Chemical Reactions and Equations

Why This Matters

Leave a glass of milk out on a summer afternoon. By evening, it turns sour. Leave an iron tawa in a damp corner of the kitchen. A week later, it has a reddish-brown coat on it. Eat a piece of bread. Within a few hours, your body turns it into the energy that lets you read this sentence.

All three of these are really the same thing happening. In each one, a substance quietly turned into a new substance. The milk changed into something different. The iron’s surface became something that is no longer iron. The bread became glucose, and then carbon dioxide, water and energy.

This “turning into something new” is what scientists call a chemical reaction.

Once you can spot reactions, name them, and write them down correctly, the rest of chemistry becomes much easier. This chapter teaches you how to do exactly that. You will learn how to see when a reaction has happened, how to write it as a balanced equation, and how to sort every reaction into a few simple groups. After this one chapter, chemistry will stop looking like magic and start looking like something you can read and understand.

The Big Idea

The whole chapter rests on a few words from Class 9 — atom, element, molecule and compound. Let us warm them up before we go further:

Here is the one big idea that the whole chapter is built on:

In a chemical reaction, atoms are never created and never destroyed. They are only rearranged into new groups.

Think of atoms like a fixed set of LEGO bricks. A reaction takes apart some shapes and builds new ones. But you always end with the exact same bricks you started with. Not one brick more, not one brick less.

This idea has a name. It is the Law of Conservation of Mass, which you met in Class 9. It says that mass cannot be created or destroyed in a chemical reaction. (“Conservation” just means nothing is lost.)

But why is this true? Why can a reaction never make an atom appear out of nowhere, or make one vanish? Figure 1.1 answers that by following the very same atoms through a reaction, so you can count them before and after:

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Here is the reason, step by step. An atom is a real, solid object — the tiniest building block of matter. A chemical reaction does just one thing: it breaks the joins (bonds) between atoms and makes new joins. That is all. It cannot crush an atom into nothing, and it cannot build a fresh atom from nothing. (Making or destroying an atom would need a nuclear reaction, which is a completely different thing — not ordinary chemistry.)

So think of it like rearranging the same handful of LEGO bricks into a new model. You take the model apart and build something else, but you still hold every brick you started with. Because the same atoms come out the other end, the number of each kind of atom on the left must equal the number on the right. This is exactly why every correct equation has to be balanced — balancing is just us making the equation honest about the atoms that were there all along.

Two important things follow from this idea, and both matter for the rest of the chapter:

1. New substances form during a reaction. So reactions show visible signs — a colour change, a gas bubbling out, a change in temperature, a new smell, or a solid appearing.
2. Atoms are never lost. So every correct chemical equation must be balanced. “Balanced” means the same number of atoms of each element on both sides of the arrow.

Keep these two ideas in mind. Everything else in this chapter builds on them.

Let’s Break It Down

How do we even know a reaction happened?

You cannot see atoms rearranging. They are far too small. So how do you know a reaction took place, and not just something simple like water drying up? You look for clues.

Here are three examples. When you burn magnesium ribbon in air, it gives a very bright white flame and leaves behind a white powder (this powder is magnesium oxide). When you drop zinc granules (small pieces of zinc) into dilute hydrochloric acid or sulphuric acid, bubbles of gas rise up and the test tube feels warm. When you mix lead nitrate solution with potassium iodide solution, a bright yellow solid suddenly forms out of two clear liquids.

Figure 1.2 shows what that burning magnesium ribbon actually looks like:

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

• a change in state (a new substance appears as a solid, liquid or gas)
• a change in colour
• a gas being given off (bubbling or fizzing)
• a change in temperature (the container becomes warmer or cooler)
• a precipitate forming (a solid that does not dissolve), or a new smell

Try spotting the signs yourself in this everyday kitchen example:

Writing a chemical reaction: from a sentence to a symbol

Writing a reaction as a full sentence takes a lot of words. For example: “When magnesium ribbon is burnt in oxygen, it changes into magnesium oxide.” That is a long sentence. Scientists make it shorter in two steps.

Step 1 — the word-equation. Write the names of the substances. Put the starting substances on the left and the new substances on the right. Draw an arrow in between to show the direction:

Magnesium + Oxygen → Magnesium oxide

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

Step 2 — the chemical equation. Now replace the names with chemical formulae (short symbols for each substance). Before we do, here is a quick reminder of how those symbols and formulae work, from Class 9:

Mg + O₂ → MgO

This is much shorter and more useful. But it has a hidden problem. Let us count the oxygen atoms. On the left there are two oxygen atoms (in O₂). On the right there is only one oxygen atom (in MgO). So one oxygen atom seems to have vanished. But the Law of Conservation of Mass says atoms can never just disappear. An equation like this — correct formulae, but an unequal number of atoms on the two sides — is called a skeletal chemical equation. We must balance it before it is correct.

Balancing equations, step by step

To balance an equation, we add coefficients. A coefficient is a big number written in front of a formula. For example, the 2 in “2H₂O” is a coefficient. We choose these numbers so that each element has the same number of atoms on both sides. Here is the golden rule:

You are allowed to change the coefficients (the big numbers in front). You must never change the subscripts (the small numbers inside a formula). Changing a subscript would turn the substance into a completely different substance.

(A “subscript” is the small lowered number inside a formula, like the 2 in H₂O. It tells you how many atoms of that element are in one unit of the substance.)

Figure 1.3 shows balancing as a weighing scale — watch how only the front numbers change:

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The usual method is called hit-and-trial. You balance one element at a time, use the smallest whole numbers you can, and then check at the end. Let us work through the classic example together.

Once the equation is balanced, we can add even more information. We write the physical state of each substance in brackets after it: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous, which means “dissolved in water”. We also write any special conditions — like heat, light, a catalyst or pressure — above or below the arrow.

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Here, the (g) written after water tells us the water is being used as steam (water in gas form). Now the equation tells the full story: what reacts, what forms, in what state, and in what amounts.

Here is a trap that catches a lot of students — see if you can explain why it is wrong:

Make sure the coefficient-vs-subscript difference is really clear in your head:

The five families of chemical reactions

During a reaction, atoms never turn into other kinds of atoms. They just move around and form new groups. The way they regroup lets us sort almost every reaction into five families (groups). Figure 1.4 sums up the main four at a glance before we take them one by one:

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1. Combination reactions — two become one

When two or more substances join together to make a single product, it is called a combination reaction. For example, when quicklime (calcium oxide) is added to water, it forms slaked lime. The mixture hisses and gives out a lot of heat:

CaO(s) + H₂O(l) → Ca(OH)₂(aq) + heat

Two more examples are the burning of coal (C + O₂ → CO₂) and the making of water (2H₂ + O₂ → 2H₂O). Notice the pattern in all of them: many reactants come in, but only one product comes out.

The slaking of lime gives out heat. Reactions that give out heat are called exothermic reactions. (An easy way to remember: “exo” sounds like “exit”, and heat exits the reaction.) Burning natural gas, respiration inside our body cells, and even vegetables rotting 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 of a combination reaction. Here, a single substance breaks down into two or more simpler substances. This always needs some energy to make it happen. The kind of energy used gives the reaction a special name:

• Thermal decomposition — broken apart by heat (“thermal” means heat). Heating limestone gives quicklime. Heating ferrous sulphate or lead nitrate gives off coloured fumes (smoke-like gases):

  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. When you pass electric current through acidified water (water with a little acid added), it splits into hydrogen and oxygen:

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

• Photolytic decomposition — broken apart by light (“photo” means light). White silver chloride turns grey when left in sunlight. This was once used in black-and-white photography:

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

Figure 1.5 shows how the electrolysis one is actually set up, with the two gases collecting in separate tubes:

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Decomposition reactions take in (absorb) energy, so they are called endothermic reactions. (“Endo” means inside — energy goes in.) This is a neat pair to remember: combination reactions often give heat out (exothermic), while decomposition reactions take energy in (endothermic).

But why does one reaction pour out heat while another keeps drinking it in? Where does that heat even come from, or go to? The answer is hidden in the bonds between atoms, and Figure 1.6 shows it as a simple money budget:

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Here is the idea in plain words. Every bond between two atoms is like a tiny stretched spring storing energy. To break a bond, you have to spend energy (you pull the spring apart). To make a new bond, energy is given back (the spring snaps shut and releases it).

A reaction does both: first it breaks the old bonds in the reactants, then it builds new bonds in the products. So it is just a budget. Add up what you spend breaking, and what you earn building.

• If the new bonds give back more energy than the old bonds cost to break, the extra energy has nowhere to stay — it escapes, usually as heat. The reaction feels warm. That is exothermic (like burning, or slaking lime).
• If breaking the old bonds costs more than the new bonds give back, the reaction is short of energy. It can only keep going if you keep feeding heat in. That is endothermic (like heating limestone — it stops the instant you stop heating).

This also clears up a confusing point. Burning needs a match to start, yet it is exothermic. The match only pays the small upfront cost of breaking the first few bonds. Once the reaction gets going, it earns back far more than that, so the flame keeps itself alive. An endothermic reaction is different — it runs at a loss the whole time, so it needs heat fed in non-stop.

This table puts combination and decomposition side by side so the contrast is easy to revise:

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

Now test both ideas together — family and energy — on one reaction:

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

In a displacement reaction, a more reactive element pushes a less reactive element out of its compound and takes its place. (“More reactive” just means it reacts more easily.) For example, dip an iron nail into blue copper sulphate solution. After some time, the blue colour fades to pale green and the nail turns brownish. This happens because iron is more reactive than copper, so iron pushes the copper out:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

You can actually watch this colour change happen, as Figure 1.7 shows:

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

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

4. Double displacement reactions — partners swap

In a double displacement reaction, two compounds in solution swap their ions with each other. (Ions are charged parts of a compound; you can think of them here as the two “partners” inside each compound.) For example, mix sodium sulphate solution with barium chloride solution. A white solid that does not dissolve drops to the bottom:

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

Here, barium and sodium have swapped partners. That solid which does not dissolve (BaSO₄) is called a precipitate. So any reaction that forms a precipitate is also called a precipitation reaction.

These two “displacement” reactions sound alike but work differently — here they are side by side:

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 some copper powder in air, and its surface turns black. The copper has gained oxygen to become copper oxide:

2Cu + O₂ →[heat] 2CuO

Now pass hydrogen gas over that hot black copper oxide. It turns brown again. This time the copper oxide loses its oxygen and goes back to being copper:

CuO + H₂ →[heat] Cu + H₂O

Here are the two key definitions. Notice that each one has two parts:

• Oxidation = gaining oxygen or losing hydrogen.
• Reduction = losing oxygen or gaining hydrogen.

Look again at that second reaction. CuO loses oxygen, so CuO is reduced. At the same time, H₂ gains oxygen, so H₂ is oxidised. One cannot happen without the other. When one substance is oxidised, another is always reduced in the same reaction. Because of this, such reactions are called oxidation–reduction reactions, or redox reactions for short (“red” from reduction + “ox” from oxidation).

But why must they always come as a pair? Why can a substance never just lose oxygen on its own? Figure 1.8 shows the reason by tracking the single oxygen atom as it moves:

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The reason is simple once you see it. When CuO is reduced, it loses an oxygen atom. But that oxygen atom does not disappear — by the very rule we just learnt, atoms are never destroyed. So it has to go somewhere. In this reaction it lands on the hydrogen, turning it into water. The hydrogen has now gained oxygen, so the hydrogen is oxidised.

See what happened? The same oxygen atom leaving one substance and joining another is one single event, looked at from two sides. The substance that drops the oxygen is reduced. The substance that catches it is oxidised. You cannot have one without the other, for the same reason you cannot pass a ball to a friend without someone catching it. That is why we give the pair one combined name — redox.

Figure 1.9 lays out both halves once more and names them:

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If you ever mix them up, this small table is all you really need to memorise:

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

Now spot the reduced substance in a reaction that has no hydrogen at all:

Oxidation in everyday life: corrosion and rancidity

Redox reactions are not just something you see in a lab. Two slow redox reactions cause huge problems in real life.

Corrosion is the slow eating-away of a metal. It happens when the metal is attacked by moisture (water in the air), air, and other chemicals around it. The rusting of iron (its reddish-brown coat), the black layer that forms on silver, and the green film on copper are all examples of corrosion. Over time, corrosion slowly weakens car bodies, bridges, railings and ships. This is why we paint, grease, or galvanise iron to protect it. (Galvanising means coating iron with a layer of zinc.) These coatings keep oxygen and water away from the metal.

Rancidity is what happens when the fats and oils in food get oxidised. The food then develops a bad smell and a bad taste — you may have noticed this in old biscuits or fried snacks. To slow this down, food makers add antioxidants (chemicals that stop oxidation), seal food in airtight packets, keep it in the fridge, and fill packets with nitrogen gas. Nitrogen does not react easily, so it keeps oxygen away from the food. (That little “puff” of gas you hear when you open a chips packet is this nitrogen escaping.)

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Practice Problems

Try each problem on your own before you tap “Show Solution”. The effort you put in is what helps you learn. 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 part of everyday chemistry comes down to one special tug-of-war between two opposite kinds of substances. This is true for many things you see daily, from the sour taste of a lemon to the relief you feel after taking an antacid tablet. In Chapter 2: Acids, Bases and Salts, you will meet these two opposites. You will learn why some things taste sour while others taste bitter and feel slippery. And you will see the same reaction types from this chapter — especially neutralisation, which is a kind of double displacement — happening in your kitchen and even in your stomach.