Chemical Reactions and Equations
Leave a glass of milk on the counter through a summer afternoon and by evening it has soured. Leave an iron nail out in the rain and within days a reddish-brown crust creeps over its surface. Bite into a week-old bag of chips and the oil tastes stale. In each case something has happened that stirring or cooling cannot undo — the identity of the substance has changed. That is the signature of a chemical reaction, and this chapter builds the toolkit for describing, writing and classifying reactions like these. The chapter opens with three simple activities — burning a magnesium ribbon, mixing lead nitrate with potassium iodide, dropping zinc into acid — and asks what tells us a reaction has occurred: a change of colour or state, a gas bubbling out, a change in temperature. From there it builds a precise written language for chemistry. A word equation such as magnesium plus oxygen gives magnesium oxide becomes the compact Mg + O2 gives MgO, and because atoms can never simply vanish or appear, that equation must be balanced so the same number of atoms of every element sits on both sides — the law of conservation of mass made visible in a formula. Once you can write and balance an equation, the chapter sorts every reaction you will meet in school chemistry into recognisable families: combination reactions that build one substance out of several, decomposition reactions that break one substance into several, displacement reactions where a reactive element evicts a less reactive one from a compound, and double displacement reactions where two compounds swap partners, often dropping an insoluble precipitate. Alongside this runs oxidation and reduction — the gain or loss of oxygen or hydrogen — which explains two things you already know without a chemistry name for them: the rusting of iron (corrosion) and the staleness of old oil (rancidity). By the end you should be able to look at any chemical change and describe it precisely, as a balanced equation with state symbols, correctly sorted into its type.
- A chemical reaction is signalled by a change of state, colour, gas evolution or temperature; substances that change are reactants, new substances formed are products.
- A chemical equation writes a reaction using formulae, reactants on the left, products on the right, joined by an arrow pointing toward the products.
- An equation is skeletal until it is balanced: the atom count of every element must match on both sides, from the law of conservation of mass — mass is neither created nor destroyed.
- State symbols (s), (l), (g), (aq) and conditions over the arrow (heat, sunlight, catalyst, pressure) turn a bare equation into a complete one.
- Combination reactions join two or more substances into one product; decomposition reactions break one substance into two or more — exact opposites of each other.
- Reactions releasing heat are exothermic (combination, respiration); reactions absorbing heat, light or electricity are endothermic (most decomposition).
- In a displacement reaction a reactive element ousts a less reactive one from its compound; in a double displacement reaction two compounds exchange ions, often dropping an insoluble precipitate.
- Oxidation is gain of oxygen or loss of hydrogen; reduction is loss of oxygen or gain of hydrogen; a reaction where both happen is a redox reaction.
- Corrosion (rusting, tarnishing) and rancidity (spoiling fats and oils) are everyday, slow oxidation reactions with real costs.
- Key idea: every chemical change is a balanced chemical equation obeying the law of conservation of mass, sorted into combination, decomposition, displacement or double displacement, and further understood through the gain or loss of oxygen and hydrogen as oxidation and reduction.
Every chemical reaction can be represented by a chemical equation that must be balanced to satisfy the law of conservation of mass, and every such reaction can be classified as combination, decomposition, displacement or double displacement while also being tracked as oxidation or reduction through the gain or loss of oxygen and hydrogen.
What this chapter covers
The chapter moves through four connected stages. First, it establishes how we recognise a chemical reaction — change of state, colour, temperature or gas evolution — using three simple activities as evidence. Second, it builds the written language of chemistry: word equation, skeletal equation and balanced equation, arriving at the hit-and-trial balancing method and the reasons an equation needs state symbols and conditions. Third, it sorts reactions into families: combination, decomposition (thermal, electrolytic, photolytic), displacement and double displacement, alongside exothermic and endothermic change and precipitation. Fourth, it introduces oxidation and reduction through the gain or loss of oxygen and hydrogen, and uses this to explain two everyday phenomena — corrosion and rancidity. Carry one habit through the chapter: whenever you see an equation, first ask whether the atom count of each element actually matches on both sides.
The core ideas
From word equations to chemical equations
A chemical reaction can always be described in a sentence — magnesium ribbon burnt in air changes into magnesium oxide — but a sentence is slow to write and awkward to check. The first shorthand is the word equation, which places reactants on the left and products on the right of an arrow that shows the direction of change.
where: the substances before the arrow, magnesium and oxygen, are the reactants; the substance after the arrow, magnesium oxide, is the product. A plus sign separates multiple reactants or products, and the arrowhead points toward the products. For anything more complex, chemists replace words with chemical formulae, giving a chemical equation such as Mg + O2 gives MgO — compact, universal, and countable atom by atom.
Balancing a chemical equation: the law of conservation of mass in symbols
Count the atoms on each side of Mg + O2 gives MgO and something is wrong: one magnesium and two oxygen atoms sit on the left, only one magnesium and one oxygen on the right. This is a skeletal chemical equation — it names the reactants and products but has not been checked for atom count. You already learned in Class IX that mass can neither be created nor destroyed in a reaction, the law of conservation of mass; since atoms too are neither created nor destroyed, the atom count of each element must be identical before and after.
where: this equality holds only when nothing enters or leaves the reaction vessel — a closed system. To fix a skeletal equation you adjust coefficients (whole numbers placed in front of a formula) until the atom count matches; you may never touch the subscripts inside a formula, because changing a subscript changes the substance itself. Doubling every coefficient in Mg + O2 gives MgO turns it into the balanced equation 2Mg + O2 gives 2MgO, where two magnesium and two oxygen atoms now sit correctly on both sides. This trial-and-improvement process is called the hit-and-trial method, and the full worked procedure for a trickier equation is carried out step by step in the first derivation below.
Writing a complete equation: state symbols and conditions
A balanced equation such as 3Fe + 4H2O gives Fe3O4 + 4H2 is correct arithmetically, but it says nothing about whether the water was liquid or steam, or whether the iron was a solid rod or a fine powder. To make an equation fully informative, chemists attach a state symbol to every formula: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous, meaning dissolved in water.
where: the (g) on H2O here is deliberate — it tells you the water is present as steam, exactly the condition under which red-hot iron reacts with water. Reaction conditions that matter — temperature, pressure, a catalyst or light — are written above or below the arrow rather than as a state symbol; methanol synthesis is written CO(g) + 2H2(g) gives CH3OH(l) with 340 atmospheres over the arrow, and photosynthesis is written with sunlight and chlorophyll over the arrow. State symbols and conditions are not decorative: a question asking for "a balanced equation with state symbols" is graded on both the atom count and the symbols.
Combination and decomposition reactions: opposite processes
When calcium oxide (quick lime) is mixed with water it reacts vigorously, releasing a great deal of heat and leaving behind calcium hydroxide (slaked lime), the substance used for whitewashing walls.
where: two reactants have combined to give a single product, calcium hydroxide. Any reaction in which two or more substances combine to give a single product is a combination reaction; burning coal (C plus O2 gives CO2) and the formation of water from hydrogen and oxygen are further examples. Because heat is released here, this is also an exothermic reaction — respiration, in which glucose combines with oxygen in your cells to release energy, is another everyday exothermic example. Decomposition reactions run in exactly the opposite direction: a single reactant breaks down into two or more simpler products. Heating limestone is the industrial example behind cement manufacture.
where: the delta above the arrow is the standard symbol for heat; a decomposition driven by heat is a thermal decomposition. Ferrous sulphate crystals lose water and break down into ferric oxide, sulphur dioxide and sulphur trioxide on strong heating, turning green to brown-black; lead nitrate decomposes to lead oxide, brown nitrogen dioxide gas and oxygen. Decomposition can also be driven by electricity, as when water splits into hydrogen and oxygen (the figure above), or by light, as when white silver chloride turns grey in sunlight, decomposing into silver and chlorine gas — the same reaction with silver bromide is the chemistry of old black-and-white film. Since decomposition needs an input of heat, light or electricity, these reactions are typically endothermic, the opposite of the exothermic combination reactions above.
Displacement and double displacement reactions
Dip a clean iron nail into blue copper sulphate solution and, after a few minutes, the nail turns brownish and the blue colour visibly fades.
where: iron has pushed copper out of its compound, forming iron sulphate in solution while depositing copper metal on the nail. A reaction where a more reactive element displaces a less reactive one from its compound is a displacement reaction; zinc displacing copper from copper sulphate, and lead displacing copper from copper chloride, follow the same pattern. Mix a colourless solution of sodium sulphate with a colourless solution of barium chloride and a white solid appears immediately, insoluble in the water.
where: here two compounds have swapped partners — sulphate pairs with barium instead of sodium, chloride pairs with sodium instead of barium. Such a reaction, exchanging ions between two reactants, is a double displacement reaction; whenever the exchange produces an insoluble solid like barium sulphate here, it is also a precipitation reaction, and the solid itself is the precipitate.
Oxidation, reduction and redox reactions
Heat copper powder in air and its shiny surface turns black, coated with copper(II) oxide; pass hydrogen gas over that same heated black coating and it turns back to brown copper metal, with water formed as the second product.
where: copper oxide has lost its oxygen atom and hydrogen has gained one. If a substance gains oxygen (or loses hydrogen), it is oxidised; if it loses oxygen (or gains hydrogen), it is reduced. Here copper oxide is reduced to copper, and hydrogen is oxidised to water; because oxidation and reduction always happen together, this whole reaction is a redox reaction. Carbon reducing zinc oxide to zinc while being oxidised to carbon monoxide, and hydrochloric acid being oxidised to chlorine while manganese dioxide is reduced to manganese chloride, are two further textbook redox reactions; the general skill, worked through in the second derivation below, is tracking oxygen and hydrogen atom by atom rather than guessing.
Oxidation in everyday life: corrosion and rancidity
Two slow, everyday oxidation reactions cost households and industries enormous sums every year. When a metal is attacked by moisture, acids or the oxygen in air, it is said to corrode, and the process is called corrosion; the reddish-brown coating on iron (hydrated iron oxide) is rusting, the best-known example, while silver tarnishes black and copper develops a green coating through the same general process. Corrosion damages car bodies, bridges, ships and iron railings, and protecting against it — by painting, oiling, galvanising or plating — is a major branch of applied chemistry. When fats and oils in food are oxidised over time, their smell and taste change and the food is said to have gone rancid; this spoilage, called rancidity, is why chip packets are flushed with unreactive nitrogen gas before sealing, and why antioxidants and airtight containers slow down the oxidation that causes it.
Two worked derivations
The goal here is to turn an unbalanced, skeletal chemical equation into a fully balanced one using only atom-counting and small whole-number coefficients, following the systematic hit-and-trial procedure the chapter uses for iron reacting with steam, Fe + H2O gives Fe3O4 + H2 — a good test case since no element is balanced by accident, so every step is genuinely needed.
Step 1. Box every formula and treat its contents as fixed; nothing inside a box may change, since a changed subscript would be a different substance. List the atoms each side: reactants 1 iron, 2 hydrogen, 1 oxygen; products 3 iron, 2 hydrogen, 4 oxygen.
Step 2. Start with the most atom-rich formula and its most mismatched element. Fe3O4's oxygen, 4 against only 1, is the biggest gap. Raise reactant-side oxygen to 4 by placing a coefficient of 4 in front of H2O, giving 4H2O, since a coefficient multiplies the whole formula without touching its internal ratio.
Step 3. Now Fe + 4H2O gives Fe3O4 + H2. Oxygen is balanced, 4 and 4, but hydrogen is 8 on the left against 2 on the right. Place a 4 in front of H2, giving 4H2, raising product-side hydrogen to 8.
Step 4. Now Fe + 4H2O gives Fe3O4 + 4H2, with only iron unmatched: 1 against 3. Place a 3 in front of Fe, giving 3Fe.
Step 5. The full equation, 3Fe + 4H2O gives Fe3O4 + 4H2, checks out: iron 3 and 3, hydrogen 8 and 8, oxygen 4 and 4. Since this happens on heating iron with steam, the complete equation is 3Fe(s) + 4H2O(g) gives Fe3O4(s) + 4H2(g). Verified: reactant atoms 3 Fe, 8 H, 4 O; product atoms 3 Fe, 8 H, 4 O — every count matches.
Go deeper · Derivation B · 1 minIdentifying which substance is oxidised and which is reduced in a redox reaction
The goal here is a reliable, step-by-step method for deciding which reactant is oxidised and which is reduced in any equation, rather than memorising one or two textbook answers — the method rests only on tracking whether each substance gains or loses oxygen (or hydrogen), applied first to the copper oxide and hydrogen reaction and then checked against an independent example.
Step 1. Write the balanced equation and identify every reactant and product: CuO(s) + H2(g) gives Cu(s) + H2O(l). Reactants: copper(II) oxide and hydrogen gas. Products: copper metal and water.
Step 2. Track the oxygen atom. It starts bonded to copper, inside CuO, and ends bonded to hydrogen, inside H2O. Copper has lost its oxygen (CuO to Cu) while hydrogen has gained one (H2 to H2O).
Step 3. Apply the definitions to what Step 2 found. A substance that loses oxygen is reduced, so copper oxide is reduced to copper. A substance that gains oxygen is oxidised, so hydrogen is oxidised to water. Both happen together, exactly why this is called a redox reaction — one reactant's gain balancing the other's loss.
Step 4. Check the method on an independent example: ZnO + C gives Zn + CO. Zinc oxide loses oxygen to become plain zinc, so it is reduced; carbon gains oxygen to become carbon monoxide, so it is oxidised. The same reasoning, extended to hydrogen (losing hydrogen means oxidised, gaining hydrogen means reduced), covers reactions where oxygen never appears, such as hydrogen sulphide oxidised to sulphur by losing its hydrogen. Verified: in CuO + H2 gives Cu + H2O, oxygen moves from Cu to H2, so CuO is reduced and H2 is oxidised; in ZnO + C gives Zn + CO, oxygen moves from Zn to C, so ZnO is reduced and C is oxidised.
Worked examples
Board level - Balancing a double displacement reaction and adding state symbols
Solutions of barium chloride and aluminium sulphate react in water to give a precipitate of barium sulphate and a solution of aluminium chloride. Write the balanced chemical equation with state symbols, and name the type of reaction.
Step 1. Write the word equation, then the skeletal formula equation: BaCl2 + Al2(SO4)3 gives BaSO4 + AlCl3.
Step 2. Count atoms each side. Reactants: 1 barium, 2 chlorine, 2 aluminium, 3 sulphur, 12 oxygen. Products (per formula unit): 1 barium, 1 sulphur, 4 oxygen, 1 aluminium, 3 chlorine. Only barium matches, so balancing is needed.
Step 3. Balance sulphate first, since 3 sulphate groups sit on the left. Placing 3 in front of BaSO4 also fixes barium, since 3 barium now needs 3BaCl2: 3BaCl2 + Al2(SO4)3 gives 3BaSO4 + AlCl3.
Step 4. Balance chlorine and aluminium last. 3BaCl2 gives 6 chlorine and Al2(SO4)3 gives 2 aluminium; matching these needs 2AlCl3. The fully balanced equation is 3BaCl2 + Al2(SO4)3 gives 2AlCl3 + 3BaSO4.
Step 5. Add state symbols: reactants dissolved, so (aq); barium sulphate is the insoluble precipitate, so (s); aluminium chloride stays dissolved, so (aq). Answer: 3BaCl2(aq) + Al2(SO4)3(aq) gives 2AlCl3(aq) + 3BaSO4(s), a double displacement (precipitation) reaction.
Board level - The thermit reaction as displacement and as a redox reaction
Iron(III) oxide reacts with aluminium powder to give aluminium oxide and molten iron, releasing enormous heat — the thermit reaction used to weld railway tracks. Verify Fe2O3(s) + 2Al(s) gives Al2O3(s) + 2Fe(s) is balanced, name the reaction type, and identify what is oxidised and what is reduced.
Step 1. Count atoms each side. Reactants: 2 iron, 3 oxygen, 2 aluminium. Products: 2 aluminium, 3 oxygen, 2 iron. All three match, so the equation is already balanced.
Step 2. Aluminium, the more reactive metal, has taken iron's place in the oxide, pushing iron out as the free metal — a displacement reaction.
Step 3. Track the oxygen: it starts on iron, in Fe2O3, and ends on aluminium, in Al2O3. Fe2O3 has lost oxygen, so it is reduced to iron; aluminium has gained oxygen, so it is oxidised to aluminium oxide. Since both happen together, this displacement reaction is also a redox reaction.
Step 4. So much heat is released that the iron produced is molten, which is why this reaction welds railway tracks on site. Answer: the equation is already balanced; it is a displacement reaction that is also a redox reaction, with Fe2O3 reduced and Al oxidised, releasing a large amount of heat.
Board HOTS / foundation JEE-NEET - Stoichiometry of a displacement reaction from atomic masses
Zinc granules are added to excess silver nitrate solution: Zn(s) + 2AgNO3(aq) gives Zn(NO3)2(aq) + 2Ag(s). Taking the atomic mass of zinc as 65 u and silver as 108 u, calculate the mass of silver deposited when 65 g of zinc reacts completely, name the reaction type, and identify the substance oxidised.
Step 1. Read the ratio from the coefficients: 1 formula unit of Zn reacts with 2 of AgNO3 to deposit 2 atoms of Ag. In mass terms, one atomic mass of zinc (65 g) corresponds to two atomic masses of silver (2 times 108 u = 216 g).
Step 2. The question gives exactly 65 g of zinc, one atomic-mass-worth, so it produces the matched amount of silver from Step 1: 216 g.
Step 3. No oxygen changes hands here, so use the broader redox view: zinc loses electrons to become Zn(NO3)2, a positive ion, while silver ions gain electrons to become neutral silver metal. Zinc is oxidised (it loses electrons); silver ion is reduced.
Step 4. Name the type: zinc, more reactive than silver, displaces silver from its compound, so this is a displacement reaction, and since electrons transfer it is simultaneously a redox reaction. Answer: 216 g of silver is deposited; the reaction is a displacement (and redox) reaction; zinc is oxidised. Timing note: aim for under three minutes. The chain is coefficients to mass ratio to scaled answer — read the coefficients first, since the whole calculation is just 65 g of Zn is to 216 g of Ag as 1 is to 2. The one trap is forgetting the factor of 2 in front of AgNO3 and Ag; skipping straight from atomic mass to atomic mass without checking coefficients is the commonest error here.
Common traps
The single most common mistake is balancing an equation by changing a subscript inside a formula — turning H2O into H2O2 to add an oxygen atom — which silently changes the substance rather than balancing anything; only coefficients in front of a formula may be adjusted. A second trap is forgetting state symbols, or assigning them without thought — (s) for a solution, (aq) for a dry solid. Students often mix up combination and decomposition when a product happens to be a compound, forgetting the real test is how many separate substances start on each side. Confusing displacement with double displacement is common too: a true displacement swaps one element, a double displacement exchanges ions between two compounds. In oxidation and reduction, a frequent slip is tracking the wrong atom instead of identifying it first. Finally, students often think oxidation means only burning with a visible flame, missing that rusting and rancidity are oxidation too, just without fire.
You should be able to…
- Recognise, from the four types of evidence (state, colour, gas, temperature), that a reaction has taken place.
- Write a word equation, with reactants on the left and products on the right.
- Convert a word equation into a skeletal chemical equation using correct formulae.
- Balance a skeletal equation by hit-and-trial, adjusting only coefficients, never subscripts.
- Add state symbols (s), (l), (g), (aq) and reaction conditions above the arrow.
- Classify a reaction as combination, decomposition, displacement or double displacement, and justify it.
- Decide whether a reaction is exothermic or endothermic from what is released or absorbed.
- Identify which reactant is oxidised and which is reduced, by tracking oxygen or hydrogen.
- Explain corrosion and rancidity as everyday oxidation reactions, with one method to slow each.
Ask the tutor: Want to practise balancing a tricky equation step by step, or work through a few more reaction-type and oxidation-reduction identification questions together?
Why can I not just add a subscript to balance an equation, since it changes the atom count just as easily as a coefficient does?
Because a subscript is part of a formula's identity, not a multiplier you are free to adjust. Changing H2O to H2O2 raises the oxygen count, but H2O2 is hydrogen peroxide, a completely different substance from water — you would no longer be describing the same reaction. A coefficient in front of a formula, like the 4 in 4H2O, means four separate, unchanged water molecules, exactly what balances atom counts without inventing new substances. Examiners check this rule first: an equation "balanced" by editing a subscript is marked wrong even if the atom count works out.
Respiration does not look like a chemical reaction to me — there is no flame or fizzing. How is it exothermic?
Respiration is glucose reacting with oxygen inside your cells, written C6H12O6(aq) + 6O2(aq) gives 6CO2(aq) + 6H2O(l) plus energy, and it satisfies every test for a chemical reaction: new substances form from old ones and energy is released, just as when magnesium burns. The difference is only how the energy is released — respiration happens gradually, at body temperature, inside cells, rather than all at once as a flame. That energy still keeps you warm and powers your body, which is exactly what "exothermic" means: energy given out with the products, seen or not.
What is the actual difference between corrosion and rancidity, since both seem to be about things going bad over time?
Both are slow oxidation reactions acting on different substances. Corrosion is the oxidation of a metal by moisture, oxygen or acids in the air — rusting of iron, tarnishing of silver, the green coating on copper — converting the metal into a compound at its surface. Rancidity is the oxidation of fats and oils in food, changing smell and taste without forming a solid crust; it happens to organic molecules, not metals. The chemistry is oxidation in both cases, but the fix differs: painting or galvanising a metal, versus antioxidants or airtight, nitrogen-flushed packaging for food.
If oxidation is gaining oxygen, how can hydrogen sulphide be oxidised in a reaction that does not involve oxygen changing hands?
The definition widens once you look closely: a substance is also oxidised if it loses hydrogen, even with no oxygen atom directly involved, since losing hydrogen and gaining oxygen are two faces of the same underlying process studied further in later grades. At Class 10 level, use whichever half of the rule fits the equation in front of you — check first whether oxygen is gained or lost, and if oxygen genuinely does not appear, check whether hydrogen is gained or lost instead. Both checks test the same underlying idea, from two angles.