CHEMISTRY FORM FOUR STUDY NOTES TOPIC 1: NON METALS AND THEIR COMPOUNDS & TOPIC 2: ORGANIC CHEMISTRY

August 14, 2017

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TOPIC 1: NON METALS AND THEIR COMPOUNDS

General Chemical Properties of Non Metals

The non-metals are very reactive and most of them react with other elements to form different compounds.

The
following are important chemical properties of non-metals which are
connected with their tendency towards electron gain in the course of
formation of compounds:

1.
The oxide of a non-metal is either acidic or neutral but never basic.
The oxide of a non-metal is a covalent compound. Being acidic, it
combines with water to form an acid, e.g.

2.
A non-metal never replaces hydrogen in an acid to form a salt. This is
because replacement of hydrogen in an acid is due to the fact that H⁺ accepts electrons supplied by a metallic atom.

A non-metal is an electron acceptor and so cannot release electrons to hydrogen ions in solution.

3. Non-metals form covalent chlorides, for example, the behaviour of phosphorus forming its chlorides is well known.

A covalent chloride like this is usually a volatile liquid, a non electrolyte, and rapidly hydrolysed by water.

These properties are characteristic of non metallic chlorides (except CCl₄ which is not hydrolysed by water).

4.
Non-metals combine with hydrogen to form many hydrides. A covalent
compound is formed by equal sharing of electrons between or among the
combining atoms. For example, methane ammonia, hydrogen chloride and
hydrogen sulphide are the covalent compounds.

5.
Non-metals are oxidizing agents As discussed early, non-metals accept
electrons from other substances. Therefore, they are called oxidizing
agents because, upon accepting the electrons, the substances donating
these electrons are oxidized. So they act as the agents for oxidation of
other substances.

The Oxidizing Properties of Non-metals

Explain the oxidizing properties of non-metals

Non-metals
react by gaining electrons to become negative ions. A univalent
non-metal accepts one electron while a divalent one accepts two
electrons. The ion formed carries the corresponding number of negative
charges, but they rarely exceed two and never exceed three.

When a substance loses electron(s), it becomes oxidized, i.e. its oxidation number increases. This is called oxidation.
Due to the fact that non-metals accept the electrons(s) donated by
other substances, particularly metals, they are, therefore, termed as oxidizing agents. This is because by accepting the electrons, they help oxidize the electron donors.

Those substances or metals which donated the electrons are called reducing agents. This is because the electrons they donate reduce the oxidation number of non-metals. This process is called reduction. In this respect, non-metals act as oxidizing agents while metals act as reducing agents.

Strong and weak oxidants

As
we have already seen, non-metals ionize by electron gain. In all cases,
the extra electron(s) accepted lead to the formation of negative ions.
The easiness of formation of negative ions depends on the ability of an
element to accept the electrons. The ability of accepting electrons is
called electronegativity of an element. Some elements are more electronegative than others.

The
order of electronegativity of some non-metals is as follows: Fluorine
< Chlorine > Bromine > Iodine > Nitrogen > Carbon

The
degree of electronegativity indicates reactivity and hence oxidizing
power of the element. Elements with higher electronegativity will
displace those elements with lower electronegativity from their
compounds.

Referring
to the series above, fluorine will displace all the rest of the
elements from their compounds as it is more electronegative than any
other element in the series. Likewise, chlorine can displace bromine,
iodine and nitrogen from their compounds. The displacement reaction
occurs in the manner:

Where X is more electronegative than Y

The
higher the electronegativity the stronger the oxidant. For example,
bromine is a stronger oxidant than iodine, nitrogen and carbon.

The Displacement of Non-metals by another Non-metal from a Compound

Describe the displacement of non-metals by another non-metal from a compound

Non-metals in the reactivity series

It is useful to placecarbonandhydrogeninto the reactivity series because these elements can be used to extract metals.

Here is the reactivity series including carbon and hydrogen:

Note
that zinc and iron can bedisplacedfrom theiroxidesusing carbon but not
using hydrogen. However, copper can be extracted using carbon or
hydrogen.

Chlorine

Chlorine
is very reactive, so it is never found as the free element in nature.
It occurs mainly as sodium chloride or rock salt. It also occurs in the
combined state as chlorides of sodium, potassium and magnesium.

How is chlorine made?

In
industry, chlorine is made by electrolysis of molten sodium chloride or
brine. Brine is a concentrated solution of sodium chloride in water.

In
the laboratory, chlorine is made by the oxidation of concentrated
hydrochloric acid. The oxidation can be brought about by a number of
oxidizing agents, for example, lead (IV) oxide, manganese (IV) oxide,
trilead tetraoxide (Pb₃O₄) or potassium manganate (VII).

The reaction equation is:

When potassium permanganate is used, reaction equation is:

The poisonous nature of chlorine

Chlorine
is a useful but dangerous gas. It is very poisonous if inhaled to even a
small extent (1 part of chlorine in 50,000 parts of air causes death).
Chlorine poisoning occurs when the gas is inhaled or swallowed. It
reacts with water inside and outside of the body (such as water in the
digestive tract and moisture on the lungs and eyes) to form hydrochloric
acid and hydrochlorous acid. Both of these substances are extremely
poisonous.

Most
incidents of chlorine poisoning are due to ingesting household cleaners
or bleach products. Due to its poisonous nature, chlorine was used in
the World War I (1914-1918) to kill people and it caused many deaths.

The Chemical Properties of Chlorine

Describe the chemical properties of chlorine

1.
When a little litmus solution is poured onto a gas jar of chlorine,
litmus immediately turns colourless. The gas also bleaches a damp litmus
paper since the gas is acidic. If blue litmus paper is used, it is
first turned red before being bleached. The bleaching action is due to
the fact that chlorine reacts with the water, forming a mixture of
hypochlorous and hydrochloric acids.

The hypochlorous acid is a very reactive compound and readily gives up its oxygen to the dye, to form a colourless compound:

Thus,
dry chlorine does not bleach. It has to be moist since the hypochlorous
acid formed by its reaction with water is the one that is responsible
for the bleaching action.

2. Chlorine reacts with many reactive metals, on heating, giving their respective chlorine salts, e.g.:

3. Some non-metals burn in chlorine to form covalent compounds:

(i) Dry, yellow phosphorous burns spontaneously in the gas to produce white fumes of chlorides of phosphorous, mainly PCl₃.

4.
When a filter paper dipped into a little turpentine is dropped into a
gas jar of chlorine, a violent reaction occurs and a black cloud of
solid particles of carbon is formed.

The
reaction also takes place with other hydrocarbons as well and as with
turpentine, hydrogen chloride gas is formed. The hydrogen chloride can
be shown to be present by passing a little ammonia gas across the top of
the jar whereby dense white fumes of ammonium chloride are observed.

5.
Hydrogen chloride gas is oxidized to elemental sulphur by chlorine gas.
When a gas jar of hydrogen sulphide is inverted over a gas jar of
chlorine, a yellow precipitate of sulphur and hydrogen chloride gas will
be formed.

6.
When a stream of chlorine gas is bubbled through a pale green iron (II)
chloride solution, a red-brown precipitate of iron (III) chloride is
formed, showing that the iron (II) chloride has been oxidized to iron
(III) chloride.

Ionically:

7. When chlorine is bubbled into the solution of sulphur dioxide in water for a few minutes, dilute sulphuric acid is obtained.

8. The reaction between chlorine and alkalis gives products which depend upon reaction conditions:

Chlorine
reacts with cold dilute aqueous solution of sodium or potassium
hydroxides, forming a pale yellow solution of the hypochlorite and
chloride of the metal.

With the hot concentrated aqueous solution, a mixture of the chloride and chlorate is formed.

A similar reaction occurs if hot concentrated calcium hydroxide solution is used.

The Uses of Chlorine

Explain the uses of chlorine

1.
Chlorine is bleaching agent and is also used in the manufacture of
other bleaches. When chlorine is added to sodium hydroxide solution,
bleach is made.

The active chemical in bleach is sodium hypochlorite (sodium chlorate (I))

If chlorine is passed for a considerable time over solid calcium hydroxide the product formed is bleaching powder.

Bleaching
power finds extensive use in dye works, and in laundries. It is used in
industries where cloth, cotton, paper, etc. need to be bleached. Many
textile industries use chlorine for bleaching purposes. Bleach is also
used to kill bacteria for example in the toilet. It will also remove
colour from the dyed materials.

2.
Chlorine is added to water supplied to homes and industries to kill
disease-causing germs like bacteria. If they were not killed, these
pathogens might cause diseases such as cholera and typhoid. It is also
used to sterilize the water in swimming pools.

3.
Chlorine is used to make some important chemicals such as hydrochloric
acid, chlorofluorocarbons (CFCs), tetrachloromethane (CCl₄), and chloroform (CHCl₃).

CFC_s are carbon compounds containing both chlorine and fluorine. An example is trichchlorofluoromethane (CCl₃F).

CFC_s are very uncreative compounds. They were used in aerosol cans and refrigerators. However, CFC_s damage the ozone layer. Consequently, their use is strongly discouraged.

4.
Chlorine is a reactant in the manufacture organo-chloro compounds which
are used to make pesticides, antiseptics (e.g. dettol), herbicides
(weed killers) and polyvinyl chloride (PVC) which is an intermediate
compound in the production of plastics.

Hydrogen Chloride

A Dry Sample of Hydrogen Chloride Gas

Prepare a dry sample of hydrogen chloride gas

Hydrogen
chloride is a gaseous compound at room temperature. It is usually
prepared in the laboratory by the reaction between concentrated
sulphuric acid and any chloride, e.g. sodium chloride. When a mixture of
the two is gently warmed, hydrogen chloride gas is formed.

Preparation of hydrogen chloride gas

The Properties of Hydrogen Chloride Gas

Explain the properties of hydrogen chloride gas

(a) Physical properties

Includes

It is a colourless gas with a choking, irritating smell and an acid taste.
It is heavier (denser) than air.
It fumes in most air due to the formation of tiny droplets of hydrochloric acid.
It is very soluble in water (450 cm³ of gas in 1 cm³ of water). The acidic properties and the solubility of hydrogen chloride gas can best be shown by the fountain experiment

b) Chemical properties

1.A
dry sample of hydrogen chloride gas has no effect on dry, blue litmus
paper but it turns moist, blue litmus paper to red. This is due to
acidic properties of hydrogen chlorine gas.

2. It does not burn and does not support combustion.

3. When a stream of hydrogen chloride gas is passed over some heated metals, the anhydrous chloride of these metals are formed:

4. It is decomposed by strong oxidizing agents such as manganese dioxide, lead dioxide and red lead to give chlorine gas.

5.
When the gas is passed though a solution of silver nitrate, acidified
with dilute nitric acid, a white precipitate of silver chloride is
formed. This is another test for hydrogen chloride gas and all soluble
chlorides.

6. It gives clouds of white fumes when brought into contact with ammonia vapour. The fumes are solid ammonium chloride (NH₄Cl).

In qualitative analysis, this is the chemical test for hydrogen chloride gas as well as for ammonia gas.

Reactions of aqueous hydrogen chloride

Hydrogen
chloride gas is very soluble in water (and in other polar solvents). In
water, an acid solution is formed, which is hydrochloric acid. In
aqueous solution, the hydrogen chloride molecule dissociates into
hydrogen ions (H⁺) and chloride ions (Cl^–):

The
solution is called hydrochloric acid. Hydrochloric acid reacts with
metals, metal oxides, hydroxides (soluble bases) and metal carbonates..

This suggests that when hydrogen chloride gas is dissolved in water an acidic solution is formed, for example:

it turns damp blue litmus paper red;
it reacts with various substances just like other acids (see table bellow); and
it conducts electricity, yielding hydrogen gas at the cathode and chlorine gas at the anode.

Reactions of aqueous hydrogen chloride

Acid reacting with	General equation
oxide (base)	acid + metal oxide ® salt + water
alkali (soluble base)	acid+metal hydroxide(akali)®salt+water
metal	acid + metal ® salt + hydrogen
metal carbonate	acid + metal carbonate®salt+water+CO₂

The Uses of Hydrogen Chloride

Explain the uses of hydrogen chloride

These are:

It
is chiefly used in the production of hydrochloric acid. When the gas is
dissolved in water in the appropriate proportions, hydrochloric acid is
formed.
Aqueous hydrogen chloride is used in qualitative and quantitative analysis.
It
is an important reagent in other industrial chemical transformations,
e.g. hydrochlorination of rubber and production of vinyl and alkyl
chlorides.
In the electronics industry, it is used to both rub semiconductor crystals and to purify silicon.
It is used in the textile industry, to separate cotton from wool and fluff.
In the laboratory, anhydrous forms of the gas are particularly useful for generating chloride-based Lewis acids.
It is used to remove rust from the oxidized metals.
It
is extensively used in the manufacture of medicines and is a key
substance utilized to turn maize and other agricultural products into
artificial sweeteners.

Sulphur

Occurrence

Sulphur
is a yellow, crystalline, non-metallic solid. Its symbol is S. It has
an atomic number of 16. Sulphur exists in nature as a free element and
in compounds.

As
a free element, sulphur is found in several countries such as Italy,
Mexico, Japan, Poland, USA and Sicily. In its combined state, sulphur is
found combined with metal ores such as galena (PbS), iron pyrites (FeS₂), Copper pyrites (CuFeS₂) and zinc blend (ZnS). It is also found in natural gas as hydrogen sulphide (H₂S₎and in crude oil as organic sulphur compounds.

The Extraction of Sulphur from Natural Deposits

Describe the extraction of sulphur from natural deposits

Sulphur
is extracted from its underground deposits by the Frasch process. The
Frasch process makes use of the low melting point (119^oC) of sulphur.

In
this process, a hole about 30 cm in diameter is bored down through the
clay, sand, and limestone to the sulphur beds. This boring is lined with
an iron pipe and inside the pipe, is sunk a device called sulphur pump.
The pump consists of three concentric pipes (cylindrical pipes with a
common centre) which end in a reservoir of a large diameter (see figure
bellow).

Extraction of sulphur

Down the outermost of the three pipes is forced a stream of water at about 170^oC.
This water must be kept at a pressure of about 10 atmospheres to
maintain it in the liquid state, i.e. it is superheated water, and it is
hot enough to melt the sulphur. The molten sulphur flows into the
reservoir at the base of the pump and is forced up to the surface. The
sulphur obtained is 99.5% pure and can be used without any purification.

The stages involved in the entire extraction process can be summarized as follows:

Superheated water (170^oC) is pumped through the outer pipe to melt the sulphur.
Hot
compressed air (10 atm) is pumped down through the inner pipe. The
combination of the hot water and the hot air melts the sulphur. The
molten sulphur, hot air and hot water form a froth.
The froth is forced to the earth’s surface through the middle pipe by the compressed air.
The molten sulphur is collected in large tanks (where the water drains off), cooled and solidified.

The Properties of Sulphur

Explain the properties of sulphur

Sulphur
is a relatively reactive element that readily reacts with other
elements to form compounds such as oxides, chlorides and sulphides.

In chemical reactions, sulphur exhibits both oxidizing and reducing properties.

1. Oxidizing properties of sulphur

Heated sulphur reacts with metals such as iron, copper, zinc and tin to give metal sulphides.

When
a mixture of iron filings and powdered sulphur, in the proportion of 56
to 32, that is 7:4 (the ratio of their relative atomic masses), is
heated, the two react in a highly exothermic reaction. The heat given
out makes the mixture to continue glowing even after the heating has
stopped. A black or dark grey iron (II) sulphide is formed.

In this reaction, sulphur acts as an oxidizing agent which oxidizes iron metal to iron (II) ion.

Sulphur is reduced to sulphur ion

Iron
(II) sulphide is not attracted by a magnet since it is not magnetic.
The magnetic property of iron is lost when this compound is formed.

The copper powder reacts with sulphur in the same way to form copper (II) sulphide.

2. Reducing properties of sulphur

Dilute
acids do not act upon sulphur. However, sulphur is oxidized by hot
concentrated sulphuric acid with the formation of sulphur dioxide.

In this reaction, the sulphur is oxidized by the acid to sulphur dioxide and the acid is reduced to the same substance.

The filter paper soaked in potassium dichromate changes from orange to green. The sulphur dioxide reduces chromium from chromium (VI) to chromium (III).

When sulphur reacts with hot concentrated nitric acid, brown fumes of nitrogen dioxide are formed.

Sulphur
is oxidized by hot concentrates nitric acid to sulphuric acid while
nitric acid is reduced to nitrogen dioxide and water.

The Uses of Sulphur

Explain the uses of sulphur

Includes

Most
of the sulphur produced in the world (90%) is used to manufacture
sulphuric acid. Sulphuric acid is an important reagent in many
industrial processes.
Sulphur is used in the manufacture of sulphur dioxide (used in the Contact Process for the manufacture of sulphuric acid).
Manufacture of calcium hydrogensulphite, Ca(HSO₃)_2,and sodium sulphite which are used for bleaching wood straw and wood pulp in the paper industry.
It
is also used for vulcanization of natural rubber. Rubber is usually
sticky and soft. When heated with sulphur (vulcanization), it becomes
hard and strong.
It is used for dusting vines to prevent growth of certain kinds of fungi and also as an insecticide.
Sulphur
is used in smaller quantities for the manufacture of dyes, explosives,
fireworks, gunpowder etc. For example gunpowder is a mixture of
potassium nitrate, carbon and sulphur.
It is used in the
manufacture of various organic compounds such as plastics and
pharmaceuticals like sulphur ointments e.g. sulphadimadine, septrin
e.tc.
Photographic chemicals such as carbon disulphide (CS₂) and sodium thiosulphate (Na₂S₂O₃) are made using sulphur as one of the raw materials.
Some
is added to cement to make sulphur concrete. Unlike ordinary cement,
this is not attacked by the acid. So it is used for walls and floors in
plants where acid is used.

Sulphur Dioxide

Preparation of sulphur dioxide

In
the laboratory, sulphur dioxide is prepared by heating a mixture of
sodium sulphite and dilute hydrochloric acid. The reaction equation is:

Alternatively,
sulphur dioxide can be prepared by heating a mixture of concentrated
sulphuric acid and copper. In this case, there is no reaction until the
mixture in the flask becomes hot. Then rapid effervescence occurs and
the gas is usually collected as shown.

A dark brown mixture is left in the flask. It contains anhydrous copper (II) sulphate and certain impurities.

The Properties of Sulphur Dioxide

Describe the properties of sulphur dioxide

Physical properties

These are:

The gas is colourless with an irritating (pungent), chocking smell.
It is denser than air. Its density is 2½ times that of air.
It is readily soluble in water and forms an acidic solution of sulphurous acid.

Acidic characteristics of sulphur dioxide

As indicated above, sulphur dioxide gas dissolves in water to form an acid solution of sulphurous acid, H₂SO₃. The solution turns blue litmus paper red. Sulphur dioxide is thus an acidic gas.

Chemical properties

The reaction characteristics described below explains the chemical properties of sulphur dioxide gas.

The reducing properties of sulphur dioxide

(i) Reduction of dyes in flower petals or colour in paper (bleaching)

Sulphur
dioxide bleaches the colours in dyes such as flower pigments. When the
flower pigments or dyes contain oxygen, they are coloured. Sulphur
dioxide reduces the dye (removes oxygen from it) and the dye, therefore,
turns colourless. This process can be summarized as follows:

In this process, sulphur dioxide dissolves in water to form sulphurous acid:

Then
the acid takes up oxygen from the dye of the flowers or paper and forms
sulphuric acid. The removal of oxygen from the dye converts it to a
colourless compound:

However, the oxygen from the air may oxidize the reduced colourless compound back to the original coloured compound.

(ii) Reduction of nitric acid

If sulphur dioxide gas is bubbled through concentrated nitric acid, brown fumes of nitrogen dioxide gas are formed:

The nitric acid is reduced to nitrogen dioxide and the sulphur dioxide is oxidized to sulphuric acid.

(iii) Reduction of acidified potassium permanganate solution

Sulphur dioxide decolourized purple potassium permanganate solution.

Sulphur dioxide is first converted to a sulphite (SO₃^2-), after reacting with water, and then oxidized to a sulphate (SO₄^2-) by the permanganate.

The ionic equation is:

The oxidation state of manganese changes from +7 to +2. This is a reduction reaction.

(iv) Reduction of acidified potassium dichromate solution

When sulphur dioxide is mixed with potassium dichromate (VI) solution the orange colour of the solution changes to green. The dichromate (VI) is reduced to chromate (III) which is green in colour when in aqueous state:

The
oxidation state of chromium changes from +6 in potassium dichromate
(VI) to +3 in chromic sulphate. On the other hand, sulphur dioxide is
oxidized by the dichromate (VI) to sulphate (SO₄²^–).

Other reduction reactions involving sulphur dioxide gas

Sulphur dioxide reduces chlorine, bromine and iodine to the hydrogen halides in the presence of water:

In all these cases, the solution changes from brown to colourless.

The oxidizing properties of sulphur dioxide

(i) Oxidation of hydrogen sulphide

Iron (II) sulphide reacts with dilute hydrochloric acid to produce hydrogen sulphide gas and iron (III) chloride:

Hydrogen
sulphide gas has a smell similar to that of a rotten egg. The hydrogen
sulphide gas is oxidized by sulphur dioxide in the gas jar to sulphur:

The
sulphur produced is a yellow residue.The reaction takes place in the
presence of moisture which acts as a catalyst. That is why water is
added in the jar.

(ii) Oxidation of magnesium

When
burning magnesium is lowered into a jar containing sulphur dioxide gas,
a white solid, magnesium oxide, and yellow pieces of sulphur are
formed:

Tests for sulphur dioxide gas

Includes

The gas can be identified by its characteristic pungent and choking smell.
It
can also be detected by putting into it a filter paper that has been
previously dipped into an acidified solution of potassium dichromate
(VI). The colour of the filter paper changes from orange to green due to the reduction of dichromate (VI) to chromate (III).
Sulphur dioxide also decolourized acidified potassium permanganate solution.

The Uses and Hazards of Sulphur Dioxide

Explain the uses and hazards of sulphur dioxide

Sulphur
dioxide has got a number of uses in daily life. However, there are also
some hazards which can be caused by the gas if its production is not
controlled.

Uses of sulphur dioxide

The following are some uses of sulphur dioxide gas:

The gas is used in the industrial manufacture of sulphuric acid in the Contact Process.
It is used as a bleaching agent for wood pulp, silk, wool and straw.
Its
poisonous nature makes it a useful fumigant. So it is used in
fumigation. The gas is poisonous to all organisms, particularly
bacteria.
It is used as a preservative and sterilizing agent in
making soft drinks and jam, and in dried fruits. A very low
concentration of the gas in food prevents fermentation as it stops the
growth of bacteria and moulds. Its reaction with oxygen prevents
oxidation of juices and other liquids to which it is added for
preservation.

Hazards of sulphur dioxide gas

The
hazards of sulphur dioxide gas are due to its effect in environmental
pollution and the health problems accompanied with that pollution.
Sulphur dioxide is a major air pollutant. The major sources of sulphur
dioxide in the air are power plants that use fossil fuels such as coal,
diesel and petrol; industrial boilers; and exhaust emissions from motor
vehicles. The gas is also produced during metal smelting and other
industrial processes.

Half
of sulphur dioxide output comes from burning coal in coal-fired power
stations. All coal contains small amounts of sulphur. So when the coal
is burnt to produce energy, the sulphur in the coal reacts with oxygen
in the air to produce sulphur dioxide

Sulphur
dioxide is a very irritating gas and is thought to be the cause of
bronchitis and other lung diseases. Exposure to higher concentrations of
the gas can cause impairment of the respiratory function and heart
diseases.

Sulphur
dioxide also causes acid rain. This occurs when the gas comes in
contact with moist air. The sulphur dioxide dissolves in water vapour
from the clouds and combines with oxygen from the atmosphere to form an
acid – sulphuric acid:

Acid
rain damages the leaves and barks of plants making them more vulnerable
to diseases, weather and insects. When acid rain reaches the lake,
river or other water bodies it makes the whole water body acidic. Even a
low concentration of acid in the water can kill fish and other marine
organisms.

Sulphuric Acid

The Contact Process for the Manufacture of Sulphuric Acid

Describe the contact process for the manufacture of sulphuric acid

Sulphuric
acid is an important laboratory and industrial reagent. It is produced
in large scale through the Contact Process. The process involves four
major stages. These are:

production of sulphur dioxide;
purification of sulphur dioxide;
catalytic conversion of sulphur dioxide (SO2) to sulphur trioxide (SO3); and
conversion of sulphur trioxide to sulphuric acid.

1. Production of sulphur dioxide

The sulphur dioxide used in the Contact Process can be obtained from different sources. These include:

(a) Burning sulphur in air:

This is the most convenient method of producing sulphur dioxide.

(b) Burning sulphide ores such as iron pyrite (FeS₂) and zinc blend (ZnS):

Sulphur dioxide gas is produced as a by-product.

2. Purification of sulphur dioxide

The
sulphur dioxide produced in the first stage is mixed with air, ready
for passing it over the catalyst. Before contact with the catalyst is
allowed, the gas mixture has to be purified to remove impurities. This
is achieved by passing the mixture through an electrostatic precipitator
to remove any dust. It is then washed with water to remove impurities
such as traces of arsenic (III) oxide (As₂O₃). The gas mixture is then passed through concentrated sulphuric acid to remove all moisture. The three impurities (As₂O₃, dust and moisture), if not removed, will poison the catalyst thereby rendering it useless.

3. Catalytic conversion of sulphur dioxide to sulphur trioxide

The
purified and dried mixture of sulphur dioxide and air is passed through
a heat exchanger to acquire the necessary heat for the conversion to
sulphur trioxide. The mixture is then taken to the conversion chamber,
which contains a catalyst. The catalyst used is finely divided vanadium
(V) oxide (vanadium pentoxide, V₂O₅) which is heated to 450^°C.

Originally,
platinized asbestos was used as a catalyst. But, compared to vanadium
(V) oxide, platinum is very expensive and easily poisoned by impurities.
So it has been replaced by vanadium (V) oxide as the usual catalyst
used in the Contact Process.

Sulphur
dioxide remains in contact with the catalyst during the conversion
process, hence the name Contact Process. The reaction that takes place
during the conversion is:

The conversion takes place at a temperature of 450^°C and normal atmospheric pressure (1 atm).

The
reaction is exothermic, which means that, as sulphur trioxide is
formed, heat energy is released. If the temperature rises above 450^°C the yield of sulphur trioxide decreases.

Once the reaction has started, no external heating is required. Thus, the heat exchanger maintains the temperature at 450^°C.
According to Le Chatelier’s principle, a lower temperature should be
used to shift the equilibrium to the right, hence increasing the
percentage yield. However, too low temperature will lower the formation
rate to an economical level. Hence, to increase the rate, high
temperature (450^°C), medium pressure (1-2 atm) and a catalyst (V₂O₅)
are used to ensure maximum yield. The catalyst only serves to increase
the rate of reaction as it does not change the position of the dynamic
equilibrium.

About 98% of the possible yield of sulphur trioxide is obtained.

4. Conversion of sulphur trioxide to sulphuric acid

The
sulphur trioxide from the conversion chamber is passed through a heat
exchanger to remove excess heat. It is then taken to an absorption tower
where it is dissolved in concentrated sulphuric acid to form oleum or fuming sulphuric acid:

Oleum is then carefully diluted with the correct amount of water to give ordinary concentrated sulphuric acid:

Sulphur
dioxide cannot be dissolved directly in water to form sulphuric acid.
The reaction is so highly exothermic that the heat produced vapourizes
the sulphuric acid formed. This makes it difficult to collect the gas
because the acid vapour (mist) produced is very stable and cannot be
condensed.

The flow diagram for the manufacture of sulphuric acid is show in the figure below

A flow diagram of the Contact Process

The Properties of Sulphuric Acid

Explain the properties of sulphuric acid

Chemical properties of dilute sulphuric acid

Dilute
sulphuric acid reacts with metals, metal oxides, metal hydroxides and
metal carbonates and hydrogencarbonates to produce salts.

Reaction with metals

Dilute sulphuric acid reacts with many metals above hydrogen in the activity series to form sulphates and hydrogen gas, e.g.

Reaction with metal oxides

Reactions
of metal oxides with dilute sulphuric acid are neutralization
reactions. Metal oxides react with dilute sulphuric acid to form a salt
(sulphate) and water, e.g.

Reaction with metals hydroxides

The
reaction between dilute sulphuric and a metal hydroxide is a
neutralization reaction. Metal hydroxides react with dilute sulphuric
acid to form a sulphate and water, e.g

Reaction with metal carbonates and hydrogencarbonates

Dilute
sulphuric acid reacts with metal carbonates and hydrogencarbonates to
give metal sulphates, carbon dioxide and water, e.g.

Chemical properties of concentrated sulphuric acid

Concentrated sulphuric acid as a dehydrating agent

As
a dehydrating agent, it will remove the elements of water (hydrogen and
oxygen) from a compound to form a new compound. It will dehydrate
sugar, paper and wood. These are all made of carbon, hydrogen and
oxygen. The acid removes the hydrogen and oxygen as water, leaving
carbon behind.

When
concentrated sulphuric acid is added to sugar, a vigorous reaction
occurs, causing the reaction mixture to rise and fill the beaker.

The
colour of the sugar changes to brown and finally black. Concentrated
sulphuric acid dehydrates sugar (glucose) by taking away the elements of
water (hydrogen and oxygen) from the sugar, leaving carbon.

The acid will also dehydrate sucrose to form carbon.

The
final product is a black mass of carbon. The reaction is highly
exothermic. The heat produced evaporates the water formed from the
reaction.

The
acid also dehydrates some hydrated salts. When concentrated sulphuric
acid is added to hydrated blue copper (II) sulphate crystals, the colour
changes from blue to white. The acid dehydrates the hydrated copper
(II) sulphate crystals to anhydrous copper (II) sulphate powder:

Concentrated sulphuric acid also dehydrates methanoic acid and ethanol to give carbon monoxide and ethene gases respectively:

Concentrated sulphuric acid as a drying agent

As
a drying agent, concentrated sulphuric acid absorbs traces of water
from substances. Because of its ability to absorb water, it is used for
drying most gases prepared in the laboratory that it would not react
with. It cannot be used for drying ammonia, carbon dioxide, hydrogen
sulphide or any gas with which it reacts.

Sulphuric acid as an oxidizing agent

Hot
concentrated sulphuric acid is a strong oxidizing agent. It oxidizes
both metals and non-metals while it is reduced to sulphur dioxide.

Concentrated
sulphuric acid oxidizes charcoal (carbon) to carbon dioxide, sulphur to
sulphur dioxide and copper to copper (II) sulphate.

With sulphur and copper, the orange colour of the dichromate (VI) paper changes to green. This confirms the presence of sulphur dioxide gas.

With
carbon, a white precipitate is formed on the glass rod when the rod
dipped in lime water (calcium hydroxide) is placed in the mouth of the
test tube. This confirms the presence of carbon dioxide gas which reacts
with the calcium hydroxide on the glass rod to produce a white
precipitate of calcium carbonate.

The Uses of Sulphuric Acid

Explain the uses of sulphuric acid

Sulphuric
acid is one of the most important industrial chemicals. It has widely
varied uses and plays some part in the production of nearly all
manufactured goods. The following are some of the uses of sulphuric
acid:

Manufacture of fertilizers
The major use of sulphuric acid is the production of fertilizers such
as ammonium sulphate and superphosphates (phosphate fertilizers).
Manufacture of chemicals
It is widely used in the manufacture of chemicals e.g. in making
hydrochloric acid, nitric acid, phosphoric acid, sulphate salts,
synthetic detergents, soap, paints and pigments, explosives, plastics
and drugs.
Refining of crude oil A large quantity of sulphuric acid is used in refining crude oil.
Extraction and manufacturing of metals
Sulphuric acid is used in the iron and steel-making industry to remove
rust and scale from the surface of the rolled iron sheets. It is also
used in processing metals e.g. in pickling (cleaning) iron and steel
before plating them with tin or zinc to produce galvanized iron.
Manufacture of alum
Sulphuric acid is used in the manufacture of aluminium sulphate, which
is used in water treatment plants to filter impurities and to improve
the taste of water. Aluminium sulphate is made by reacting bauxite with
sulphuric acid.
Manufacture of natural and man-made fibres
Sulphuric acid is used for making natural and synthetic (artificial)
fibres. For example, the textile called rayon is made from cellulose
fibres derived from wood. These fibres are dissolved in a solution of
tetraamminecopper (II) sulphate to produce a thick liquid which is then
injected into sulphuric acid to form rayon fibres.
Other uses:Sulphuric
acid is used as (i) an electrolyte in lead-acid batteries, which are
used in cars, to allow the flow of electrons between the plates in the
battery. The sulphuric acid used in this way is called battery acid;
(ii) as a general dehydrating agent in its concentrated form in tanning
leather; and (iii) in waste water treatment.

Nitrogen

Occurrence

Nitrogen
makes about 78% of the air by volume. The element also occurs combined
with other compounds in the form of sodium nitrate, Chile saltpetre,
NaNO₃ (as a mineral deposit in Chile), and in the soil as ammonium sulphate, sodium nitrate, potassium nitrate and calcium nitrate.

Combined
nitrogen also occurs as a constituent of all living matter of plants
and animals in the form of proteins, enzymes, alkaloids.

Preparation of a sample of Nitrogen in the Laboratory

Prepare a sample of nitrogen in the laboratory

Nitrogen can be prepared in the laboratory by either separating the gas from the air or by heating ammonium nitrite.

(i) Preparation of nitrogen by the action of heat on ammonium nitrite

The reaction between sodium nitrite and ammonium chloride gives sodium chloride and ammonium nitrite.

A solution of ammonium nitrite readily decomposes, on slight warming, to give nitrogen gas.

As the solution becomes warm, rapid effervescence occurs as more nitrogen is given off.

The
nitrogen evolved may be collected over water because the gas is only
slightly soluble in water, at ordinary temperature, and slightly denser
than air.

Tests for nitrogen gas

Aim: To test for the presence of nitrogen.

Procedure

Using
the four gas jars of nitrogen gas collected in the previous experiment
(Experiment 1.16), carry out the following tests for nitrogen gas and
write down your observations:

Remove the cover from the first jar and smell the gas. Observe the colour of the gas and identify its smell.
Remove the cover from the second jar and put in it a piece of damp universal indicator.
Place a lighted splint into the third gas jar.
To the fourth jar, add some calcium hydroxide solution (lime water) and shake.

Observations and inferences

The gas is colourless and odourless. This distinguishes it from gases such as sulphur dioxide, ammonia, hydrogen chloride, etc.
The colour of the indicator does not change. This shows that nitrogen is a neutral gas.
The
lighted splint is extinguished and the gas does not burn. It can not,
therefore, be any gas which supports combustion, e.g. oxygen, dinitrogen
oxide, or any combustible gas, e.g. hydrogen sulphide, carbon monoxide,
hydrogen, etc.
After the above tests, the only gas with which
nitrogen may be confused is carbon dioxide. To distinguish it from
carbon dioxide, the gas is dissolved in calcium hydroxide solution.
Nitrogen leaves the calcium hydroxide unchanged while carbon dioxide
turns the solution milky (due to formation of a precipitate of CaCO₃).

Properties of nitrogen

Physical properties

Nitrogen
is colourless and odourless. It is slightly less dense than air and
sparingly insoluble in water. The gas is neutral to litmus.

Chemical properties

Under
ordinary conditions, nitrogen gas is inert. However, the gas only takes
part in reactions at very high temperature as follows:

1. It reacts with red hot metals, like magnesium and calcium to form nitrites

2.
Nitrogen does not burn nor does it support combustion. When heated, the
gas combines with oxygen to form nitrogen monoxide gas:

3. At ordinary pressure and moderately high temperatures, nitrogen reacts with hydrogen in the Haber Process to produce ammonia:

The Uses of Nitrogen

Explain the uses of nitrogen

Here are some uses of Nitrogen:

Manufacture of fertilizers:
Nitrogen is used to manufacture nitrogenous fertilizers. These include
diammonium phosphate (DAP), calcium ammonium nitrate (CAN), ammonium
superphosphate (ASP), ammonium nitrate (AN), ammonium phosphate sulphate
(APS), ammonium sulphate nitrate (ASN), ammonium sulphate (AS) and
urea.
Refrigeration: Because of its
low boiling point (-196oC), liquid nitrogen is used as a refrigerant for
storing organs in research laboratories, bull semen for artificial
insemination, eggs and other cells for medical research and reproductive
technology, etc. It is also used for preservation of food products and
for their transportation.
Processing reactive substances:
Some reactions require an inert atmosphere in order to proceed as
desired. Because of its low reactivity, nitrogen is used to provide an
inert atmosphere for storing and processing reactive substances.
Manufacture of nylon:
Nitrogen is used in the manufacture of synthetic fibres such as
polyamides. Polyamides are commonly known as nylons. Nylons are
chemically inert and are stronger than natural fibres. They are used in
making fishing nets, clothes, ropes, and many other items.
Manufacture of ammonia:
Nitrogen is used in the manufacture of ammonia through the Haber
Process.In the Haber Process, ammonia is manufactured by direct
combination of nitrogen and hydrogen.Nitrogen and hydrogen are mixed in
the ratio of 1:3. The gases are then reacted together at a temperature
of about 450°C and a pressure of 250 atmospheres in the presence of
finely divided iron as a catalyst. N2(g) + 3H2(g) ⇔2NH3(g) + heat The
gases are cooled while still under pressure and ammonia is removed as a
liquid.
Manufacture of nitric acid: The ammonia
gas manufactured using nitrogen in the Haber Process is used in the
manufacture of nitric acid by catalytic oxidation.
Plant nutrition:
When atmospheric nitrogen is fixed into the soil by bacterial actions,
it becomes a nutrient to plants. Nitrogen fixation refers to the
conversion of atmospheric nitrogen, by certain species of bacteria, into
nitrites.

Ammonia

Preparation of a Dry Sample of Ammonia Gas in the Laboratory

Prepare a dry sample of ammonia gas in the laboratory

Ammonia
is a binary compound of nitrogen and hydrogen. Ammonia gas is
colourless and has a strong pungent and choking smell. It does not occur
free in air, but exists in nitrogenous organic materials such as hoofs
and horns of animals. The gas can be released by heating or burning
these materials.

Laboratory preparation of ammonia

Ammonia
is prepared in the laboratory by heating any ammonium salt with an
alkali. In most cases, ammonium chloride and calcium hydroxide (the
cheaper alkali) are used. Both are solid so they must be thoroughly
ground first to give a very fine mixture so that the reaction can occur
efficiently.

Ca(OH)_2(s) + 2NH₄Cl_(s)→CaCl2_(s)+ 2H₂O_(l) + 2NH_3(g)

Calcium
hydroxide reacts with ammonium chloride to produce ammonia gas, calcium
chloride and water. The flask is tilted to prevent any condensed water
formed from running back into the hot flask, which might break it.

Ammonia
gas is dried by passing it over quicklime because it reacts with all
the usual drying agents. Concentrated sulphuric acid is acidic and would
absorb the gas forming a salt e.g. 2NH_3(g)+ H₂SO_4(l) → (NH₄)₂SO_4(s)

It reacts with calcium chloride, forming solid complex compounds, e.g. 8NH_3(g) + CaCl_2(s)→ CaCl₂.8NH_3(s)

Ammonia
is less dense than air and very soluble in water, so it is collected as
shown by upward delivery (or downward displacement of air).

Ammonia is an alkaline gas and it turns wet, red litmus paper blue.

The Properties of Ammonia

Describe the properties of ammonia

Physical Properties of Ammonia

It is a colourless poisonous gas with a strong chocking smell.
It is less dense than air.
It is easily liquefied by cooling to -33°C or by compressing it.
It turns wet, red litmus paper blue as it is the only common alkaline gas
It
is very soluble in water forming alkaline solution. Ammonia has the
highest solubility of all known gases. About 800 volumes of the gas
dissolve in 1 volume of water at 15°C. The fountain experiment below
demonstrates this solubility.

Chemical properties of ammonia

Reaction with air and oxygen

Ammonia will not normally burn in air but can be made to do so in two ways:

Burning ammonia in an oxygen-rich atmosphere; and
Use of a catalyst.

Burning ammonia in oxygen-rich air

When ammonia is burned in air enriched with oxygen, the products formed are nitrogen and water. 4NH_3(g) + 3O_2(g)→ 2N_2(g) + 6H₂O_(l)

Using a catalyst

The
source of ammonia gas in this experiment is concentrated ammonia
solution which gives off fumes of ammonia gas. If some concentrated
ammonia solution is left in a stoppered flask for a few minutes, the
flask will quickly become full of ammonia gas by diffusion.

The
catalyst used is the metal platinum. The platinum coil is heated in a
bunsen flame until it is red hot. It is then lowered into a flask
containing ammonia. The coil continues to glow even though it is not
being heated. This indicates that a chemical reaction is taking place.
Near the top of the flask brown fumes can be seen.

The products of this reaction are nitrogen monoxide (nitrogen (II) oxide), which is a colourless gas, and water.

4NH_3(g) + 5O_2(g) → 4NO_(g) + 6H₂O_(g)

However,
as the nitrogen monoxide gas moves towards the neck of the flask and
comes into contact with the air, it reacts with the oxygen in the air to
form nitrogen dioxide gas which is brown in colour. Hence, the brown
fumes are seen at the neck of the flask.

2NO_(g) + O_2(g) → 2NO_2(g)

Reaction with copper

When
ammonia gas is passed over heated copper (II) oxide, the gas is
oxidized to nitrogen and water by the hot copper (II) oxide while the
oxide is reduced to copper.

3CuO_(s) + 2NH_3(g) → 3Cu_(s) + 3H₂O_(l) + N_2(g)

Copper (II) oxide is black. During the reaction it is oxidized to copper which is reddish brown in colour.

Reaction with hydrogen chloride gas

When ammonia and hydrogen chloride are mixed, dense white fumes of ammonium chloride are formed.

NH_3(g)+ HCl_(g) → NH₄Cl_(s)

This
test is simply performed by dipping a glass rod into concentrated
hydrochloric acid then holding the glass rod at the mouth of a gas jar
containing ammonia gas.

The Uses of Ammonia

Explain the Uses of Ammonia

Manufacture of fertilizer

Since ammonia solution is alkaline, it reacts with acids to form ammonium salts which are used as fertilizers.

Other
nitrogenous fertilizers manufactured using ammonia include ammonium
sulphate nitrate (ASN), diammonium phosphate (DAP) and calcium ammonium
nitrate (CAN)

Ammonia itself can be used directly as a fertilizer but has three disadvantages:

Ammonia comes as a gas or a concentrated solution which is less easy to store; it is easier to store the solid ammonium salts.
Ammonia is alkaline and can affect the natural pH of the soil.
Ammonia easily evaporates if directly applied to the soil.

Manufacture of nitric acid

A
lot of the ammonia from the Haber Process is used to make nitric acid.
The raw materials for manufacture of nitric acid are ammonia, air and
water.

The overall result is that ammonia is oxidized to nitric acid.

Cleaning

Ammonia
solution is very useful in cleaning. This is because it softens water
in homes and laundries and neutralizes acid stains caused by
perspiration, hence making washing easier. It is also used as a grease
solvent.

Refrigeration

Liquid
ammonia is used in large-scale refrigerating plants such as in ships
and warehouses. It used to be used in domestic refrigerators, but has
now been replaced by non-toxic, non-corrosive chlorofluorohydrocarbons
(CFHs).

Manufacture of synthetic fibres

In the textile industry, ammonia is used in the manufacture of synthetic fibres such as nylon and rayon.

Carbon

The Forms in which Carbon Exists

Name the forms in which carbon exists

Although
carbon forms less than one percent of the earth’s crust by weight, it
is the most interesting of all elements. This is because of the
following reasons:

All living things are made of carbon and its compounds.
Over ¾ of the world’s power is obtained from carbon and its compounds.
Over ¾ of all substances in the world are made of carbon.

Carbon
is not commonly found in a free state. The free element is mainly found
as graphite and diamond. Carbon occurs in a number of other forms, e.g.
wood charcoal, animal charcoal, coke, soot (lampblack).

Carbon
compounds are found in many naturally occurring substances such as
coal, petroleum, wood, coal gas, natural gas, carbonates, shells,
organic matter of all kinds, all living things, and occurs in the air to
a small, but very important extent (0.03 – 0.04% by volume) as carbon
dioxide. The carbohydrates, proteins and lipids in all living things
contain carbon.

Pure
carbon is found in the form of diamond in Tanzania (Mwadui), Sierra
Leone, India, South Africa, Russia and South America; and impure carbon,
as graphite, is found in Sri Lanka.

Allotropic Forms of Carbon

Describe allotropic forms of carbon

An
element that exists in more than one form is said to exhibit allotropy.
Allotropy is the existence of an element in more than one form (without
change of state). The various forms are known as allotropes.

Allotropes
of a given element differ in their physical properties and may differ
in some chemical properties as well. The allotropic forms of carbon are:

graphite;
diamond; and
amorphous carbon.

These
allotropes have got different molecular structures. The structural
differences are mainly due to the way their atoms are packed.

Graphite

Graphite
has a layer structure. Figure 1.18 illustrates one layer of the
structure of graphite. Each layer consists of carbon atoms covalently
bonded together into hexagonal rings. These rings form flat parallel
layers, one over the other. The force that hold the carbon atoms
together are very strong. Adjacent layers are held together by weak van
der Waals’ forces as shown in figure 1.19. The layers readily slide over
one another, accounting for the soft and greasy texture of graphite.

Carbon atoms in one graphite layer

Graphite structure

Carbon
atom has got four electrons in its outer shell. Each carbon atom forms
three covalent bonds to other carbon atoms. Thus, three of its four
outermost electrons are paired up to form covalent bonds. The fourth
electron is not attached to any particular atom (delocalized) and is
free to move anywhere along the layers. Graphite conducts electricity in
the plane of the layers but not at right angles to them.

Because
conduction of electricity involves movement of unshared electrons from
one atom to another atom, graphite is a good conductor of electricity
since the hexagonal layers permit this movement. It is also a good
conductor of heat for a similar reason.

Physical properties of graphite

It is a black, soft and slippery substance. It feels soapy and greasy. It has a metallic lustre and is opaque to light.
It has low relative density (2.3) as compared to diamond (3.5)
Graphite is a good conductor of heat and electricity due to the delocalized electrons.
It
has a very high melting point (3730°C) and boiling point (4830°C). The
melting and boiling points are high because of strong covalent bonds
between the carbon atoms which require more energy (heat) to break in
order to melt graphite.

Uses of graphite

It
is used as an electrode in electrolysis and as a positive terminal in
dry cells. The use of graphite as electrode in electrolysis has an
advantage because it does not react readily with most substances (it is
an inert electrode).
It is used as a lubricant, particularly
when high temperatures are involved, where the usual lubricating oils
easily decompose due to extreme heat. It is a lubricant for dynamos,
electric motors and fast-moving parts of machinery.
The major
use of graphite is in making lead pencils of different hardness, by
mixing it with different proportions of clay. The weakly held layers of
carbon atoms in graphite easily slide over each other and are left
behind on paper as black marks.
Being resistant to chemicals and
having a high melting point and also because it is a good conductor of
heat, graphite is used to make crucibles.
Graphite has the
ability to absorb fast-moving neutrons, thus, it is used in nuclear
reactors to control the speed of the nuclear fission reaction.

Diamond

The
basic unit of the diamond structure is shown in figure 1.20. Each
carbon atom is covalently bonded to four other carbon atoms. This basic
unit is repeated in three dimensions as shown in figure 1.21 to form a
giant tetrahedral structure of millions of carbon atoms, all forming
four covalent bonds to each other. The melting point of diamond is high.
This is because of the strong covalent bonds between carbon atoms,
which require a large amount of heat energy to break up.

Physical properties of diamond

It
is the hardest natural substance known. This due to the strong covalent
bonds between the carbon atoms in diamond. Again the compact
tetrahedral arrangement of carbon atoms contributes to its hardness.
It has the highest melting point (3550°C) and boiling point (4289°C).
It has a high relative density (3.5) compared to graphite (2.3)
It
is a poor conductor of heat and electricity. This is because there are
no free electrons to conduct electricity. All electrons are firmly held
in covalent bonds.
It is colourless, transparent and has a dazzling (amazing) brilliant lustre and radiance.
It
has a high refractive index of 2.5. The high refractive index results
in high dispersion of light, making it suitable for use in jewellery.

Uses of diamond

It is used in making jewellery.
Due
to its extreme hardness, it is used to make glass cutters, drilling
devices, rock borers, and as an abrasive for smoothing very hard
materials.

Comparison of the properties of diamond and graphite

Diamond	Graphite
1. Colourless, transparent andglittering	Black, opaque with metallic lustre
2. Hardest natural substanceknown, used to cut glass andin drills	Soft to touch, greasy or soapy, can be used as a lubricant and in making lead pencils

3. High relative density (3.5)	Low relative density (2.3)
4. Non-conductor	Good conductor of heat and electricity
5. Burns in air least readily (atabout 900°C)	Burns in air readily (at 700°C)
6. Have strong C-C covalentbonds arranged octahedrally toform a giant molecular crystal	Have strong C-C bonds within the hexagonal rings in the sheets but weak Van der Waal’s forces in between layers.
7. Its cleavage is difficult and itoccurs along the boundaries ofthe octahedral crystal unit	Cleavage easy, and occurs along the sheets or the layers.
8. Prepared from graphite at very high pressure and temperature	Prepared from coke and silica mixture at high temperature
9. Not attacked by potassiumchlorate and nitric acid together	Attacked by these reagents

A
proof that diamond and graphite are both allotropic forms of carbon can
be shown experimentally by burning equal masses of each allotrope in a
stream of oxygen. It is found that equal masses of each allotrope
produce equal masses of carbon dioxide. C_(s) + O_2(g) → CO_2(g)

The
weight of carbon dioxide produced can be determined by allowing it to
be absorbed in a weighed amount of potassium hydroxide solution. 2KOH_(aq) + CO_2(g) → K₂CO_3(aq) + H₂O_(l). Then, the weight of carbon dioxide is obtained by calculations based on the equation above.

Amorphous carbon

Amorphous
carbon is carbon that does not have any clear shape, form or
crystalline structure. Amorphous carbon is made of tiny bits of graphite
with varying amounts of other elements considered as impurities. It is
formed when a materials containing carbon is burned in a limited supply
of oxygen, resulting in incomplete combustion.

Amorphous carbon exists in many forms. The following are the major forms in which it occurs:

Charcoal
Lampblack (soot) or carbon black
Coke

Charcoal

Charcoal
is made by heating organic material (animal or plant parts) to a high
temperature in the absence of air or in the presence of limited amounts
of oxygen or other reagents, catalysts, or solvents. This process is
called destructive distillation. There are three categories of charcoal,
namely, wood charcoal, animal charcoal and sugar charcoal.

Wood charcoal

Wood
charcoal is made by heating wood or other vegetable matter (for example
coconut shells) in the almost complete absence of air. Wood charcoal is
light, porous and is a remarkably good absorbent for liquid or gases (1
cm³ of wood charcoal will absorb 100 cm3 of ammonia gas at 0°C).

Uses of wood charcoal

Because
of its ability to absorb large amounts of gas or liquid, it is used in
gas masks to absorb poisonous gases in air in industrial process to
recover volatile materials from waste gases.
Wood charcoal can be used in metal refining instead of coke.
Wood charcoal is a good source of energy. Thus, it is used as fuel for cooking and heating in homes.

Animal charcoal

Animal
charcoal is made by heating animal bones in the absence of air. Its
main component is calcium phosphate, Ca3(PO4)2, and carbon constitutes
about 10% of the components.

Uses of animal charcoal: It is used in sugar industries to remove the colouring matter from brown sugar to make it white

Sugar charcoal

Sugar
charcoal is a very pure form of carbon, and is made by removing the
elements of water (oxygen and hydrogen) from sugar. This is achieved by
using strong dehydrating agents such as concentrated sulphuric acid or
concentrated nitric acid, which removes water from sugar.

Uses of sugar charcoal: Sugar charcoal is chiefly used in the manufacture of artificial diamonds.

Lampblack/soot/carbon black

This
is produced by burning petroleum or any hydrocarbon in a limited supply
of air (or in chlorine). It can also be produced from the burning of
organic material at home. It is commonly found in the kitchen chimneys,
lamps and on bases of cooking pans and pots.

Uses of lampblack/soot/carbon

It is used in making printers’ ink, shoe polish and carbon papers.
It is an important industrial material in the manufacture of tyres. It is used as a filler material in tyres.

Coke

Coke
is made by destructive distillation of coal. During the process, coal
is obtained as the major product. The other products formed when coal is
destructively distilled are coal gas, coal tar, coal oil, gas carbon
and ammonia liquor.

Uses of coke

Coke is used as a fuel and as a reducing agent in the extraction of iron, lead and zinc. It is also used as a fuel in boilers.
It is used in the manufacture of producer gas and water gas.

Chemical properties of carbon

Carbon burns in excess oxygen to form carbon dioxide C_(s) + O_2(g) → CO_2(g).In insufficient oxygen, carbon monoxide is formed.2C_(s) + O_2(g) → 2CO_(g)
Carbon
has got reducing properties and thus acts as a reducing agent. Carbon
reduces oxides of metals below it in the electrochemical and activity
series to their respective metals. This occurs on strong heating, and
this reaction is used industrially for extraction of metals from their
ores:PbO_(s) + C_(s) → Pb_(s) + CO_(g);Fe₂O_3(s) + 3C_(s) → 2Fe_(s) + 3CO_(g);ZnO_(s) + C_(s) → Zn_(s) + CO_(g).
Sulphur vapours react with red hot carbon to give carbon disulphide. C_(s) + 2S_(g) → CS_2(l)
Carbon dioxide is reduced by red hot carbon to carbon monoxide C_(s) + CO_2(g) → 2CO_(g);This reaction is used in the industrial manufacture of producer gas.

Carbon Dioxide

Preparation of a Dry Sample of Carbon Dioxide Gas in the Laboratory

Prepare a dry sample of carbon dioxide gas in the laboratory

Carbon
dioxide is one of the oxides of carbon. The gas is present in the air
at a level of approximately 0.03% by volume. It is also found dissolved
in water. The gas is one of the by-products of all decaying organic
matter. Without carbon dioxide there is no life on earth. It is used by
all plants in the process of photosynthesis and both plants and animals
evolve carbon dioxide in respiration.

Laboratory preparation of carbon dioxide

Carbon dioxide is prepared in the laboratory by the action of dilute hydrochloride acid on marble (calcium carbonate).

When
dilute hydrochloric acid is poured on marble chips, effervescence
occurs. Dilute hydrochloric acid reacts with the marble chips to give
calcium chloride, water and carbon dioxide.

CaCO_3(s) + 2HCl_(aq)→ CaCl_2(g) + H₂O_(l) + CO_2(g)

The Properties of Carbon Dioxide

Analyse the properties of carbon dioxide

Physical properties

Carbon dioxide is a colourless and odourless gas.
It is denser than air.
When
the gas is cooled to –78°C, it turns straight into the solid (it
sublimes). Sublimation is the change of a gas straight into a solid or
change of a solid straight into a gas. Solid carbon dioxide is called
dry ice. It sublimes when it is heated or exposed to air.
It has a melting point of –199°C and boiling point of –91.5°C.
Carbon dioxide does not support combustion. This is why it is used in fire extinguishers.

Chemical properties

Reaction of carbon dioxide with lime water (Test for carbon dioxide)

When
a little carbon dioxide gas is bubbled into lime water (calcium
hydroxide solution), the solution turns milky. This is due to the
formation of a white precipitate of insoluble calcium carbonate.

Ca(OH)_2(aq) + CO_2(g) → CaCO_3(s) + H₂O_(l)

This
is a confirmatory test for the presence of carbon dioxide. The test
serves to distinguish carbon dioxide from any other gas.

When
excess carbon dioxide is bubbled into the lime water, the white
perceptible dissolves to form a clear solution of soluble calcium
hydrogen carbonate: CaCO_3(s) + H₂O_(l) + CO_2(g)→ Ca(HCO₃)_2(aq)

Barium hydroxide can also be used to test for carbon dioxide as it forms a precipitate of barium carbonate: Ba(OH)_2(aq) + CO_2(g)→ BaCO3(s) + H₂O_(l)

Reaction of carbon dioxide with magnesium

When
a burning magnesium ribbon is lowered into a gas jar containing carbon
dioxide gas, it continues to burn for a short time with a spluttering
flame. A white ash of magnesium oxide and black specks of carbon are
formed. The black specks of carbon can be seen on the sides of the gas
jar.

2Mg_(s) + CO_2(g)→ 2MgO_(s) + C_(s)

This clearly shows that carbon dioxide contains carbon and oxygen.

Reaction of carbon dioxide with water

Carbon
dioxide reacts with water to form a weak carbonic acid. When carbon
dioxide is bubbled into water, it dissolves to form a weakly acidic
solution of carbonic acid:

H₂O_(l) + CO_2(g) ⇔ H₂CO_3(aq)

The
solution turns a blue litmus paper pink. This indicates that the
solution is slightly acidic and hence too weak to turn the blue litmus
paper to red (as strong acids do). The solution has no effect on red
litmus paper.

The Uses of Carbon Dioxide

Explain the uses of carbon dioxide

Uses of Carbon Dioxide include:

Fire extinguisher:
Carbon dioxide is inert (i.e. it does not burn). It is dense than air
and does not support combustion. Hence, it is a very useful
fire-fighting chemical. When applied to fire, it forms a blanket over
the burning material. Thus, it prevents air (oxygen) from reaching the
burning material and therefore, extinguishing the flames.
Manufacture of aerated (fizzy) drinks:
Soda water and mineral water contain carbon dioxide dissolved under
pressure. Because the gas is only slightly soluble, it is bubbled into
these drinks under pressure to make more of it dissolve. When the
bottles are opened, the gas escapes and it causes the “fizzy”.Dissolved
carbon dioxide is responsible for the pleasant taste of soft drinks such
as lemonade, Coca cola, Pepsi cola and other aerated drinks and mineral
waters. Other beverages like beers are also bottled together with
carbon dioxide.
Refrigeration: Carbon dioxide
is used for refrigeration purposes (i.e. in the deep-freezing of foods).
The gas liquefies at ordinary pressure to form dry ice which sublimes
at -78°C. Dry ice is a good refrigerant because it leaves no liquid
after sublimation as is the case with ordinary ice.
Manufacture of sodium carbonate by the Solvary Process:Carbon
dioxide is used in the manufacture of anhydrous sodium carbonate in the
Solvary Process. The sodium carbonate produced is used in the
manufacture of glass.
Manufacture of baking soda:
Carbon dioxide is used in making baking soda (sodium bicarbonate).
Baking soda is prepared by passing carbon dioxide into cold concentrated
sodium hydroxide solution: CO_2(g) + 2NaOH_(aq) → Na₂CO_3(aq) + H₂O_(l).Further
addition of carbon dioxide produces sodium bicarbonate which, at
sufficiently high concentration, will precipitate out of the solution as
a solid: Na₂CO_3(aq) + CO_2(g) + H₂O_(l) → 2NaHCO_3(s)
Yeast and sodium bicarbonate (hydrogencarbonate) are important in the
baking industry. Thus in baking of bread, yeast is added to flour, sugar
and water (forming the dough). In the making of cakes, baking powder (a
mixture of bicarbonate and an acid) is used instead of yeast.
Rain making:
When pieces of dry ice (solid carbon dioxide) are dropped into clouds,
the temperature of the clouds is lowered to such an extent that rain
precipitates out.
Photosynthesis: Plants use carbon dioxide from the air to manufacture their own food through the process of photosynthesis

TOPIC 2: ORGANIC CHEMISTRY

Introduction to Organic Chemistry

Organic
chemistry is the chemistry of carbon compounds. Due to the ability of
carbon to form chains of atoms, and for other atoms or groups of atoms
to be attached to these chains, there are a huge number of carbon
compounds. All organic compounds contain carbon together with one or
more other elements such as hydrogen, oxygen, nitrogen, sulphur and the
halogens.

Normally
every compound of carbon is an organic compound. Even after discovering
that these compounds could be synthesized in the laboratory, the
definition that they are organic (of organic nature, that is,
they originate from living things) has been retained. However, for
conventional and historic reasons, some of the simpler compounds such as
carbon dioxide (CO₂) carbonates, carbon monoxide (CO₃) are usually studied with other non-carbon compounds in Inorganic Chemistry.

Difference between Organic from Inorganic Chemistry

Distinguish organic from inorganic chemistry

Organic
chemistry is the chemistry of carbon compounds. All organic compounds
contain carbon and other elements such as H, O, N, S and the halogens.
Normally every compound of carbon is an organic compound. Examples of
organic compounds/substances are plastics, milk, carbohydrates, lipids,
proteins, sugar and hydrocarbons. Inorganic substances includes table
salt, CO2, diamond, iron and water.

The Importance of Organic Chemistry in Life

Explain the importance of organic chemistry in life

Carbon
is the most unusual atom. It has the ability to join up to itself and
form very long chains of atoms. Without this ability, life on Earth
would not exist because the molecules that make our bodies contain
mostly long chains of carbon atoms.

All
living things contain carbon compounds. Raw materials such as oil and
coal, derived from living things, are also based on carbon. Our modern
society is very much dependent on organic chemistry to make the fuels
and materials that we use in every day of our lives. In particular,
polymers, large molecules obtained from alkanes, have very widespread
use. Without alkanes from crude oil our transport system would be
impossible.

We
also need various fractions obtained from crude oil (petrol, diesel,
kerosene, oil, natural gas, etc.) to run motor vehicles and other
machines to simplify our work and life.

In brief, the compounds obtained from crude oil have thousands of different uses, for example:

some are used as fuel or converted into fuels;
some
are used for making detergents, dyes, drugs, paints, and cosmetics; and
yet some are used for making polyethene, polyvinyl chloride (PVC) and
other plastics.

The Origin of Organic Compounds

Explain the origin of organic compounds

Fossil
fuels were formed in the Earth’s crust from material that was once
living. Coal comes from fossils of plant material. Crude oil and natural
gas are formed from bodies of marine microorganisms. The formation of
these fuels took place over geological periods of time (many millions of
years).

Crude
oil is one of the world’s major natural resources. The oil is the
result of a process that began up to 400 million years ago. Prehistoric
marine organisms died and sunk to the sea bed and were covered by mud
and sand. The change into crude oil and natural gas was brought about by
high pressure, high temperature and bacteria acting over millions of
years. The original organic material broke down into hydrocarbons.
Compression of the mud above the hydrocarbon mixture transformed it into
shale. Then geological movements and pressure changed this shale into
harder rocks, squeezing out the oil and gas. The oil and gas moved
upwards through the porous rocks, moving from high- pressure to
low-pressure conditions. Sometimes they reached the surface, but often
they became trapped by a layer of non porous rock. Reservoirs of oil and
gas were created. These reservoirs are not lakes of oil or pockets of
gas. Instead, the oil or gas is spreadthroughout the pores in coarse
rocks such as sandstone or limestone the same as water is held in a
sponge.

The Fractional Distillation of Crude Oil

Describe the fractional distillation of crude oil

Crude
oil from an oilfield is separated from impurities such as sand and
water and is pumped through pipelines to the refinery. At the refinery,
fractional distillation is used to separate the crude oil into several
fractions, each fraction containing several hydrocarbons which boil
within a certain range of temperatures. These different boiling points
are roughly related to the number of carbon atoms in the hydrocarbon
(Table 2.1)

Separation
of the hydrocarbons takes place in a fractionating column
(fractionating tower). At the start of the refinery process, crude oil
is preheated to a temperature of 350–400^°C and pumped in at
the base of the tower. As it boils, the vapour passes up the tower. It
passes through a series of bubble caps and cools as it rises further up
the column. The different fractions cool and condense at different
temperatures, and therefore at different heights in the column. The
fractions condensing at the different levels are collected on trays.

Thus,
vapour is rising and liquid falling at each level in the tower. As a
result very efficient fractionation occurs. Liquid is taken off at
several different levels, the higher the level, the lower the boiling
point of the fraction removed. Figure 2.2 shows the process of
separation of crude oil into different fractions.

After
fractional distillation, impurities are removed. The commonest impurity
is sulphur, which is removed and used to manufacture sulphuric acid. If
petrol (gasoline) containing sulphur is not purified before it is used
in an internal combustion engine, the exhaust fumes will contain oxides
of sulphur (SO₂ and SO₃). These are poisonous gases and will pollute the environment.

Uses of different petroleum fractions

1.
Natural gas (refinery gas). The gas fractions consist of mainly
methane, ethane, propane, and butane. The methane and ethane are usually
burnt as fuel. The propane and butane are liquefied and are distributed
in high pressure gas cylinders and tanks to be used for lighting and
heating purposes.

2.
Petrol (motor gasoline). It is mainly used as a fuel in internal
combustion engines in motor vehicles. It is also used as a solvent for
grease stains and paints.

3. Naphtha. It is used as a raw material for making many chemicals and plastics.

4.
Kerosene (paraffin). It is used in homes as a fuel in paraffin lamps
and stoves for heating, lighting and cooking food. However, in addition
to its domestic use, it is used as a fuel for jet engines in aeroplanes.
It is also used as a detergent.

5. Diesel oil. It is used as a fuel in diesel engines (e.g. in diesel train engines, tractors, lorries, diesel car engines).

6.
Lubricating oil. It is used to make petroleum jelly (e.g. Vaseline). It
is also used as oil for lubricating moving parts of cars and other
machines.

7. Fuel oil. It is used as a fuel for power stations, ships and factories.

8. Paraffin waxes. They are used to make candles, polishes and waxed papers. They are also used in water proofing and as grease.

9.
Asphalt, bitumen. They are used to make protective coatings for road
surfaces and concrete roof tops, and also as binding agents for roofing
sheets.

Hydrocarbons

Hydrocarbons
are compounds containing hydrogen and carbon and no other element. That
is, a hydrocarbon has the molecular formula CxHy, x and being whole
numbers. For example, methane (CH4) ethane (C2H2) and benzene (C6H6) are
hydrocarbons.

The three Families of Hydrocarbons

Classify the three families of hydrocarbons

In
order to simplify the study of these compounds, chemists have grouped
them into families. The members of each family have characteristic
chemical properties and graded physical properties, such as boiling
point, etc. The main family members of hydrocarbons that we shall study
at this level are the alkanes, the alkenes and the alkynes.

Alkanes

The
members of this group of hydrocarbons are distinguished by possessing
the general molecular formula Cn H2n+2, where is 1, 2, 3, etc., for
successive members of the group. The first member of the series (n = 1)
is methane (CH4) and the second (n = 2) is ethane (C2H6). Both are gases
at room temperature and pressure. This general formula can be used to
work out the formula of any other alkane if we know the number of carbon
atoms it consists of. The following table gives the molecular formula
and name of the first few alkanes, plus an indication of some of their
physical properties.

Table 2.2 Formulae and physical properties of some alkanes

name	Molecular formula	Melting point (oC)	Boiling point (oC)	Density (gcm–3)
Methane	CH4	–183	–162	gas
Ethane	C2H6	–172	–89	gas
Propane	C3H8	–188	–42	gas
Butane	C4H10	–135	–1	gas
Pentane	C5H12	–130	36	0.626
Hexane	C6H14	–95	69	0.659
Heptane	C7H16	–91	98	0.684
Octane	C8H18	–57	126	0.703
Nonane	C9H20	–54	151	0.718
Decane	C10H22	–30	174	0.730

Structure of alkanes

In
the alkanes, all carbon atoms show a covalency of four. Table 2.3 shows
the structures of the first five members of the alkanes.

The Homologous Series of the Three Families of Hydrocarbons

Write the homologous series of the three families of hydrocarbons

A
homologous series is a series of compounds with similar chemical
properties which differ by –CH2. Such series has the following
characteristics:

All members obey the general molecular formula, e.g. for alkanes Cn H2n+2, alkenes Cn H2n and alkynes, Cn H2n-2.
Each member differs, in molecular formula, from the next by CH2, e.g. alkanes are CH4, C2H6, C3H8, and so on.
All members show similar chemical properties.
The
general properties of members change gradually in the same direction
along the series, e.g. in alkanes, boiling points and freezing points
rise with increase in the number of carbon atoms (CH4 – a gas; C5H12 – a
liquid; C20H42 – a solid at ordinary temperatures and pressure).
General
methods of preparation are known which can be applied to any member of
the series. Other homologous series are alcohols, CnH2n+1 and carboxylic
acids, CnH2n+1COOH.

Nomenclature of alkanes

The
term nomenclature means naming. Organic compounds are named using IUPAC
system of nomenclature. IUPAC stands for International Union of Pure
and Applied Chemistry.

Rules of Nomenclature

1. Name the longest unbranched carbon chain. For example, consider the following compound:

In this compound, the longest unbranched chain is

The longest chain is, therefore, butane.

2. Name the functional groups. Functional groups are branches of the main carbon chain. These are sometimes called alkyl groups.

Examples of alkyl groups are as follows:

CH3 – methyl
CH3CH2 – ethyl
CH3CH2CH2 – propyl
Cl – chloro
Br – bromo
I – iodo
F – flouro

3. Give the position(s) of the functional group(s) using lower numbers when possible. Consider the following example:

There are two possibilities of naming the above compound but only one is correct.

2 – methylbutane
3 – methylbutane

The first name is correct because it bears the lowest number (2). The second name is wrong because it bears a large number (3).

4. Names of the functional groups are named in alphabetical order in the final name.

Example

Consider the following compound:

The name of the compound is: 3–bromo–2–choloropentane

Note: We start with bromo and not chloro because b comes before c in alphabet (a, b, c, d …)

5.
If there are identical functional groups, use the prefixes di (2), tri
(3), tetra (4), penta (5), etc. These prefixes do not account for
alphabetical order.

Examples:

6. Commas separate numbers and hyphens separate numbers from words.

The Concept of Isomerism

Explain the concept of isomerism

Isomers are organic compounds with the same molecular formula but different structural formulae.

Isomerism
is the occurrence of two or more compounds with the same molecular
formula but different molecular structures (structural formulae).

Isomers of the same molecular formula have different physical and chemical properties because of structural differences.

Structural Formulae of all Isomers of Hydrocarbons up to Five Carbon Atoms

Write structural formulae of all isomers of hydrocarbons up to five carbon atoms

Example 1

1.
Butane The two isomers of butane, with molecular formula C4H10, are as
shown below. Isomers can be presented either in an open formula
(structure) as follows.

Example 2

Pentane

The three isomers of pentane showa open formula asfollows:

Isomers of Hydrocarbons up to Five Carbon Atoms

Name the isomers of hydrocarbons up to five carbon atoms

Example 3

alkanes showing isomerism

1.
Butane The two isomers of butane, with molecular formula C4H10, are as
shown below. Isomers can be presented either in an open or condensed
formula (structure).

Note:
Remember that CH4, C2H6 and C3H8 have no isomers because they cannot
form alkyl groups, that is, they cannot be branched. Isomerism in
alkanes starts from an alkane with four carbon atoms (butane) in its
structure (C4H10).

2. Pentane

The three isomers of pentane are shown as follows:

Alkenes

Alkanes
are members of a homologous series of hydrocarbons with general
molecular formula Cn H2n where n = 2, 3, 4, etc. The common structural
feature is the presence of a carbon–carbon double bond:

There
is no member of this series for n = 1 since there needs to be at least
two carbon atoms present to have a double bond. Table 2.3 shows the
first five members of alkenes.

Structure of alkenes

Alkanes
are characterized by possessing one or more carbon to carbon double
bond(s) (C=C) in the carbon chain. Alkenes are unsaturated hydrocarbons
because it is possible to break the double bond and add extra atoms to
the parent molecule.

Table 2.4 shows the structures of the first five members of the homologous series of alkenes.

Nomenclature of alkenes

Rules for naming alkenes

Names of alkenes have the ending –ene.
The double bond is given the number of the carbon atom where it begins.
In
addition to the rules discussed earlier, the position of the double
bond must be included in the name .Examples: CH3–CH=CH3 prop-1-ene ,
CH3–CH=CH–CH3 but-2-ene
Double bonds are given the lowest number possible, usually lower than the functional groups.

Counting
from the left to right, the double bond (=) is on carbon number 4 and
from the right to left, it is on carbon number 3. Because the double
bonds should be given the lowest umber possible, the name of the
compound is, therefore, hept-3-ene instead of hept-4-ene.

Isomerism

Alkenes,
like alkanes show branching isomerism, for exampleCH3–CH=CH–CH3
(pent-1-ene) can be branched into (3-methylbut-1-ene) as follows:

They
also show positional isomerism due to different positions of the double
bond in the molecule, for example, pent-1-ene and pent-2-ene are both
the isomers of pentene (C5H10):

CH2=CH–CH2–CH2–CH3 pent-1-ene

CH3–CH=CH–CH2–CH3 pent-2-ene

Other examples include:

Alkynes

These
are homologous series of hydrocarbons with the general formula CnH2n-2.
The common structural feature is the presence of a carbon-carbon triple
bond (CC). There is no member for n = 1 because, as for the alkenes,
there needs to be at least two carbonatoms present to have a triple
bond. Table 2.5 shows the first five members of the alkynes.

Nomenclature of alkynes

Rules for naming alkynes

Alkynes have the ending –yne
Triple bond is given the lowest number possible, usually lower than the functional groups.
As for the alkenes the position of the triple bond must be included in the name.

Isomerism

Alkynes
show branching isomerism and positional isomerism. The alkyne with the
molecular formula, C5H8, shows chain branching isomerism, for instance,

A General Formula to Identify the Families of Hydrocarbons

Apply a general formula to identify the families of hydrocarbons

Activity 1

Apply a general formula to identify the families of hydrocarbons

Properties of Hydrocarbons

The Physical Properties of Lower Hydrocarbons; Alkanes, Alkenes and Alkynes

Explain the physical properties of lower hydrocarbons; alkanes, alkenes and alkynes

Alkanes

The first four alkanes are gases at room temperature, the next twelve (C5–C17) are liquids, and the rest are solids.
The
boiling and melting points of unbranched alkanes increase as the molar
masses increase. The larger the molecule, the higher the boiling point
or melting point.
Branched alkanes have lower boiling points than the unbranched alkanes,
Alkanes are very sparingly soluble in water but they easily dissolve in organic solvents.
In an alkane molecule, each carbon atom forms four single covalent bonds. This means that alkanes are saturated hydrocarbons.

Alkenes

Here again we find that alkenes posses physical properties that arethe same as those of alkanes, i.e.

They are insoluble in water but soluble in organic solvents suchas benzene, ether and chloroform.
They are less dense than water.
The boiling point increases with increase in molecular weighti.e., increase in the number of carbon atoms.

Alkynes

The physical properties of the alkynes are essentially the same as those of the alkanes and alkenes.

The Concept of Saturated and Unsaturated Hydrocarbons

Explain the concept of saturated and unsaturated hydrocarbons

Hydrocarbons
are classified into two distinct categories: saturated and unsaturated.
Saturated hydrocarbons contain only one carbon-carbon single bond.
Unsaturated hydrocarbons contain at least one carbon-carbon double or
triple bond. The unsaturated hydrocarbons are more reactive and they
contain fewer hydrogen atoms bonded to the carbon atoms than saturated
hydrocarbons.

The Chemical Properties of Alkanes, Alkenes and Alkynes

Compare the chemical properties of alkanes, alkenes and alkynes

Alkanes

Alkanes are not as reactive as other hydrocarbons. However, they exhibit the following chemical properties:

Combustion: alkanes burn in a sufficient amount of air to form carbon dioxide and steam.

Cracking: this is a process in which large alkanes are broken down into smaller hydrocarbons.

Halogenation:
under special conditions (sunlight or 300°C), halogens react with
alkanes in a substitution reaction. A substitution reaction is a
reaction in which one atom or group of atoms replaces another.

Alkenes

Alkenes
are much more reactive than alkanes. This is because thedouble bond can
break to form single bonds and add on otheratoms. Because the double
bond allows them to add on moreatoms, alkenes are said to be
unsaturated. The alkanes don‟t havedouble bonds and can‟t add on more
atoms, so they are saturated.

Combustion: alkenes burn in sufficient supply of air (explodes)to form carbon dioxide and steam.

Addition
reactions: Alkenes give a number of addition reactions in which two
hydrogen atoms (or their equivalent) are taken intocombination per
molecule to form a single product. The following are some of the
addition reactions shown by alkenes:

Alkynes

Addition: having carbon–carbon trip bonds, alkynes are unsaturated compounds. They give a number of addition reactions:

With
acidified potassium manganate (VII) solution: At room temperature, with
shaking, ethyne quickly decolourizes this solution (i.e. reduces it)
with formation of ethanedioc acid

Polymerization

Combustion:

Salt formation:

This
is a test for terminal alkynes, i.e. those with C≡C at the end. For
those with C≡C at the centre, no reactions take place e.g.
CH3–CH2–C≡C—CH2–CH3 + AgNO3→ No reaction!

Alcohols

Alcohols
form a homologous series of general molecular formula, CnH2n+2, where n
= 1, 2, 3, etc. for successive members of the group. The hydroxyl
group, OH, is the characteristic functional group of the alcohols.

Preparation of Ethanol in the Laboratory

Prepare ethanol in the laboratory

Laboratory
preparation of an alcohol involves heating a haloalkane with an aqueous
alkali. Ethanol is prepared by heating a mixture of bromoethane and
sodium hydroxide solution. The reaction proceeds as follows:

The ethanol, as indicated in the equation will be in solution. It canbe isolated by distillation.

Ethanol
can also be prepared by fermentation of glucose wherebyyeast produces
an enzyme called zymase which acts as an organiccatalyst to catalyse the
degradation of glucose.

For
substantial fermentation to take place, a mixture of glucose and yeast
is left in a warm place for a considerable length of time.

The equation for the reaction taking place is:

The Homology of Alcohols up to Five Carbon Atoms

Write the homology of alcohols up to five carbon atoms

Alcohols
are named as if they are derived from alkanes by the replacement of
hydrogen atom by the hydroxyl group (–OH). The lowest two members
containing one and two carbon atoms in the molecule, are, respectively,
methanol and ethanol. Their structural formulae are shown below:

Other members of the series are as shown below:

Number of carbons per molecule (n)	Name	Formula
1	Methanol	CH3OH
2	Ethanol	C2H5OH
3	Propanol	C3H7OH
4	Butanol	C4H9OH
5	Pentanol	C5H11OH
6	Hexanol	C6H13OH
7	Heptanol	C7H15OH
8	Octanol	C8H17OH
9	Nonanol	C9H19OH
10	Decanol	C10H21OH

The successive members of the series have molecular formulae which differ by CH2.

Like
other series, alcohols exhibit isomerism. For instance, propanol
(C3H7OH) can have two different structural formulae as exemplified
below:

The first is called propan-1-ol because the OH group is on the first carbon atom.

The second is called propan-2-ol because the OH group is on the second carbon atom.

Structure of all Isomers of Saturated Alcohols up to Five Carbon Atoms

Write structure of all isomers of saturated alcohols up to five carbon atoms

Nomenclature of branched alcohols

Rules:

For branched alcohols, the–OH group is given the lowest number than the alkyl group.
The
alcohol group (–OH) or the alkyl groups can be attached to different
positions of the carbon chain. It is the different positions of these
groups that result to the different names of the alcohol.
The carbon chain may have several branches of different alkyl groups.

Examples:

Note: Both the methyl (–CH3) and hydroxyl (–OH) groups are on carbon number 2. The longest unbranched carbon chain is

Isomers of Alcohols up to Five Carbon Atoms

Name all isomers of alcohols up to five carbon atoms

Alcohols
exhibit chain branching isomerism as with all the other homologous
series. In addition, they show an isomerism based on the position of the
hydroxyl group (–OH) in the molecule.

Alcohols
that have the formula CH3OH and CH3CH2OH have only one possible
structure. They do not have isomers. The –OH group in ethanol, CH3CH2OH,
is always on the last carbon atom.

With
the formula CH3CH2CH2OH, however, the –OH group can either be attached
to the first or second carbon of the carbon chain to give two isomers
namely, CH3CH2CH2OH (propan-1-ol) and

Isomers of other members in the series are as follows:

Butanol: C4H9OH or CH3CH2CH2CH2OH

Isomers:

Pentanol: C5H11OH or CH3CH2CH2CH2CH2OH

ItsIsomers:

The Properties of Alcohol

Describe the properties of alcohol

Ethanol
is the best known and most important of all alcohols. It has several
different uses. Due to its diverse use, its name is used interchangeably
with that of “alcohol” as if it represents all other classes of
alcohols.

Physical properties of ethanol (alcohol)

It is a clear, colourless liquid with a boiling point of 78°C.
It is readily soluble in water and it mixes completely with it (miscible).
It evaporates very easily when exposed to air (it is a very volatile liquid).

Chemical properties of ethanol

Reaction with oxygen (combustion)

It burns well in oxygen giving carbon dioxide and water, and plenty of heat.

This reaction is the basis of its use as a fuel.

Reaction with sodium metal

Ethanol
reacts vigorously with sodium to produce sodium ethoxide and hydrogen.
However, the reaction is not as vigorous as the reaction between sodium
and water.

Methanol reacts similarly to produce sodium methoxide and hydrogen:

Reaction with ethanoic acid (esterification)

Alcohols react with organic acids to form sweet-smelling oily liquids known as esters.

The
symbol, H+, indicates that the reaction takes place in acidified
conditions. Normally dilute sulphuric acid is added as a catalyst for
this esterification reaction.

Reaction with concentrated sulphuric acid

Concentrated sulphuric acid dehydrates ethanol to give ethene. The reaction takes place at 180°C.

Reaction with phosphorus pentachloride

Ethanol reacts vigorously with phosphorus pentachloride to give chloroethane, phosphorus oxychloride and hydrogen chloride.

Hydrogen chloride is evolved as white fumes, and this reaction is used to test for alcohols and other hydroxyl compounds.

Oxidation

Ethanol
is oxidized by strong oxidizing agents such as warm acidified potassium
dichromate or acidified potassium permanganate to ethanoic acid.

The Uses of Alcohol

Explain the uses of alcohol

Alcohol
is used as an alcoholic beverage. Different alcohol drinks contain
different amounts of ethanol. Beer and wines contain 8–13% of ethanol.
Sprits such as whisky, gin, brandy, rum, etc. contain about 35–40% of
ethanol. However, these beverages must only be drunk with moderation as
too much consumption of alcohol can lead to heath problems.

Ethanol as fuel.

Ethanol
burns with a clear flame, giving out quite a lot of heat.C2H5OH(l) +
3O2(g) → 2CO2(g) + 3H2O(g) + heat On a small scale, ethanol can be used
as methylated spirit (ethanol mixed with methanol or other compounds) in
spirit lamps and stoves.
Ethanol produced by fermentation of
sugar from sugar cane has been used as an alternative fuel to gasoline
(petrol), or mixed with gasoline to produce “gasohol”.

Ethanol as a solvent.

Alcohol
is a good solvent, dissolving many compounds that are insoluble in
water. The fact that it evaporates easily makes it a good solvent in
products like glues, paints, varnishes, deodorants (perfumes), and
aftershaves.

It is used in the manufacture of many chemicals

For
example, it is used in making sweet–smelling liquids called esters
which are used as solvents, in food flavourings, and as fragrance in
beauty products (cosmetics).

Ethanol
is also used as an ingredient in iodine tincture (a mixture of alcohol
and iodine) which is largely used in hospitals for treatment of wounds.

The Harmful Effects of Alcohol

Explain the harmful effects of alcohols

In
spite of its widest use as a useful product, ethanol, taken as an
alcoholic beverage, has a lot of detrimental health effects. Some of the
effects of alcohol are as follows:

Even
just one drink impairs coordination and judgement. It is a factor among
the many causes of road accident as it leads to blurred vision.
It
leads to lack of muscular control (e.g. drunken stagger) and ultimately
to coma, the state in which a person is said to be dead–drunk.
Prolonged
consumption of too much alcohol causes liver deterioration (cirrhosis
of the liver) which can cause liver failure and death. Also heavy
drinking eventually damages the muscle tissue of the heart. There may
well be some long–term damage to the brain. All these health effects
accelerate death.
Alcohol is a depressive drug and can be
addictive. Occasional drinking may lead to alcohol addiction, a
condition during which a person is said to be alcoholic. When a drunkard
finds him/herself in such a condition it is very difficult to go
without drinking. This can eventually lead to poverty as a drunkard
spends most of his/her time and money drinking (buzzing).
Alcohol can make someone aggressive. This accounts for many arrests, conviction and jail sentences.
Excessive drinking causes depression and other mental disorders.
It
can lead to gastric ulcers, high blood pressure and cancer of the
mouth, throat, and gullet. People who smoke as well are at greater risk
from thesecancers.

Women
are at greater risk than men from the harmful effects of alcohol. One
reason is that men‟s bodies have higher water content. So alcohol in
their body fluids is more diluted.

Drinking during pregnancy can damage the baby‟s brain and heart, and slow down its growth so that it is born underweight.

Carboxylic Acids

Carboxylic
acids form a homologous series of the general formula CnH2n+1COOH (or
CnH2n+1CO2H), where n = 1, 2, 4, etc. for successive members of the
group. All these acids have the characteristic functional (carboxyl)
group, –COOH, attached to a hydrocarbon chain.

Natural Sources of Organic Acids

Identify natural sources of organic acids

There are various natural sources of organic acids. Some of these sources are:

milk (lactic acid)
citrus fruits (citric acid);
tobacco (nicotinic acid); and
tea (tartaric acid).

The Oxidation of Ethanol to Ethanoic Acid

Explain the oxidation of ethanol to ethanoic acid

When exposed to open air, ethanol is oxidized (by oxygen of the air) to ethanoic acid. The reaction for the process occurs thus:

The Structures of the Homologues of Carboxylic Acids up to Five Carbon Atoms

Write the structures of the homologues of carboxylic acids up to five carbon atoms

Carboxylic
acids are named as if they are derived from alkanes by the replacement
of one hydrogen atom by the –COOH group. The two lowest members,
containing one atom and two carbon atoms respectively, are:

The other members of the homologous series are as shown below:

The
successive members of the series have molecular formulae which differ
by –CH2 It is important to remember that every carboxylic acid molecule
contains the functional group –COOH which is called the carboxyl group.

The Isomers of Carboxylic Acids up to Five Carbon Atoms

Name the isomers of carboxylic acids up to five carbon atoms

Like
other organic compounds, carboxylic acids also exhibit isomerism.
Isomers of carboxylic acid are a result of branching of the hydrocarbon
end (R) rather than the position of the carboxyl group in a molecule of
the carboxylic acid. More isomers of the carboxylic acids can be created
by branching the hydrocarbon end in as many different ways as possible.

Rules

The carbon of the carboxyl group (–COOH) is considered as carbon atom number 1.

Identify the positions of the alkyl group(s) attached to the (longest) acid chain. For example, in a molecule,

the alkyl group is methyl (–CH3) and it is attached to carbon number 2.

Name
the branched alkyl group, followed by the name of the acid to which the
alkyl group is attached. For example, in the case above (rule no.2):

the alkyl group is methyl;
it is attached to carbon number and
the acid to which it is attached is butanoic acid,CH3CH2CH2COOH

Therefore, the name of the compound is 2-methylbutanoic acid.

In
case there occurs more than one alkyl groups in the compound the
prefixes di(2), tri(3), tetra(4) etc (as it was the case in alkanes) may
be used. For, example in the compound

there are two methyl groups, one attached to carbon number 2 and the other to carbon number 3; and
they are both attached to butanoic acid chain.

Therefore, the name of the compound is 2,3-dimethylbutanoic acid.

Isomerism and nomenclature

Branching
isomerism is found in this homologous series. Isomerism in carboxylic
acids begins from butanoic acid, C3H7COOH. The first three members of
the series do not show isomerism because their hydrocarbon ends do not
form branches. The following are the structures and names of the isomers
of carboxylic acids up to five carbon atoms:

Butanoic acid, C3H7COOH or C3H7CO2H or CH3CH2CH2COOH

Isomers:

Pentanoic acid, C4H9COOH

Isomers

The Properties of Carboxylic Acids

Explain the properties of carboxylic acids

Carboxylic acids are weak acids. They are slightly ionized in dilute solutions.

Like
inorganic acids, their solutions contain H+ ions. The presence of H+
ions give the solutions acidic behaviour, that is, their solutions
affect indicators, just like the inorganic acids do.

Neutralization

Like
inorganic acids, carboxylic acids react with metals, alkalis,
carbonates, and hydrogen carbonates to form salts. For example:

Esterification

The
reaction between carboxylic acids and alcohols is called
esterification. The acids will react reversibly with alcohols to form
sweet–smelling esters. Concentrated sulphuric acid is a catalyst for the
reaction.

The
reaction can be reversed to recover an acid and alcohol again by
boiling the products (an ester + water) with a mineral acid (HCl or
H2SO4) or with an aqueous alkali (KOH or NaOH) as a catalyst.

Esters
are manufactured for use as solvents, food flavourings, and fragrance
for perfumes and beauty products. Ethyl ethanoate is just one example of
many esters. The esters usually have strong and pleasant smells. Many
of these compounds occur naturally. They are responsible for the
flavours in fruits and for the scents of flowers. Fats and oils are
naturally occurring esters used for energy storage in plants and
animals. Some of the naturally occurring esters include:

vegetable oils e.g. palm oil, groundnut oil, cashewnut oil, olive oil, sunflower oil, etc; and
animal fats.

All esters contain the functional group,

, where R is any alkyl group.

Preparation of Soap from Animal Fats or Vegetable Oil

Prepare soap from animal fats or vegetable oil

Vegetable
oils are formed from fatty acids and an alcohol called glycerol (also
called glycerine). Fatty acids are carboxylic acids with long chains of
carbon atoms. They are called “fatty” because the long chains repel
water, making them immiscible with water. Glycerol or glycerine (or
propane–1,2,3-triol) has three –OH groups. This is how fatty acids and
glycerol react:

Preparation of soap from oils

Soap
is made by heating animal fats or vegetable oils with sodium hydroxide
solution. The oils react with the solution of sodium hydroxide and break
down to form glycerol and the sodium salts of their fatty acids. These
salts are used as soap. The reaction equation is:

This
process is known as saponification. The soap you buy is made from a
blend of different oils. When soap dissolves in water it ionizes thus:

The cleansing agent in soap is the ion, RCOO–

CHEMISTRY FORM FOUR ALL TOPICS
CHEMISTRY FORM FOUR TOPIC 1 & 2.
CHEMISTRY FORM FOUR TOPIC 3.
CHEMISTRY FORM FOUR TOPIC 4.
CHEMISTRY FORM FOUR TOPIC 5.

O’LEVEL CHEMISTRY
CHEMISTRY STUDY NOTES, FORM FOUR.
CHEMISTRY STUDY NOTES, FORM THREE.
CHEMISTRY STUDY NOTES, FORM TWO.
CHEMISTRY STUDY NOTES FORM ONE.

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CHEMISTRY FORM FOUR STUDY NOTES TOPIC 1: NON METALS AND THEIR COMPOUNDS & TOPIC 2: ORGANIC CHEMISTRY

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