 (s) —–> CaSiO3 (l)
The product is calcium silicate. This melts and trickles to the bottom of the furnace
as a molten slag, which floats on top of the molten iron as it is less dense, and can
be tapped off separately. Slag is used

Uses of iron

 Cast iron
Molten iron straight from the furnace can be cooled rapidly and solidified by
running it into sand moulds. This is known as pig iron. If the pig iron is remelted and
cooled under controlled conditions, cast iron is formed. This is very impure iron,
containing about 4% carbon as its main impurity. Although cast iron is very hard, it
is also very brittle and tends to shatter if it is hit hard. It is used for things like
manhole covers, gutterings and drainpipes, and cylinder blocks in car engines.
 Mild steel
Mild steel is iron containing up to about 0.25% carbon. This small amount of carbon
increases the hardness and strength of the iron. It is used for (among other things)
wire, nails, car bodies, ship building, girders and bridges.

 Wrought iron
This is pure iron. It was once used to make decorative gates and railings but has
now been largely replaced by mild steel. The purity of the iron makes it very easy
to work because it is fairly soft, but the softness and lack of strength mean that it
isn’t useful for structural purposes.

 High-carbon steel
High carbon steel is iron containing up to 1.5% carbon. Increases the carbon
content makes the iron harder, but at the same time it gets more brittle. High-
carbon steel is used for cutting tools and masonry nails. Masonry nails are designed
to be hammered into concrete blocks or brickwork where a mild steel nail would
bend. If you miss-hit a masonry nail, it tends to break into two bits because of its
increased brittleness.

 Stainless steel
Stainless steel is an alloy of iron with chromium and nickel. Chromium and nickel
form strong oxide layers in the same way as aluminium, and these oxide layers
protect the iron as well. Stainless steel is therefore very resistance to corrosion.
Obvious uses include kitchen sinks, saucepans, knives and forks, and gardening
tools. But there are also major uses for it in the brewing, dairy and chemical
industries where corrosion- resistant vessels are essential
Types of iron Iron mixed with Some uses
Decorative work such as
Wrought iron (pure iron)
gates and railings
Nails, car bodies, ship
Mild steel Up to 0.25% carbon
building, girders
Cutting tools, masonry
High-carbon steel 0.25-1.5% carbon
nails
 Manhole covers, guttering,
Cast iron About 4% carbon
engine blocks
Cutlery, cooking utensils,
Stainless steel Chromium and nickel
kitchen sinks
Physical Properties of Iron

1. Pure iron is a grey metal with density 7.8g cm-3

2. It melts at 1535oC and boils at 2800oC
3. It is very malleable and ductile
4. It is a good conductor of heat and electricity
5. It can easily be magnetized

Chemical Properties of Iron

 Reaction With Air: It burns in air to form reddish hydrated iron (III) oxide of
variable composition
4Fe(s) + 3O2(g) + 2xH2O(l) ——-> 2Fe2O3.xH2O

 When finely divided iron is heated in air it burns at high temperature to form
magnetic iron oxide, which behaves like a compound oxide
3Fe(s) + 2O2(g) ——> Fe3O4 (s)

 Reaction With Steam: When steam is passed over red-hot iron filings, tri iron
tetraoxide and hydrogen are produced and the reaction is reversible
Fe(s) + 4H20 <——> Fe3O4(s) + 4H2(g)
Order of reactivity Symbol Method of Extraction
Potassium K
Electrolysis

The metal compound is:

1. Melted, then
2. Has electricity passed through it
Assessment

1. What are the uses of the following Iron

– Cast Iron
– Stainless steel
– Wrought Iron
2. Stainless steel is an alloy of iron with ………. and ………
3. Mention 2 common ores of iron ……… and ………
 SS 3
SECOND TERM NOTES ON
CHEMISTRY
 TABLE OF CONTENT
SECOND TERM
WEEK 1: NUCLEAR CHEMISTRY (I)

WEEK 2: NUCLEAR CHEMISTRY (II)

WEEK 3: SIMPLE MOLECULES AND THEIR SHAPES

WEEK 4: PETROLEUM OR CRUDE OIL

WEEK 5: VOLUMETRIC (QUANTITATIVE) ANALYSIS

WEEK 6: QUALITATIVE ANALYSIS

 WEEK 1
TOPIC: NUCLEAR CHEMISTRY
Nuclear Reactions

Rutherford in 1919 transmitted nitrogen isotope into an oxygen isotope. The nitrogen was subjected to
the action of swift alpha – particles derived from radium salt.

Transmutation is the process by which radioactive elements change into different elements.

Nuclear reaction is a process in which two nuclei or nuclear particles collide, to produce different
products than the initial particles.

Nuclear fission and nuclear fusion both are nuclear phenomena that release large amounts of energy,
but they are different processes which yield different products. Learn what nuclear fission and nuclear
fusion are and how you can tell them apart.

Nuclear Fission

Nuclear fission takes place when an atom’s nucleus splits into two or more smaller nuclei. These smaller
nuclei are called fission products. Particles (e.g., neutrons, photons, alpha particles) usually are released,
too. This is an exothermic process releasing kinetic energy of the fission products and energy in the form
of gamma radiation. Fission may be considered a form of element transmutation since changing the
number of protons of an element essentially changes the element from one into another.

Nuclear Fission Example:

235
92U + 10n → 9038Sr + 14354Xe + 310n

Nuclear Fusion

Nuclear fusion is a process in which atomic nuclei are fused together to form heavier nuclei. Extremely
high temperatures (on the order of 1.5 x 107°C) can force nuclei together. Large amounts of energy are
released when fusion occurs.

Nuclear Fusion Examples

The reactions which take place in the sun provide an example of nuclear fusion:
1
1H + 21H → 32He
3
2He + 32He → 42He + 211H
1
1H + 11H → 21H + 0+1β
Comparison between Nuclear Fission and Fusion

Nuclear Fission Nuclear Fusion

Fission is the splitting of a large atom into two or more Fusion is the fusing of two or more lighter at
Definition:
smaller ones. into a larger one.

Natural occurrence
Fission reaction does not normally occur in nature. Fusion occurs in stars, such as the sun.
of the process:

Few radioactive particles are produced by fu

Byproducts of the
Fission produces many highly radioactive particles. reaction, but if a fission “trigger” is used,
reaction:
radioactive particles will result from that.

of the substance and high-speed neutrons are High density, high temperature environment
Conditions:
required. required.

Extremely high energy is required to bring tw

Energy Takes little energy to split two atoms in a fission
more protons close enough that nuclear forc
Requirement: reaction.
overcome their electrostatic repulsion.

The energy released by fission is a million times greater

The energy released by fusion is three to fou
Energy Released: than that released in chemical reactions; but lower
greater than the energy released by fission.
than the energy released by nuclear fusion.

One class of nuclear weapon is the hydrogen

One class of nuclear weapon is a fission bomb, also
Nuclear weapon: which uses a fission reaction to “trigger” a fu
known as an atomic bomb or atom bomb.
reaction.

Comparison of Nuclear Reaction and Ordinary Chemical Reaction

Nuclear Reaction Ordinary Chemical Reaction

During chemical reactions, elements do not lose their id

During nuclear reactions, the nuclei of atoms undergo change
In these reactions, only the electrons in the outermost s
and therefore new elements are formed as a result of such
atoms participate whereas the nuclei of atoms remain
reactions.
unchanged.

Reactivity of an element towards nuclear reactions is nearly

Reactivity of an element towards chemical reactions dep
independent of oxidation state of the element. For example, Ra
upon the oxidation state of the element. In ordinary che
element or Ra2+ ion in RaC2 behave s similarly during nuclear
reactions, Ra and Ra2+ behave quite differently.
reactions.

In nuclear reactions, isotopes behave quite differently. For

Different isotopes of an element have nearly same chem
example, U-235 undergoes fission quietly readily but U-238 does
reactivity.
not.

Rate of a nuclear reaction is independent of temperature and Rate of a chemical reaction is largely affected by temper
pressure. and pressure.
A nuclear reaction cannot be reversed. A chemical reaction can be reversed.

Chemical reactions are accompanied by relatively small

Nuclear reactions are accompanied by large energy changes.
changes.

Assessment

What is the most significant difference between Nuclear fission and fusion?

What does Transmutation mean?

 WEEK 2
TOPIC: NUCLEAR CHEMISTRY (II)
Introduction to Radioactivity

Antoine Henri Becquerel, a French scientist, 1896, was conducting an experiment which started with the
exposure of a uranium-bearing crystal to sunlight. Once the crystal had sat in the sunshine for a while,
he placed it on a photographic plate. As he had anticipated, the crystal produced its image on the plate.
Becquerel theorized that the absorbed energy of the sun was being released by the uranium in the form
of x-rays. The husband and wife team of Pierre and Marie Curie in 1898 became interested in
Becquerel’s discovery. While experimenting with their own uranium-containing ore, they came up with
the term “radioactivity” to describe the spontaneous emissions that they studied. This word is still used
today to describe this special characteristic of some elements (radioisotopes).

While comparing the activity of pure uranium to a uranium ore sample, they found that the ore was
significantly more radioactive than the pure material. They concluded that the ore contained additional
radioactive components besides the uranium. This observation led to the discovery of two new
radioactive elements which they named polonium and radium.

Radioactivity is the spontaneous disintegration of atomic nuclei. In other word, Radioactivity is the
energy and mass released by spontaneous changes in the nucleus of an atom.

Radiation is energy that travels (‘radiates’) as waves or particles. Heat, light, sound, microwaves, radar,
radio waves, X-rays, alpha and beta particles, and gamma rays are all forms of radiation.

All substance is made of atoms. These have electrons (e) around the outside, and a nucleus in the
middle. The nucleus consists of protons (p) and neutrons (n), and is extremely small. In some types of
atom, the nucleus is unstable, and will decay into a more stable atom.

Unstable atomic nuclei will spontaneously decompose to form nuclei with a higher stability. The
decomposition process is called radioactivity. The energy and particles which are released during the
decomposition process are called radiation. When unstable nuclei decompose in nature, the process is
referred to as natural radioactivity. When the unstable nuclei are prepared in the laboratory, the
decomposition is called induced radioactivity.

Types and Nature of Radiation

There are three different types of radiation: the alpha, beta and gamma. They have different
penetrating powers.
Alpha Radiation

Alpha radiation consists of a stream of positively charged particles, called alpha particles (α), which have
an atomic mass of 4 and a charge of +2 (a helium nucleus). When an alpha particle is ejected from a
nucleus, the mass number of the nucleus decreases by four units and the atomic number decreases by
two units. For example:
238
92U → 42He + 23490Th

The helium nucleus is the alpha particle.

Properties of Alpha Particles (α)

1. They are positively charged

2. They have low penetrating power (low energy) and because of their longer wavelengths, they
cannot pass through a piece of paper

3. They have very powerful ionizing effects upon any gas through which they pass

4. They also cause some substances like zinc sulphide to fluorescence.

Beta Radiation

Beta radiation is a stream of electrons, called beta particles (β). When a beta particle is ejected, a
neutron in the nucleus is converted to a proton, so the mass number of the nucleus is unchanged, but
the atomic number increases by one unit. For example:
234 0
90Th→ -1e + 23491Pa

Properties of Beta Particles (β)

1. The electron is the beta particle.

2. These rays are fast-moving stream of electrons

3. They are negatively charged

4. They are therefore deflected to the positive plate in an electric field

5. They have higher penetrating power than alpha particles with their shorter wave-length. They
would not only pass through a piece of paper but also through a piece of aluminium which are
metallic.

6. They have low ionizing effect on gases

7. They cause fluorescence in substances like anthracene but not zinc sulphide

Gamma Radiation

Gamma rays are high-energy photons with a very short wavelength (0.0005 to 0.1 nm). The emission of
gamma radiation results from an energy change within the atomic nucleus. Gamma emission changes
neither the atomic number nor the atomic mass. Alpha and beta emission are often accompanied by
gamma emission, as an excited nucleus drops to a lower and more stable energy state.
Properties of Gamma Rays (ϓ)

1. They are not particles but electromagnetic waves similar to light but are of shorter wavelength
than light rays and have higher frequency.

2. They have neither mass nor charge

3. They are therefore not deflected in an electric field

4. They travel at the speed of light

5. They have very high penetrating power. Out of the three types of radioactive emission, they are
the most penetrating.

6. They have the least ionizing power

Penetration of radiation

Detection of Radiation

 Diffusion Cloud Chamber

The Diffusion Cloud Chamber is used to view high energy alpha particles, lower energy beta particles,
and electrons produced by gamma rays interacting with gas molecules. The Chamber allows for the
viewing of cosmic rays without the need for dry ice or external illumination. Historically the Cloud
Chamber was the first particle detector for making ionizing particles visible. Its working principle is
based on supersaturated vapour. This vapour shows tracks of condensated alcohol while being
penetrated by ionizing particles. Vapour clouds build up along the particle tracks, which are sometimes
thin and long, sometimes thick and round or bulbous. They may appear gradually or pop up all of a
sudden or move very fast like a projectile splintering into all directions.
  Geiger – Muller Counter

The Geiger-Muller counter is a particle detector designed to detect ionizing radiation, such as alpha and
beta as well as gamma radiation (although with significantly lower sensitivity than other types of
detectors). It was invented by the German physicist Hans Geiger (co-discoverer of the atom nucleus) and
later improved by his student Walther Muller, therefore the name Geiger-Muller counter. It is probably
one of the most famous radiation detectors, mostly due to its simplicity and the distinctive audible clicks
produced with the detection of individual particles.

Alternatively, except for its use in the detection of ionizing radiation, the Geiger counter is also used as a
random number generator.

The main element of a Geiger counter is the Geiger-Muller tube, which is basically a chamber filled with
inert gas or a mix of organic vapor and halogens. The tube contains two electrodes, the anode and the
cathode, which are usually coated with graphite. The anode is represented by a wire in the center of the
cylindrical chamber while the cathode forms the lateral area. One end of the cylinder, through which the
radiation enters the chamber, is sealed by a mica window.

As ionizing radiation coming from the surrounding medium passes through the mica window and enters
the Geiger-Muller tube, it ionizes the gas inside, transforming it into positively charged ions and
electrons. The electrons eventually migrate towards the anode of the tube detector, while the positively
charged ions accelerate towards the cathode. As the positive ions move towards the cathode, they
collide with the remaining inert gas thus producing more ions through an avalanche effect. When this
happens an electrical current is established between the two electrodes. The current is amplified and
the resulting flow of electrons can be used to produce sound, light flashes or meter readings.

 Scintillation Counter

Certain minerals, such as zinc sulphide fluoresce or glow when exposed to radiation. The glow is made
up of tiny flashes of light or scintillation and these may be seen under a microscope or counted with
suitable device.

Radioactive Decay

Radioactive decay is the spontaneous radioactive disintegration of an atomic nucleus, resulting in the
release of energy. Some atoms are stable. Others are unstable and ‘decay’, emitting radiation to achieve
a stable state. The emissions from an unstable atom’s nucleus, as it decays, can be in the form of alpha,
beta or gamma radiation.
 When an atom decays, it changes into another isotope, or form, of the same element or into a
completely different element, in a process called transmutation. Different isotopes of the same element
differ in the number of neutrons in their nuclei. Some elements reach stability via a series of steps
through several isotopes, or ‘daughter products’.

One example is uranium-238 (U-238), which, through the process of radioactive decay, will eventually
become a stable isotope of lead. However, this process takes billions of years. Along the way, as the U-
238 isotope’s initial energy declines, it will transmute via a series of elements, each more stable than the
last – thorium, radium, radon, polonium and bismuth – before it stabilizes as lead.

Alpha decay

In alpha decay, a positively-charged particle is emitted from the nucleus of an atom. This alpha particle
consists of two protons and two neutrons (the same structure as a helium-4 nucleus). Although alpha
particles are normally highly energetic, they travel only a few centimeters in air and are stopped by a
sheet of paper or the outer layer of dead skin.

Beta decay
In beta decay, a particle is emitted from the nucleus of an atom. This beta particle is an electron with
either negative or positive electric charge. Beta particles may travel metres in air and several millimetres
into the human body. Most beta particles may be stopped by a small thickness of a light material such as
aluminium or plastic.

Gamma decay

Gamma decay occurs because the nucleus of an atom is at too high an energy state. The nucleus ‘falls
down’ to a lower energy state, emitting a high energy photon known as a gamma particle in the process.
Gamma particles travel in a wave-like pattern at the speed of light. They can only be stopped by a dense
material such as lead, steel, concrete or several metres of water.

Half Life of Radioactive Elements

The half-life of a radioactive element is the time that it takes for one half of the atoms of that substance
to disintegrate into another nuclear form. The decay of an isotope can be measured by its half life.
These can range from mere fractions of a second, to many billions of years.

Element Most Stable Isotope Half-life of Most Stable Isotope

Polonium Po-209 102 years

Astatine At-210 8.1 hours

Radon Rn-222 3.82 days

Radium Ra-226 1600 years

Thorium Th-229 7.54 x 104 years

Uranium U-236 2.34 x 107 years

Protactinium Pa- 234 1.18 minutes

Example Rate of Radioactive Decay Problem

226
88Ra, a common isotope of radium, has a half-life of 1620 years. Knowing this, calculate the first order
rate constant for the decay of radium-226 and the fraction of a sample of this isotope remaining after
100 years.

Solution

The rate of radioactive decay is expressed by the relationship: k = 0.693/t1/2

Where k is the rate and t1/2 is the half-life.

Plugging in the half-life given in the problem: k = 0.693/1620 years = 4.28 x 10-4/year

Radioactive decay is a first order rate reaction, so the expression for the rate is:

log10 X0/X = kt/2.30

Where X0 is the quantity of radioactive substance at zero time (when the counting process starts) and X
is the quantity remaining after time t. k is the first order rate constant, a characteristic of the isotope
that is decaying. Plugging in the values:

log10 X0/X = (4.28 x 10-4/year)/2.30 x 100 years = 0.0186

Taking antilogs: X0/X = 1/1.044 = 0.958 = 95.8% of the isotope remains

Uses of Radioactivity

1. Radioactivity tracers are commonly used in the medical field and also in the study of plants and
animals.

2. Radiation is used and produced in nuclear reactors, which controls fission reactions to produce
energy and new substances from the fission products.

3. Radiation is also used to sterilize medical instruments and food.

4. Radiation is used by test personnel who monitor materials and processes by non-destructive
methods such as x-rays.

Assessment

1. Radiation is detected by the following except

a. diffusion cloud chamber
b. scintillation counter
c. Geiger-Muller counter
d. pulse current

2. ……….. they came up with the term “radioactivity” to describe the spontaneous emissions that
they studied.
a. Becquerel
 b. Pierre and Marie Currie
c. Geiger Muller
d. Marsden

3. Radioactive decay is expressed in terms of

a. rate of radioactive absorption
b. rate of radioactive stability
c. position of element in the periodic table
d. half life

4. The following characteristics belong to one of the basic nuclear particles (i) – Low penetrating (ii)
– Powerful ionizing power on gases (iii) – Particulate
a. Alpha particles
b. Beta particles
c. Gamma rays
d. X-ray

5. Radioisotopes have wide application in

a. medicine
b. industries
c. agriculture
d. all of the above

Answers

1. D

2. B

3. B

4. A

5. D
 WEEK 3
TOPIC: SIMPLE MOLECULES AND THEIR SHAPES
Polar and non-polar molecules

Chemical bonding is the result of either an atom sharing one or more outer orbit electrons with another
atom or an atom taking outer orbit electrons from the atom with which it is bonding. Normally, an atom
has an even distribution of electrons in the orbits or shells, but if more end up on one side that the other
in a molecule, there can be a resulting electrical field in that area.

Water is a polar molecule because of the way the atoms bind in the molecule such that there are excess
electrons on the Oxygen side and a lack or excess of positive charges on the Hydrogen side of the
molecule.

Water is a polar molecule with positive charges

on one side and negative on the other.

Examples of polar molecules of materials that are gases under standard conditions are:

 Ammonia (NH3)

 Sulfur Dioxide (SO2)

 Hydrogen Sulfide (H2S).

Also, Ethanol is polar, since its oxygen molecule draws electrons towards it due to its high electro-
negativity, causing a negative charge around itself.

Non-polar molecules

A non-polar molecule is one that the electrons are distributed more symmetrically and thus does not
have an abundance of charges at the opposite sides. The charges all cancel out each other.
The electrical charges in non-polar Carbon Dioxide are evenly distributed

Examples of non-polar liquids

Most hydrocarbon liquids are non-polar molecules. Examples include:

 Toluene

 Gasoline

Alkynes are non-polar because they cannot be dissolved in water, as do polar molecules. However,
alkynes but do dissolve in other non-polar substances. A rule is that like substances dissolve in like
substances.

Examples of non-polar gases

Common examples of non-polar gases are the noble or inert gases, including:

 Helium (He)

 Neon (Ne)

 Krypton (Kr)

 Xenon (Xe)

Other non-polar gases include:

 Hydrogen (H2)

 Nitrogen (N2)

 Oxygen (O2)

 Carbon Dioxide (CO2)

 Methane (CH4)

 Ethylene (C2H4)

Since Chloroform is more soluble in fats than in water, it is also classified as non-polar.

Rule for Solutions

The rule for determining if a mixture becomes a solution is that polar molecules will mix to form
solutions and non-polar molecules will form solutions, but a polar and non-polar combination will not
form a solution.
Water is a polar molecule and oil is a non-polar molecule. Thus they won’t form a solution. On the other
hand, since alcohol is a polar molecule, it will form a solution with water.

From these examples, we note that three factors determine whether a molecule is polar and non-polar.
These are

1. The polarity of the covalent bond

2. The number and position of lone pairs

3. The shape of the molecule

General Structures of Crystals

The external shape of a crystal is the result of the spatial arrangement of the particles that make up the
crystal. The three-dimensional arrangement of the particles is called the crystal lattice. The crystal lattice
may be considered to be made up of unit cells which are each the smallest portion of the crystal lattice
which shows the complete pattern of the particles in their relative positions

There are three kinds of unit cells based on the cubic structure. These are as follow:

1. Simple cubic: These are structures in which the particles are placed one at each corner of the
cube.

2. Face-centred cubic: These are structure in which there is a particle at each corner and one in the
centre of each face of the cube

3. Body-centred cubic: These are structure in which there is a particle at each corner and one at
the centre of the cube.

Covalent Compounds – Simple Molecules

Form between non-metal atoms. Each bond consists of a shared pair of electrons, and is very strong.
Covalently bonded substances fall into two main types:

1. Simple molecules and

2. Giant covalent structures.

Simple molecules

These contain only a few atoms held together by strong covalent bonds. An example is carbon dioxide
(CO2), the molecules of which contain one atom of carbon bonded with two atoms of oxygen.

Properties of Simple Molecular Substances

 Low melting and boiling points – This is because the weak intermolecular forces break down
easily.

 Non-conductive – Substances with a simple molecular structure do not conduct electricity. This
is because they do not have any free electrons or an overall electric charge.

Hydrogen, ammonia, methane and water are also simple molecules with covalent bonds. All have very
strong bonds between the atoms, but much weaker forces holding the molecules together. When one of
these substances melts or boils, it is these weak ‘intermolecular forces’ that break, not the strong
covalent bonds. Simple molecular substances are gases, liquids or solids with low melting and boiling
points.

Covalent bonding – Giant Covalent structures

Giant covalent structures contain a lot of non-metal atoms, each joined to adjacent atoms by . The
atoms are usually arranged into giant regular lattices – extremely strong structures because of the many
bonds involved. The graphic shows the molecular structure of diamond and graphite: two allotropes of
carbon, and of silica (silicon dioxide).

From left to right – graphite, diamond, silica

Properties of Giant Covalent Structures

  Very high melting points – Substances with giant covalent structures have very high melting
points, because a lot of strong covalent bonds must be broken. Graphite, for example, has a
melting point of more than 3,600ºC.

 Variable conductivity – Diamond does not conduct electricity. Graphite contains free electrons,
so it does conduct electricity. Silicon is semi-conductive – that is, midway between non-
conductive and conductive.

Graphite

Graphite is a form of carbon in which the carbon atoms form layers. These layers can slide over each
other, so graphite is much softer than diamond. It is used in pencils, and as a lubricant. Each carbon
atom in a layer is joined to only three other carbon atoms. Graphite conducts electricity.

Diamond

Diamond is a form of carbon in which each carbon atom is joined to four other carbon atoms, forming a
giant covalent structure. It is the best example of covalent solid, which is usually octahedral in shape. As
a result, diamond is very hard and has a high melting point. It does not conduct electricity. Diamond is
the hardest substance known.

Silica

Silica, which is found in sand, has a similar structure to diamond. It is also hard and has a high melting
point, but contains silicon and oxygen atoms, instead of carbon atoms.

The fact that it is a semi-conductor makes it immensely useful in the electronics industry: most
transistors are made of silica.

Molecular Shape

While Lewis dot structures can tell us how the atoms in molecules are bonded to each other, they don’t
tell us the shape of the molecule. In this section, we’ll discuss the methods for predicting molecular
shape. The most important thing to remember when attempting to predict the shape of a molecule
based on its chemical formula and the basic premises of the VSPER model is that the molecule will
assume the shape that most minimizes electron pair repulsions. In attempting to minimize electron pair
repulsions, two types of electron sets must be considered: electrons can exist in bonding pairs, which
are involved in creating a single or multiple covalent bond, or non-bonding pairs, which are pairs of
electrons that are not involved in a bond, but are localized to a single atom.

The VSPER Model—Determining Molecular Shape

Total number of single bonds, double bonds, and lone pairs on the central Structural pair
Shape
atom geometry

2 Linear

3 Trigonal planar

4 Tetrahedral
5 Trigonal bipyramidal

6 Octahedral

The above table represents a single atom with all of the electrons that would be associated with it as a
result of the bonds it forms with other atoms plus its lone electron pairs. However, since atoms in a
molecule can never be considered alone, the shape of the actual molecule might be different from what
you’d predict based on its structural pair geometry. You use the structural pair geometry to determine
the molecular geometry by following these steps:

1. Draw the Lewis dot structure for the molecule and count the total number of single bonds,
multiple bonds, and unpaired electrons.

2. Determine the structural pair geometry for the molecule by arranging the electron pairs so that
the repulsions are minimized (based on the table).

3. Use the table above to determine the molecular geometry.

For instance, look at methane, which is CH4:

So as you can see, lone pairs have more repulsive force than do shared electron pairs, and thus they
force the shared pairs to squeeze more closely together.

As a final note, you may remember that we mentioned before that only elements with a principal
energy level of 3 or higher can expand their valence and violate the octet rule. This is because d
electrons are necessary to make possible bonding to a fifth or sixth atom. In XeF4, there are two lone
pairs and four shared pairs surrounding Xe, and two possible arrangements exist:
In the axial arrangement, shared pairs are situated “top and bottom.” In the equatorial arrangement,
shared pairs surround Xe. The equatorial arrangement is more stable since the lone pairs are 180˚ apart
and this minimizes their repulsion. In both molecular arrangements, the electronic geometry is
octahedral, with 90˚ angles. The top figure has a molecular geometry known as “seesaw,” while the
bottom figure has a molecular geometry that is more stable, known as square planar.

Example

Draw the dot formula for SeF4 and determine the hybridization at Se.

Explanation

First determine the number of valence electrons this molecule has: SeF4 has 6 + 4(7) = 34 valence
electrons, which is equal to 17 pairs of electrons.

Selenium is surrounded by four fluorines and a lone pair of electrons. That’s five sites of electron
density, which translates into sp3d hybridization. Se is from the fourth period, so it may have an
expanded octet.

So, to recap, focus on the number of binding “sites” or areas of concentrated electron density:

Two areas of electron density: linear, planar molecule

Three areas of electron density: trigonal planar molecule

Four areas of electron density: tetrahedral molecule

Five areas of electron density: trigonal bipyramidal molecule

Six areas of electron density: octahedral molecule

Molecular Polarity

In chemical bonds, polarity refers to an uneven distribution of electron pairs between the two bonded
atoms—in this case, one of the atoms is slightly more negative than the other. But molecules can be
polar too, and when they are polar, they are called dipoles. Dipoles are molecules that have a slightly
positive charge on one end and a slightly negative charge on the other. Look at the water molecule. The
two lone electron pairs on the oxygen atom establish a negative pole on this bent molecule, while the
bound hydrogen atoms constitute a positive pole. In fact, this polarity of water accounts for most of
water’s unique physical properties. However, molecules can also contain polar bonds and not be polar.
Carbon dioxide is a perfect example. Both of the C—O bonds in carbon dioxide are polar, but they’re
oriented such that they cancel each other out, and the molecule itself is not polar.

Assessment
 1. These are factors which determine whether a molecule is polar or non-polar except
a. shape of the molecule
b. number of lone pairs in an atom
c. number of electronegative elements present
d. polarity of covalent bond

2. These are examples of unit cell cubic structure

a. simple cubic structure
b. non-linear cubic structure
c. face-centered cubic structure
d. body centered cubic structure

3. These are intermolecular bonds or attractions except

a. metallic bond
b. dipole – dipole attractions
c. van-der Waal’s forces
d. hydrogen bond

4. List the properties of simple molecular substances

Answers

1. C

2. B

3. A
 WEEK 4
TOPIC: PETROLEUM OR CRUDE OIL
Crude Oil

Crude oil or Petroleum is the chief source of hydrocarbons. Petroleum which means rock oil in Latin
occurs as a dark, sticky, viscous liquid. It is found in huge underground deposits in many parts of the
world. Natural gas is usually found together with it. Petroleum is a mixture of gaseous liquid and solid
alkanes, alkenes, cycloalkanes, aromatic hydrocarbons and others. Natural gas consists mainly of
methane. Crude oil is a mixture of hydrocarbons. It exists in liquid phase in natural underground
reservoirs and remains liquid at atmospheric pressure after passing through surface separating facilities.

Crude oil occurs in large quantities in Nigeria, especially in Bayelsa, Edo, Imo, rivers, Delta, Abia, Ondo
and Cross river state. It is dark brown in colour though its composition and consistency vary from place
to place. In fact, different oil producing areas yield significantly different varieties of crude oil. We have
light and heavy crude oil. The light one has low metal and sulphur content, light in colour and flows
easily. It is very expensive. The heavy one has high metal and sulphur content and must be heated to
become fluid. It is less expensivePetroleum is a naturally occurring, yellow-to-black liquid found
in geological formations beneath the Earth’s surface, which is commonly refined into various types
of fuels. Components of petroleum are separated using a technique called fractional distillation.

It consists of hydrocarbons of various molecular weights and other organic compounds

Origin of Crude Oil and Natural Gas

They are formed from the remains of marine algae and animals. When these tiny aquatic organisms
died, their remains gradually settled on the seabeds. Over the years, the remains became covered by
mud, silt and other sediments. As the sediments piled up, their mass exerted a great pressure on the
lower layers, changing them to hard sedimentary rocks. During this process, bacterial activity, heat and
pressure probably changed the plant and animal remains into crude oil and natural gas.

Refining of Crude Oil

Petroleum or crude oil occurs naturally. it contains many useful products also called fractions. These are
separated by the method of fractional distillation. This process of obtaining useful fractions from
petroleum is called refining.

The process of dividing petroleum into fractions with different boiling range volatilities and free from
impurities is called refining.
The process of turning petroleum into a useful form is done in a crude oil refinery. The steps for making
crude oil into oil, petrol or whatsoever are fractional distillation, cracking and reforming.

Petroleum is refined by fractional distillation. The process of separating a mixture into a series of
fractions of different volatilities by means of distillation is known as fractional distillation.
In the process of fractional distillation, a mixture of different liquids is evaporated followed by
condensation. Different liquids are evaporated according to their boiling point and they are collected in
different chambers of distillation tower
Fractional Distillation of Crude Oil

Fractional distillation differs from distillation only in that it separates a mixture into a number of
different parts, called fractions. A tall column is fitted above the mixture, with several condensers
coming off at different heights. The column is hot at the bottom and cool at the top. Substances with
high boiling points condense at the bottom and substances with low boiling points condense at the top.
Like distillation, fractional distillation works because the different substances in the mixture have
different boiling points.

Fractions of petroleum from refining are petroleum gases (Methane, butane, etc.), petrol or gasoline,
kerosene, diesel, lubricating oil and bitumen (asphalt). Because they have different boiling points, the
substances in crude oil can be separated using fractional distillation. The crude oil is evaporated and its
vapours allowed to condense at different temperatures in the fractionating column. Each fraction
contains hydrocarbon molecules with a similar number of carbon atoms.

The mixture is inserted at the bottom, where mostly everything will condense as the temperature is
350°C and more. The condensed crude oil will rise to the next fraction above, which has a very high
temperature as well, but a slightly smaller one. Only the part of the mixture, which boiling point is under
the temperature of the fraction, will condense and rise to the next fraction. The part of the mixture,
whose boiling point is higher than the temperature inside the distillation fraction, will stay there and be
pumped out.

Crude oil is heated until it boils and then the hydrocarbon gases are entered into the bottom of the
fractionating column. As the gases go up the column the temperature decreases.
The hydrocarbon gases condense back into liquids and the fractions are removed from the sides of the
column. The different fractions have different uses.The smaller the hydrocarbon molecule, the further it
rises up the column before condensing.
The fractionating column operates continuously. The temperatures shown are approximate. A sample of
crude oil may be separated in the laboratory by fractional distillation. The collection vessel is changed as
the temperature rises to collect the different fractions.

Cracking

After the fractional distillation process, the separated mixtures have to be cracked down. This means
that a long molecule will be split up in smaller parts.
Firstly, single bonds will be broken down. This results into some lone electrons.
The lone electrons form double bonds. Thus, hydrogen will disconnect from the carbon atom. Hydrogen
(H2) remains as a side product. The loss of hydrogen in these smaller organic molecules is logical,
because when they are lost, more lone electrons remain with what the previous lone electrons can
make a bond.

Reforming

After cracking, the molecules are ready to undergo the reforming process.
This is given by the octane number. The octane number is very important in petrol. It tells what the
percentage of pure heptane (in the earlier days it was octane – that is why it is called octane number) in
petrol is. This is of great importance for the chemical behaviour. The quality of petrol is improved by
adding mixtures to pure heptane. The chains of heptane are heated up (where platinum is used as a
catalyst). So they can change. After the heating process, it shows a higher amount of branched chains.
This increases the octane number.

Example: When the petrol you buy has a 98 in the name, then it means that 98% are branched chains
and 2% of the mixture is pure heptane (or another pure molecule).

Octane Number
The octane number or octane rating of petrol is a measure of the proportion of branched chain
hydrocarbons in a given blend of gasoline (petrol).

In other words, Octane number is a standard which determines the knocking ability and quality of
gasoline. Higher is the octane number of a gasoline, lower is the knocking it produces.

Gasoline is composed of C7 – C9 hydrocarbons i.e. heptanes, octane and nonane. These hydrocarbons
are present in their straight chain or branched chain isomers. It has been shown that straight chain
hydrocarbons burn too rapidly in the car engine thus, causing irregular motion of the pistons which
results in rattling noise. This rattling noise is known as “Knocking”.

Knocking

Knocking is a sharp metallic sound produced in the internal combustion engine. Knocking is caused by
the low octane number of gasoline

Synthetic Petrol

Synthetic petrol is made from materials such as coal, coke and hydrogen which do not occur in crude
petroleum. Synthetic petrol can be gotten from two sources

1. From coal: When powdered coal is heated with hydrogen in the presence of iron or tin as
catalyst at 500oC and 200 atmospheric pressure, it is converted into an oily mixture of
hydrocarbons. The mixture is separated by distillation into a petrol fraction boiling at 200oC and
a heavy oily residue which can be further treated with fresh coal to obtain more petrol

2. From Coke: When steam is passed over heated coke at 1000 oC, a mixture containing equal
volumes of carbon (II) oxide and hydrogen known as water gas is obtained

1. C(s) + H2O(l) —–> CO(g) + H2(g)

The water gas can be hydrogenated to a mixture of hydrocarbons by adding hydrogen and passing it
over finely divided nickel as catalyst at 200oC. About half of the product is petrol, the less volatile
fraction is used as fuel for diesel engines

Fractions of Crude Oil

1. Petroleum gases – These gases are liberated at the top of the column at temperatures below
40o They are mainly hydrocarbons consisting of one of four carbon atoms per molecule, e.g
methane, ethane, propane, butane and ethane. They are used mainly as fuels for homes and
industries and for manufacture of products like hydrogen, sulphide and ethyne

2. Petroleum ethers and ligroin – This is the mixture of very volatile liquids consisting of
hydrocarbons with five or seven carbon atoms per molecule. The petroleum ether or light
petroleum distils over between 2oC to 60oC, while the ligroin or light naphtha distils over
between 60oC to 100o

3. Petrol – Petrol or gasoline is a mixture of hydrocarbons consisting of four to twelve carbon

atoms per molecule. It is a volatile liquid which distils over in the temperature range of 40 oC to
200o Petrol is used as a fuel for aeroplanes and motor vehicles. It is a good solvent for paints and
grease
 4. Kerosene – Kerosene contains hydrocarbons with twelve to eighteen carbon atoms per
molecule. It distils over in the temperature of 200oC to 250o it is a fairly volatile liquid and is
used as a fuel for lighting and heating and for driving tractors and some modern jet engines. It is
a good solvent for paint and grease

5. Gas oil and Diesel oil – They have twenty to twenty five carbon atoms per molecule. They distil
over in the temperature of 250oC to 350o They are mainly used as fuel for heating and diesel
engines and as raw materials of the cracking process

6. Lubricating oils – They are heavy oils. They consist of hydrocarbons with more than twenty five
carbon atoms per molecule. They distil over in the temperature of 350oC to 500o They are
viscous liquids used for lubricating moving parts of engines and machines. Grease belongs to this
group. Vaseline and paraffin wax are solid long chain hydrocarbons which are obtained from re-
distillation of heavy oils.

7. Bitumen and other residues – these are solid residues left behind at temperatures above
500o Bitumen are used for surfacing roads and airfields while other residues may be used as
fuel, in protective paints and water proofing roofs.

Assessment

1. ……… is the chief source of hydrocarbons

2. ………. of petrol is a measure of the proportion of branched chain hydrocarbons in a given blend
of gasoline (petrol).

3. ……….. involves the separation of a mixture into a number of different parts, called fractions.

4. ……….. is a sharp metallic sound produced in the internal combustion engine

5. Mention 4 fractions of crude oil

Answers

1. Crude oil/Petroleum

2. Octane Number

3. Fractional Distillation

4. Knocking

5. Kerosene, Petrol, Bitumen, Diesel, Gas oil, Lubricating oil

 WEEK 5
TOPIC: VOLUMETRIC ANALYSIS
Introduction

Volumetric analysis is an analytical method or procedure for working out the titre or concentration of an
analyte in a solution. This is done by measuring the volume of a standard solution of an appropriate
reagent whose precise concentration is already known.

Preparing A Standard Solution

1. A standard solution is a solution in which its concentration is known.

2. The steps taken in preparing a standard solution are:

 Determine the volume and concentration that you want to prepare.

 Calculate the mass of solute needed to give the required volume and concentration.

 Weigh the solute

 Dissolve the solute completely dissolved in distilled water and then transfer it to a volumetric
flask partially filled with distilled water.

 Add distilled water to the calibration mark of the volumetric flask.

 Invert the flask and shake it to make sure thorough mixing.

Acid Base Titration

Some materials used during acid – base titration and precautions in using some of them

1. weighing balance

2. chemical balance

3. pipette

4. burette

5. retort stand

6. filter paper

7. funnel

8. white tile

9. standard volumetric flask

10. conical flask

Pipette

 rinse the pipette with the solution it will be measured with e.g. base

 avoid air bubbles in the pipette

 make sure the mark to be read is at same level with your eye

 do not blow the last drop on the burette

Burette

 rinse the burette with acid or allow it to drain after rinsing

 make sure the burette jar is filled

 avoid air bubbles in the burette

 make sure that burette is not leaking4remove the funnel befor taking your reading

 avoid inconsistent burette reading

Conical flask

 do not rinse with any of the solutions used in the titration but with distilled water

 wash down with distilled water any drop of the solution that stick by the sides of the conical
flask

Concentration of a Solution
 1. The concentration of a solution tells you how much solute is dissolved in 1 unit volume of
solution.

2. The volume of a solution is measured in dm³ (litres) 1 dm³ = 1000 cm³.

3. The amount of solute can be measured in grams or moles.

4. 2 units of concentration used in chemistry are g dm-3 and mol dm-3

Concentration in g dm-3

1. Concentration is the number of moles of solute per liter of solution.

2. A concentration of 10 g dm-3 means there is 10 g of solute dissolved in1 dm3 of solution.

Concentration=Mass of solute (g) / Volume of solution (dm3)

Example 1:
Calculate the concentration of the solution if 28g of NaOH is dissolve in 250cm 3 of water.
Answer:
Mass of solute = 28g
Volume of solvent = 250cm³ = 0.25dm³

Concentration = Mass / Volume

=28g / 0.25dm3

=112g/dm3

Concentration in mol dm-3 (Molarity)

1. Molarity is probably the most commonly used unit of concentration. It is the number of moles of
solute per liter of solution.

2. A concentration of 2 mol dm-3 means there are 2 moles of solute dissolved in 1 dm3 of solution.

Molarity=Mole of solute(mol) / Volume of solution(dm3)

Example 2:
What is the molarity of a solution made when water is added to 0.2 mol of CaCl 2 to make 100 cm³ of
solution? [RAM: Ca = 40; Cl = 35.5]

Answer:
Number of mole of solute = 0.2 mol
Volume of solvent = 100 cm³ = 0.1 dm³

Molarity = Number of Mole / Volume

= 0.2mol / 0.1dm3

=2mol/dm3

Conversion of Concentration Unit

 1. The chart above shows how to convert the units of concentration from g dm-3 to mol dm-3 and
vice versa.

2. The molar mass of the solute is equal to the relative molecular mass of the solute.

Example 3:
The concentration of a Potassium chloride solution is 14.9 g dm -3. What is the molarity ( mol dm-3) of the
solution? [Relative Atomic Mass: Cl = 35.5; K = 39]

Answer:
Relative Formula Mass of Potassium Chloride (KCl)
= 39 + 35.5 = 74.5

Molar Mass of Potassium Chloride = 74.5 g/mol

Molarity of Potassium Chloride

Molarity = ConcentrationMolar Mass=14.9gdm−374.5gmol−1=0.2mol/dm3

Molarity and Number of Moles

Number of mole of solute in a solution can be calculated by using the following formula

n=MV / 1000

where
n = number of mole of solute
M = molarity of the solution
V = volume of the solution in cm3

Example 4
How many moles of zinc sulphate is present in 200cm3 of 0.1 mol dm-3 zinc sulphate solution?
Answer:
Molarity, M = 0.1 mol dm-3
Voloume, V = 200cm3

n=MV / 1000
n= (0.1)(200) / 1000 = 0.02mol

Example 5
A solution of barium hydroxide have molarity 0.1 mol dm-3. What is the concentration of the solution in
g dm-3? [Relative Atomic Mass: Ba = 137; O = 16; H = 1]
Answer:
Relative Formula Mass of barium hydrokxide, Ba(OH)2
= 137 + 2(16+1) = 171

Molar Mass of Potassium Chloride = 171 g/mol

Concentration = Molarity × Molar Mass

=0.1mol/dm3 × 171gmol−1 = 17.1gmol−1

Solubility of substance

The term solubility refers to the maximum amount of material that will dissolve in a given amount of
solvent at a given temperature to produce a stable solution.

Calculations on Solubility

Examples 1: If 12.2 g of lead(II) trioxonitrate(V) were dissolved in 21 g of distilled water at 20 oC,

calculate the solubility of the solute in mol dm-3.

Solution:

Molar mass of Pb(NO3)2 = 331 g

12.2 g of Pb(NO3)2 = 12.2/331 = 0.037 mole

21 g of water at 20oC dissolved 0.037 mole of Pb(NO3)2

1000 cm3 of water at 20 oC dissolved 1000 x 0.037 / 21

= 1.76 moles of Pb(NO3)2

At 20oC, the solubility of Pb(NO3)2 in water is 1.76 mol dm-3

Using formula Method

Solubility = mass/molar mass x 1000/vol

= 12.2/331 x 1000/21

= 1.7751 mol dm-3

= 1.76 mol dm-3

Example 2: The solubility of potassium trioxonitrate (V) is exactly 1800 g per 1000 g water at 83oC and
700 g per 1000 g water at 40oC. Calculate the mass of potassium trioxonitrate (V) that will crystallize out
of solution if 155 g of the saturated solution at 83 oC is cooled to 40oC.

Solution:
 Solvents + Solution = Solution

Masses involved at 83oC 1000 g 1800 g = 2800 g

Masses involved at 40oC 1000 g 700 g = 1700 g

Solute deposited on cooling from 83oC to 40oC = 2800 – 1700 = 1100 g

On cooling from 83oC to 40oC,

2800 g of saturated solution deposit 1100 g of solute

155 g of saturated solution will deposit = 1100 x 155 / 2800

= 60.89 g of solute

A titration calculation

The method for titration calculations is the one you have used for mole calculations already:

1. Write a balanced equation for the reaction.

2. Find the number of moles of the known substance.

3. Use the balanced equation to find the number of moles of the unknown substance.

4. Work out the mass, concentration or volume of the unknown.

The formula:

MAVA/MBVB = NA/NB

MA = Molarity of acid in mol/dm3

VA = Volume of the acid in cm3

MB = Molarity of base in mol/dm3

VB = Volume of the base in cm3

NA = Number of moles of acid

NB = Number of moles of base

Example 1: Calculate:

(a) the mass of anhydrous Na2CO3 present in 300cm3 of 0.1M

(b) the number of Na2CO3 particles present in the solution (Na = 23, C = 12, O = 16)

Solution:

(a) Molarity of the Na2CO3 solution = 0.1 M

Molar mass of Na2CO3 = 23 x 2 + 12 + 16 x 3 = 106g/mol

Concentration (g/dm3) = Molarity x molar mass

0.1 x 106

= 10.6g/dm3

This means 1000cm3 of 0.1M solution contain 10.6g of Na2CO3

300cm3 of 0.1 M solution will contain: 300 x 10.6 / 1000

= 3.18g of Na2CO3

(b) Number of Na2CO3 particles = molarity x 6.02 x 1023

0.1 x 6.02 x 1023

= 0.602 x 1023

This means 1000cm3 of 0.1 M solution contain 0.602 x 1023 Na2CO3 particles

300cm3 of 0.1M solution contain

Example 2: 20.30cm3 of hydrochloric acid solution was titrated against 25cm3 of 0.1M sodium hydroxide
solution. Calculate:
(i)
the concentration of the acid in mol/dm3

(ii) the concentration of the acid in g/dm3

Solution:

Equation of reaction

HCl (aq) + NaOH (aq) ——–> NaCl (aq) + H2O (l)

1 : 1

(i) MA =?, MB = 0.1 M, VA = 20.30cm3, VB = 25cm3, NA = 1, NB = 1

MAVA/MBVB = NA/NB

MA = MBVB NA / VA NB

MA = 0.1 x 25 x 1 / 20.30 x 1

MA = 0.123 mol/dm3

(ii) Molarmass of Acid = 36.5 g/mol, Concentration of Acid in mol/dm3 = 0.123 mol/dm3

Concentration of Acid in g/dm3 = Concentration of Acid in mol/dm3 x Molarmass

= 0.123 x 36.5

= 4.49 g/dm3
Assessment

1. Calculate the concentration of the solution if 25g of NaOH is dissolve in 500cm3 of water.

2. 32.30cm3 of hydrochloric acid solution was titrated against 25cm3 of 0.2M sodium hydroxide
solution. Calculate: a. the concentration of the acid in mol/dm3 b. the concentration of the acid
in g/dm3

3. If 21.2 g of Calcium Trioxocarbonate were dissolved in 15 g of distilled water at 30oC, calculate

the solubility of the solute in mol dm-3.

4. The concentration of a Potassium chloride solution is 28.9 g dm-3. What is the molarity ( mol dm-
3
) of the solution? [Relative Atomic Mass: Cl = 35.5; K = 39]
 WEEK 6
TOPIC: QUALITATIVE ANALYSIS

Identification of Ions

There are 10 cations and 4 anions to be studied:

Cation

Sodium Na+ Iron (II) Fe2+

Calcium Ca2+ Iron (III) Fe3+

Magnesium Mg2+ Lead(II) Pb2+

Aluminium Al3+ Copper (II) Cu2+

Zinc Zn2+ Ammonium NH4+

Anion

Chloride ion Cl–

sulphate ion SO42-

nitrate ion NO3–

carbonate ion CO3–

Steps in qualitative analysis

Colour of Ions

Salt or metal oxide Solid Aqueous solution

Salt of Sodium, Calcium, Magnesium, Aluminium, zinc, Lead,
White Colourless
ammonium

Salt of Chloride, sulphate, nitrate, carbonate White Colourless

Salt of Copper(II) –

Copper(II) Carbonate Green –

Copper(II) sulphate, Copper(II) nitrate, Copper(II) chloride Blue Blue

Copper(II) oxide Black –

Salt of Iron (II):

Green Green
Iron(II) sulphate; Iron(II) nitrate; Iron(ID chloride

Salt of Iron (III):

Brown Brown
Iron(III) sulphate; Iron(III) nitrate; Iron(III) chloride

Yellow when it is
Zink oxide hot and white when –
it is cold.

Brown when it is
Lead(II) oxide- hot and yellow –
when it is cold.

Magnesium oxide, Aluminium oxide White Insoluble

Potassium oxide, Sodium oxide, Calcium oxide White Colourless

Heating Effect on Carbonate Salts

1. All carbonates salts except potassium carbonate and sodium carbonate can be decomposed by
heat to produce carbon dioxide gas.

2. Table below shows the effect of heating on metal carbonate.

Carbonate Salt Equation of The Reaction

Potassium carbonate
Will not decompose by heat
Sodium carbonate

Calcium carbonate
Magnesium carbonate
Aluminium carbonate Calcium carbonate
Zinc carbonate
Iron (III) carbonate
Lead(II) carbonate
Copper(II) carbonate

CaCO3 → CaO + CO2

Aluminium carbonate

Al2(CO3)3 → Al2O3 + 3CO2

Copper carbonate

CuCO3 → CuO + CO2Mercury(II) carbonate

Silver carbonate
Aurum(II) carbonateSilver carbonate

2Ag2CO3 → 2Ag + 2CO2 + O2Ammonium carbonate(NH4)2CO3 → 2NH3 + 2CO2 + H2O

Heating Effect on Nitrate Salts

1. All nitrates salts decompose when heated.

2. Table below shows the products formed when different nitrate salts are heated.

Nitrate Salt Equation of The Reaction

Ammonium nitrate decompose to nitrogen monoxide and water vapour

Ammonium nitrate
when heated.

NH4NO3 → N2O + 2H2OPotassium nitrate

Sodium nitratePotassium nitrate

2KNO3 → 2KNO2 + O2

Sodium nitrate

2NaNO3 → 2NaNO2 + O2Calcium nitrate

Magnesium nitrate
Aluminium nitrate
Zink nitrate
Iron (III) nitrate
Lead(II) nitrate
Copper(II) nitrateMagnesium nitrate

2Mg(NO3)2 → 2MgO + 4NO2 + O2

Iron(III) nitrate

4Fe(NO3)3 → 2Fe2O3 + 12NO2 + 3O2

Lead(II) nitrate
2Pb(NO3)2 → 2PbO + 4NO2 + O2Mercury(II) nitrate
Silver(I) nitrate
Aurum(II) nitrateSilver nitrate

2AgNO3 → 2Ag + 2NO2 + O2

Heating Effect on Sulphate Salts

1. Most sulphate salts do not decompose by heat. For instance, sodium sulphate, potassium
sulphate, and calcium sulphate are not decomposable by heat.

2. Only certain sulphate salts are decomposed by heat when heated strongly.

For instance:

Strong heating of green crystal iron (II) sulphate will release steam, sulphur dioxide, sulphur trioxide and
leave behind a reddish solid iron (III) oxide residue. The steam released comes from the hydrated water
of the crystallize salt.

2FeSO4.7H2O → Fe2O3(p) + SO2(g) + SO3(g) + 14H2O(g)

Meanwhile, zinc sulphate, copper (II) sulphate, and iron (III) sulphate decompose when heated strongly
to evolve sulphur (VI)oxide gas and form a metal oxide.
Example
Zinc sulphate

ZnSO4 → ZnO + SO3

Copper (II) sulphate

CuSO4 → CuO + SO3

Iron (III) sulphate

Fe2(SO4)3→ Fe2O3 + 3SO3

When ammonium sulphate is heated strongly, this white solid sublimate and is decomposed to form
ammonia gas and sulphuric acid.

(NH4)2SO4 → NH3 + H2SO4

Identifying Gases

 Oxygen (Glowing wooden splinter)

A glowing wooden splinter is inserted into the test tube that contain the gas.

The gas rekindles the glowing wooden splinter.

 Hydrogen (Lighted Wooden Splinter)

A lighter wooden splinter is brought close to the mouth of the test tube that contain the gas.

A “pop” sound is produced.

  Carbon dioxide (Lime Water)

The gas is directed to flow through lime water.

The lime water turn chalky.

 Sulphur Dioxide (Potassium Dichromate(VI))

The gas is directed to flow through potassium dichromate(VI) solution.

The orange colour of potassium dichromate(VI) solution become green.

 Chlorine (Moist Litmus Paper)

Moist blue litmus paper is inserted into the test tube that contain the gas.

The blue litmus paper turn red and then white.

 Ammonia Gas (Moist litmus paper)

Moist red litmus paper is inserted into the test tube that contain the gas.

The red litmus paper turn blue.

Identifying Anions – Carbonate

The carbonate salts react with acid produce carbon dioxide and water. This chemical property of
carbonate is used to test the presence of carbonate in a salt.

CO32– + 2H+ → H2O + CO2

 During the test, some dilute hydrochloric acid / nitric acid /sulphuric acid is added to
the carbonate salt.

 If the salt contain carbonate, effervescence occurs.

 If the gas given off is passed through lime water, the lime water will turns chalky.

 This indicates that the gas is carbon dioxide, and hence the salt contain carbonate.

Identifying Anions – Sulphate

All salts of sulphate are soluble in water, except lead(II) sulphate, barium sulphate and calcium sulphate.

The insolubility of barium sulphate is used to test for the presence of sulphate in a salt.

1. 2 cm3 of dilute hydrochloric / nitric acid is added to 2 cm3 of sulphate solution. This is to check
whether carbonate is presence in the solution or not because carbonate may give the same
result as sulphate.

2. If not effervescence, then 2 cm3 of barium chloride is added into the solution.

3. If sulphate ions are presence, a white precipitate will form. The precipitate is barium sulphate.

4. This is actually the double decomposition reaction that you have learned in preparation of
insoluble salt.

Ba2+ + SO42– → BaSO4

Identifying Anions – Chloride

All the salts of chloride are soluble in water except lead(II) chloride, silver(I) chloride and mercury
chloride.The insolubility if silver(I) chloride is used in the test of presence of chloride.

1. 2 cm³ of dilute nitric acid is added to 2 cm3 solution of chloride ions. This is the check if
carbonate ions are presence because carbonate ions may give the same result.

2. If there is no effervescence, 2 cm³ of silver nitrate solution is then added into the mixture.

3. A white precipitate will form if chloride is presence in the salt.

4. The precipitate is silver chloride

Ag+ + Cl– → AgCl

Identifying Anions – Nitrate

Test 1

Add dilute sodium hydroxide and a little aluminium powder. If a nitrate is present, ammonia gas is
produced. The gas can be identified as it turns moist red litmus paper blue.

Test 2

1. About 2cm3 of dilute sulphuric acid is added into the solution that wants to be tested and then
followed by 2cm3 iron (II) sulphate solution.

2. A few drops of concentrated sulphuric acid are carefully drop through the inclined side of the
test tube without shaking the test tube.

3. A brown ring will form in the middle of the solution.

4. Explanation: Iron (II) sulphate reduce nitric acid (from the reaction between nitrate ion and
concentrated sulphuric acid) to nitrogen monoxide. Afterwards, nitrogen monoxide combines
with iron (II) sulphate to form the compound FeSO4.NO which is brown in colour (brown ring).
Identifying Cations – Test with Sodium Hydroxide and Ammonia Solution

 Cations can be identified by their reaction with aqueous sodium hydroxide and aqueous
ammonia.

 Sodium hydroxide and aqueous ammonia produce hydroxide ion which will react with most
anion to form precipitate.

NaOH + H2O → Na+ + 2OH– + H+

NH3 + H2O → NH4+ + OH–

 Different cations like aluminium Al3+, calcium Ca2+, copper(II) Cu2+, iron(II) Fe2+, iron(III) Fe3+,
lead(II) Pb2+, zinc Zn2+ produce different coloured precipitates, which may or may not dissolve in
excess alkali.

 Zn(OH)2, Al(OH)3 and Pb(OH)3 dissolve in excess NaOH solution, this is because Zn(OH)2,
Al(OH)3 and Pb(OH)3 are amphoteric, they can react with NaOH to form salt and water.
Zn(OH)2 + 2NaOH → Na2ZnO2 + 2H2O
Al2(OH)3 + 3NaOH → Na3Al2O3 + 3H2O
Pb(OH)2 + 2NaOH → Na2PbO2 + 2H2O

 Table below shows the summary of the precipitate form by different cation.

NaOH(ak) NH3(ak)

Na+ – –

Ca2+ White precipitate. –

Mg2+ White precipitate. White precipitate.

Al3+ White precipitate.

Dissolve in excess NaOH solution.White precipitate.Zn2+White precipitate.

Dissolve in excess NaOH solution.White precipitate.

Dissolve in excess NH3 solution.Pb2+White precipitate.

Dissolve in excess NaOH solution.White precipitate.Fe2+Dirty green precipitate.Dirty green

precipitate.Fe3+Red brown precipitate.Red brown precipitate.Cu2+Blue precipitate.Blue precipitate.

Dissolve in excess NH3 solution and form a blue solution.NH4+—

Identifying Cations – Test with Chloride Ions

1. Out of the 10 cations, only lead(II) ions will form a precipitate with chloride ions.

2. This is because lead(II) chloride is insoluble in water.

3. The chemical reaction is a double decomposition reaction.

Pb2+ + 2Cl– → PbCl2

 1. Lead(II) chloride will dissolve in hot water.

HCl or NaCl

Na+ –

Ca2+ –

Mg2+ –

Al3+ .-

Zn2+ –

White precipitate.
Pb2+
Dissolve in hot water

Fe2+ –

Fe3+ –

Cu2+ –

NH4+ –

Identifying Cations – Test with Sulphate Ions

1. Out of the 10 cations, only calcium ions and lead(II) ions will form a precipitate with sulphate
ions.

2. This is because both calcium sulphate and lead(II) sulphate are insoluble in water.

3. The chemical reaction is a double decomposition reaction.

Pb2+ + SO42- → PbSO4

Ca2+ + SO42- → CaSO4

H2SO4 or Na2SO4

Na+ –

Ca2+ White precipitate.

Mg2+ –

Al3+ –

Zn2+ –

Pb2+ White precipitate.

Fe2+ –
Fe3+ –

Cu2+ –

NH4+ –

Identifying Cations – Test with Carbonate Ions

1. All ions, except sodium ions and ammonium ions will form precipitate with carbonate.

2. This is because sodium carbonate and ammonium carbonate are soluble in water.

Na2CO3

Na+ –

Ca2+ White precipitate.

Mg2+ White precipitate.

Al3+ White precipitate.

Zn2+ White precipitate.

Pb2+ White precipitate.

Fe2+ Green precipitate.

Fe3+ Brown precipitate.

Cu2+ Blue precipitate.

NH4+ –

Identifying Cations – Test with Iodide Ions

1. Iodide ions will form precipitate with lead(II) ions and copper(II) ions.

2. You only need to know the reaction between lead(II) ions and iodide ions.

3. The yellow precipitate formed will dissolve in hot water.

Pb2+ + 2I– → PbI2

KI

Na+ –

Ca2+ –

Mg2+ –

Al3+ –
Zn2+ –

Pb2+ Yellow precipitate. Dissolve in hot water

Fe2+ –

Fe3+ A red brown solution formed.

Cu2+ White precipitate form in brown solution

NH4+ –

Identifying Cations – Tests to Distinguish Iron(II) and Iron(III) ions.

Some Tests to Distinguish Fe2+ ion From Fe3+ Ion

1. The presence of Fe2+ ion and Fe3+ ion in a salt can be confirmed by using solution of
potassium hexacyanoferrate (II), solution of potassium hexacyanoferrate (III) and potassium
thiocyanate.

2. Table below shows the observation of the tests.

Reagent Observation Ion presents

Solution of potassium Light blue precipitate Fe2+

hexacyanoferrate (II) Dark Blue precipitate Fe3+

Solution of potassium Dark blue precipitate Fe2+

hexacyanoferrate (III) Greenish brown solution Fe3+

Pinkish solution Fe2+

Potassium thiocyanate
Blood red solution Fe3+

Assessment

1. How can the following gas be tested for

a. Oxygen
b. Hydrogen
c. Chlorine
d. Hydrogen Sulphide

2. What are the heating effects of the following

a. Carbonate salts
b. Nitrate salts