Chapter 20 – Gases of the Atmosphere
After completing this chapter, you should be able to:
 describe how oxygen, carbon dioxide and simple
oxides of nitrogen are made in the laboratory
 explain how oxygen and nitrogen are extracted by
fractional distillation of liquid air
 describe how carbon dioxide is produced
commercially
 describe a laboratory test for oxygen and carbon
dioxide
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draw the structure of oxygen, nitrogen and carbon
dioxide molecules
recall the properties and uses of nitrogen, oxygen or
carbon dioxide
explain the physical properties, including solubility,
of each gas studied in terms of its structure
describe how oxides of nitrogen are produced in the
atmosphere.
Nitrogen
Nitrogen, like oxygen and carbon dioxide, plays a part in sustaining life in the biosphere. It is also widely used in industry. Nitrogen
makes up almost 80%, by volume, of the atmosphere.
Properties of Nitrogen:
 a colourless and odourless gas
 insoluble in water
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generally unreactive (Its low reactivity can be
related to the strong triple bond between the two
nitrogen atoms in each N2 molecule.)
Uses of nitrogen –
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Plants need a continual supply of nitrogen because nitrogen atoms are present in key compounds in the cells of all living
organisms. However, since they cannot use nitrogen gas directly from the atmosphere plants depend on nitrogen-fixing
bacteria to convert nitrogen into the useful NO3and NH ions. These bacteria live in the soil or in nodules in the roots of
legumes, such as peas, beans, clover and lucerne.
Nitrogen-based fertilisers are compounds of nitrogen added to soils to complement bacterial nitrogen fixation and so
maximise crops. The fertiliser industry is the largest consumer of nitrogen gas, using it to manufacture compounds such as
ammonium sulfate, ammonium nitrate and ammonia.
The low reactivity of nitrogen makes it useful when an unreactive, oxygen-free environment is required, such as in the
electronics industry, glass-making and welding. It is used to remove flammable vapours from oil pipelines and tanks and
the refuelling tubes of aircraft during flight.
The low boiling temperature of liquid nitrogen (196°C) makes it suitable for the storage of frozen eggs, sperm and
embryos. It is also used for freeze-drying foods. Its intense cold allows for very rapid cooling, which results in minimal
damage to the food and so improved appearance and nutritional value.
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Commercial production of nitrogen
Industry values show that nitrogen, like oxygen, is one of the most important
chemicals in terms of quantity produced each year. Nitrogen is produced by the
fractional distillation of liquid air. As liquid air at the bottom of the distillation
column evaporates, nitrogen, which has a lower boiling temperature than the
other gases in the air, rises the most in the column and can be collected.
Oxides of nitrogen
Nitrogen(II) oxide (NO) forms whenever nitrogen and oxygen are mixed
at extremely high temperatures. High temperatures can be generated in the
atmosphere by lightning flashes and meteor trails.
This colourless gas is very difficult to isolate because it reacts readily with oxygen in the atmosphere, at lower temperatures, to form
the brown poisonous gas nitrogen(IV) oxide (NO2).
Laboratory preparation of nitrogen(II) oxide
Nitrogen(II) oxide can be made in the laboratory from the reaction between 50% nitric acid solution and copper. Nitrogen(II) oxide
is only slightly soluble in water, so it can be collected by the displacement of water.
Laboratory preparation of nitrogen(IV) oxide
Nitrogen(IV) oxide can be made in the laboratory from the reaction
between concentrated nitric acid and copper. Since the gas is soluble in
water but much denser than air, it is collected by the upward displacement
of air.
Oxygen
Oxygen is the most abundant element in the Earth’s crust and comprises
about 21% by volume of the atmosphere. The importance of oxygen to
life cannot be overstated. Most of the energy needed by living things is
obtained from respiration, the chemical reaction of oxygen with glucose
produced from the digestion of food. Without a constant and adequate supply of oxygen, most living things could not exist. In
addition, the energy released by reactions of oxygen with fossil fuels is the source of most of the heat and electrical energy needed
to run our world.
Properties of oxygen
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is a colourless, odourless gas
is slightly soluble in water. One litre of water will dissolve 0.03 L of oxygen at 20°C and a pressure of 1 atm. Although the
solubility of oxygen is low, enough of it will dissolve in water to support aquatic life
supports combustion
reacts with most metals and non-metals.
Many oxides of non-metals are acidic oxides and dissolve in water to form acids. The reaction of oxygen with metals in the
environment is known as corrosion. This is one of the disadvantages of the reactivity of oxygen. Many of the oxides of metals are
basic oxides. Basic oxides react with acids. Although oxygen reacts with a wide range of metals and non-metals, it is less reactive
than ozone.
Test for oxygen - If you place a glowing splint of wood or piece of string into a test-tube of oxygen, it will re-ignite. This test is
often used to identify oxygen.
Uses of oxygen
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Manufacture of steel. Steels contain small amounts of carbon, sulfur, silicon and phosphorus impurities. Oxygen is used to
oxidise these impurities and so reduce their amount in the steel.
Manufacture of chemicals. Some chemicals that are produced in large quantities, such as sulfuric acid and methanol, also
require large quantities of oxygen.
Medical uses. Because an adequate supply of oxygen is essential for respiration, oxygen is used in hospitals and by
paramedics to help patients with respiratory problems. Pure oxygen gas or oxygen-enriched air is given to people suffering
from heart or lung problems, such as pneumonia, suffocation and exposure to toxic fumes.
A large number of industrial processes rely on the oxygen in air rather than
pure oxygen:
 Sewage treatment. Air is bubbled through sewage in treatment plants to oxidise organic matter.
 Combustion processes. Large quantities of oxygen are required to support the combustion of fossil fuels in power stations
and in industry.
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When natural gas burns in an excess of oxygen, such as in a Bunsen burner with the air hole open, then the reaction can be
represented by the equation.
If, however, the air hole is closed, the supply of oxygen is
reduced and the combustion of methane is incomplete. The yellow colour of the flame indicates the presence of unburnt carbon.
Carbon monoxide may also be produced in this process. The incomplete combustion of fuels in car engines results in significant
environmental pollution. The amount of oxygen in the atmosphere is critical for the combustion of materials. Experiments show that
an atmosphere with about 21% oxygen is about the safe upper limit for life. If the concentration were, say, 25%, natural fires would
be far more frequent, with even damp vegetation burning easily. If the
concentration were lower, say 15%, natural fires would be very difficult to start.
Commercial production of oxygen
The commercial production of oxygen is one of the largest chemical industries.
Nearly all of the oxygen for commercial use is obtained by separating it from
the other gases in the air by fractional distillation of liquid air
Laboratory preparation of oxygen
Oxygen can be prepared in the laboratory by the decomposition of hydrogen
peroxide solution (H2O2), using manganese dioxide (MnO2) as a catalyst.
Ozone
Ozone (O3) is an allotrope of oxygen. Ozone is an extremely reactive gas and a very strong oxidant. Its ability to react with the very
stable CFCs and halons is evidence of this high reactivity. It is this high reactivity that makes ozone a useful sterilising agent. Ozone
destroys viruses and bacteria in water supplies and large swimming pools. The ozone is converted into harmless oxygen. One of the
problems associated with this method of sterilisation, however, is that the water supply is not protected from recontamination since
no ozone remains in the water after the initial treatment.
Carbon dioxide
Although carbon dioxide occupies only about 0.035% of the total volume of the atmosphere, this concentration is sufficient to
sustain life on the Earth. The total amount of carbon dioxide in the oceans, lakes and rivers is about 20 times that in the atmosphere.
Dissolved carbon dioxide supports photosynthesis by aquatic plants, which in turn support aquatic animals.
Properties of carbon dioxide
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is a colourless and odourless gas
will not support combustion unless the fire is hot enough to break the carbon–oxygen double bonds. Because of the strong
double bonds, carbon dioxide is a stable molecule. Magnesium is one of the few substances that will continue to burn in
carbon dioxide.
is denser than air. This accounts for the ‘blanketing’ effect carbon dioxide has on a fire.
when cooled at atmospheric pressure, it does not form a liquid but instead turns into a white solid at 78°C. This solid form
of carbon dioxide is known as dry ice.
is slightly soluble in water - 1.5 g of carbon dioxide will dissolve in 1 litre of water. As the solubility increases under
pressure, carbon dioxide is suit able for use in carbonated drinks, such as lemonade and mineral water. When the pressure
is released by opening a bottle, the solubility is reduced and so the extra carbon dioxide, which has dissolved under
pressure, bubbles out of the solution.
forms slightly acidic solutions. When carbon dioxide dissolves in water, most of it remains as CO 2(aq). A small amount,
however, will react with water to form carbonic acid (H2CO3): CO2(aq) H2O(l)
H2CO3(aq) For this reason,
carbon dioxide is classified as an acidic oxide. It will react with bases to form salts, many of which are insoluble.
Test for carbon dioxide
If you blow through a straw into a test-tube of limewater (calcium hydroxide), a white precipitate forms. This precipitate is calcium
carbonate: CO2(g) Ca(OH)2(aq)
CaCO3(s) H2O(l) This test is often used to identify carbon dioxide.
If you continue to pass carbon dioxide into the test-tube, the precipitate
dissolves to give a colourless solution. Soluble calcium hydrogen carbonate
has been formed: CaCO3(s) CO2(g) H2O(l)
Ca2(aq) 2HCO3-(aq)
Commercial production of carbon dioxide
In industry, carbon dioxide can be obtained as a by-product of the fermentation of sugars to alcohol:
2C2H5OH(aq) 2CO2(g)
C6H12O6(aq)
glucose
ethanol
Carbon dioxide can also be obtained by heating calcium carbonate in the form of limestone in a lime kiln:
CaCO3(s)
CaO(s) CO2(g)
calcium oxide
(quick lime)
Laboratory preparation of carbon dioxide
Carbon dioxide is usually prepared in the laboratory by the reaction of dilute hydrochloric acid with calcium carbonate in the form
of marble chips. The gas can be collected over water. The equation for the reaction is: CaCO (s) 2HCl(aq)
CaCl (aq) CO (g) H O(l)
3
2
2
2
Uses of carbon dioxide
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Carbon dioxide is the ultimate source of carbon for all living things. Through photosynthesis, in the presence of
chlorophyll, a plant pigment and sunlight, green plants convert carbon dioxide and water to glucose and oxygen. The
glucose then undergoes condensation polymerisation to form starches and cellulose. In addition to providing food for
ruminants, cellulose molecules are the structural unit in cotton fibres. Cellulose molecules provide plants with support.
Carbon dioxide is used to make carbonated drinks. Carbon dioxide is dissolved in the drink at high pressure. It comes out
of solution as bubbles when the bottle or can is opened.
Solid carbon dioxide, known as dry ice, is used as a refrigerant. Dry ice sublimes at 78°C. It changes directly from a solid
to a gas without going through a liquid phase.
Carbon dioxide is used in fire extinguishers. It is denser than air and so blankets the flames, reducing the access of oxygen
to the fuel.
Chapter 20 Questions:
Q2, 3, 4, 5, 6, 8, 9, 12, 15, 16, 19