Non metals

Chapter 22: Non-metals
• Describe the physical and chemical properties of
non – metals.
• Describe the industrial preparation of chlorine,
sulphuric acid and ammonia.
• List the uses of the non – metals: carbon, sulphur,
phosphorus, chlorine, nitrogen, silicon and their
compounds.
Learning Outcomes
You should be able to:

Metals
Non-metals
Semi-metal
Non-metals like Hydrogen (H), carbon (C), Nitrogen
(N) are found on the right side of the Periodic Table.
C N
H

Metals Non-Metals
Malleable and ductile Brittle, neither ductile nor
malleable
Good conductors of
electricity and heat
Poor conductors of heat and
electricity except graphite
Lustrous and can be
polished
Non-lustrous and cannot be
polished, except graphite and
iodine which are lustrous non-
metals
Solids at room
temperature, except
mercury
May be solids, liquids or gases at
room temperature
Differences between metals and non-metals

Metals Non-Metals
Strong, tough and have
high tensile strength,
except mercury and zinc
Not strong and have low tensile
strength, except diamond and
carbon fibre
Hard and have high
density, except sodium
and potassium
Generally soft and have low
density, except diamond
High melting and boiling
points
Low melting and boiling points,
except diamond and graphite
Differences between metals and non-metals

Mainly occurs in a combined state in compounds
such as water, acids and many organic substances.
Elemental hydrogen exists as diatomic molecules, H2,
which has the following properties:
•Colourless, odourless and neutral gas
•Non-conductor of electricity
•Low melting point (–259 °C) and low boiling point
(–253 °C)
Hydrogen (H)

Chlorine (Cl)
Highly reactive so it never occurs in the uncombined
form in nature, mainly occurs as sodium chloride or
rock salt
Elemental chlorine exists as diatomic molecules, Cl2,
• Greenish-yellow, poisonous gas
• Non-conductor of electricity
• Denser than air
• Low melting point (–101 °C) and low boiling point (–35 °C)

Oxygen (O)
Nearly half the mass of the Earth’s crust comprises
oxygen in a combined state in compounds such as
water, silicates, oxides and salts. Elemental form
exists in the air, forming 21% of air by volume.
Elemental oxygen exists mainly as diatomic
molecules, O2, which has the following properties:
• Colourless, odourless and neutral gas
• Low melting point (–218 °C) and low boiling point
(–183 °C)

Carbon (C)
Found in the form of diamond (India, South Africa) and
graphite (Sri Lanka), main constituent of numerous
naturally occurring compounds such as coal, mineral
oils, carbonates, organic matter and carbon dioxide gas.
Graphite
• Black
• Soft
• Good electrical conductor
• Very high melting point (3652 °C)
• Very high boiling point (4200 °C)
Diamond
• Colourless, transparent and has a very
high refractive index
• Hardest known natural substance
• Good thermal conductor
• Very high melting point (3550 °C)
• Very high boiling point (4827 °C)

Sulphur (S)
Exists as natural deposits of elemental or native sulphur,
compounds such as sulphur dioxide, zinc blende, pyrite
and gypsum. Present as hydrogen sulphide in petroleum
gases.
• Light yellow powdery solid
• Allotropes – rhombic and monoclinic sulphur
• Rhombic sulphur has a melting point of 113 °C and
boiling point of 445 °C

Nitrogen (N)
Occurs combined in compounds such as sodium nitrate,
calcium nitrate and ammonium sulphate and as an
important constituent of protein. Exists as elements in
79% of the air by volume.
Elemental nitrogen exists as diatomic molecules, N2,
• Colourless, odourless and neutral gas
• Low melting point (–210 °C)
• Low boiling point (–196 °C)

Quick Check 1
List the differences between metals and non-metals.
Solution

Solution to Quick Check 1
Strong, tough and have high tensile strength,
except mercury and zinc
Not strong and have low tensile strength, except
diamond and carbon fibre
Hard and have high density, except sodium
and potassium
Generally soft and have low density, except diamond
High melting and boiling points Low melting and boiling points, except diamond and
graphite
Metals Non-Metals
Malleable and ductile Brittle, neither ductile nor malleable
Good conductors of electricity and heat Poor conductors of heat and electricity except graphite
Lustrous and can be polished Non-lustrous and cannot be polished, except graphite
and Iodine which is lustrous non-metals
Solids at room temperature, except mercury May be solids, liquids or gases at room temperature
Return

Apart from carbon, other non-metals like
oxygen, nitrogen and sulphur have
allotropes as well.
Diatomic
oxygen (O2)
Triatomic
oxygen (O3)
Sulphur has 2 allotropes : rhombic and monoclinic sulphur. Rhombic
sulphur changes to monoclinic sulphur and vice versa at temperatures
above 96 oC and below 96 oC respectively.
Monoclinic sulphurRhombic sulphur

Metals Non-Metals
Have 1–3 electrons in the outermost
shell
Have 4–8 electrons in the outermost
shell
Lose valence electron(s) to form
cations
Gain electron(s) to form anions
or share valence electrons to form
covalent molecules
Electropositive Electronegative
Lose electrons in the valence shell
(oxidised) and make good reducing
agents
Gain electrons from other elements
(reduced) and make good oxidising
agents
Cationic metals are discharged at the
cathode during electrolysis
Anionic non-metals are discharged at
the anode during electrolysis

Metals Non-Metals
Many metals displace hydrogen
from dilute acids
Do not react with dilute acids
and do not displace hydrogen
from dilute acids
Form chlorides which are
electrolytes but non-volatile
Form covalent chlorides which
are non-electrolytes but volatile
Do not combine with hydrogen,
except the reactive elemental metals
which form metal hydrides
React with hydrogen to form
stable covalent hydrides
Form basic oxides, except Cr2O3
(acidic), and Al, Zn, Pb (amphoteric)
Form acidic or neutral oxides

Hydrogen: 2H2(g) + O2(g) → 2H2O(g)
Colourless,
stable liquid
Carbon: C(s) + O2(g) → CO2(g)
Colourless,
odourless gas and
an acidic oxide
In the presence of
insufficient oxygen,
carbon monoxide is
formed instead.

2Cl2(g) + O2(g) → 2Cl2O(g)
Chlorine
Cl2(g) + 2O2(g) → 2ClO2(g)
The chloride oxides are acidic and
highly unstable, reacts with alkalis to
form salt.
Anhydride of hypochlorous
acid. It is an orange gas.
Reddish yellow gas, a
useful oxidising agent.

Sulphur S(s) + O2(g) → SO2(g)
Nitrogen
N2(g) + O2(g) → 2NO (g)
2NO(g) + O2(g) → 2NO2(g)
Acidic oxide, colourless gas with a
pungent smell that dissolves in
water to form acid.
Colourless gas
and a toxic air
pollutant.
Reddish brown toxic gas, which has a
sharp biting odour and is an air pollutant.

Reaction with metals are always a redox reaction where
the non-metal is the oxidising agent (electron acceptor)
and the metal is the reducing agent (electron donor).
Hydrogen H2(g) + 2Na(s) → 2NaH2(s)
Sulphur
Nitrogen N2(g) + 3Ca(s) → Ca3N2(s)
S(S) + Zn(s) → ZnS(s)
Hydrogen + Alkali metals → Metal hydrides
Nitrogen + Reactive metals → Metal nitrides
Sulphur + Metals → Metal sulphides

Chlorine
Oxygen
3Cl2(g) + 2Al(s) → 2AlCl3(s)
O2(g) + 2Ca(s) → 2CaO(s)
3O2(g) + 4Al(s) → 2Al2O3(s)
Aluminium oxide is amphoteric and forms an
impervious layer on the aluminium metal,
protecting it from corrosion.
Chlorine + Metals → Metal chlorides
Oxygen + Metals → Metal oxides

Metals
Non-metals
Semi-metal
Oxidising power increases
Electronegativityincreases
Electronegativity increases
Oxidisingpowerincreases
Nitrogen, oxygen, bromine, chlorine and
fluorine are all good oxidising agents
with fluorine being the strongest.

Across the Periodic Table, the atomic radius decreases, ionisation
energy increases, and thus the electronegativity increases. The increase
in electronegativity also reflects an increase in oxidising power.
(most
(least

A more reactive non-metal would be able to displace a less reactive non-
metal from its salts in aqueous solutions. For example, chlorine replaces
bromine from a solution containing bromide ions.
Cl2(g) + 2Br-(aq) → Br2(g) + 2Cl-(aq)
Chlorine oxidises the bromide ions by
removing an electron from it. Bromine gas
is then liberated from the solution and is
detected by its reddish-brown colour.

There are several methods to collect gases from experiments
depending on the solubility and density of the gas.
Method Suitable gases
Downward displacement of water For gases which are insoluble or slightly soluble in
water
(e.g. oxygen, hydrogen, nitrogen, carbon dioxide)
Downward displacement of air /
upward delivery
For gases which are less dense than air
(e.g. hydrogen, ammonia)
Upward displacement of air /
downward delivery
For gases which are denser than air
(e.g. chlorine, carbon dioxide, sulphur dioxide)
Using a gas syringe (for any gas) For any gas

Moderately active metals, such as zinc, are used to react with mineral
acids to produce hydrogen gas.
The liberated hydrogen gas
is then bubbled through a
solution of concentrated
sulphuric acid.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
Concentrated sulphuric acid
acts as a drying agent as it
removes any water
molecules present in the gas.

Chlorine can be prepared by
removing hydrogen from
hydrochloric acid using an
oxidising agent, e.g.,
manganese dioxide.
Concentrated hydrochloric
acid is added to manganese
dioxide and the mixture is
heated.
MnO2(s) + 4HCl(l) → MnCl2(aq) + 2H2O(l)
+ Cl2(g)
Since chlorine is denser than air, it is
collected using downward delivery.

When hydrogen peroxide is added to manganese dioxide, it
decomposes at room temperature, liberating oxygen gas.
2H2O2(l) → 2H2O(l) + O2(g)
This method of collecting
oxygen is known as
downward displacement of
water.

Carbon dioxide is normally produced by the action of dilute hydrochloric
acid on marble chips.
CaCO3(s) + 2HCl(aq) → CaCl2(aq) +H2O(l) + CO2(g)
The gas can be collected by the downward delivery method.

Sulphur dioxide can be prepared by treating sodium sulphite with
dilute sulphuric acid or hydrochloric acid.
Na2SO3(s) + 2HCl(aq) → 2NaCl (aq) +H2O(l) + SO2(g)
The SO2 produced is passed
through a gas washing bottle
containing concentrated
sulphuric acid. The gas is
dried and collected by
downward delivery.

Ammonia can be prepared by heating an ammonium salt with a strong
base. In the lab, ammonia is prepared by heating a mixture of solid
ammonium chloride and calcium hydroxide.
2NH4Cl(s) + Ca(OH)2(aq) → CaCl2(s) +2H2O(l) + 2NH3(g)
Ammonia gas is
collected by upward
delivery as it is less
dense than air. Since
water is also
produced, the gas is
dried by passing
through a drying
tower filled with a
drying agent.

CH4(g) + H2O(g) → CO(g) + 3H2(g)
Steam is mixed with methane (the main constituent of natural gas),
with a nickel catalyst at a temperature of 1200 °C and 50 atm
pressure to produce hydrogen gas.

• Chlorine is produced by the electrolysis of a concentrated aqueous
solution of sodium chloride known as brine.
• The anode is made of graphite (or titanium) while the cathode is made
of mercury.
• Chloride ions migrate to the anode and are discharged.
• Sodium ions are preferentially discharged to form sodium.
• Sodium combines with the mercury cathode to form sodium amalgam.
• The amalgam is treated with water to produce sodium hydroxide and
hydrogen gas. The mercury is thus freed up for use as cathode again.
• Since chlorine reacts with sodium hydroxide, they must be produced
in separate chambers and kept apart. This is achieved by the use of a
circulating mercury cathode.
The mercury cell process
At the anode: 2Cl- (aq) – 2e- Cl2 (g)
At the cathode: Na+ (aq) + e- Na (l)
Chlorine formed

• The anode and cathode are separated by an ion-
exchange membrane.
• The membrane allows sodium ions and water to
pass through, but not chloride ions.
The membrane cell process
At the anode: 2Cl- (aq) – 2e- Cl2 (g)
At the cathode: 2H+ (aq) + 2e- H2 (g)
Chlorine formed

• Most of the sulphuric acid in the world today is
manufactured by the Contact Process.
• The Contact Process involves the catalytic
oxidation of sulphur dioxide, SO2, to sulphur
trioxide, SO3.
The Contact Process

Sulphur dioxide
(from burning of sulphur or
roasting of iron sulphide)
Sulphur dioxide
+
Excess Air
450 oC
1 – 2 atm
Vanadium(V) oxide
catalyst
Sulphur trioxide
dissolved in conc. sulphuric
acid
OleumSulphuric acid
Water
unreacted
sulphur dioxide

Step 1: Conversion of Sulphur to Sulphur Dioxide
Sulphur dioxide is obtained from the burning of sulphur.
Most of the sulphur is obtained as a by-product of petroleum
refining. Sulphur burns in air to form a colourless pungent
gas called sulphur dioxide:
S(s) + O2(g)  SO2(g)
In some factories, sulphur dioxide is obtained as a by-
product from the roasting of iron pyrites in the extraction of
iron.
4FeS2(s) + 11O2  2Fe2O3(s) + 8SO2(g)

Step2: Conversion of Sulphur Dioxide to Sulphur Trioxide
The sulphur dioxide is mixed with excess air and passed through a
filter to remove impurities and particles before entering the reaction
chamber.
The reaction chamber (converter) contains finely divided vanadium(V)
oxide as a catalyst at a temperature of about 450 °C – 500 °C. The
conversion of sulphur dioxide to sulphur trioxide occurs.
2SO3(g) ∆H = -189 kJmol-1
The heat evolved in the exothermic reaction maintains the
temperature of the catalyst. By using several converters in series and a
slight excess of oxygen, about 98% conversion of sulphur dioxide into
sulphur trioxide is achieved.
2SO2(g) + O2(g)

Step 3: Conversion of Sulphur Trioxide into Oleum
The sulphur trioxide is dissolved in concentrated sulphuric acid to form
a fuming liquid called oleum:
SO3(g) + H2SO4(l)  H2S2O7(l)
Sulphur trioxide is not dissolved directly in water because the reaction is
extremely vigorous and will result in the production of a mist of fine
sulphuric acid particles which is damaging to health.
Step 4: Conversion of Oleum to Sulphuric Acid
The oleum obtained is carefully diluted with the correct amount of water
to form concentrated sulphuric acid:
H2S2O7(l) + H2O(l)  2H2SO4(l)

Optimal Conditions for the Conversion
of Sulphur Dioxide to Sulphur Trioxide
The oxidation of sulphur dioxide to sulphur trioxide is a
reversible reaction, and is therefore affected by the
experimental conditions.
2SO2(g) + O2(g) 2SO3(g) ∆H = -189 kJmol-1
- As it is an exothermic reaction, lower temperatures would
favour the production of more sulphur trioxide and result in a
higher yield. A temperature of 450 °C is favourable.
- Higher pressure would favour the yield of the product.
However, extreme high pressure is not necessary as the yield is
already very high (98%) at a pressure of 1 – 2 atm.
- A vanadium catalyst, vanadium(V) oxide, is also used in this
reaction to speed up the rate of the reaction.

The Haber Process
• The process for manufacturing ammonia from
nitrogen is called the Haber Process, named
after its inventor, Fritz Haber.
The Haber Process
450 – 500 oC

• Hydrogen is obtained by the action of steam on natural
gas or from the cracking of petroleum fractions.
• A mixture of three parts by volume of hydrogen to one
part of nitrogen is compressed to a pressure of about
200 atm and passed over an iron catalyst heated to a
temperature of about 450 oC.
• A yield of 17% – 20 % of ammonia is formed because the
reaction is reversible.
• Under these conditions, the nitrogen reacts with the
hydrogen to form ammonia according to the equation:
2N2(g) + 3H2(g) 2NH3(g) ∆H = -92.4 kJ mol-1

World’s Production of Ammonia
• The annual production of ammonia has been increasing rapidly
since the end of World War II (Fig 24.6).
• 140 million tonnes of ammonia are produced per year.
• Most of the ammonia is used in the manufacturing of fertilisers.
• The use of fertilisers has increased the yield of food crops
which in turn supports a continuing rise in world population.
Fig 22.5 World production of ammonia is rising rapidly

Uses of Ammonia
• The ammonia manufactured in the Haber process is
used in the industry for many purposes.
• Large quantities of ammonia are used in the
manufacture of fertilisers like ammonium nitrate,
ammonium sulphate and urea.
• Ammonia is also used in making nitric acid. This is
done by the catalytic oxidation of ammonia into
nitrogen oxide which is then made into nitric acid. Nitric
acid can be used for making explosives and textiles.
• Ammonia solution is commonly used as a cleaning
agent for dry cleaning and making window cleaners.
Ammonium fertiliser Nitric acid Window cleaner

• The most important use is for making sulphuric acid.
• It is used for bleaching wool and silk as it is a mild
reducing agent and would not damage the material.
• It is used for bleaching wood pulp for paper-making.
• It is used as a preservative for wine and other foodstuff
such as jams, tomato sauces and dried fruits. Sulphur
dioxide kills bacteria in the food and helps to maintain
the appearance of the foodstuff.
Uses of Sulphur Dioxide

Preparations and collectives of
Non-metals
Chapter 22
Quick Check 2
1. Why are some gases prepared in the laboratory passed
through concentrated sulphuric acid?
2. Why is quick lime (calcium oxide) used to dry
ammonia instead of concentrated sulphuric acid?
Solution

Solution to Quick Check 2
1. To remove water vapour from the gas
2. Because alkaline ammonia reacts with sulphuric acid
to form ammonium sulphate.
2NH3 + H2SO4 → (NH4)2SO4
Return

References
Chemistry for CSEC Examinations
by Mike Taylor and Tania Chung
Longman Chemistry for CSEC
by Jim Clark and Ray Oliver

Non metals

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