Chemical Reactions and Equations

Why This Matters

Leave a glass of milk out on a summer afternoon and by evening it has turned sour. Leave an iron tawa in a damp corner of the kitchen and a week later it wears a reddish-brown coat. Bite into a piece of bread and within hours your body has turned it into the energy that lets you read this sentence.

Every one of these is the same kind of event wearing a different costume: a substance has quietly turned into a new substance. The milk’s chemistry changed. The iron’s surface became something that is no longer iron. The bread became glucose, then carbon dioxide, water and energy.

That “turning into something new” is what chemists call a chemical reaction. And once you can spot reactions, name them, and write them down precisely, you hold the grammar of all of chemistry. This chapter teaches you that grammar — how to see when a reaction has happened, how to write it as a balanced equation, and how to sort every reaction into a handful of simple families. Master this one chapter and the rest of chemistry stops looking like magic spells and starts looking like sentences you can read.

The Big Idea

Here is the single idea the whole chapter hangs on:

In a chemical reaction, atoms are never created and never destroyed — they are only rearranged into new combinations.

Think of atoms like a fixed set of LEGO bricks. A reaction takes apart some structures (breaks bonds) and builds new ones (makes bonds), but you finish with the exact same bricks you started with — not one brick more, not one less. This is the Law of Conservation of Mass you met in Class 9: mass can neither be created nor destroyed in a chemical reaction.

Two consequences flow from this, and they power everything ahead:

1. Because new substances form, reactions announce themselves with visible signs — a colour change, a gas bubbling out, a change in temperature, a new smell, or a solid appearing.
2. Because atoms are conserved, every correct chemical equation must be balanced — the same number of atoms of each element on both sides.

Keep these two ideas in your pocket. Everything else is detail.

Let’s Break It Down

How do we even know a reaction happened?

You can’t see atoms rearranging. So how do you know a reaction took place and not just, say, water evaporating? You watch for tell-tale signs.

When magnesium ribbon is burned in air, it flares with a dazzling white flame and leaves behind a white powder (magnesium oxide). When zinc granules meet dilute hydrochloric or sulphuric acid, bubbles of gas fizz up and the tube feels warm. When lead nitrate solution meets potassium iodide solution, a bright yellow solid suddenly appears out of two clear liquids.

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So a chemical reaction has probably taken place if you notice any of these:

• a change in state (solid ↔ liquid ↔ gas as a new substance)
• a change in colour
• the evolution of a gas (bubbling, fizzing)
• a change in temperature (the container warms up or cools down)
• the formation of a precipitate (an insoluble solid) or a new smell

Writing a chemical reaction: from a sentence to a symbol

Describing a reaction in words is clumsy. “When magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide” is a whole sentence. Chemists shorten this in two steps.

Step 1 — the word-equation. Write the names, reactants on the left, products on the right, an arrow showing the direction:

Magnesium + Oxygen → Magnesium oxide

The substances that get used up (magnesium, oxygen) are the reactants; the new substance formed (magnesium oxide) is the product. The ”+” means “reacts with” on the left and “and” on the right. The arrow ”→” means “gives” or “produces.”

Step 2 — the chemical equation. Replace names with chemical formulae:

Mg + O₂ → MgO

This is shorter and far more useful — but right now it has a hidden problem. Count the oxygen atoms: two on the left (O₂), but only one on the right (MgO). Atoms have gone missing, which the Law of Conservation of Mass forbids. An equation like this — correct formulae but unequal atoms — is called a skeletal chemical equation. We must balance it.

Balancing equations, step by step

Balancing means putting coefficients (the big numbers in front of formulae) so that every element has equal atoms on both sides. The golden rule:

You may change the coefficients (the numbers in front). You may never change the subscripts (the small numbers inside a formula) — that would turn the substance into a different substance.

The standard method is hit-and-trial: balance one element at a time, using the smallest whole numbers, then check. Let’s walk through the classic example.

Once balanced, we make the equation even more informative by adding the physical states in brackets: (s) solid, (l) liquid, (g) gas, and (aq) for aqueous (dissolved in water). Conditions like heat, light, a catalyst or pressure go above or below the arrow.

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Here (g) on water tells us it is used as steam. Now the equation tells the full story: what reacts, what forms, in what state, and in what amounts.

The five families of chemical reactions

Atoms never change into other atoms during a reaction; they just regroup. The way they regroup lets us sort almost every reaction into five families.

1. Combination reactions — two become one

When two or more substances join to form a single product, it’s a combination reaction. When quicklime (calcium oxide) meets water, it forms slaked lime, hissing and releasing a lot of heat:

CaO(s) + H₂O(l) → Ca(OH)₂(aq) + heat

Other examples: burning coal (C + O₂ → CO₂) and the formation of water (2H₂ + O₂ → 2H₂O). Notice the pattern — many reactants, one product.

Reactions like the slaking of lime release heat. Reactions that release heat are called exothermic reactions. Burning natural gas, respiration in our cells, and even the rotting of vegetables into compost are all exothermic:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) (burning natural gas — exothermic)

C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + energy (respiration)

2. Decomposition reactions — one becomes many

A decomposition reaction is the exact opposite: a single substance breaks down into two or more simpler substances. It always needs an energy “push,” and the form of that push gives it a special name:

• Thermal decomposition — broken apart by heat. Heating limestone gives quicklime; heating ferrous sulphate or lead nitrate gives coloured fumes:

  CaCO₃(s) →[heat] CaO(s) + CO₂(g) 2Pb(NO₃)₂(s) →[heat] 2PbO(s) + 4NO₂(g) + O₂(g) (brown fumes of NO₂)

• Electrolytic decomposition — broken apart by electricity. Passing current through acidified water splits it into hydrogen and oxygen:

  2H₂O(l) →[electricity] 2H₂(g) + O₂(g)

• Photolytic decomposition — broken apart by light. White silver chloride turns grey in sunlight (used in black-and-white photography):

  2AgCl(s) →[sunlight] 2Ag(s) + Cl₂(g)

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Because decomposition reactions absorb energy, they are endothermic. That’s a neat mirror image: combination often gives heat out (exothermic), decomposition takes energy in (endothermic).

Feature	Combination	Decomposition
What happens	Two or more → one product	One reactant → two or more products
Energy	Usually releases energy (often exothermic)	Needs energy: heat, light or electricity (endothermic)
Example	CaO + H₂O → Ca(OH)₂	CaCO₃ → CaO + CO₂

3. Displacement reactions — the stronger metal pushes out the weaker

When a more reactive element kicks a less reactive element out of its compound, that’s a displacement reaction. Dip an iron nail in blue copper sulphate solution: the blue fades to pale green and the nail turns brownish, because iron (more reactive) displaces copper (less reactive):

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

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Zinc and lead can do the same to copper, because both are more reactive than copper:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

4. Double displacement reactions — partners swap

When two compounds in solution exchange their ions, it’s a double displacement reaction. Mix sodium sulphate and barium chloride solutions and a white insoluble solid drops out:

Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)

The barium and sodium have swapped partners. The insoluble solid (BaSO₄) that appears is called a precipitate, so any reaction that produces a precipitate is also called a precipitation reaction.

Feature	Displacement	Double displacement
What swaps	One element replaces another	Two compounds exchange ions (partners swap)
Reactants	An element + a compound	Two compounds (usually in solution)
Example	Fe + CuSO₄ → FeSO₄ + Cu	Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl

5. Oxidation and reduction — the gain and loss game

Heat copper powder in air and its surface turns black — copper has gained oxygen to become copper oxide:

2Cu + O₂ →[heat] 2CuO

Now pass hydrogen gas over that hot black copper oxide and it turns brown again — the copper oxide loses oxygen back to copper:

CuO + H₂ →[heat] Cu + H₂O

Here are the definitions, and they’re broader than you might first think:

• Oxidation = gain of oxygen or loss of hydrogen.
• Reduction = loss of oxygen or gain of hydrogen.

In that second reaction, CuO is reduced (loses oxygen) while H₂ is oxidised (gains oxygen). One can’t happen without the other — so these are called oxidation–reduction or redox reactions. An easy memory hook is OIL RIG: Oxidation Is Loss, Reduction Is Gain (of hydrogen — and the reverse for oxygen).

Process	Oxygen	Hydrogen
Oxidation	Gains oxygen	Loses hydrogen
Reduction	Loses oxygen	Gains hydrogen

Oxidation in everyday life: corrosion and rancidity

Redox reactions aren’t just lab curiosities — two slow ones cost the world a fortune.

Corrosion is the slow eating-away of a metal when it’s attacked by moisture, air and chemicals around it. The rusting of iron (a reddish-brown coat), the black tarnish on silver, and the green film on copper are all corrosion. Corrosion silently weakens car bodies, bridges, railings and ships — which is why we paint, grease, or galvanise iron to keep oxygen and water away.

Rancidity is what happens when the fats and oils in food get oxidised: they develop a foul smell and taste. To slow it down, manufacturers add antioxidants, seal food in airtight packets, refrigerate it, and flush packets with nitrogen — an unreactive gas that keeps oxygen out (that “puff” of gas when you open a chips packet is nitrogen doing its job).

Common Mistakes

Quick Check

Practice Problems

Try each one yourself before tapping “Show Solution” — that struggle is where the learning happens. These problems are written by Curriv and are completely free.

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Summary

After this chapter, you should be able to explain each of these in your own words:

• A chemical reaction makes new substances by rearranging atoms; atoms are conserved, so mass is conserved (Law of Conservation of Mass).
• You can tell a reaction happened by signs like a change in state, colour, temperature, the evolution of a gas, or the formation of a precipitate.
• A reaction can be written as a word-equation, then a chemical equation, which must be balanced by adjusting coefficients (never subscripts).
• Physical states (s, l, g, aq) and conditions (heat, light, catalyst) make an equation more informative.
• The five families: combination (many → one), decomposition (one → many, by heat/light/electricity), displacement (a more reactive element pushes out a less reactive one), double displacement (compounds swap ions, often making a precipitate), and oxidation–reduction (redox).
• Exothermic reactions release heat; endothermic reactions absorb it.
• Oxidation = gain of oxygen / loss of hydrogen; reduction = loss of oxygen / gain of hydrogen — and they always happen together.
• Slow redox in daily life shows up as corrosion (e.g. rusting) and rancidity (spoiling of fats and oils), which we fight with paint, antioxidants, airtight packing and nitrogen flushing.

What’s Next

You now know that substances react and how to write those reactions down. A huge share of everyday chemistry — from the tang of a lemon to the relief of an antacid tablet — comes down to one special tug-of-war between two opposite kinds of substances. In Chapter 2: Acids, Bases and Salts, you’ll meet those opposites, learn why some things taste sour and others bitter and slippery, and watch the very same reaction types from this chapter (especially neutralisation, a double displacement) play out in your kitchen and your stomach.