Unit 1 — Chemical Reactions and Their Environmental Impacts

1.1 Starting point: what is a chemical reaction?

• Look at two experiments that both show bubbles.
• In the first, dilute hydrochloric acid is poured onto a piece of marble (calcium carbonate).
• In the second, the same acid is poured onto a piece of zinc.
• Both show effervescence, yet they are not the same reaction.
• The gas from the marble is carbon dioxide; the gas from the zinc is hydrogen.
• In one case the acid reacts with a carbonate (a salt); in the other it reacts with a metal.
• The key question of this lesson is: how can we sort the huge number of reactions into a few families so that we can predict their products?

• You already know the basic definition.
• A chemical reaction is a process in which the bonds in the molecules of the reactants are broken and new bonds form, giving products that are new substances with different properties.

How do we know a reaction has happened? We look for one or more signs:

• a change in colour (for example blue copper sulphate solution fading);
• a change in odour;
• evolution of a gas (bubbles);
• formation of a precipitate (an insoluble solid that appears in a solution);
• release of heat and light (burning).

• Scientists classify reactions according to how the atoms or ions rearrange and re-bond.
• This classification is not just filing; it lets us predict the products of reactions we have never seen.
• The four main types in this lesson are:

1. Direct combination reactions
2. Decomposition reactions
3. Single replacement reactions
4. Double replacement reactions (which include neutralisation and precipitation)

• Later in the lesson we meet a different way of looking at reactions, oxidation–reduction (redox), which describes what happens to electrons.
• Combustion, studied in Lesson 2, is another important type.

1.2 Direct combination reactions

A direct combination reaction is a reaction in which two or more substances join to form one new, more complex substance.

General pattern: A + B → AB

Think of two separate objects snapping together. Typical examples:

• Metal + oxygen → metal oxide
• Non-metal + oxygen → non-metal oxide
• Metal + non-metal → salt

Activity: reaction of metals with oxygen. Using tongs, a student heats three samples in a Bunsen flame for the same time: a sanded magnesium ribbon, a piece of steel wool (iron) and a copper wire.

Observations and conclusions:

• Conclusion: most metals combine with oxygen when heated in air, but they do so with very different vigour.
• Magnesium reacts most readily, iron less so, and copper least.
• So magnesium is more reactive than iron, and iron is more reactive than copper.

• Safety note. Burning magnesium gives out very intense light and ultraviolet radiation that can injure the retina.
• Never stare at it; wear safety goggles.

Notice how each equation is balanced. In 2Mg + O₂ → 2MgO there are two Mg atoms and two O atoms on each side. In 3Fe + 2O₂ → Fe₃O₄ there are three Fe and four O atoms on each side. In 2Cu + O₂ → 2CuO there are two Cu and two O on each side. The oxygen molecule O₂ always contains two atoms, so we must produce an even number of oxygen atoms on the left.

Other direct combination examples you can understand with the same pattern: 2Na + Cl₂ → 2NaCl (a metal and a non-metal give a salt), and Fe + S → FeS (heating iron and sulphur). Also, magnesium combines directly with bromine: Mg + Br₂ → MgBr₂.

1.3 Decomposition reactions

• A decomposition reaction is the opposite of combination.
• One compound breaks down into two or more simpler substances, usually when energy such as heat (thermal decomposition) or light is supplied.

General pattern: AB → A + B

Examples:

1. Red mercury(II) oxide decomposes when heated into silvery liquid mercury and oxygen gas. 2HgO → 2Hg + O₂ (Check: 2 Hg and 2 O on each side.) This is the experiment in which Joseph Priestley obtained oxygen.

2. Calcium carbonate (limestone, marble, chalk) decomposes on strong heating into calcium oxide (quicklime) and carbon dioxide. CaCO₃ → CaO + CO₂

3. In a matchstick, potassium chlorate decomposes by heat to give oxygen: 2KClO₃ → 2KCl + 3O₂. (More on this in Lesson 2.)

Do not confuse the two: a combination reaction has one product from several reactants; a decomposition reaction has one reactant that yields several products.

Check your understanding. Magnesium ribbon burning in air is a combination reaction (Mg + O₂ → MgO, two reactants, one product). The heating of mercury oxide is a decomposition reaction (one reactant, two products).

1.4 Single replacement reactions

• A single replacement reaction is one in which a free element takes the place of another element in a compound.
• The more reactive element pushes out the less reactive one.

General pattern: A + BC → AC + B

The element that is replaced is released in its free form. Before we study the examples we need a tool: the chemical reactivity series.

1.4.1 Replacement of hydrogen of an acid by a metal

Activity.

• Place a piece of sanded magnesium ribbon in a test tube and add about 5 mL of dilute hydrochloric acid.
• At once bubbles of gas appear.
• Close the mouth of the tube for a few seconds, then bring a burning splint near: the gas burns with a "pop".
• That is the test for hydrogen.

Mg + 2HCl → MgCl₂ + H₂↑

• The magnesium has replaced the hydrogen of the acid: hydrogen was released as a gas and the salt (magnesium chloride) was formed.
• In general:

Reactive metal + dilute acid → salt + hydrogen gas

• Do all metals react with the same speed?
• Place equal masses of zinc, iron and copper in equal volumes of the same dilute hydrochloric acid and compare the rate of bubbling:

• Zinc: steady bubbles. Zn + 2HCl → ZnCl₂ + H₂↑
• Iron: slower bubbling. Fe + 2HCl → FeCl₂ + H₂↑
• Copper: no bubbles at all.

• Magnesium is faster than zinc, zinc is faster than iron, and copper does not react.
• The rate of reaction with acids (and with oxygen and water) is a measure of the metal's chemical reactivity.

1.4.2 The chemical reactivity series

• The chemical reactivity series is the arrangement of the metal elements in descending order of their chemical reactivity.
• Hydrogen, which is not a metal, is placed in the series as a reference point.

Most reactive (top) to least reactive (bottom):

Potassium (K) – Sodium (Na) – Lithium (Li) – Barium (Ba) – Calcium (Ca) – Magnesium (Mg) – Aluminium (Al) – Zinc (Zn) – Iron (Fe) – Tin (Sn) – Lead (Pb) – Hydrogen (H) – Copper (Cu) – Mercury (Hg) – Silver (Ag) – Platinum (Pt) – Gold (Au)

Important rules from the series:

• Metals above hydrogen replace hydrogen from dilute acids (and the most reactive ones also from water).
• Metals below hydrogen (Cu, Hg, Ag, Pt, Au) do not react with dilute acids or with water, because they cannot replace hydrogen.
• A metal higher in the series replaces a lower metal from the solution of its salt.
• Platinum and gold are inactive elements. This is why gold and platinum stay shiny and are used in jewellery.

• Enrichment: why does aluminium seem slow? Aluminium is above hydrogen, yet it does not react with dilute hydrochloric acid immediately.
• The reason is the thin, tough layer of aluminium oxide (Al₂O₃) on its surface, which acts as an insulating layer and prevents the acid from reaching the metal.
• When this layer is slowly eaten away by the acid, bubbles of hydrogen start to appear.
• So the delay does not mean aluminium is unreactive.

• Example. If small pieces of copper, zinc, silver and iron are put separately in dilute hydrochloric acid, hydrogen is released only from zinc and from iron, because both are above hydrogen.
• Copper and silver are below hydrogen, so nothing happens.

1.4.3 Replacement of hydrogen of water by a metal

Highly reactive metals at the top of the series react with water and replace hydrogen, which is released as a gas that burns with a pop.

Activity (done only by the teacher).

• Half-fill a glass trough with water.
• Using forceps, take a very small piece of sodium (the size of a grain of rice) from the kerosene, dry it with filter paper, and place it on the water.

Observations: the sodium floats because its density is less than that of water; it darts about quickly across the surface because it reacts instantly and violently; a yellowish-orange flame may form; the solution becomes alkaline (it turns red litmus blue).

2Na + 2H₂O → 2NaOH + H₂↑

Sodium replaces hydrogen from water, forming sodium hydroxide (an alkali) and hydrogen gas.

• Safety. A great deal of heat is released, which can ignite the hydrogen or cause a small explosion.
• Water must never be used to put out a sodium fire; dry sand is kept ready.
• Sodium is stored under kerosene so that it does not react with air and moisture.

Do all metals react with water equally? No. Compare sodium and magnesium: sodium reacts instantly with cold water, while magnesium ribbon does not react with cold water and needs thermal energy (hot water) before bubbles form: Mg + 2H₂O(hot) → Mg(OH)₂ + H₂↑ This proves magnesium is less reactive than sodium. Metals lower in the series such as copper do not react with water at all.

1.4.4 Replacement of one metal by another in its salt solution

A more reactive metal can displace a less reactive metal from the solution of its salt.

Activity. Pour blue copper(II) sulphate solution into a beaker. Clean a zinc rod with sandpaper and immerse it for about ten minutes.

Observations: the blue colour of the solution gradually fades, and a reddish-brown deposit of copper forms on the rod.

Zn + CuSO₄ → ZnSO₄ + Cu (blue) (colourless) (red-brown solid)

• Zinc is higher than copper in the series, so zinc replaced copper.
• Had the experiment been done the other way round (a copper rod in zinc sulphate), nothing would happen, because copper cannot displace the more reactive zinc.

Rule: a metal that comes earlier in the series replaces any metal that follows it from its salt solution.

• Life application: cleaning silverware. Silverware slowly develops a black layer of silver sulphide (tarnish).
• Place it in boiling water with sodium bicarbonate and a piece of aluminium foil.
• Aluminium is more reactive than silver, so it takes the place of silver in the tarnish, and the silver is restored to its shiny state.

• Cross-cutting concept: patterns. Reactions are not random.
• They follow predictable patterns (combination, decomposition, replacement), and the reactivity series is a pattern that lets us predict what will happen before we do the experiment.

1.5 Double replacement reactions

Revise first: an ionic compound such as sodium chloride (NaCl) consists of a positive ion (a cation, for example Na⁺) and a negative ion (an anion, for example Cl⁻).

A double replacement reaction is a reaction between two ionic compounds, usually dissolved in water, in which the cations and anions exchange partners to make two new compounds.

General pattern: AB + CD → AD + CB

Three important kinds are studied.

1.5.1 Neutralisation: acid + alkali

Activity.

• Put a little dilute sodium hydroxide solution in a test tube and add drops of universal indicator.
• Then add dilute hydrochloric acid drop by drop.

The universal indicator gives a colour that depends on how acidic or alkaline the solution is:

• strong alkali: violet (dark blue/purple),
• weak alkali: blue,
• neutral: green,
• weak acid: yellow, then orange,
• strong acid: red.

• Observations: NaOH solution is colourless but turns the indicator violet.
• As acid is added gradually the colour changes through blue to green, the neutralisation point.
• If we continue to add acid the colour moves to yellow, orange and finally red.

NaOH + HCl → NaCl + H₂O

• The acid and the alkali exchange partners: Na⁺ combines with Cl⁻ (salt) and H⁺ combines with OH⁻ (water).
• A reaction between an acid and an alkali that gives a salt and water only is a neutralisation reaction.

1.5.2 Acid + salt

Activity.

• Put sodium carbonate powder in a flask.
• Add dilute hydrochloric acid slowly from a funnel with a stopcock.
• Pass the gas through clear limewater using a delivery tube.

Observation: effervescence occurs and limewater turns milky (turbid).

Conclusion: the gas is carbon dioxide. The reaction is:

Na₂CO₃ + 2HCl → 2NaCl + H₂CO₃ H₂CO₃ → H₂O + CO₂↑ (carbonic acid is unstable and breaks down)

Overall: Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂↑

• The ions exchanged positions: sodium joined chloride, hydrogen joined carbonate.
• The carbonic acid formed is unstable, so it immediately decomposes into water and carbon dioxide.

This also explains the first experiment of the lesson: marble (calcium carbonate) with dilute hydrochloric acid gives carbon dioxide: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂↑ So the marble reaction is a double replacement reaction, whereas the zinc reaction is a single replacement reaction. Not every reaction of dilute hydrochloric acid produces a gas: with sodium hydroxide, for example, the products are a salt and water and no gas is produced.

1.5.3 Precipitation: salt solution + salt solution

Activity. Put sodium chloride solution in a test tube and add a few drops of silver nitrate solution.

Observation: a white precipitate of silver chloride forms immediately. It slowly turns violet/grey in light.

NaCl + AgNO₃ → AgCl↓ + NaNO₃

• The two solutions are exchanging partners: Na⁺ with NO₃⁻ (stays dissolved) and Ag⁺ with Cl⁻ (insoluble solid).
• A reaction between two soluble salts in solution that forms an insoluble solid is a precipitation reaction.

• Check your understanding. Two test tubes each contain silver nitrate solution.
• Sodium chloride solution is added to the first, and a piece of aluminium is put in the second.
• In the first, a white precipitate shows that a reaction happened.
• In the second, a grey/silvery deposit of silver forms on the aluminium (aluminium is more reactive than silver and replaces it), and the colourless solution slowly changes because aluminium nitrate forms.

Life application: treating hard water. Underground water may contain a lot of calcium and magnesium salts; this is hard water. It leaves hard deposits (scale) inside pipes, boilers and washing machines and finally damages them. At treatment plants, sodium carbonate is added. Carbonate ions combine with calcium and magnesium ions to make insoluble carbonates, which can be filtered or settle out: Ca²⁺ + CO₃²⁻ → CaCO₃↓ Mg²⁺ + CO₃²⁻ → MgCO₃↓ Equivalent full equation: CaCl₂ + Na₂CO₃ → CaCO₃↓ + 2NaCl.

• Model activity. In class you can design models of the different types of reactions using coloured paper clips: each colour is an atom, and linked clips are bonded atoms.
• Taking apart and re-joining the clips shows how combination (join), decomposition (split), single replacement (one clip changes partner) and double replacement (two pairs swap partners) rearrange atoms.

1.6 Oxidation–reduction (redox) reactions

In earlier years you learned a definition based on oxygen and hydrogen:

• Oxidation: increase in the percentage of oxygen, or decrease in the percentage of hydrogen, in a substance.
• Reduction: increase in the percentage of hydrogen, or decrease in the percentage of oxygen.

• But in many reactions (for example zinc in copper sulphate solution) there is no oxygen or hydrogen at all, yet clearly something is oxidised and something reduced.
• We need a definition in terms of electrons.

Look again at: Zn + CuSO₄ → ZnSO₄ + Cu. The ionic form is

Zn + Cu²⁺ → Zn²⁺ + Cu

(The sulphate ion SO₄²⁻ is unchanged and is called a spectator ion.)

Before the reaction, zinc consists of neutral atoms and copper is present as positive ions (Cu²⁺) in the solution. During the reaction:

• Oxidation half-reaction: a zinc atom loses two electrons and becomes a zinc ion: Zn → Zn²⁺ + 2e⁻.
• Reduction half-reaction: a copper ion gains those two electrons and becomes a neutral copper atom: Cu²⁺ + 2e⁻ → Cu.

Therefore:

• Oxidation is a process in which an atom or ion loses one or more electrons.
• Reduction is a process in which an atom or ion gains one or more electrons.
• A reducing agent is the substance that loses electrons (it is itself oxidised and it causes the other substance to be reduced).
• An oxidising agent is the substance that gains electrons (it is itself reduced and it causes the other to be oxidised).

• Oxidation and reduction always occur together and in equal amounts of electrons; one cannot happen without the other.
• Hence the name "redox".

In the reaction above, zinc is oxidised and is the reducing agent; copper ions (the copper sulphate) are reduced and copper sulphate is the oxidising agent.

A useful memory aid: OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

Worked example 1: magnesium and bromine

Mg + Br₂ → MgBr₂

Magnesium becomes Mg²⁺ (loses 2 electrons): Mg → Mg²⁺ + 2e⁻. Each bromine atom gains one electron: Br₂ + 2e⁻ → 2Br⁻.

• Oxidised: magnesium.
• Reduced: bromine.
• Oxidising agent: bromine (it gains electrons).
• Reducing agent: magnesium (it loses electrons).

Worked example 2: magnesium and hydrochloric acid

Mg + 2HCl → MgCl₂ + H₂

Magnesium atoms become Mg²⁺ ions by losing two electrons: oxidation, so Mg is the reducing agent. The hydrogen ions of the acid gain electrons and become hydrogen molecules: 2H⁺ + 2e⁻ → H₂ : reduction, so the acid (the hydrogen ions) is the oxidising agent.

Worked example 3: sodium and water

2Na + 2H₂O → 2NaOH + H₂. Sodium atoms each lose one electron (Na → Na⁺ + e⁻): sodium is oxidised, the reducing agent. Hydrogen of water gains electrons to make H₂: it is reduced; water is the oxidising agent.

Note that every single replacement reaction of a metal is a redox reaction, because a metal atom loses electrons while another ion gains them. Combination reactions between a metal and a non-metal, and the combustion of a metal in oxygen (2Mg + O₂ → 2MgO: Mg loses electrons, oxygen gains) are also redox.

• Why the reactivity series matters. The more reactive a metal, the more readily its atoms lose electrons.
• That is why potassium and sodium react so violently, and why gold hardly reacts at all.

1.7 Summary of Lesson 1

Key ideas: the reactivity series predicts single replacement; neutralisation gives salt and water; precipitation gives an insoluble solid; oxidation is loss of electrons and reduction is gain of electrons; they occur together.

Environmental and economic link.

• Corrosion (for example rusting of iron) is slow oxidation of a metal.
• Because iron is rather high in the series, it corrodes, and corrosion costs economies enormous sums every year.
• Using less reactive metals, painting, galvanising (coating iron with zinc) or oiling can save money and resources.
• This links to responsible consumption.

2.1 Why do some things burn easily?

• Look at a forest fire: the leaves catch fire easily while a thick branch takes much longer.
• And the smoke affects the whole atmosphere.
• This lesson answers: what is combustion?
• Why do some materials burn faster?
• What are the products?
• And how can we reduce the pollution that burning causes?

2.2 The combustion triangle

Combustion (burning) is a chemical reaction between a substance and oxygen that releases heat, light, or both, and forms new substances that depend on what is burned.

For combustion to occur, three things must be present together. They form the combustion triangle (fire triangle):

1. Fuel: a flammable material.
2. Oxygen (from the air).
3. Heat energy: an ignition source that brings the fuel to its ignition point.

• Remove any one of the three and the fire stops.
• This is the principle of fire extinguishing: water removes heat (by absorbing it), a lid or fire blanket removes oxygen, and clearing away trees or switching off a gas supply removes fuel.

2.3 Ignition point and flammability

Activity.

• Take two identical paper strips, one dry and one wet.
• Hold each, in turn, 2 cm above the tip of a candle flame and time how long until it catches fire.
• The dry strip ignites fast.
• The wet one does not burn easily because the water absorbs the heat of the flame and delays the paper from reaching its ignition temperature.

(Safety: wear goggles and heat-resistant gloves; work on a stable heat-resistant board away from flammable objects; have an extinguisher available.)

• Ignition point is the lowest temperature to which a material must be heated so that it starts burning and continues to burn by itself.
• Flammability expresses how easily and quickly a substance ignites at the surrounding temperature. Substances with a low ignition point, such as petrol and alcohol, are highly flammable and their containers carry warning signs.

This answers the forest question: thin dry leaves have a large surface area and little mass, so they heat up and reach their ignition point quickly; a thick branch takes much longer to heat to that temperature.

2.4 Life application: how a match lights

• The match head contains an oxidising agent, potassium chlorate (KClO₃), together with easily burned substances, glue and powdered glass.
• The striking strip on the side of the box contains red phosphorus, glue and glass powder.

When you strike a match, friction heat turns a tiny amount of red phosphorus into white phosphorus, which ignites spontaneously in air. This heat decomposes potassium chlorate: 2KClO₃ → 2KCl + 3O₂ The oxygen released makes the match head and the wood burn faster and more intensely.

2.5 Combustion of metals and non-metals

Metals. Magnesium burns with a dazzling white light and forms white magnesium oxide: 2Mg + O₂ → 2MgO Most metals burn to give basic oxides. Some dissolve in water to form alkaline solutions (which turn litmus blue). For example, MgO + H₂O → Mg(OH)₂ (slightly soluble); Na₂O + H₂O → 2NaOH.

Non-metals. Coal (carbon) glows red and releases carbon dioxide: C + O₂ → CO₂ Most non-metals form acidic oxides, which dissolve in water giving acids (which turn litmus red). Examples: CO₂ + H₂O → H₂CO₃ (carbonic acid), S + O₂ → SO₂ and SO₂ + H₂O → H₂SO₃ (sulphurous acid).

So: metal oxides are usually basic; non-metal oxides are usually acidic.

2.6 Complete and incomplete combustion

Activity with a Bunsen burner.

• Put water in a test tube.
• First open the air valve fully and light the burner: the flame is blue.
• Heat the tube for a minute; its bottom remains clean.
• Then close the air valve partially: the flame turns yellow.
• Heat the tube again: a black deposit (soot) forms on its bottom.

• Blue flame: plenty of oxygen entering through the air inlet; the gas burns completely.
• Yellow flame: air inlet narrow, little oxygen; the gas burns incompletely and unburned carbon (soot) is formed.

• Cumulative knowledge. Bunsen burners and home gas stoves use hydrocarbons, compounds of hydrogen and carbon only.
• Natural gas is mostly methane (CH₄).
• Butagas (cooking gas) is mostly propane and butane.

Complete combustion of a hydrocarbon

With plenty of oxygen, the products are carbon dioxide and water vapour, with a large amount of heat:

CH₄ + 2O₂ → CO₂ + 2H₂O + heat

Balancing check: carbon 1 = 1; hydrogen 4 = 4; oxygen 4 = 2 + 2 = 4. Two oxygen molecules are required for each methane molecule.

Incomplete combustion of a hydrocarbon

With limited oxygen, the products depend on how much oxygen is available.

Producing carbon monoxide (CO): 2CH₄ + 3O₂ → 2CO + 4H₂O + heat (C: 2 = 2; H: 8 = 8; O: 6 = 2 + 4 = 6.)

Producing soot (carbon, C): CH₄ + O₂ → C + 2H₂O + heat (C: 1 = 1; H: 4 = 4; O: 2 = 2.)

Extra: with even less oxygen, a mixture of CO and C forms.

Important note: burning the same mass of fuel incompletely releases less heat than burning it completely, so incomplete combustion is wasteful as well as polluting.

Worked example 1: oxygen needed

How many oxygen molecules are needed to burn two methane molecules in each case?

• Complete (to CO₂): 2CH₄ + 4O₂ → 2CO₂ + 4H₂O : 4 O₂ molecules.
• Incomplete to CO: 2CH₄ + 3O₂ → 2CO + 4H₂O : 3 O₂ molecules.
• Incomplete to carbon: 2CH₄ + 2O₂ → 2C + 4H₂O : 2 O₂ molecules.

Conclusion: the less oxygen available, the more carbon is left in a dangerous or useless form.

Worked example 2: 4 methane molecules

Complete combustion of 4 methane molecules requires 4 × 2 = 8 oxygen molecules, forming 4 CO₂ and 8 H₂O.

Life application: flame colour on a stove

• A blue flame on a gas stove means complete combustion: it gives carbon dioxide, water vapour and a lot of heat, so cooking is quick and less gas is used.
• A yellow or orange flame means incomplete combustion because of a lack of oxygen: it produces toxic carbon monoxide and black soot that blackens pots, and heating is less efficient.
• The solution is to clean the burner openings or adjust the air intake.
• Never run a gas heater in a closed room.

2.7 Detecting the products of combustion

Activity. Burn natural gas (hydrocarbon) in a Bunsen burner and pass the products through an apparatus: first a U-tube surrounded by ice that contains light blue anhydrous cobalt chloride paper, then through clear limewater.

Observations:

• The ice cools the gases so the water vapour condenses; the blue cobalt chloride paper turns pink, showing that water is formed.
• Limewater turns milky, showing that carbon dioxide is formed.

So the complete combustion of a hydrocarbon forms water and carbon dioxide, as the equation says.

2.8 Properties of the combustion products

Equations: S + O₂ → SO₂; 2NO + O₂ → 2NO₂.

• Cross-cutting concept: stability and change. In every combustion the substances change but the total mass of products equals the total mass of the reactants (conservation of matter).
• Yet the products can change the environment and human health over time.

2.9 Environmental pollution from combustion

Burning agricultural waste: the "black cloud"

• At the end of the harvest (usually October–November) some farmers burn rice straw to clear the fields quickly.
• Piled straw receives too little oxygen, so combustion is mainly incomplete, releasing carbon monoxide, carbon dioxide and soot (fine carbon particles).
• These form dense smoke called the black cloud, which reduces visibility, causes breathing difficulties and increases asthma and chest allergies.
• The alternative is to use the straw for compost, animal feed, paper or biofuel and not burn it.

Car engines

• Burning petrol or diesel in the engine gives the heat that becomes motion.
• The exhaust contains carbon monoxide, carbon dioxide, soot, nitrogen oxides, and (if the fuel has sulphur) sulphur dioxide.
• It is a main cause of urban air pollution.
• When too little air enters the engine, or the engine is inefficient, burning is incomplete and the harmful emissions rise.
• Reducing exhaust pollution:

1. Regular engine maintenance to keep efficiency high.
2. Cleaner fuels such as natural gas, or electric cars.
3. Public transport that does not harm the environment.

Technology: catalytic converter. A catalytic converter in the exhaust contains catalysts (substances that speed up a reaction without being used up). They speed up redox reactions that turn toxic CO and nitrogen oxides into safer carbon dioxide and nitrogen. For example: 2CO + 2NO → 2CO₂ + N₂.

Industry

• Metal smelting, ceramics, cement, glass, oil distillation and many power stations burn coal, oil or natural gas.
• Their chimneys can emit large amounts of pollutants.
• Reduction measures:

1. Use natural gas, which causes less pollution.
2. Fit filters to clean the gases before leaving chimneys.
3. Use renewable energy (solar and wind) for electricity.

• Why is natural gas preferred to coal? Burning natural gas gives less carbon dioxide for the same heat, and very little sulphur dioxide and soot.
• Coal gives much more carbon dioxide, more sulphur oxides and more particles.

Technology: acid-gas scrubber. A scrubber is placed in factory chimneys. The polluted gases pass through a spray of an alkaline solution. Acidic gases such as sulphur dioxide are neutralised (for example SO₂ + Ca(OH)₂ → CaSO₃ + H₂O) and form harmless solid salts, reducing the chance of acid rain and smog.

Acid rain

• forms when sulphur dioxide and nitrogen oxides dissolve in rain water to give acids.
• It damages buildings, harms plants and kills life in lakes.

Cross-cutting concept: cause and effect. More combustion processes (cause) lead to emissions that disturb the ecological balance (effect).

Be an environmental engineer. Choose one problem (the black cloud, car exhaust, or burning waste) and propose a practical solution: a filter that converts harmful gases, a recycling scheme for straw, a better maintained engine.

2.10 Summary of Lesson 2

• Combustion needs fuel, oxygen and heat (fire triangle).
• Ignition point and flammability explain why some fuels burn more easily.
• Metals give basic oxides, non-metals acidic oxides.
• Complete combustion of a hydrocarbon gives CO₂ and H₂O; incomplete gives CO and/or soot, plus water, and releases less energy.
• Products can be identified using cobalt chloride paper (water) and limewater (carbon dioxide).
• Burning waste, vehicles and industry pollute the air; catalytic converters, scrubbers, clean fuel and renewable energy reduce pollution.

3.1 A behaviour that threatens public health

• Tobacco smoking harms the smoker and the people around him, and it harms the environment.
• In this lesson you learn what tobacco contains, what chemicals form when it burns, how these chemicals harm the body's systems, and how to avoid smoking.

3.2 The tobacco plant and the products of tobacco combustion

• The tobacco plant grows mainly in warm, tropical regions.
• It has a thick upright stem, large green leaves and tubular flowers that may be white, pink or red.
• The dried leaves are used to make cigarettes.

A simple experiment to collect cigarette smoke.

• A transparent plastic bottle with a hole near the bottom (sealed by a stopper) is two-thirds filled with water.
• A cigarette is fixed in a hole in the bottle cap, and a piece of white cotton is pressed into the mouth of the bottle.
• The cigarette is lit and the stopper removed.
• As water runs out, air is drawn through the cigarette, carrying its smoke into the bottle.
• The white cotton slowly turns brown/dark as products of combustion are deposited among the fibres.
• This shows that the smoke contains solid and sticky particles, not just gas.

3.3 What tobacco contains

1. Nicotine. A natural organic compound containing carbon, hydrogen and nitrogen, with an alkaline nature. Heat from the burning cigarette vaporises nicotine, and the smoker inhales it with the smoke. Nicotine is the substance that causes addiction.
2. Carbohydrates and proteins (sugars, plant proteins and fibres). When burned, especially incompletely, they produce carbon monoxide and tar, a thick sticky dark liquid made of the combustion products of these compounds.
3. Mineral elements. Potassium, magnesium and calcium (harmless as minerals), but also toxic heavy elements such as cadmium, lead and arsenic.
4. Radioactive elements. Small amounts, for example polonium-210, which can damage the lungs by radiation.

Burning tobacco releases a mixture of harmful gases:

• toxic gases: carbon monoxide (CO) and hydrogen cyanide (HCN);
• irritant gases: nitrogen oxides (NO, NO₂) and ammonia (NH₃).

A burning cigarette is therefore not just paper and dry leaves: complex reactions produce several thousand chemicals, many of them toxic or carcinogenic (cancer-causing).

What happens to minerals and radioactive elements?

1. Volatilisation: some heavy metals and radioactive elements vaporise when the tobacco ignites and become gases or stick to fine tar particles.
2. Inhalation: the atoms enter the lungs and deposit in the branches of the bronchi.
3. Accumulation: the build-up causes damage to the DNA of lung cells, which raises the risk of cancerous tumours.
4. Ash: the leftover ash contains the non-volatile minerals (calcium, potassium), giving it a white or grey colour.

According to the World Health Organization (WHO), tobacco causes more than 8 million deaths every year.

3.4 Harmful effects on health

Respiratory system

• Normal lungs are pink and elastic.
• A smoker's lungs become dark or black because tar and fine particles are deposited in the tissue.
• The lungs lose elasticity and capacity, so the smoker has chronic shortness of breath, persistent cough, reduced ability to exercise and a higher risk of lung cancer.

• Structure and function. The alveoli are small, thin-walled and very numerous (structure) and give a large surface area for exchanging oxygen and carbon dioxide between lungs and blood (function).
• Damage to this structure damages the function.

Emphysema.

• Sticky tar condenses on the inner walls of the alveoli, which lose their elasticity.
• As tar accumulates, the walls rupture and many small sacs turn into a few large, flabby cavities.
• The surface area available for gas exchange drops sharply, so less oxygen reaches the blood.
• The patient suffers severe, chronic difficulty in breathing.

Circulatory system (heart and blood vessels)

Carbon monoxide

• binds with the haemoglobin of the blood much more readily than oxygen does, so the blood carries less oxygen, and the body (heart, brain and muscles) is short of oxygen.
• Nicotine raises the heart rate and blood pressure and narrows (constricts) blood vessels, which raises the risk of heart disease and atherosclerosis (fatty deposits in the arteries).

Nervous system and senses

Toxins cause nervous disorders: nicotine can increase anxiety, disturb sleep and vision, and in the long term reduce brain function and concentration, especially in teenagers.

Fertility and reproduction

Smoking reduces fertility in men and women and increases the risk of abnormalities in the fetus.

Immune system and appearance

Smoking weakens immunity; it also causes pale skin and hair loss.

Addiction

• When someone smokes, nicotine reaches the brain quickly.
• It stimulates the release of substances that give a brief feeling of well-being.
• This effect fades quickly, creating a craving for another cigarette.
• Repeated smoking leads to addiction, so the habit becomes harmful and difficult to quit.
• This is why many smokers who want to stop find it so hard.

• Worked reasoning example 1. Why does a smoker get out of breath quickly when climbing stairs?
• Carbon monoxide reduces the oxygen carried by the blood; tar damages the alveoli and cuts the surface for gas exchange.
• Muscles receive less oxygen than they need, so the person breathes hard and tires easily.

• Worked reasoning example 2. Why is the gas X (colourless, odourless, toxic) in cigarette smoke carbon monoxide, and Y (dark, sticky) tar?
• CO is toxic, colourless and odourless; tar is the dark, viscous liquid made of combustion products of sugars, proteins and fibres.

3.5 Environmental damage caused by smoking

1. Air pollution: smoking releases large amounts of carbon dioxide, adding to global warming (the rise in Earth's temperature caused by greater proportions of certain gases in the atmosphere). WHO estimates that tobacco releases tens of millions of tons of carbon dioxide every year.
2. Water and soil pollution: cigarette butts are among the commonest forms of litter. They contain harmful chemicals that leach into soil and water and harm living organisms. Several trillion cigarette filters reach oceans, rivers, streets and beaches each year.

3.6 Types of smoking and ways of exposure

Types of smoking

1. Traditional smoking: burning tobacco leaves in cigarettes (or other tobacco products).
2. Vaping (electronic smoking): heating a liquid that contains nicotine with an electronic device to produce an aerosol (vapour) that is inhaled.

• There is a common misconception that vaping is less harmful.
• It has been shown that both cause severe and direct damage to the lungs and heart.
• Vaping liquids often contain high concentrations of nicotine, which causes rapid addiction and may cause acute heart and nervous problems.

Exposure to smoking products

1. Active smoking: the smoker deliberately inhales smoke from burning tobacco or the aerosol from a vaping device.
2. Passive smoking: a non-smoker inhales the smoke exhaled by a smoker, or the smoke that rises from the burning tip of the cigarette. That smoke has not passed through the filter, so it can contain higher concentrations of toxic and carcinogenic gases. Non-smokers are therefore exposed to the same heart and respiratory diseases as active smokers.

• Comparison of situations: a person who smokes (active), a non-smoker sitting beside him (passive), and a person who vapes (active, with aerosol released into the air).
• All three inhale harmful substances; the differences are the method (burning versus heating a liquid) and whether the person chose to inhale.

3.7 Preventing smoking: how to say no

• Decide to refuse: practise saying "No" clearly to anyone who offers you a cigarette.
• Choose positive friends: stay away from peers who encourage smoking and make friends who care about their health.
• Stay busy with positive activities: sport, reading, hobbies.
• Avoid smoking areas to protect yourself from passive smoking.
• Ask for help from a parent or teacher if you feel pressure.
• Learn about the dangers of smoking; knowledge strengthens your decision.

• Awareness campaigns (for example World No Tobacco Day each year, and posters by health ministries) help the community.
• Governments also ban smoking in public places.

• Issue for discussion: Is vaping less risky than smoking?
• The scientific answer in this book: no, it still harms the heart and lungs and causes quick addiction.

3.8 Summary of Lesson 3

• Tobacco contains nicotine, carbohydrates and proteins, minerals (some toxic heavy metals) and radioactive elements.
• Its combustion gives CO, HCN, NO/NO₂, NH₃, tar, and thousands of other chemicals.
• Harm: lungs (tar, emphysema, cancer), heart and blood (CO, nicotine), nerves, fertility, immunity, and addiction.
• Smoking pollutes air (CO₂), water and soil (butts).
• Types: traditional and vaping; exposure: active and passive.
• Prevention begins with a personal decision and a supportive environment.