Gases

The atmosphere is the layer of air that covers the Earth all around. It is a mixture of many gases.

About 99% of the air is nitrogen (78%) and oxygen (21%). The last 1% is mostly argon, with a little carbon dioxide and traces of neon, helium, methane, krypton, hydrogen, ozone and others. Air also holds water vapor, from almost none in a dry winter to about 4% on a humid monsoon day.

Gases are not only in the air. They are locked up in compounds in plants, animals, water, food and rocks: oxygen in water and in limestone, hydrogen in water and in every sugar, nitrogen in every protein. By mass, about 96% of the human body is just four elements, oxygen (about 65%), carbon (18%), hydrogen (10%) and nitrogen (3%); three of them are gases as elements. Living things use gases in respiration and photosynthesis, and to build compounds.

We get the gases we need from the atmosphere, or by making them in factories. For scientific study, we prepare them in the laboratory.

Hydrogen is the lightest gas there is, and the most common element in the universe: the Sun and the stars are mostly hydrogen. On Earth there is very little free hydrogen gas; it is found in compounds: water, acids, bases, hydrocarbons (like the methane in biogas), carbohydrates, proteins and fats.

Hydrogen has three isotopes: atoms of the same element, with the same atomic number but different mass numbers.

Isotope	Protons	Neutrons	Mass number
protium, ¹₁H	1	0	1
deuterium, ²₁H	1	1	2
tritium, ³₁H	1	2	3

Almost all hydrogen is protium; deuterium is rare, and tritium is radioactive.

In 1766 the English chemist Henry Cavendish made hydrogen by pouring acids on metals, and showed that it is a separate gas that burns; he called it “inflammable air”. The French chemist Antoine Lavoisier named it hydrogen, from the Greek for “water-former”, because it burns in oxygen to make water.

Symbol	Atomic number	Atomic mass	Molecular formula	Density
H	1	1	H₂	0.09 kg/m³ (air: about 1.29 kg/m³)

Laboratory preparation of hydrogen gas

Active metals like zinc and magnesium displace hydrogen from dilute acids. So in the laboratory, hydrogen is made by the reaction of granulated zinc with dilute hydrochloric acid (dilute sulphuric acid also works):

$$\text{Zn(s)} + 2\text{HCl(aq)} \rightarrow \text{ZnCl}_2\text{(aq)} + \text{H}_2\uparrow$$

(The arrow ↑ after H₂ means it escapes as a gas.) Very pure zinc reacts only slowly; impure (commercial, granulated) zinc reacts much faster, because the traces of other metals, such as copper, in it speed up the reaction.

Precautions

1. Wash all the apparatus well.
2. Use impure (granulated) zinc, so that the reaction is quick.
3. Make every joint airtight, so that no gas escapes.
4. Use dilute acid. Concentrated sulphuric acid and nitric acid give other gases instead of hydrogen (sulphur dioxide and oxides of nitrogen).
5. The lower end of the thistle funnel must dip under the acid, so the gas can’t escape through it; the end of the delivery tube inside the bottle must stay above the acid, so only gas goes into it.
6. Fill the gas jar completely with water, with no air bubble left in it.
7. Keep flames away: hydrogen mixed with air explodes when lit.

Test for hydrogen gas

Bring a burning splint (or a lit candle) to the mouth of a test tube or jar of hydrogen. The hydrogen burns with a “pop” sound and a pale blue flame, and the splint itself goes out if pushed inside the jar. The pop is a tiny explosion of hydrogen mixing with the oxygen of the air: 2H₂ + O₂ → 2H₂O. No other common gas does this.

Physical properties of hydrogen gas

1. It is a colorless, odorless and tasteless gas.
2. It is the lightest gas: about 14 times lighter than air.
3. It hardly dissolves in water.
4. It has no effect on litmus paper: it is a neutral gas.
5. It turns into a liquid at −253 °C and freezes to a solid at −259 °C.

Chemical properties of hydrogen gas

1. With oxygen it forms water. A mixture of hydrogen and oxygen explodes when lit:

$$2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$$
2. With some non-metals, at high temperatures:

$$\text{C} + 2\text{H}_2 \rightarrow \text{CH}_4 \ \text{(methane)}$$
$$\text{S} + \text{H}_2 \rightarrow \text{H}_2\text{S} \ \text{(hydrogen sulphide)}$$
3. With very active metals (sodium, potassium, calcium), heated, it forms solid hydrides, which react with water to give hydrogen back:

$$2\text{Na} + \text{H}_2 \rightarrow 2\text{NaH} \ \text{(sodium hydride)}$$
$$2\text{K} + \text{H}_2 \rightarrow 2\text{KH} \ \text{(potassium hydride)}$$
$$\text{Ca} + \text{H}_2 \rightarrow \text{CaH}_2 \ \text{(calcium hydride)}$$
4. It takes oxygen away from the oxides of less reactive metals. When hydrogen is passed over the red-hot oxide of a metal like iron, copper or lead, it combines with the oxygen to form water and leaves the free metal:

$$\text{Fe}_2\text{O}_3 + 3\text{H}_2 \rightarrow 2\text{Fe} + 3\text{H}_2\text{O}$$
$$\text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O}$$
$$\text{PbO} + \text{H}_2 \rightarrow \text{Pb} + \text{H}_2\text{O}$$
With copper oxide you can watch it happen: the black powder turns into shiny, pinkish-brown copper.

The removal of oxygen from a compound (or the addition of hydrogen to it) is called reduction, and these are reduction reactions. Hydrogen is a reducing agent.

Uses of hydrogen gas

1. To get metals from their oxides (reduction).
2. To make ammonia, by combining it with nitrogen in the Haber process; most ammonia becomes fertilizer.
3. To turn liquid vegetable oils into solid fat, vanaspati ghee, by hydrogenation. (Partly hydrogenated fats contain trans fats, which are bad for the heart, so many countries now limit them.)
4. As rocket fuel: liquid hydrogen burning with liquid oxygen gives a huge amount of energy, which launches rockets.
5. For welding and cutting metals: hydrogen burning in oxygen gives the oxy-hydrogen flame, at about 2,800 °C.
6. To fill weather balloons, and as a clean fuel in some cars and buses, whose only exhaust is water.

Living things need oxygen for respiration. It makes up about 21% of the air by volume, and is found in compounds everywhere: in organic compounds such as carbohydrates, proteins, fats and wood, and in inorganic compounds such as calcium carbonate (limestone), silica (sand) and water. By mass, nearly half the Earth’s crust is oxygen.

In 1774 the English scientist Joseph Priestley made oxygen by heating red mercuric oxide (HgO) with sunlight focused through a large lens; the Swedish chemist Carl Scheele had made it a little earlier, but published later. Lavoisier showed what it really is, and named it oxygen.

Symbol	Atomic number	Atomic mass	Molecular formula	Density
O	8	16	O₂	1.43 kg/m³

Laboratory preparation of oxygen gas

A. By heating

Some compounds rich in oxygen give it off when heated. In the laboratory, oxygen is made by heating potassium chlorate (KClO₃) with a little manganese dioxide (MnO₂) as a catalyst:

$$2\text{KClO}_3 \xrightarrow[\Delta]{\text{MnO}_2} 2\text{KCl} + 3\text{O}_2\uparrow$$

The manganese dioxide is not used up; it makes the potassium chlorate give its oxygen at a much lower temperature, and more steadily.

Precautions

1. Clamp the hard-glass test tube with its mouth slightly lower than its closed end, as in the figure. Any moisture in the chemicals that condenses near the mouth then runs out, instead of running back onto the hot glass and cracking it.
2. Make every joint airtight.
3. Fill the gas jar completely with water, leaving no air.
4. When you finish, take the delivery tube out of the water before you stop heating. Otherwise, as the tube cools, water is sucked back into the hot test tube and cracks it.

B. Without heating

Oxygen can also be made at room temperature by the decomposition of hydrogen peroxide (H₂O₂), with manganese dioxide as a catalyst:

$$2\text{H}_2\text{O}_2 \xrightarrow{\text{MnO}_2} 2\text{H}_2\text{O} + \text{O}_2\uparrow$$

Precautions

1. Make every joint airtight.
2. Pour in enough hydrogen peroxide to cover the end of the thistle funnel, so no oxygen escapes through it.
3. Use pure manganese dioxide.

Test for oxygen gas

Put a glowing splint (a wooden splint lit and then blown out, so that only its tip glows red) into a jar of the gas. If it bursts into flame again, the gas is oxygen. A burning candle or splint put into oxygen burns much more brightly, with a bigger flame, than in air.

Physical properties of oxygen gas

1. It is a colorless, odorless and tasteless gas.
2. It is slightly heavier than air.
3. It is slightly soluble in water: just enough for fish to breathe.
4. It has no effect on litmus paper: it is a neutral gas.
5. It turns into a pale blue liquid at −183 °C and freezes at about −219 °C.
6. It doesn’t burn itself, but it supports burning: things burn much more fiercely in it than in air.

Chemical properties of oxygen gas

1. Metals burn in oxygen to form metal oxides. Metal + oxygen → metal oxide:

$$4\text{Na} + \text{O}_2 \rightarrow 2\text{Na}_2\text{O} \ \text{(sodium oxide)}$$
$$4\text{K} + \text{O}_2 \rightarrow 2\text{K}_2\text{O} \ \text{(potassium oxide)}$$
$$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO} \ \text{(magnesium oxide)}$$

2. Non-metals burn in oxygen to form non-metal oxides. Non-metal + oxygen → oxide of the non-metal:

$$2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O} \ \text{(water)}$$
$$\text{C} + \text{O}_2 \rightarrow \text{CO}_2 \ \text{(carbon dioxide)}$$
$$\text{S} + \text{O}_2 \rightarrow \text{SO}_2 \ \text{(sulphur dioxide)}$$
$$4\text{P} + 5\text{O}_2 \rightarrow 2\text{P}_2\text{O}_5 \ \text{(phosphorus pentoxide)}$$
3. Hydrocarbons burn in oxygen to give carbon dioxide, water and energy. Hydrocarbon + oxygen → carbon dioxide + water + energy:

$$\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} + \text{energy} \ \text{(methane)}$$
$$\text{C}_2\text{H}_4 + 3\text{O}_2 \rightarrow 2\text{CO}_2 + 2\text{H}_2\text{O} + \text{energy} \ \text{(ethene)}$$
4. Glucose reacts with oxygen to give carbon dioxide, water and energy. This is what happens in respiration in living things:

$$\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy}$$
5. Alcohol burns in oxygen, giving carbon dioxide, water and energy:

$$\text{C}_2\text{H}_5\text{OH} + 3\text{O}_2 \rightarrow 2\text{CO}_2 + 3\text{H}_2\text{O} + \text{energy} \ \text{(ethyl alcohol)}$$

Uses of oxygen gas

1. All living things (except a few microbes) use oxygen for respiration, to get energy from food.
2. For welding and cutting metals, with the oxy-acetylene flame (about 3,200 °C) or the oxy-hydrogen flame.
3. In hospitals, for patients who have difficulty breathing.
4. In cylinders, where there isn’t enough of it to breathe: for mountaineers, miners, divers, astronauts and in submarines.
5. In making steel: oxygen blown through molten iron burns away its impurities.
6. With liquid hydrogen or kerosene, as the oxidizer in rocket engines.

Ozone is a form of oxygen whose molecule has three oxygen atoms, O₃ (ordinary oxygen is O₂). It is a pale blue gas with a sharp smell; you may notice it near a photocopier, or in the air after a thunderstorm. As a liquid it is deep blue.

High in the atmosphere, ultraviolet (UV) rays from the Sun split some oxygen molecules into single oxygen atoms (sometimes called nascent oxygen). An oxygen atom then joins an oxygen molecule to make ozone:

$$\text{O}_2 \xrightarrow{\text{UV rays}} \text{O} + \text{O}$$
$$\text{O}_2 + \text{O} \rightarrow \text{O}_3 \ \text{(ozone)}$$

The ozone made this way forms a layer in the stratosphere, the part of the atmosphere above the clouds and the weather. This is the ozone layer. It lies roughly 15 to 35 km above the ground, with the most ozone at about 20 to 25 km. (Books give slightly different heights for it, from about 15 to 50 km; the layer has no sharp edges.)

Importance of the ozone layer

The ozone layer absorbs most of the ultraviolet rays from the Sun, so these harmful rays can’t reach the Earth’s surface. That is why it is called the protective layer of the Earth.

• a.It helps keep the Earth’s climate and temperature in balance (by absorbing UV, it warms the stratosphere, and so shapes how the atmosphere moves).
• b.It protects living things on land and in water, and the ecosystems they form.
• c.It keeps ecosystems from being destroyed, and keeps them in balance.
• d.It protects people, animals and plants from the diseases and damage that UV radiation causes.

Depletion of the ozone layer

The thinning of the ozone layer in the stratosphere is called ozone layer depletion. The chief chemicals that deplete it are chlorofluorocarbons (CFCs), once used in refrigerators, air conditioners, aerosol sprays and foam. Others are halons (used in fire extinguishers), carbon tetrachloride, methyl chloroform, methyl bromide (used to kill pests in soil) and oxides of nitrogen (especially nitrous oxide, N₂O, from nitrogen fertilizers).

These gases are very stable in the lower air, so they last for years and slowly drift up into the stratosphere. There, UV rays break them up and set free atoms of chlorine or bromine. Each chlorine atom destroys ozone, and is set free again to destroy more:

$$\text{CFCl}_3 \xrightarrow{\text{UV rays}} \text{CFCl}_2 + \text{Cl}$$
$$\text{Cl} + \text{O}_3 \rightarrow \text{ClO} + \text{O}_2$$
$$\text{ClO} + \text{O} \rightarrow \text{Cl} + \text{O}_2$$

Because the chlorine atom isn’t used up, one chlorine atom can destroy many thousands of ozone molecules before it is finally locked away in another compound. Continued depletion thins the layer so much that, over Antarctica every spring (September to November), a hole opens in it.

Effects of ozone layer depletion

When the ozone layer thins, more UV rays from the Sun reach the Earth’s surface. They harm both living things and the environment:

• a.On human health: UV burns and irritates the eyes and causes cataracts (the lens of the eye goes cloudy); it causes sunburn, itching and blistering of the skin, and skin cancer; and it weakens the body’s immune system. UV also damages DNA, which can cause harmful changes (mutations).
• b.On plants: UV slows down the growth of plants and the rate of photosynthesis, so crops give less.
• c.On the ecosystem: UV harms the tiny plankton at the surface of the sea that feed fish, and the eggs and young of many animals, so whole food chains and ecosystems are thrown out of balance.
• d.On materials: UV makes plastics, paints and rubber crack and fade faster.

Ways to protect the ozone layer

• a.Stop using devices and products that release CFCs and other ozone-depleting chemicals; dispose of old refrigerators and air conditioners safely, so their gases are collected, not let out.
• b.Use substitutes that don’t harm the ozone layer. Hydrofluorocarbons (HFCs) replaced CFCs because they contain no chlorine, but they are strong greenhouse gases, so under a 2016 change to the Montreal Protocol (the Kigali Amendment) they too are being phased out, in favor of gases such as isobutane, now used in many new refrigerators.
• c.Use nitrogen fertilizers only as much as crops need, because the extra produces nitrous oxide, which also depletes ozone.
• d.Use fewer petrol and diesel vehicles, which give out oxides of nitrogen, and more electric vehicles. Nepal, which makes its electricity mostly from rivers, is a good place for electric vehicles.

In 1772 the Scottish scientist Daniel Rutherford found that air in which a candle had burned out and mice had breathed still had a gas left in it that would support neither burning nor life. In 1790 the French chemist Jean-Antoine Chaptal named it nitrogen, “niter-former”, because it is part of niter (saltpeter).

Nitrogen is much less reactive than most gases: under ordinary conditions it hardly reacts with any element or compound. The reason is the very strong triple bond holding its two atoms together (N≡N). Yet it is everywhere in living things, in proteins, enzymes, DNA and RNA, and in some rocks and minerals, as nitrates.

Laboratory preparation of nitrogen gas

In the laboratory, nitrogen is made by gently heating a solution of sodium nitrite (NaNO₂) and ammonium chloride (NH₄Cl):

$$\text{NaNO}_2 + \text{NH}_4\text{Cl} \xrightarrow{\Delta} \text{NaCl} + \text{N}_2\uparrow + 2\text{H}_2\text{O}$$

Precautions

1. Fit the apparatus airtight, and heat the mixture gently: once it starts, the reaction gives out heat and can become violent.
2. Make the solution before heating. Dry ammonium chloride, when heated, turns straight from a solid into a gas (it sublimes) instead of reacting.

Test for nitrogen gas

1. A burning splint goes out in nitrogen, and the gas doesn’t burn: nitrogen neither burns nor supports burning. (Carbon dioxide does the same, but nitrogen doesn’t turn lime water milky.)
2. A strongly burning magnesium ribbon, lowered into a jar of nitrogen, keeps burning and leaves a yellowish powder, magnesium nitride. Add a few drops of water to the powder: if you smell ammonia, the gas was nitrogen.

$$3\text{Mg} + \text{N}_2 \rightarrow \text{Mg}_3\text{N}_2 \qquad \text{Mg}_3\text{N}_2 + 6\text{H}_2\text{O} \rightarrow 3\text{Mg(OH)}_2 + 2\text{NH}_3$$

Physical properties of nitrogen gas

1. It is a colorless, odorless and tasteless gas.
2. It is slightly lighter than air.
3. It is sparingly soluble in water.
4. It is neutral: it doesn’t affect litmus paper.
5. It neither burns nor supports burning.
6. It turns into a liquid at about −196 °C and freezes at −210 °C.

Chemical properties of nitrogen gas

Nitrogen is unreactive under ordinary conditions, but under special conditions it reacts with several substances.

1. With hydrogen, it forms ammonia, at about 450 °C and a high pressure (about 200 times the pressure of the air; books give 200 to 500 atm), with an iron catalyst. This is the Haber process. (Many books add “a little molybdenum as a promoter” and write Fe/Mo over the arrow, and an exam accepts it; today’s plants help the iron with potassium and aluminium oxides instead.)

$$\text{N}_2 + 3\text{H}_2 \xrightarrow[\text{Fe/Mo}]{450\,°\text{C},\ 200\ \text{atm}} 2\text{NH}_3 \ \text{(ammonia)}$$
2. With oxygen, it forms nitric oxide, only at about 2,000 to 3,000 °C, in an electric arc, or in a flash of lightning:

$$\text{N}_2 + \text{O}_2 \xrightarrow{\text{electric arc}} 2\text{NO} \ \text{(nitric oxide)}$$
3. With burning magnesium, calcium or aluminium, it forms metal nitrides:

$$3\text{Mg} + \text{N}_2 \xrightarrow{\Delta} \text{Mg}_3\text{N}_2 \ \text{(magnesium nitride)}$$
$$3\text{Ca} + \text{N}_2 \xrightarrow{\Delta} \text{Ca}_3\text{N}_2 \ \text{(calcium nitride)}$$
$$2\text{Al} + \text{N}_2 \xrightarrow{\Delta} 2\text{AlN} \ \text{(aluminium nitride)}$$

Uses of nitrogen gas

1. To make ammonia and nitric acid, which are used to make nitrogen fertilizers such as urea.
2. To fill packets of food, like chips and biscuits: with no oxygen inside, the food doesn’t go stale, and the gas cushions it.
3. To fill electric bulbs (usually mixed with argon), so that the hot tungsten filament doesn’t burn away.
4. As liquid nitrogen (at −196 °C), to freeze and store things: blood, cells, and the frozen semen used by vets in Nepal for artificial insemination of cows and buffaloes; and to cool equipment in laboratories.

1. Choose the correct answer

2. Give reasons

3. Answer the following questions

4. Complete and balance these equations