Lesson 1: Types of Chemical Reactions

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.