Download Topic 16 Some non-metals and their compounds notes

Hydrogen and Nitrogen
Hydrogen is the most common element in the universe. It is a colourless, odourless
gas which is less dense than air and extremely reactive. Unfortunately it is so light
that there is none of it in our atmosphere. We can, however, produce hydrogen
fairly easily by the electrolysis of water (H2O is broken down into the elements
hydrogen and oxygen) or by the reaction of methane with steam. Nitrogen on the
other hand makes up 78% of the air around us and can be obtained by fractional
distillation of air or by removing the 21% oxygen from the air by combustion.
Nitrogen is also a colourless, odourless gas which is very unreactive.
Hydrogen and nitrogen can react together to make ammonia
The Haber Process
Air contains 78% nitrogen, and the nitrogen can be used to manufacture several
important chemicals, including nitrogen-based fertilisers. Nitrogen-based
fertilisers are important in agriculture for increasing the yields of crops. The
starting compound for the manufacture of nitrogen-based fertilisers is ammonia,
NH3. Ammonia can also be used to manufacture nitric acid.
The process for the manufacture of ammonia
was devised by the German physical chemist,
Fritz Haber and was put into practice by the
German chemical engineer, Carl Bosch.
Haber
Bosch
The raw materials for the process are nitrogen and hydrogen. Nitrogen is obtained
by the fractional distillation of liquid air. Hydrogen is obtained from the reaction
of natural gas (methane) and steam.
H2O(g) + CH4(g)
3H2(g) + CO(g)
The purified gases are passed over a catalyst of iron at a high temperature and
pressure. About 15% of the nitrogen and hydrogen react to produce ammonia. The
gases are cooled to liquefy the ammonia, which is removed. Then the nitrogen and
hydrogen are re-circulated through the reactor.
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The reaction between nitrogen and hydrogen is reversible, this means that
ammonia also breaks down again into nitrogen and hydrogen.
3H2(g) + N2(g)
2NH3(g)
H = -92 kJ.mol-1
Economics of the Haber Process
The optimum conditions for this equilibrium reaction, which produce a reasonable
yield of ammonia quickly, are decided as follows:
Pressure
An increase in pressure will shift the equilibrium to the right hand side, because
this has fewer moles of gas. Therefore, the yield of ammonia will increase.
In addition to its effect on the equilibrium position, increasing the pressure will:

increase the reaction rate, establishing equilibrium faster. Since the
molecules of gas are forced closer together, there will be more collisions in a
given time.

increase the cost. As pressure increases, the cost of industrial plant
increases because, for example, vessels and pipes have to be thicker.
A compromise has to be struck between increasing rate and yield, which are
beneficial to the economics of the process, and increasing cost, which is not. In
practice, a pressure of about 250 atmospheres is used.
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Temperature
An increase in temperature will displace the equilibrium to the left hand side,
because the backward reaction is endothermic. Therefore, the yield of ammonia
will decrease.
In addition to its effect on the equilibrium position, increasing the temperature
will:

increase the reaction rate, establishing equilibrium faster. Since the
molecules of gas have more energy, a higher proportion will have energy
greater than or equal to the activation energy and undergo successful
collisions.

increase the cost. Higher temperatures incur greater fuel costs.
A compromise has to be struck between increasing rate, which is beneficial to the
economics of the process, and increasing cost & decreasing yield, which are not. In
practice, a temperature of about 450oC is used.
Catalyst
Addition of a catalyst has no effect on the equilibrium position and therefore does not
affect the yield of ammonia.
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A catalyst provides an alternative reaction path with a lower activation energy.
Therefore, it increases the reaction rate, establishing equilibrium faster.
Using a catalyst leads to a small increase in cost.
In practice, an iron catalyst containing promoters is used.
USES OF AMMONIA
Ammonia is a colourless gas with a very characteristic smell. It is an alkaline gas
which turns damp red litmus paper blue.
Manufacture of Fertilisers
Ammonia is used in the production of fertilisers, which increase the yield of crops.
The most common compounds used as fertilisers are:

ammonium sulphate, which is made by the neutralisation of ammonia with
sulfuric acid
2NH3 + H2SO4
(NH4)2SO4

ammonium nitrate, which is made by the neutralisation of ammonia with nitric
acid
NH3 + HNO3
NH4NO3

urea, which is made by the reaction of ammonia and carbon dioxide
2NH3 + CO2
CO(NH2)2 + H2O
PROBLEMS WITH THE USE OF FERTILISERS
Although artificial fertilisers have done much to increase world food production
and reduce starvation, there are problems associated with their use. For example,
the run-off of fertilisers from fields into water courses is a problem because:

Too much fertiliser in water can lead to excessive algal growth
(eutrophication). The algae form a thick mat on the surface of the water,
which blocks out light and kills aquatic plants. This reduces oxygen levels in
the water. When the mat starts to die and falls to the bottom, it removes
oxygen from the water as it decomposes. As a consequence of this, fish are
killed.
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
The presence of nitrate in drinking water can cause health problems,
particularly blue baby syndrome. In new-born babies, the presence of nitrate
affects their ability to carry oxygen in the blood.
Some changes to farming practice could reduce these problems. Farmers could, for
example:
 not apply fertilisers when rain is expected, so that they are not washed into
water courses before the plants have had a chance to absorb them
 apply fertilisers several times in small doses rather than in fewer larger
doses
Attempts are being made to reduce the need for artificial fertilisers by developing
strains of the most important food crops which can fix their own nitrogen.
Manufacture of Nitric Acid
Nitric acid is made from ammonia in the Ostwald Process. Ammonia is oxidised by
oxygen in the air in the presence of a hot platinum catalyst:
4NH3 + 5O2
4NO + 6H2O
The nitrogen oxide (NO) formed in the above reaction is cooled and reacted with
water and more oxygen to form nitric acid.
4NO + 2H2O + 3O2
4HNO3
Apart from its use in the manufacture of ammonium nitrate, nitric acid is also used
in the manufacture of dyes, explosives and nylon
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SULFUR
You will see sulfur written as sulphur in
textbooks – it is the same thing. Sulfur is a
yellow non-metallic element. As an element it
exits as small molecules of 8 sulfur atoms
arranged in a ring. The molecule can be written as
S8.
Like carbon, which has two allotropic forms
(diamond and graphite), sulfur has two allotropes
called rhombic and monoclinic. Allotropes are
different forms of the same element
Most sulfur is obtained from compounds found in natural gas. Natural gas contains
hydrogen sulfide, which smells of rotten eggs,
2H2S + O2

2S + 2H2O
Sulfur can also be extracted from underground sulfur beds. In this case, hot water
is pumped underground and the sulfur melts and is carried back to the surface.
Sulfur is used to produce sulfuric acid though it can be used to toughen rubber
tyres (vulcanizing), make drugs, cosmetics and even to modify the properties of
concrete. Sulfur burns in air to produce sulfur dioxide which causes acid rain.
Sulfur dioxide is a colourless gas with a choking smell
S + O2

SO2
The SO2 produced dissolves in water (clouds) to for an acidic solution.
SO2 + H2O

H2SO3
Sulfur dioxide does, however, have its uses. It can be used to bleach wool and
wood pulp for making paper. It also kills bacteria and can therefore be used to
preserve food.
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THE MANUFACTURE OF SULFURIC ACID
(The Contact Process)
Concentrated sulfuric acid is manufactured from the raw materials sulfur, water
and air.
Step 1:
Sulfur is burned in oxygen to form sulfur dioxide.
S(s) + O2(g)
SO2(g)
H = -297 kJ/mol
This reaction, which is exothermic, is a redox reaction. Sulfur is oxidized from an
oxidation state of 0 to an oxidation state of +4.
Step 2:
The sulfur dioxide is then mixed with oxygen, and the mixture is passed over a
catalyst of vanadium(V) oxide (V2O5) at a temperature of 450oC and a pressure of
2 atmospheres. This reaction, in which sulfur trioxide is formed, is exothermic. It
is also a redox reaction, because sulfur is oxidized from an oxidation state of +4 to
an oxidation state of +6.
2SO2(g) + O2(g)
2SO3(g)
H = -192 kJ/mol
The optimum conditions for this equilibrium reaction are decided as follows:
Pressure
An increase in pressure will shift the equilibrium to the right hand side, because
this has fewer moles of gas. Therefore, the yield of sulfur(VI) oxide will increase.
When pure oxygen is used, the volume reduction is from 3 volumes to 2. If
air is used instead of oxygen, one volume of oxygen will be accompanied by 4
volumes of unreactive nitrogen. Therefore the volume reduction will be smaller
(from 7 volumes to 6), and pressure will have a smaller effect on the equilibrium
position. Even at low pressures, the equilibrium lies well to the right hand side, so
the use of expensive high pressures is not justified. A compromise has to be struck
between increasing rate and yield, which are beneficial to the economics of the
process, and increasing cost, which is not. In practice, a pressure of about 2
atmospheres is used.
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Temperature
An increase in temperature will displace the equilibrium to the left hand side,
because the backward reaction is endothermic. Therefore, the yield of sulfur(VI)
oxide will decrease.
In addition to its effect on the equilibrium position, increasing the temperature
will:
 increase the reaction rate
 increase the cost
A compromise is needed between increasing rate, which helps the economics of the
process, and increasing cost & decreasing yield, which do not. In practice, a
temperature of about 450oC is used.
Catalyst
Addition of a catalyst has no effect on the equilibrium position and therefore does
not affect the yield of sulfur(VI) oxide.
Using a catalyst leads to a small increase in cost. In practice, a catalyst of
vanadium(V) oxide is used.
Step 3:
The sulfur trioxide is dissolved in concentrated sulfuric acid to form fuming
sulfuric acid (oleum, H2S2O7). Water is then carefully added to the oleum to
produce concentrated sulfuric acid (98%).
This reaction is very exothermic and has to be cooled by passing cold water
through pipes. The heated water can then be used to generate electricity for the
plant.
USES OF SULFURIC ACID
Sulfuric acid is used:
 as car battery acid
 in the manufacture of fertilizers
 in the manufacture of soaps & detergents
 Concentrated sulfuric acid is also a dehydrating agent– ie it can ‘remove
water’
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Carbon
The element carbon is found in the earth’s crust in two allotropic forms. Diamond is
a clear, hard solid which does not conduct electricity, whilst graphite is a black,
slippery solid which does conduct electricity. There are millions of compounds
containing carbon. Most living things are made up of compounds containing carbon.
Carbon cycle
The carbon cycle describes the movement of carbon through the
environment. Carbon exists in the ground, in the sea, in plants and animals, and in
the air. The natural processes shown in the picture below can release carbon
dioxide into the atmosphere (by combustion, decay and respiration) or remove
carbon dioxide from the atmosphere (by dissolving in the sea and by
photosynthesis. Carbon can also be "locked away" in the ground as a fossil fuel or
by forming carbonate rocks.
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Carbon dioxide
Carbon dioxide is released into the atmosphere when we burn fossil fuels or
we thermally decompose carbonate rocks (see p12). Both of these processes take
place daily on an industrial scale.
CH4 + 2O2
CaCO3
2H2O + CO2


CaO + CO2
Carbon dioxide is heavier than air and does not support combustion. It is therefore
used in fire extinguishers.
Carbon monoxide
Carbon monoxide (CO) is also produced when fuels burn in too little oxygen.
We call this incomplete combustion. Carbon monoxide is an odourless, poisonous gas
which binds to haemoglobin in blood cells preventing them from carrying oxygen
around the body. Carbon monoxide poisoning therefore is very often fatal.
CH4 + 1½ O2

2H2O + CO
Carbonates
You have already met decomposing carbonates in Topic 13 (behaviour of
metals). Carbonates decompose when heated to form oxides and CO2. Only the
carbonates of the most reactive metals (Na2CO3 and K2CO3) as so stable that they
do not decompose when heated.
Organic compounds
Methane (CH4) is a colourless, odourless gas made up of the elements carbon
and hydrogen. It is found trapped below impermeable rock under the sea floor
where it formed millions of years ago. Some animals produce methane during the
digestion of plant matter (cows and sheep) due to the presence of certain bacteria
in their gut. Methane is also a greenhouse gas which contributes to global warming.
There are millions of organic compounds such as proteins, carbohydrates,
fats, pharmaceutical drugs, fuels and plastics which are all made of carbon bonded
to other elements. You have already met these in Topics 17 & 18 (organic
chemistry).
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Global Warming
Carbon dioxide, methane, water vapour and nitrous oxides absorb infra-red
radiation in the same range that is given off by the Earth at night. More carbon
dioxide in the atmosphere means more heat is absorbed, and less can escape at
night into space. The overall effect is that the Earth receives more heat in the day
than it can lose at night, so the temperature rises. The same thing happens in a
greenhouse where glass instead of carbon dioxide is used to prevent the heat
escaping at night. So, carbon dioxide, along with methane and water vapour have
been called greenhouse gases.
Climate change
Carbon dioxide is causing the global temperature of the Earth to rise.
This rise in temperature affects the climate all over the planet in ways which are
hard to predict. Some areas may become wetter or dryer, other areas may become
hotter. The main (predictable) directions of wind and ocean currents on the planet
may slow down, speed up, or change direction. This could have catastrophic
consequences for life on planet Earth, including rising sea levels because of ice
melting at the poles and more extreme weather because of more convection in the
hotter wetter atmosphere. Finally, some species may not be able to adapt to the
different climate and may die out altogether.
If we are causing global warming by burning fossil fuels we can do something
about it. Cutting back on the burning of fossil fuels will help as will switching to
more environmentally friendly sources of power such as wind, tidal and solar power.
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LIMESTONE
Limestone is formed when sea creatures die and fall to the bottom of the sea floor. Over
millions of years this material builds up and is pressed together to form limestone rock.
These rocks are slowly pushed upwards and sea levels have fallen exposing the
limestone. Limestone contains mostly calcium carbonate (CaCO3) which is an important
raw material. Limestone is obtained by quarrying.
Malham cove (limestone)
Dover cliffs (chalk)
marble quarry
Thermal Decomposition
When some substances are heated, they decompose (break down) into simpler
substances. This process is called THERMAL DECOMPOSITION. Calcium carbonate
undergoes thermal decomposition when heated in a kiln to about 1100 oC, forming
calcium oxide (quicklime) and carbon dioxide.
Other metal carbonates decompose in a similar way. The lower the metal is in the
reactivity series, the lower the temperature at which this decomposition takes place.
CaCO3(s)
CaO(s) + CO2(g)
Hydration
Quicklime reacts with water in a very exothermic reaction, forming calcium hydroxide
(slaked lime).
CaO(s) + H2O(l)
Ca(OH)2(s)
The Limestone Cycle
H2O
LIMESTONE
(calcium carbonate)
CO2
CO2
QUICKLIME
(calcium oxide)
SLAKED LIME
(calcium hydroxide)
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H2O
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Uses of Calcium Carbonate & Calcium Hydroxide
Building Material
Limestone is used as a building material. A form of limestone, called Portland stone,
was used in the construction of St Paul’s Cathedral in London. Limestone is a hard rock,
but it erodes badly in polluted urban environments, because it dissolves readily in acids.
Marble is used in sculpture and as a building material.
St. Paul’s Cathedral
marble bust
Removal of Acidity
Powdered limestone is used to neutralise acidity in lakes and in light, sandy soils.
Slaked lime (calcium hydroxide) is used to neutralise acidity in heavy soils and in
drinking water.
Cement Manufacture
Cement is made by roasting
 powdered limestone
 powdered clay
 gypsum in a rotary kiln.

Cement is then used to make concrete.
Flue Gas Desulfurisation is a process used to prevent SO2 escaping into the
atmosphere. Waste gases containing SO2 are passed through a flue (chimney)
containing calcium hydroxide (Ca(OH)2) which absorbs the SO2 producing calcium
sulphite (CaSO3) and water.
Ca(OH) 2
+
SO2
CaSO3 +
H2O
This can easily be oxidised to to make hydrated calcium sulphate (CaSO 4. 2H2O), also
known as gypsum, which is used to make the cement and also plasterboard for the
building industry.
2CaSO3 + O2 +
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4H2O
CaSO4. 2H2O (gypsum)
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Glass Manufacture
Glass is made by heating a mixture of:
 limestone
 sand
 sodium carbonate
Iron production
Limestone is added to the Blast Furnace, in the production of iron, to remove acidic
impurities present in iron ore.
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

NEUTRALISATION


GLASS
heavy soils
drinking water
SLAKED
LIME
lakes
light soils
add water
QUICKLIME
heat + sand &
sodium carbonate
heat
LIMESTONE
CHALK
MARBLE
CALCIUM CARBONATE
heat + powdered clay
BUILDING
removal of impurities
CEMENT
IRON & STEEL
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CONCRETE
mix with
 water
 sand
 crushed rock
Topic 16 Non-metals and their compounds
Summary questions
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(d)
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