9.6 Nitrogen and Sulphur
Learning outcomes:
Candidates should be able to:
(a) Explain the lack of reactivity of nitrogen
(b) Describe:
(i) The formation and structure of the ammonium ion
(ii) The displacement of ammonia from its salts
(c) describe the Haber process for the manufacture of ammonia from its
elements, giving essential operating conditions, and interpret these
conditions (qualitatively) in terms of the principles of kinetics and equilibria
(d) Understand the industrial importance of ammonia and nitrogen compounds
derived from ammonia
(e) Understand the environmental consequences of the controlled use of nitrate
fertilisers
(f) Understand and explain the occurrence and catalytic removal of oxides of
nitrogen
(g) Explain why atmospheric oxides of nitrogen are pollutants, including their
catalytic role in the oxidation of atmospheric sulphur dioxide
(h) Describe the formation of atmospheric sulphur dioxide from the combustion of
sulphur contaminated carbonaceous fuels
(i) State the role of sulphur dioxide in the formation of acid-rain and describe the
main environmental consequences of acid-rain
(j) State the main details of the Contact process for sulphuric acid production
(k) Understand the industrial importance of sulphuric acid
(l) Describe the use of sulphur dioxide in food preservation
(a) Explain the lack of reactivity of nitrogen:
Group V
Nitrogen is a gas making up approximately
78% of the atmosphere with a triple bond
energy of 945 KJmol-1 hence explaining the
lack of reactivity compared to other non-metals
:N
N:
This results in a large activation energy, Ea
It does react with the s-block metals, e.g.
3Mg(s) + N2(g)
Mg3N2(s)
Magnesium nitride
(b) Describe:
The formation and structure of the ammonium ion
The displacement of ammonia from its salts
Ammonia, is produced commercially by the Haber Process.
N2(g) + 3H2(g)
2NH3(g)
DH = -92 KJmol-1
It can be acidified to give the ammonium ion:
NH3(aq) + H+(aq)
NH4+(aq)
Ammonia is alkaline (a weak base) in aqueous solution and a typical reaction
would be using dilute hydrochloric acid:
NH3(aq) + HCl(aq)
NH4+(aq) + Cl-(aq)
Ammonia is clearly a base as it is a proton acceptor according to the BronstedLowry definition of acids and bases.
Ammonia dissociates in water to a small extent (1%)
NH3(aq) + H20(l)
NH4+(aq) + OH-(aq)
Can you remember the conjugate pairs (of acids and bases)?
Structure of the ammonium ion
It has a tetrahedral shape similar to methane (CH4)
In the gas phase, ammonia gas can react with hydrogen chloride:
NH3(g) + HCl(g)
NH4Cl(s)
NH3(g)
HCl(g)
cotton wool soaked
in conc. ammonia
cotton wool soaked
in conc. HCl(aq)
Displacement of ammonia from its salts
If we add sodium hydroxide to ammonium chloride, the characteristic
smell of ammonia is produced:
NaOH(aq) + NH4Cl(aq)
NH3(aq) + NaCl(aq) + H2O(l)
The same is also true of organic amine salts:
NaOH(aq) + CH3CH2NH3Cl(aq)
CH3CH2NH2(aq) + NaCl(aq) + H2O(l)
The group I metal hydroxide is clearly the stronger base!
(c) Describe the Haber process for the manufacture of ammonia from its
elements, giving essential operating conditions, and interpret these
conditions (qualitatively) in terms of the principles of kinetics and equilibria
Ammonia, NH3, is produced commercially by the Haber Process.
N2(g) + 3H2(g)
Fe
460 oC
200 atm
2NH3(g)
DH = -92 KJmol-1
Haber Process
N2(g) + 3H2(g)
Iron
catalyst
2NH3(g)
Haber Process (contd.)
•
•
•
•
•
Operating Conditions
N2 and H2 are pumped into a chamber.
The pre-heated gases are passed through a heating coil to
the catalyst bed.
The catalyst bed is kept at 460 - 550 C under high
pressure.
The product gas stream (containing N2, H2 and NH3) is
passed over a cooler to a refrigeration unit.
In the refrigeration unit, ammonia liquefies, but not N2 or
H2.
Haber Process (contd.)
1. Kinetics
A catalyst (iron) is used to speed up the rate of
reaction and to lower the high activation energy in
breaking the N2 triple bond. However, if the
temperature is too high, it begins to get destroyed and
must be changed more regularly.
Effect of temperature and pressure
on Haber Process
An example:
In the Haber Process for the production of ammonia,
based on the reversible reaction:
N2(g) + 3H2(g)
2NH3(g)
it is observed that:
• As the total pressure increases, the amount of ammonia
present at equilibrium increases.
• As the temperature decreases, the amount of ammonia at
equilibrium increases.
Haber Process
N2(g) + 3H2(g)
2NH3(g)
(d) Understand the industrial importance of ammonia and nitrogen
compounds derived from ammonia
Most of the ammonia produced is used to make nitric acid and fertilizers.
1. Ammonia and oxygen are passed over a hot platinum catalyst:
4NH3(g) + 5O2(g)
Pt
4NO(g) + 6H2O(l) DH = -909 KJmol-1
Conditions: 1100K and 4-10 atm. pressure
About 95% of the ammonia is converted into nitrogen monoxide
2. The gases are cooled and mixed with more air. Then under pressure,
they are passed up a tower that has water trickling down it and nitric acid
is produced:
4NO(g) + 3O2(g) + 2H2O(l)
Overall:
NH3(g) + 2O2(g)
4HNO3(aq)
HNO3(aq) + H2O(l)
The conversion of NO into nitric acid actually occurs in three stages:
a. 2NO(g) + O2(g)
b. 2NO2(g)
DH = -115 KJmol-1
2NO2(g)
DH = -58 KJmol-1
N2O4(g)
c. 3N2O4(g) + 2H2O(l)
4HNO3(aq)
+
2NO(g)
Commercial uses of nitric acid:
Nitric acid is used to make fertilizers, explosives and the polyamide, Nylon
Fertilizers 化肥
Ammonium salts are used as fertilizers because they contain nitrogen in a form
that plants can use. The fertilizer Nitram is ammonium nitrate and is made from
a solution of ammonia in water and nitric acid in an acid-base reaction:
NH3(aq) + HNO3(aq)
NH4NO3(aq)
Similar neutralization reactions with phosphoric(V) and sulphuric acids:
2NH3(aq) + H2SO4(aq)
(NH4)2SO4(aq)
(e) Understand the environmental consequences of the controlled use of
nitrate fertilisers
Ammonium nitrate is a good fertilizer: it contains 35% by mass of
nitrogen all of which is available to plants as it dissolves easily.
However, some nitrate ions are washed into rivers and streams which
encourages rapid growth of algae. However, when algae die they are
decomposed by bacteria and this uses up oxygen dissolved in water.
The fish begin to suffocate and die from a lack of oxygen.
Nitrates also get into drinking water as they are not purified and at
high levels can cause problems for small babies by affecting haemoglobin
(used in respiration and transport of O2 in the blood).
In today’s lesson…..
(d) Understand the industrial importance nitrogen compounds derived
from ammonia (revision from yesterday)
(e) Understand the environmental consequences of the controlled
use of nitrate fertilisers (revision from yesterday)
(f) Understand and explain the occurrence and catalytic removal of
oxides of nitrogen (new)
(g) Explain why atmospheric oxides of nitrogen are pollutants, including
their catalytic role in the oxidation of atmospheric sulphur dioxide (new)
We can test some of our knowledge in a past-paper style question on manufacture
of nitric acid
We can test some of our knowledge in a past-paper style question on catalytic
removal of NO in car engines
(d) Understand the industrial importance of ammonia and nitrogen
compounds derived from ammonia
Most of the ammonia produced is used to make nitric acid and fertilizers.
1. Ammonia and oxygen are passed over a hot platinum catalyst:
4NH3(g) + 5O2(g)
Pt
4NO(g) + 6H2O(l) DH = -909 KJmol-1
Conditions: 1100K and 4-10 atm. pressure
About 95% of the ammonia is converted into nitrogen monoxide
2. The gases are cooled and mixed with more air. Then under pressure,
they are passed up a tower that has water trickling down it and nitric acid
is produced:
4NO(g) + 3O2(g) + 2H2O(l)
Overall:
NH3(g) + 2O2(g)
4HNO3(aq)
HNO3(aq) + H2O(l)
The conversion of NO into nitric acid actually occurs in three stages:
a. 2NO(g) + O2(g)
b. 2NO2(g)
DH = -115 KJmol-1
2NO2(g)
DH = -58 KJmol-1
N2O4(g)
c. 3N2O4(g) + 2H2O(l)
4HNO3(aq)
+
2NO(g)
Properties of nitric(v) acid
Concentrated nitric acid (66%) is a strong acid:
HNO3(aq)
H+(aq) + NO3-(aq)
It is also an oxidising agent and reacts with most metals to form nitrates
and nitrogen oxides:
Cu(s) + 2HNO3(aq)
Cu(NO3)2(aq) + H2O(l) + NO2(g)
Commercial uses of nitric acid:
Nitric acid is used to make fertilizers, explosives and the polyamide, Nylon
Fertilisers 化肥
Ammonium salts are used as fertilisers because they contain nitrogen in a form
that plants can use. The fertiliser Nitram is ammonium nitrate and is made from
a solution of ammonia in water and nitric acid in an acid-base reaction:
NH3(aq) + HNO3(aq)
NH4NO3(aq)
Similar neutralization reactions with phosphoric(V) and sulphuric acids:
2NH3(aq) + H2SO4(aq)
(NH4)2SO4(aq)
(e) Understand the environmental consequences of the controlled use of
nitrate fertilisers
Ammonium nitrate is a good fertiliser: it contains 35% by mass of
nitrogen all of which is available to plants as it dissolves easily.
However, some nitrate ions are washed into rivers and streams which
encourages rapid growth of algae 藻类. However, when algae die they
are decomposed by bacteria and this uses up oxygen dissolved in water.
The fish begin to suffocate and die from a lack of oxygen.
Nitrates also get into drinking water as they are not purified and at
high levels can cause problems for small babies by affecting haemoglobin
血红蛋白 (used in respiration and transport of O2 in the blood).
(f) Understand and explain the occurrence and catalytic removal
of oxides of nitrogen
The combustion engines in cars produce a number of polluting
gases, including a small amount of NOx:
Carbon monoxide, which is poisonous.
Nitrogen oxides (NO, NO2 and N2O4) which cause acid rain and
help to destroy the ozone (O3) layer in the upper atmosphere.
N2(g)+ O2(g)
2NO(g) : the excess heat in an engine produces NO
Unburnt hydrocarbons may cause cancer.
Unburnt hydrocarbons react with nitrogen oxides in sunlight to form
smogs (containing ozone) at low levels in the atmosphere, especially
in large hot cities.
Catalytic removal of nitrogen oxides
These poisonous gases may be removed by the use of a catalytic converter
In car exhaust systems. It is honeycomb in shape to provide a large
surface area and is coated with platinum and rhodium metal catalysts:
Carbon monoxide + nitrogen oxide
Hydrocarbons + nitrogen oxide
CO
ON NO-CO
N2 CO2 CO2
Pt or Rh
1100K
Inside a catalytic converter
nitrogen + carbon dioxide
nitrogen + carbon dioxide + water
(g) Explain why atmospheric oxides of nitrogen are pollutants,
including their catalytic role in the oxidation of atmospheric
sulphur dioxide
Nitrogen oxides (NO, NO2 and N2O4) which cause acid rain and
help to destroy the ozone (O3) layer in the upper atmosphere.
In more detail:
Nitric oxide (NO) reacts with oxygen to produce nitrogen dioxide:
2NO(g) + O2(g)
2NO2(g)
NO2 dissolves in rain to form a mixture of nitric(v) and nitric(III) acids:
2NO2(g) + H2O(l)
HNO3(aq) + HNO2(aq)
What type of redox reaction is this? disproportionation
These acids are help to form acid rain; they are corrosive and damage construction
materials that are attacked by acids especially marble (大理石), so some ancient
monuments lose their features.
NO and NO2 have a catalytic role in the conversion of atmospheric sulphur
dioxide to sulphur trioxide:
2NO(g) + O2(g)
SO2(g) + NO2(g)
2NO2(g)
SO3(g) + NO(g)
Sulphur trioxide dissolves in rain water to form sulphuric acid and aquatic
life may be destroyed in lakes and it also destroys limestone and marble.
Trinitrotoluene
Did you know?
The explosive TNT uses fuming
nitric acid (98%) in its manufacture
To summarize…..
(d) Understand the industrial importance nitrogen compounds derived
from ammonia
(e) Understand the environmental consequences of the controlled
use of nitrate fertilisers
(f) Understand and explain the occurrence and catalytic removal of
oxides of nitrogen
(g) Explain why atmospheric oxides of nitrogen are pollutants, including their
catalytic role in the oxidation of atmospheric sulphur dioxide
In tomorrow’s lesson we will study sulphur chemistry, in particular
the contact process for formation of SO3 and H2SO4
In today’s lesson…..
(f) Understand and explain the occurrence and catalytic removal of
oxides of nitrogen (revision)
(g) Explain why atmospheric oxides of nitrogen are pollutants, including
their catalytic role in the oxidation of atmospheric sulphur dioxide (rev.)
(h) Describe the formation of atmospheric sulphur dioxide from
the combustion of sulphur contaminated carbonaceous fuels (new)
(i) State the role of sulphur dioxide in the formation of acid-rain and
describe the main environmental consequences of acid-rain (new)
(j) State the main details of the contact process for
sulphuric acid production (also see equilibria hand-out) (new)
We will test our knowledge using an exam style question to check our revision
of nitrogen chemistry, nitric oxides and catalytic role with SO2
(f) Understand and explain the occurrence and catalytic removal
of oxides of nitrogen
The combustion engines in cars produce a number of polluting
gases, including a small amount of NOx:
Carbon monoxide, which is poisonous.
Nitrogen oxides (NO, NO2 and N2O4) which cause acid rain and
help to destroy the ozone (O3) layer in the upper atmosphere.
N2(g)+ O2(g)
2NO(g) : the excess heat in an engine produces NO
Unburnt hydrocarbons may cause cancer.
Unburnt hydrocarbons react with nitrogen oxides in sunlight to form
smogs (containing ozone) at low levels in the atmosphere, especially
in large hot cities.
Catalytic removal of nitrogen oxides
These poisonous gases may be removed by the use of a catalytic converter
In car exhaust systems. It is honeycomb in shape to provide a large
surface area and is coated with platinum and rhodium metal catalysts:
Carbon monoxide + nitrogen oxide
Hydrocarbons + nitrogen oxide
CO
ON NO-CO
N2 CO2 CO2
Pt or Rh
1100K
Inside a catalytic converter
nitrogen + carbon dioxide
nitrogen + carbon dioxide + water
(g) Explain why atmospheric oxides of nitrogen are pollutants,
including their catalytic role in the oxidation of atmospheric
sulphur dioxide
Nitrogen oxides (NO, NO2 and N2O4) which cause acid rain and
help to destroy the ozone (O3) layer in the upper atmosphere.
In more detail:
Nitric oxide (NO) reacts with oxygen to produce nitrogen dioxide:
2NO(g) + O2(g)
2NO2(g)
NO2 dissolves in rain to form a mixture of nitric(v) and nitric(III) acids:
2NO2(g) + H2O(l)
HNO3(aq) + HNO2(aq)
What type of redox reaction is this? disproportionation
These acids are help to form acid rain; they are corrosive and damage construction
materials that are attacked by acids especially marble (大理石), so some ancient
monuments lose their features.
NO and NO2 have a catalytic role in the conversion of atmospheric sulphur
dioxide to sulphur trioxide:
2NO(g) + O2(g)
SO2(g) + NO2(g)
2NO2(g)
SO3(g) + NO(g)
Sulphur trioxide dissolves in rain water to form sulphuric acid and aquatic
life may be destroyed in lakes and it also destroys limestone and marble.
Trinitrotoluene
Did you know?
The explosive TNT uses fuming
nitric acid (98%) in its manufacture
Sulphur (Group VI)
(h) Describe the formation of atmospheric sulphur dioxide from
the combustion of sulphur contaminated carbonaceous fuels
Coal, an impure form of graphite is used as a fossil fuel and is still used in
boilers to heat water. It was more popular in England more than one
hundred years ago in the Victorian Days.
However, small amounts of sulphur are oxidized to corrosive and foul
smelling sulphur dioxide, which contributes to pollution:
S(s) + O2(g)
SO2(g)
structure of SO2
(i) State the role of sulphur dioxide in the formation of acid-rain and describe
the main environmental consequences of acid-rain
Sulphur dioxide released into the atmosphere dissolves in rain and forms
sulphuric (IV) acid more commonly known as sulphurous acid; some SO2 comes
from natural sources such as volcanoes, where hot sulphur reacts with O2:
SO2(g) + H2O(l)
H2SO3(aq)
formation of acid-rain
This is corrosive and damages construction materials that are attacked by acids
especially marble (大理石), so some ancient monuments lose their features.
Acid-rain also effects forests and kills some plants, algae and fish in natural lakes,
and the more acid-rain produced the worse the effect.
(j) State the main details of the contact process for
sulphuric acid production (see equilibria hand-out)
Sulphuric acid, H2SO4 is produced from sulphur, water and oxygen by the
Contact process.
1. Sulphur is burnt to sulphur dioxide:
S(s) + O2(g)
SO2(g)
2. Sulphur dioxide is converted to suphur trioxide by reaction with more
oxygen using a catalyst of vanadium(v) oxide:
V2O5
710K
2SO2(g) + O2(g)
2SO3(g)
DH = -197 KJmol-1
Effect of temperature and pressure
on the contact process
2SO2(g) + O2(g)
V2O5
710K
2SO3(g)
DH = 197 KJmol-1
 Temperature
Le Chateliers principle predicts that for the exothermic forward reaction
cooling will produce the maximum yield (95%): the equilibrium is shifted
to the right. The SO3 is removed by dissolving it in conc. H2SO4. However,
the rate of reaction to reach equilibrium increases with temperature increase.
The yield can be increased to 99.5% by passing the remaining reactant gases
over the hot catalyst (on a fourth catalyst bed).
 Pressure
Although increasing the pressure will shift the equilbrium to the right
as fewer gas molecules are produced, there are more operational energy
costs and in practice the process is run at near atmospheric pressure (only an
extra 0.5% conversion achieved!)
A flow diagram
Catalytic
chamber
Air
Sulphur
burner
Air
Purifier
and
drier
Purifier
and
drier
Heat
exchanger
drySO2+air
SO3
water
Absorption
oleum H2SO4 store
tower
98% H2SO4
c.H2SO4
3. Formation of H2SO4
The absorbed SO3 reacts with the water in sulphuric acid to form more
sulphuric acid so that the acid becomes even more concentrated:
SO3(g) + H2O(l)
H2SO4(l)
SO3(g) + H2SO4(l)
H2S2O7(l)
H2S2O7(l) + H2O(l)
oleum
Properties of H2SO4: Viscous liquid with hydrogen bonding
1. Strong acid:
H2SO4(l)
HSO4-(aq)
H2O
H+(aq) + HSO4-(aq)
H+(aq) + SO42-(aq)
2. Dehydrating agent:
CuSO4.5H2O(s)
H2SO4
CuSO4(s) + 5H2O(l)
3. Oxidizing agent
8HI(g) + H2SO4(l)
4H2O(l) + H2S(g) + 4I2(s)
2H2SO4(l)
Practice questions
Sulphuric acid is manufactured from sulphur in four stages
0
√ S(s) + O2(g)
√ 2SO2(g) + O2(g)
+VI
+IV
SO2(g)
+VI
2SO3(g)
+VI
SO3(g) + H2SO4(l)
H2S2O7(l)
H2S2O7(l) + H2O(l)
2H2SO4(l)
(a) Give the oxidation number of sulphur in each species.
(b) Which two steps are redox reactions?
To summarize what we have learnt today…..
(h) Describe the formation of atmospheric sulphur dioxide from
the combustion of sulphur contaminated carbonaceous fuels
(i) State the role of sulphur dioxide in the formation of acid-rain and
describe the main environmental consequences of acid-rain
(j) State the main details of the Contact process for
sulphuric acid production
Tomorrow we will understand:
(k) Understand the industrial importance of sulphuric acid
(l) Describe the use of sulphur dioxide in food preservation
Revise nitrogen and sulphur
In today’s lesson…..
(h) Describe the formation of atmospheric sulphur dioxide from
the combustion of sulphur contaminated carbonaceous fuels (rev.)
(i) State the role of sulphur dioxide in the formation of acid-rain and
describe the main environmental consequences of acid-rain (rev.)
(j) State the main details of the contact process for
sulphuric acid production (also see equilibria hand-out) (rev.)
(k) Understand the industrial importance of sulphuric acid
(new)
(l) Describe the use of sulphur dioxide in food preservation (new)
Keywords: SO2, acid-rain, SO3, Contact Process, H2SO4,
SO2-food preservation
Sulphur (Group VI)
(h) Describe the formation of atmospheric sulphur dioxide from
the combustion of sulphur contaminated carbonaceous fuels
Coal, an impure form of graphite is used as a fossil fuel and is still used in
boilers to heat water. It was more popular in England more than one
hundred years ago in the Victorian Days.
However, small amounts of sulphur are oxidized to corrosive and foul
smelling sulphur dioxide, which contributes to pollution:
S(s) + O2(g)
SO2(g)
structure of SO2
(i) State the role of sulphur dioxide in the formation of acid-rain and describe
the main environmental consequences of acid-rain
Sulphur dioxide released into the atmosphere dissolves in rain and forms
sulphuric (IV) acid more commonly known as sulphurous acid; some SO2 comes
from natural sources such as volcanoes, where hot sulphur reacts with O2:
SO2(g) + H2O(l)
H2SO3(aq)
formation of acid-rain
This is corrosive and damages construction materials that are attacked by acids
especially marble (大理石), so some ancient monuments lose their features.
Acid-rain also effects forests and kills some plants, algae and fish in natural lakes,
and the more acid-rain produced the worse the effect.
(j) State the main details of the contact process for
sulphuric acid production (see equilibria hand-out)
Sulphuric acid, H2SO4 is produced from sulphur, water and oxygen by the
Contact process.
1. Sulphur is burnt to sulphur dioxide:
S(s) + O2(g)
SO2(g)
2. Sulphur dioxide is converted to suphur trioxide by reaction with more
oxygen using a catalyst of vanadium(v) oxide:
V2O5
710K
2SO2(g) + O2(g)
2SO3(g)
DH = -197 KJmol-1
Effect of temperature and pressure
on the contact process
2SO2(g) + O2(g)
V2O5
710K
2SO3(g)
DH = 197 KJmol-1
 Temperature
Le Chateliers principle predicts that for the exothermic forward reaction
cooling will produce the maximum yield (95%): the equilibrium is shifted
to the right. The SO3 is removed by dissolving it in conc. H2SO4. However,
the rate of reaction to reach equilibrium increases with temperature increase.
The yield can be increased to 99.5% by passing the remaining reactant gases
over the hot catalyst (on a fourth catalyst bed).
 Pressure
Although increasing the pressure will shift the equilbrium to the right
as fewer gas molecules are produced, there are more operational energy
costs and in practice the process is run at near atmospheric pressure (only an
extra 0.5% conversion achieved!)
A flow diagram
Catalytic
chamber
Air
Sulphur
burner
Air
Purifier
and
drier
Purifier
and
drier
Heat
exchanger
drySO2+air
SO3
water
Absorption
oleum H2SO4 store
tower
98% H2SO4
c.H2SO4
3. Formation of H2SO4
The absorbed SO3 reacts with the water in sulphuric acid to form more
sulphuric acid so that the acid becomes even more concentrated:
SO3(g) + H2O(l)
H2SO4(l)
SO3(g) + H2SO4(l)
H2S2O7(l)
H2S2O7(l) + H2O(l)
oleum
Properties of H2SO4: Viscous liquid with hydrogen bonding
1. Strong acid:
H2SO4(l)
HSO4-(aq)
H2O
H+(aq) + HSO4-(aq)
H+(aq) + SO42-(aq)
2. Dehydrating agent:
CuSO4.5H2O(s)
H2SO4
CuSO4(s) + 5H2O(l)
3. Oxidizing agent
8HI(g) + H2SO4(l)
4H2O(l) + H2S(g) + 4I2(s)
2H2SO4(l)
(k) Understand the industrial importance of sulphuric acid
Sulphuric acid is an important industrially produced chemical.
Its main uses are: Paints and pigments 染料 涂料
 Fertilisers 化肥
 Fibres and dyes 县委 染料
 Soaps and detergents 香皂 洗涤剂
 Cleaning metals 除锈剂
 Chemicals and plastics 化学品 塑料
 Tanning leather 皮革
(l) Describe the use of sulphur dioxide in food preservation (new)
Sulphur dioxide is poisonous to bacteria and is used as a food preservative
(E220 on food labels in the U.K.) especially in the preparation of fruit juices
and soft drinks and to kill bacteria in wine-making as it is slightly soluble in
water.
SO2(g) + H2O(l)
H2SO3(aq)
H+(aq) + HSO3-(aq)
Sodium metabisulphite (Na2S2O5) is another preservative (E222) and liberates
SO2(g) upon acidification.
In today’s lesson we learnt:
(k) Understand the industrial importance of sulphuric acid
(l) Describe the use of sulphur dioxide in food preservation