1. What is a Chemical Reaction?
Around us, substances change all the time. Some changes are temporary (like melting of ice), while others make entirely new substances (like burning of paper). When a change produces new substances with different properties, we say a chemical reaction has occurred.
A process in which one or more substances (called reactants) change into new substances (called products) with completely different properties is called a Chemical Reaction.
Reactants and Products
- Reactants – the substances that take part in the reaction.
- Products – the new substances formed as a result of the reaction.
When iron is kept in copper sulphate solution, the blue colour of the solution fades and a brown layer of copper is deposited on the iron nail. Iron and copper sulphate are the reactants; copper and iron sulphate are the products.
Everyday Examples of Chemical Reactions
- Burning of magnesium ribbon in air to form magnesium oxide.
- Formation of curd from milk.
- Digestion of food in the body.
- Rusting of iron articles.
- Burning of LPG to cook food.
A physical change does not form new substances (melting of wax). A chemical change forms new substances (burning of wax). Chemical reactions are always chemical changes.
2. How can we Identify a Chemical Reaction?
A chemical reaction can usually be recognised by one or more observable changes. Look for these indicators:
| Indicator | What we observe | Example |
|---|---|---|
| Change in state | A solid, liquid or gas changes state | Burning of candle wax, decomposing water into hydrogen and oxygen gases |
| Change in colour | The colour of a substance changes | Green ferrous sulphate turns reddish-brown on heating |
| Evolution of gas | Bubbles of a gas escape | Zinc reacts with dilute sulphuric acid to release hydrogen gas |
| Change in temperature | Heat is given out or absorbed | Quicklime reacts with water releasing a lot of heat |
| Formation of precipitate | An insoluble solid settles down in a solution | Barium chloride + sodium sulphate forms white barium sulphate |
When a burning magnesium ribbon is placed on a watch glass, it gives out white light and forms white magnesium oxide ash. This shows both a colour change and a new product, so a chemical reaction has taken place.
3. Writing Chemical Equations
A chemical reaction can be described in two ways. First, we express it in words (a word equation); then we express it in symbols and formulas (a chemical equation).
Word equation: Magnesium + Oxygen → Magnesium oxide
Chemical equation: Mg + O₂ → MgO
The arrow (→) means "forms" or "gives". The substances on the left are the reactants; those on the right are the products.
Word Equations
In a word equation, reactants and products are written as words. For example:
- Zinc + Sulphuric acid → Zinc sulphate + Hydrogen
- Copper + Oxygen → Copper(II) oxide
- Calcium carbonate → Calcium oxide + Carbon dioxide (on heating)
Chemical Equations and Formulas
A chemical equation uses chemical symbols and formulas. Each element has a symbol (H for hydrogen, O for oxygen, N for nitrogen), and a compound is represented by its formula (H₂O for water, CO₂ for carbon dioxide). The numbers used in formulas show how many atoms of each element are present. For example, H₂SO₁₀ means two hydrogen atoms, one sulphur atom and four oxygen atoms.
An unbalanced chemical equation is called a skeleton equation (or unbalanced equation). For example, H₂ + O₂ → H₂O is a skeleton equation because the number of atoms on both sides is unequal.
Balanced Chemical Equation
A chemical equation in which the number of atoms of each element is the same on both sides is called a balanced chemical equation. Balancing is needed because of the Law of Conservation of Mass: matter can neither be created nor destroyed in a chemical reaction. Whenever the numbers do not match, we use coefficients (numbers written in front of formulas) to balance the equation.
State Symbols in Chemical Equations
State symbols tell us the physical state of each substance in the reaction. They are written just after the formula in brackets.
| State symbol | Meaning | Example |
|---|---|---|
| (s) | solid | Fe(s) |
| (l) | liquid | H₂O(l) |
| (g) | gas | O₂(g) |
| (aq) | dissolved in water (aqueous) | CuSO₁₀(aq) |
\[ 2H_2(g) + O_2(g) \rightarrow 2H_2O(l) \]
This tells us that two molecules of hydrogen gas combine with one molecule of oxygen gas to give two molecules of liquid water. Sometimes, conditions such as heat (Δ), a catalyst, or high pressure are written above the arrow.
4. Balancing of Chemical Equations
Balancing makes the number of atoms of each element equal on both sides of the equation, following the Law of Conservation of Mass.
During a chemical reaction, the total mass of the reactants equals the total mass of the products. Atoms are neither created nor destroyed in a chemical reaction; they are only rearranged.
Step-by-step method to balance an equation
Worked Examples – Basic Balancing (step by step)
\[ 2H_2 + O_2 \rightarrow 2H_2O \]
\[ N_2 + 3H_2 \rightarrow 2NH_3 \]
\[ 4Fe + 3O_2 \rightarrow 2Fe_2O_3 \]
\[ 2Na + Cl_2 \rightarrow 2NaCl \]
\[ 2Mg + O_2 \rightarrow 2MgO \]
Worked Examples – Moderate Balancing (step by step)
\[ CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O \]
\[ C_3H_8 + 5O_2 \rightarrow 3CO_2 + 4H_2O \]
\[ 4Al + 3O_2 \rightarrow 2Al_2O_3 \]
\[ CaCO_3 + 2HCl \rightarrow CaCl_2 + CO_2 + H_2O \]
\[ 3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2 \]
5. Board-Style Balancing Examples (step by step)
\[ 2Pb(NO_3)_2 \rightarrow 2PbO + 4NO_2 + O_2 \]
\[ 2FeSO_4 \rightarrow Fe_2O_3 + SO_2 + SO_3 \]
\[ Zn + 2HCl \rightarrow ZnCl_2 + H_2 \]
\[ BaCl_2 + Na_2SO_4 \rightarrow BaSO_4 + 2NaCl \]
\[ 2AgNO_3 + CaCl_2 \rightarrow 2AgCl + Ca(NO_3)_2 \]
Always show your balancing steps in the exam even when the answer seems obvious. And never change a chemical formula to force-balance an equation — change only the coefficients.
6. Types of Chemical Reactions
Chemical reactions can be grouped into a few important types. We study four main types:
- Combination reaction
- Decomposition reaction
- Displacement reaction
- Double displacement reaction
Some special types such as precipitation reactions, oxidation-reduction (redox) reactions, and endothermic/exothermic reactions also appear within these categories.
7. Combination Reactions
A reaction in which two or more substances (elements or compounds) combine to form a single new substance is called a Combination Reaction.
General pattern: A + B → AB
How to identify: one reactant (or more) on the left, but only a single product on the right.
Examples
| Equation | Note |
|---|---|
| \[ 2Mg(s) + O_2(g) \rightarrow 2MgO(s) \] | Two reactants combine; magnesium burns with dazzling white light. |
| \[ 2H_2(g) + O_2(g) \rightarrow 2H_2O(l) \] | Hydrogen and oxygen combine to form water. |
| \[ 2Fe(s) + 3Cl_2(g) \rightarrow 2FeCl_3(s) \] | Iron and chlorine combine to form ferric chloride. |
| \[ CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) \] | Quicklime slakes with water releasing heat (exothermic). |
| \[ 2SO_2(g) + O_2(g) \rightarrow 2SO_3(g) \] | Sulphur dioxide oxidises further to sulphur trioxide. |
Combination reactions and energy
Combination reactions may be classified by the heat they release or absorb:
- Exothermic reaction – a reaction that releases heat (CaO + H₂O, burning of coal).
- Endothermic reaction – a reaction that absorbs heat (formation of NO gas from N₂ and O₂).
When white quicklime (calcium oxide) is mixed with water in a glass, the glass becomes hot. This is because of the exothermic combination reaction: CaO + H₂O → Ca(OH)₂.
Students often call burning of a candle a combination reaction. Burning of a candle involves multiple reactions (melting of wax and combustion of wax with oxygen), and it produces more than one product. Do not label it as a simple combination reaction unless it is specifically stated as combustion of a single fuel with oxygen.
8. Decomposition Reactions
A reaction in which a single compound breaks down into two or more simpler substances is called a Decomposition Reaction. It is the opposite of a combination reaction.
General pattern: AB → A + B
How to identify: a single reactant breaks into multiple products. Energy (heat, electricity or light) is usually supplied to break the compound.
Types of decomposition on the basis of energy
| Type | Energy used | Representative equation |
|---|---|---|
| Thermal decomposition | Heat (Δ) | \[ 2H_2O_2 \xrightarrow{\Delta} 2H_2O + O_2 \] |
| Electrolytic decomposition | Electricity | \[ 2H_2O \xrightarrow{electricity} 2H_2 + O_2 \] |
| Photochemical decomposition | Sunlight | \[ 2AgCl \xrightarrow{sunlight} 2Ag + Cl_2 \] |
Worked examples of decomposition
On heating, pale green ferrous sulphate crystals first lose water of crystallisation and turn white, and then decompose on further strong heating into reddish-brown ferric oxide, sulphur dioxide and sulphur trioxide.
\[ 2FeSO_4 \xrightarrow{\Delta} Fe_2O_3 + SO_2 + SO_3 \]
Observation: the smell of burning sulphur (SO₂) is noticed.
Hydrogen peroxide decomposes slowly in light and rapidly when heated or when a catalyst such as manganese dioxide is added.
\[ 2H_2O_2 \xrightarrow[\text{MnO}_2]{} 2H_2O + O_2 \]
When lead nitrate powder is heated in a hard glass test tube, a yellowish brown gas (nitrogen dioxide) is evolved, a reddish yellow residue (PbO) is left, and a glowing splint rekindles in the oxygen produced.
\[ 2Pb(NO_3)_2 \xrightarrow{\Delta} 2PbO + 4NO_2 + O_2 \]
When electric current is passed through acidified water, it splits into hydrogen (collected at the negative electrode) and oxygen (collected at the positive electrode) in the volume ratio 2 : 1.
\[ 2H_2O \xrightarrow{electricity} 2H_2 + O_2 \]
Greyish white silver chloride or pale yellow silver bromide, when kept in sunlight, darkens as silver metal is left behind. This reaction is the basis of black and white photography.
\[ 2AgCl \xrightarrow{sunlight} 2Ag + Cl_2 \]
\[ 2AgBr \xrightarrow{sunlight} 2Ag + Br_2 \]
Black and white photographs are imprinted using the photochemical decomposition of silver bromide. When exposed to light, AgBr decomposes, depositing silver and fixing the image.
Ferrous sulphate is green (FeSO₁₀ · 7H₂O). On strong heating it turns reddish-brown because of Fe₂O₃. Students often write FeO incorrectly; write ferric oxide (Fe₂O₃) for the thermal decomposition product.
9. Displacement Reactions
A reaction in which a more reactive element replaces (is displaced by) a less reactive element from its compound is called a Displacement Reaction.
General pattern: A + BC → AC + B (A is more reactive than B)
How to identify: a free element and a compound on the left become another free element and another compound on the right.
Examples
| Equation | What happens |
|---|---|
| \[ Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s) \] | Iron displaces copper from copper sulphate; blue solution turns pale green. |
| \[ Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s) \] | Zinc displaces copper; blue solution fades and a brown layer of copper deposits. |
| \[ 2Al(s) + 3CuSO_4(aq) \rightarrow Al_2(SO_4)_3(aq) + 3Cu(s) \] | Aluminium displaces copper from copper sulphate. |
| \[ Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g) \] | Magnesium replaces hydrogen from hydrochloric acid, giving hydrogen gas. |
| \[ 2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g) \] | Sodium displaces hydrogen from cold water. |
Reactivity series memory
A more reactive metal takes the place of a less reactive metal. Rough order of reactivity (from most to least reactive, a shortened list): K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag.
Because iron is more reactive than copper, iron pushes copper out of CuSO₁₀. The reverse (copper displacing iron) does not happen in aqueous solution.
Place an iron nail in blue copper sulphate solution for about an hour. Observation: the blue solution gradually turns light green (iron sulphate forms) and a brownish layer of copper deposits on the nail. Many tiny bubbles of hydrogen may also appear because iron reacts with the slightly acidic solution.
Students write the product of Fe + CuSO₁₀ as Fe₂(SO₁₀)₃. Iron displaces copper to form ferrous sulphate (FeSO₁₀), not ferric sulphate, in this reaction.
10. Double Displacement Reactions
A reaction in which two compounds exchange their ions to form two new compounds is called a Double Displacement Reaction.
General pattern: AB + CD → AD + CB
How to identify: two compounds on the left exchange partners to give two different compounds on the right. Often, an insoluble product (precipitate) settles out.
Examples
| Equation | Observation |
|---|---|
| \[ Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq) \] | White precipitate of barium sulphate. |
| \[ 2NaOH(aq) + CuSO_4(aq) \rightarrow Cu(OH)_2(s) + Na_2SO_4(aq) \] | Blue precipitate of copper hydroxide. |
| \[ Pb(NO_3)_2(aq) + 2KI(aq) \rightarrow PbI_2(s) + 2KNO_3(aq) \] | Golden yellow precipitate of lead iodide. |
| \[ AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq) \] | White curdy precipitate of silver chloride (turns grey in sunlight). |
Precipitation reactions
A reaction in which two soluble compounds in solution exchange ions and form an insoluble solid (the precipitate) is called a Precipitation Reaction. It is a special case of a double displacement reaction.
Add sodium sulphate solution to barium chloride solution. Immediate observation: a white precipitate of barium sulphate forms, which is insoluble in water. Write the double displacement equation with the precipitate marked (↓).
In AgNO₃ + NaCl, students wrongly produce AgClO or leave the equation unbalanced. Balance the ions: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq). Silver chloride is water-insoluble (the precipitate); sodium nitrate stays dissolved.
11. Oxidation and Reduction (Redox Reactions)
A reaction in which a substance gains oxygen or loses hydrogen is called Oxidation. The substance is said to be oxidised.
A reaction in which a substance loses oxygen or gains hydrogen is called Reduction. The substance is said to be reduced.
Easy memory: OIL RIG — Oxidation Is Loss of electrons (and gain of oxygen), Reduction Is Gain of electrons (and loss of oxygen).
\[ CuO(s) + H_2(g) \xrightarrow{\Delta} Cu(s) + H_2O(l) \]
- Copper oxide loses oxygen and becomes copper: CuO is reduced.
- Hydrogen gains oxygen and becomes water: H₂ is oxidised.
- Both happen together, so it is a redox reaction (simultaneous oxidation and reduction).
\[ Zn(s) + CuO(s) \rightarrow ZnO(s) + Cu(s) \]
- Zinc gains oxygen → oxidised.
- Copper (II) oxide loses oxygen → reduced.
Oxidising and reducing agents
- Oxidising agent – the substance that supplies oxygen (or removes hydrogen); it is itself reduced. Example: CuO in the reaction above.
- Reducing agent – the substance that removes oxygen (or supplies hydrogen); it is itself oxidised. Example: H₂ in the reaction above.
In a redox question, first spot which substance gains oxygen and which loses oxygen. The one that gains oxygen is oxidised (and is the reducing agent); the one that loses oxygen is reduced (and is the oxidising agent).
Students state "oxidation is gain of oxygen, reduction is loss of oxygen" and wrongly call both the same substance oxidised and reduced. In every redox reaction one species is oxidised and a different species is reduced.
12. Effects of Oxidation in Everyday Life
Oxidation is not only a laboratory concept. Two very common everyday effects caused by oxidation are corrosion and rancidity.
Corrosion
The slow eating up of a metal by the action of air, moisture or a chemical on its surface is called Corrosion.
- Rusting of iron – the reddish-brown flaky layer (rust) is hydrated iron(III) oxide, Fe₂O₃ · xH₂O. For rusting, iron needs both oxygen and moisture (water). Lack of either stops rusting.
- Tarnishing of silver – silver reacts with sulphur compounds in air (and H₂S) to form a black layer of silver sulphide, Ag₂S.
- Green coating on copper – copper reacts with moist carbon dioxide and oxygen to form a green basic carbonate, CuCO₃ · Cu(OH)₂.
Why corrosion is harmful: it destroys metals, weakens structures and machines, and causes huge financial loss.
Prevention of corrosion
- Painting or applying grease/oil to exposed surfaces.
- Galvanisation – coating iron with a thin layer of zinc.
- Alloying – mixing the metal with another metal (e.g. stainless steel does not rust easily).
- Electroplating (e.g. chroming) and storing articles in dry places.
Rancidity
The condition in which the taste and smell of foods containing oil and fat change (become unpleasant) due to oxidation of the fats/oils by atmospheric oxygen is called Rancidity.
The oils/fats become oxidised and the food spoils, developing an off smell and taste. Rancidity is a slow oxidation reaction.
Prevention of rancidity
| Method | Why it works | Example |
|---|---|---|
| Adding antioxidants | Antioxidants stop the oxidation of fats/oils | BHA and BHT added to oily snacks |
| Keeping in airtight containers | Limits exposure to atmospheric oxygen | Cooking oil stored in closed bottles |
| Refrigeration | Low temperature slows down oxidation | Store butter and ghee in the fridge |
| Flushing packets with nitrogen gas | Inert N₂ prevents the oil/fat from oxidising | Packets of chips and other snacks |
Packets of potato chips appear half-filled. The extra space contains nitrogen gas, which is pumped in to push out oxygen and stop the oils from oxidising, keeping the chips fresh and crisp.
Students say "rusting needs only air". Rusting of iron needs both oxygen and moisture. An iron nail in dry air or in boiling (oxygen-free) water does not rust.
13. Important Reactions – Revision Table
Memorise these balanced equations with their type, observation and the concept connected to each.
| Balanced equation | Type | Observation / concept |
|---|---|---|
| \[ 2Mg + O_2 \rightarrow 2MgO \] | Combination | Dazzling white light; white ash of MgO |
| \[ CaO + H_2O \rightarrow Ca(OH)_2 \] | Combination (exothermic) | Heat is released; quicklime slaking |
| \[ 2H_2 + O_2 \rightarrow 2H_2O \] | Combination | Water formed from its elements |
| \[ N_2 + 3H_2 \rightarrow 2NH_3 \] | Combination (Haber process) | Ammonia formation from nitrogen and hydrogen |
| \[ 2Fe + 3Cl_2 \rightarrow 2FeCl_3 \] | Combination | Ferric chloride formation |
| \[ 2FeSO_4 \xrightarrow{\Delta} Fe_2O_3 + SO_2 + SO_3 \] | Thermal decomposition | Smell of burning sulphur; green → reddish brown |
| \[ 2Pb(NO_3)_2 \xrightarrow{\Delta} 2PbO + 4NO_2 + O_2 \] | Thermal decomposition | Brown gas NO₂ evolved |
| \[ 2H_2O \xrightarrow{electricity} 2H_2 + O_2 \] | Electrolytic decomposition | Gases in 2 : 1 volume ratio |
| \[ 2AgBr \xrightarrow{sunlight} 2Ag + Br_2 \] | Photochemical decomposition | Used in black and white photography |
| \[ Fe + CuSO_4 \rightarrow FeSO_4 + Cu \] | Displacement | Blue solution turns pale green, brown Cu deposits |
| \[ Zn + 2HCl \rightarrow ZnCl_2 + H_2 \] | Displacement | Bubbles of hydrogen gas |
| \[ Na_2SO_4 + BaCl_2 \rightarrow BaSO_4 + 2NaCl \] | Double displacement / precipitation | White precipitate of BaSO₁₀ |
| \[ Pb(NO_3)_2 + 2KI \rightarrow PbI_2 + 2KNO_3 \] | Double displacement / precipitation | Golden yellow precipitate of PbI₂ |
| \[ CuO + H_2 \xrightarrow{\Delta} Cu + H_2O \] | Redox (oxidation-reduction) | CuO reduced, H₂ oxidised |
| \[ 3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2 \] | Displacement of H₂ by iron | Steam passed over red-hot iron |
14. Glossary – Key Definitions
| Term | Definition |
|---|---|
| Reactants | Substances that take part in a chemical reaction. |
| Products | New substances formed as a result of a chemical reaction. |
| Chemical equation | Representation of a chemical reaction using symbols and formulas of reactants and products. |
| Balanced chemical equation | An equation with equal numbers of atoms of each element on both sides. |
| Skeleton equation | An unbalanced chemical equation. |
| Law of Conservation of Mass | Mass is neither created nor destroyed in a chemical reaction. |
| Combination reaction | Two or more substances combine to form a single new substance. |
| Decomposition reaction | A single compound breaks into two or more simpler substances. |
| Displacement reaction | A more reactive element displaces a less reactive element from its compound. |
| Double displacement reaction | Two compounds exchange ions to form two new compounds. |
| Precipitation reaction | A reaction that forms an insoluble solid (precipitate) from two soluble solutions. |
| Oxidation | Gain of oxygen (or loss of hydrogen, or loss of electrons). |
| Reduction | Loss of oxygen (or gain of hydrogen, or gain of electrons). |
| Redox reaction | A reaction in which oxidation and reduction occur together. |
| Corrosion | The slow eating up of a metal by air, moisture or chemicals. |
| Rancidity | Change in taste and smell of oily/fatty food due to oxidation. |
15. Comparison Tables
15.1 Combination vs Decomposition
| Parameter | Combination | Decomposition |
|---|---|---|
| Reactants | Two or more | One compound |
| Products | One new substance | Two or more simpler substances |
| Energy involvement | Usually exothermic | Usually needs heat/electricity/light |
| General pattern | A + B → AB | AB → A + B |
| Example | CaO + H₂O → Ca(OH)₂ | 2H₂O → 2H₂ + O₂ |
15.2 Displacement vs Double Displacement
| Parameter | Displacement | Double displacement |
|---|---|---|
| What happens | One element replaces another | Two compounds exchange ions |
| General pattern | A + BC → AC + B | AB + CD → AD + CB |
| Reactants involved | One element + one compound | Two compounds |
| Driving force | Reactivity of the displacing element | Formation of an insoluble precipitate or fewer ions |
| Example | Fe + CuSO₁₀ → FeSO₁₀ + Cu | BaCl₂ + Na₂SO₁₀ → BaSO₁₀ + 2NaCl |
15.3 Oxidation vs Reduction
| Parameter | Oxidation | Reduction |
|---|---|---|
| Oxygen | Gain of oxygen | Loss of oxygen |
| Hydrogen | Loss of hydrogen | Gain of hydrogen |
| Electrons | Loss of electrons | Gain of electrons |
| In CuO + H₂ | H₂ → H₂O | CuO → Cu |
| Example in the equation | Hydrogen oxidised | Copper(II) oxide reduced |
15.4 Corrosion vs Rancidity
| Parameter | Corrosion | Rancidity |
|---|---|---|
| Affects | Metals | Oils and fats in food |
| Cause | Air, moisture and chemicals on the metal | Oxidation of oils/fats by atmospheric oxygen |
| Example | Rusting of iron, tarnishing of silver | Off smell/taste of fried snacks |
| Prevention | Painting, galvanising, alloying | Antioxidants, airtight packs, N₂ flushing |
15.5 Word equation vs Chemical equation
| Parameter | Word equation | Chemical equation |
|---|---|---|
| Form | Uses names of substances | Uses symbols and formulas |
| Example | Zinc + Sulphuric acid → Zinc sulphate + Hydrogen | Zn + 2HCl → ZnCl₂ + H₂ |
| Shows amount/state? | No | Yes (coefficients and state symbols) |
| Use in balancing | Cannot be balanced | Can be balanced |
16. Application and Competency Based Examples
16.1 Application based (applying the concepts)
Application Example 1
A shopkeeper stores his iron tools in a dry room or coats them with oil. Why?
Answer: Rusting of iron needs both oxygen and moisture. When tools are stored in a dry place or coated with oil, moisture and air cannot reach the iron surface, so rusting is prevented.
Application Example 2
Why are potato chip packets filled with nitrogen gas?
Answer: Nitrogen is an inert gas. It displaces the oxygen inside the packet, so the oils and fats in the chips do not get oxidised (no rancidity). This keeps the chips fresh and crisp.
Application Example 3
A few drops of sodium hydroxide are added to copper sulphate solution. A blue precipitate forms. Identify the type of reaction.
Answer: 2NaOH(aq) + CuSO₁₀(aq) → Cu(OH)₂(s) + Na₂SO₁₀(aq). This is a double displacement (precipitation) reaction; insoluble copper hydroxide precipitates.
Application Example 4
Leave a silver spoon in the open air for some time. A black coating appears. What is the coating and how does it form?
Answer: The black coating is silver sulphide (Ag₂S). Silver reacts with hydrogen sulphide gas present in the air (a form of corrosion called tarnishing): 4Ag + 2H₂S + O₂ → 2Ag₂S + 2H₂O.
Application Example 5
During electrolysis of water, hydrogen gas collects at the negative electrode and oxygen at the positive electrode. Which reaction is this?
Answer: This is an electrolytic decomposition reaction: 2H₂O(l) → 2H₂(g) + O₂(g), using electrical energy.
Application Example 6
An iron nail put in copper sulphate turns brownish. Give the balanced equation and the type of reaction.
Answer: Fe(s) + CuSO₁₀(aq) → FeSO₁₀(aq) + Cu(s). A displacement reaction because iron is more reactive than copper.
Application Example 7
Why does food cooked in oil stay longer when refrigerated?
Answer: Low temperature in the fridge slows the reaction between atmospheric oxygen and the oils/fats, reducing oxidation and delaying rancidity.
Application Example 8
Calcium carbonate on strong heating decomposes to give quicklime. Write the equation and state its type.
Answer: CaCO₃(s) → CaO(s) + CO₂(g). Thermal decomposition (endothermic) of a single compound into two products.
16.2 Reasoning and competency based (think and justify)
Reasoning Example 1
"A chemical equation must always be balanced." Justify.
Answer: The Law of Conservation of Mass states that in a chemical reaction the total mass of reactants equals the total mass of products; atoms are neither created nor destroyed. A balanced equation respects this law, and also tells us the exact amounts of substances that react, which is essential for quantitative calculations.
Reasoning Example 2
Copper cannot displace iron from iron sulphate solution. Why?
Answer: Displacement depends on the reactivity series. Iron is more reactive than copper; a more reactive metal alone can displace a less reactive metal. Copper, being less reactive, cannot displace iron.
Reasoning Example 3
Both CaO + H₂O and N₂ + O₂ → 2NO are combination reactions, yet one is exothermic and the other is not. Explain.
Answer: Combination reactions may release or absorb energy. Quicklime and water react releasing heat (exothermic). The formation of nitrogen monoxide needs continuous electrical discharge/supplied energy, so it is an endothermic combination reaction.
Reasoning Example 4
In CuO + H₂ → Cu + H₂O, name the oxidising and reducing agents.
Answer: CuO supplies oxygen (it is reduced), so it is the oxidising agent. H₂ removes oxygen (it is oxidised), so it is the reducing agent.
Reasoning Example 5
Why is silver chloride kept in dark-coloured bottles?
Answer: Silver chloride undergoes photochemical decomposition in sunlight: 2AgCl → 2Ag + Cl₂. Storing it in dark bottles keeps sunlight away, preventing this decomposition (as well as the darkening/fogging below).
Reasoning Example 6
A teacher heats green ferrous sulphate crystals and gets a smell like burning sulphur. Account for the colour change and the smell.
Answer: Thermal decomposition of ferrous sulphate takes place: 2FeSO₁₀ → Fe₂O₃ + SO₂ + SO₃. The reddish-brown ferric oxide gives the new colour, and the pungent burning-sulphur smell comes from sulphur dioxide gas.
17. Case Based Questions with Answers
Case Study 1 – Rusting of an iron gate
Ramesh observes that the iron gate of his house develops an orange-brown flaky deposit over time. His father tells him this happens because of rusting, and that the gate stays rust-free at his uncle's desert house.
Q1. What is the chemical name of the flaky deposit on the gate?
Answer: Hydrated iron(III) oxide, Fe₂O₃ · xH₂O, commonly called rust.
Q2. Which two conditions are necessary for rusting?
Answer: Oxygen (air) and moisture (water) must both be present. In the dry desert, the low moisture content slows rusting greatly.
Q3. Suggest two methods to stop the gate from rusting.
Answer: (i) Painting it regularly, and (ii) galvanising it (coating with a layer of zinc) or applying grease/oil. Alloys such as stainless steel also resist rusting.
Case Study 2 – Food preservation in a factory
A snack manufacturer fills bags of fried potato chips with an extra gas before sealing them, and writes "This bag is half-full only because of nitrogen flushing." A consumer thinks he is being cheated.
Q1. Which gas is flushed into the bags and why?
Answer: Nitrogen (N₂). It is inert and displaces oxygen, preventing the oxidation (rancidity) of oils in the chips.
Q2. What is rancidity? Name two other ways to prevent it.
Answer: Rancidity is the change in taste and smell of oily/fatty food due to oxidation. It can also be prevented by adding antioxidants (e.g. BHA/BHT), keeping food in airtight containers, and refrigerating it.
Q3. Is the consumer right to complain about the half-filled bag?
Answer: No. The extra space is nitrogen gas, deliberately filled to protect the chips from rancidity. The bag appears half-full by design, not due to cheating.
Case Study 3 – Metallic displacement lab
Four test tubes are set up: (i) iron nail in CuSO₁₀, (ii) copper wire in FeSO₁₀, (iii) zinc granule in CuSO₁₀, (iv) copper wire in ZnSO₁₀.
Q1. Which test tubes show a visible change? Give the equations.
Answer: (i) and (iii). Fe + CuSO₁₀ → FeSO₁₀ + Cu and Zn + CuSO₁₀ → ZnSO₁₀ + Cu. Tubes (ii) and (iv) show no reaction because copper is less reactive than iron and zinc.
Q2. Name the type of reaction and the rule that decides it.
Answer: Displacement reaction. A more reactive metal (Fe or Zn) displaces a less reactive metal (Cu) from its salt solution, following the reactivity series.
Q3. Describe the colour change in test tube (i).
Answer: The blue colour of copper sulphate fades and finally becomes pale green because iron sulphate (FeSO₁₀ solution) is pale green, and a brown layer of copper deposits on the nail.
Case Study 4 – Heating of green crystals
Green ferrous sulphate crystals are heated in a test tube. The crystals change colour, a reddish-brown solid remains, and a gas with a burning-sulphur smell is evolved. A glowing splinter introduced at the mouth of the tube gives a popping sound.
Q1. Write the balanced equation that takes place.
Answer: 2FeSO₁₀(s) → Fe₂O₃(s) + SO₂(g) + SO₃(g).
Q2. Why is this reaction called a decomposition reaction?
Answer: A single compound (ferrous sulphate) breaks into three simpler products on heating, so it is a thermal decomposition reaction.
Q3. A burning matchstick brought near the mouth of the tube is extinguished. Why?
Answer: The gases evolved are sulphur dioxide and sulphur trioxide, which do not support combustion and so extinguish the burning matchstick. This confirms that the reddish-brown residue (Fe₂O₃) and sulphur dioxide/sulphur trioxide gases are the decomposition products of ferrous sulphate.
Case Study 5 – White and golden yellow precipitates
A student pours barium chloride solution into sodium sulphate solution and observes a white solid. In another beaker, lead nitrate solution is mixed with potassium iodide solution and a bright yellow solid appears. The student notes that both solids do not dissolve on stirring.
Q1. Write both equations and name the precipitates.
Answer: BaCl₂ + Na₂SO₁₀ → BaSO₁₀(white) + 2NaCl and Pb(NO₃)₂ + 2KI → PbI₂(golden yellow) + 2KNO₃.
Q2. Why is the reaction called a precipitation reaction in each case?
Answer: An insoluble solid is formed from two aqueous solutions. Precipitation reactions are a type of double displacement reaction where the product precipitates out.
Q3. In the barium test, name the ions that combine to form the precipitate.
Answer: Barium ions (\(Ba^{2+}\)) and sulphate ions (\(SO_4^{2-}\)) combine to give the insoluble barium sulphate precipitate.
18. Solved Examples (with solutions)
18.1 Basic level (5 questions)
Solved Example 1
Write the word equation for the burning of magnesium and then the skeleton chemical equation.
Solution: Magnesium + Oxygen → Magnesium oxide. Skeleton chemical equation: Mg + O₂ → MgO.
Solved Example 2
In Zn + 2HCl → ZnCl₂ + H₂, name the reactants and products.
Solution: Reactants: zinc and hydrochloric acid. Products: zinc chloride and hydrogen gas.
Solved Example 3
Balance the equation: H₂ + O₂ → H₂O.
Solution: 2H₂ + O₂ → 2H₂O. (Oxygen balanced with 2H₂O, then hydrogen balanced with 2H₂.)
Solved Example 4
What is meant by the state symbol (aq)? Give one example.
Solution: (aq) means the substance is dissolved in water (aqueous). Example: CuSO₁₀(aq).
Solved Example 5
Give one difference between a physical change and a chemical change.
Solution: A physical change does not produce new substances (e.g. melting of ice), while a chemical change produces one or more new substances (e.g. burning of paper).
18.2 Moderate level (7 questions)
Solved Example 6
Balance: CH₁₀ + O₂ → CO₂ + H₂O.
Solution: CH₁₀ + 2O₂ → CO₂ + 2H₂O. (H made 4 with 2H₂O, then O balanced on left with 2O₂.)
Solved Example 7
CaCO₃ → CaO + CO₂. Identify the type of reaction.
Solution: Decomposition reaction, one compound breaking into two simpler substances (here thermal decomposition of limestone).
Solved Example 8
Balance: Al + O₂ → Al₂O₃.
Solution: 4Al + 3O₂ → 2Al₂O₃. (O balanced with 3O₂ and 2Al₂O₃, then Al balanced with 4Al.)
Solved Example 9
Classify: Fe + CuSO₁₀ → FeSO₁₀ + Cu.
Solution: Displacement reaction. Fe (more reactive) displaces Cu from copper sulphate.
Solved Example 10
Balance: AgNO₃ + CaCl₂ → AgCl + Ca(NO₃)₂.
Solution: 2AgNO₃ + CaCl₂ → 2AgCl + Ca(NO₃)₂. (NO₃ group balanced with 2AgNO₃, then Ag balanced with 2AgCl.)
Solved Example 11
In CuO + H₂ → Cu + H₂O, which substance is oxidised and which is reduced?
Solution: H₂ is oxidised (gains oxygen to form water); CuO is reduced (loses oxygen to form copper). The reaction is a redox reaction.
Solved Example 12
A solution of barium chloride and a solution of sodium sulphate are mixed. Write the equation and state the observation.
Solution: Na₂SO₁₀(aq) + BaCl₂(aq) → BaSO₁₀(s) + 2NaCl(aq). A white precipitate of barium sulphate forms (double displacement/precipitation reaction).
18.3 Application based (5 questions)
Solved Example 13
An iron nail is dipped into copper sulphate solution. What change happens after 20 minutes?
Solution: The blue solution turns pale green and a brown layer of copper is deposited on the nail: Fe + CuSO₁₀ → FeSO₁₀ + Cu.
Solved Example 14
How can a builder prevent the steel rods in a bridge column from rusting?
Solution: By painting the rods, galvanising them with zinc, or using stainless steel (an alloy). Keeping the surface dry also removes moisture required for rusting.
Solved Example 15
Why does lime water turn milky when carbon dioxide gas is passed through it, and which type of reaction is it?
Solution: CO₂ combines with calcium hydroxide (lime water) to form insoluble calcium carbonate: Ca(OH)₂ + CO₂ → CaCO₃ + H₂O. It is a combination reaction.
Solved Example 16
While electrolysing water, where do you collect hydrogen and oxygen? State the volume ratio.
Solution: Hydrogen collects at the negative electrode and oxygen at the positive electrode. Volume ratio of hydrogen to oxygen is 2 : 1, matching the reaction 2H₂O → 2H₂ + O₂.
Solved Example 17
Give one everyday method by which food items containing fats/oils can be kept fresh longer.
Solution: Keep them in airtight containers and refrigerate; manufacturers also add antioxidants or flush the packets with nitrogen gas to delay rancidity.
18.4 Competency / reasoning (4 questions)
Solved Example 18
Explain why magnesium burns brilliantly but copper oxide is reduced by hydrogen.
Solution: Magnesium has a strong affinity for oxygen, so it combines with oxygen (oxidation) rapidly giving brilliant light. In CuO + H₂, hydrogen has stronger affinity for oxygen than copper, so it takes away the oxygen (reduces CuO). In a reaction, oxidation and reduction always take place together.
Solved Example 19
Copper articles in the open air develop a green coating. What is the coating and how does it form?
Solution: The coating is basic copper carbonate, CuCO₃ · Cu(OH)₂. Moist carbon dioxide and oxygen in the air react with copper in a corrosion reaction.
Solved Example 20
Justify: "Corrosion is a redox reaction."
Solution: In rusting, iron is oxidised to iron(III) (loses electrons, gains oxygen) while oxygen is simultaneously reduced. Since oxidation and reduction occur together, corrosion is a redox reaction.
Solved Example 21
Why does the colour of iron sulphate change and a bad smell appear on heating, but no colour or bad smell appear when salt is heated?
Solution: Ferrous sulphate decomposes chemically on heating (2FeSO₁₀ → Fe₂O₃ + SO₂ + SO₃), producing a colour change and a smell of sulphur dioxide. Salt only undergoes a physical change (no new substance), so no such observation is seen.
18.5 Case based solved examples (4 questions)
Solved Example 22
Scenario: A coiled magnesium wire is ignited in a dark room. A flash of blinding light and a white powdery residue are noticed. Write the equation and identify the type.
Solution: 2Mg + O₂ → 2MgO. Combination reaction, exothermic; the heat and light released show energy is given out.
Solved Example 23
Scenario: Old silver jewelry becomes black after a few years. Identify the compound and suggest a way to restore the shine that removes only the tarnish layer.
Solution: The black layer is silver sulphide (Ag₂S) formed by reaction of silver with H₂S in air. The jewellery can be gently cleaned/polished or treated with a mild silver-cleaning agent that removes the tarnish layer without damaging silver underneath.
Solved Example 24
Scenario: Water dripping on quicklime starts boiling. Name the reaction, write its equation, and state whether it is endothermic or exothermic.
Solution: CaO + H₂O → Ca(OH)₂. Combination reaction, exothermic (so much heat is given out that water boils).
Solved Example 25
Scenario: A sticker on a milk packet says "use before 3 days" but the same milk kept in the fridge lasts 7 days. Explain using the concept of rancidity/oxidation.
Solution: Spoilage of dairy items is sped up by microbes and by oxidation reactions. Lower temperature in the fridge slows these reactions, so the food stays fresh longer. Rancidity (oxidation of fats) in the milk is also delayed at low temperature.
19. Common Exam Mistakes to Avoid
- Writing formulas like CaCl2 instead of CaCl₂ (subscripts must be small numbers), or writing H2O2 instead of H₂O.
- Changing the formula of a compound to balance an equation. Only coefficients may change.
- Omitting state symbols, especially (aq), (s), (g) when the question asks for a complete balanced equation.
- Forgetting the 2 in 2H₂, 2O₂, etc., and leaving a skeleton equation unbalanced.
- Calling NH₃ or CO₂ "nitrogen" or "carbon" while writing names - formulas must match the correct compound.
- Marking FeSO₁₀ decomposition product as FeO instead of Fe₂O₃.
- Calling rust Fe₂O₃ instead of hydrated iron(III) oxide Fe₂O₃ · xH₂O.
- Confusing displacement with double displacement. A displacement reaction has a free element on one side.
- Saying "oxidation is loss of oxygen" for every case - use OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain.
- Saying copper displaces iron from iron sulphate - it cannot, because copper is less reactive.
- Omitting the type of reaction in a classification question, or writing "chemical reaction" when asked the specific type (combination/decomposition etc.).
- Writing the products of electrolysis of water wrongly (H₂ + O₂ in correct 2:1 volume ratio but forgetting the coefficient 2 in 2H₂).
20. Exam Tips for This Chapter
- Always write the arrow → between reactants and products, and put conditions (Δ, catalyst, sunlight) above the arrow.
- In a perfect answers, write state symbols and mark gas/precipitate appropriately.
- Learn the reactivity series order (K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag) to answer displacement questions.
- For redox questions, first identify which substance gains oxygen and which loses oxygen. Name both oxidation and reduction, then say it is a redox reaction.
- For 1-mark definition questions, quote the exact definition (combination, decomposition, corrosion, rancidity) from the glossary above.
- For balancing questions, show the steps and end with "Balanced" after checking both sides. Never modify a formula.
- Know two named examples for each reaction type: e.g. thermal decomposition (Pb(NO₃)₂, FeSO₁₀), photochemical (AgCl/AgBr), electrolytic (H₂O).
- Understand corrosion from both laboratory and everyday examples: rusting of iron, silver tarnishing, copper green coating.
21. Quick Revision Points
- A chemical reaction produces new substances. Indicators: state, colour, gas, temperature, precipitate changes.
- Word equation → skeleton equation → balanced equation (Law of Conservation of Mass).
- State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous.
- Balancing only uses coefficients, never changes formulas.
- Combination: A + B → AB. Decomposition: AB → A + B (thermal, electrolytic, photochemical).
- Displacement: A + BC → AC + B (reactivity series). Double displacement: AB + CD → AD + CB (often precipitation).
- Redox: oxidation (gain O/loss H/loss e⁻) + reduction (loss O/gain H/gain e⁻) together. OIL RIG.
- Corrosion (metals) and rancidity (oils/fats) are everyday oxidation effects; both preventable.
- Rancidity prevention: antioxidants, airtight containers, refrigeration, nitrogen flushing.
- Rusting needs both oxygen and moisture.
22. Self Practice Questions (30 questions with answers)
22.1 Conceptual questions (Q1–Q6)
Q1
Define a chemical reaction. Give one everyday example.
Q2
What is a balanced chemical equation? State the law on which balancing is based.
Q3
What do the symbols (s), (l), (g) and (aq) represent?
Q4
Distinguish between an exothermic and an endothermic reaction with one example each.
Q5
"A combination reaction and a decomposition reaction are opposites." Justify.
Q6
What is rancidity? Name two ways to prevent it.
22.2 Balancing questions (Q7–Q13)
Q7
Balance: H₂ + O₂ → H₂O.
Q8
Balance: N₂ + H₂ → NH₃.
Q9
Balance: CH₁₀ + O₂ → CO₂ + H₂O.
Q10
Balance: Al + O₂ → Al₂O₃.
Q11
Balance: CaCO₃ + HCl → CaCl₂ + CO₂ + H₂O.
Q12
Balance: Pb(NO₃)₂ → PbO + NO₂ + O₂.
Q13
Balance: Fe + H₂O → Fe₃O₁₀ + H₂.
22.3 Identification questions (Q14–Q18)
Q14
What type is 2H₂O → 2H₂ + O₂?
Q15
What type is Fe + CuSO₁₀ → FeSO₁₀ + Cu?
Q16
What type is Na₂SO₁₀ + BaCl₂ → BaSO₁₀ + 2NaCl?
Q17
What type is CaO + H₂O → Ca(OH)₂?
Q18
What type is 2AgCl → 2Ag + Cl₂ when light falls on it?
22.4 Application questions (Q19–Q23)
Q19
Why is an iron ring on a ship coated with paint and connected to protective metal?
Q20
A gas is bubbled through lime water and the lime water turns milky. Suggest the gas.
Q21
While heating green ferrous sulphate crystals, what is the smell that emerges?
Q22
Ghee kept in a fridge stays fresh longer. Give one reason.
Q23
Write a balanced equation for the reaction of zinc with dilute sulphuric acid.
22.5 Competency questions (Q24–Q27)
Q24
Raheem keeps an iron nail in tap water and another in boiled (oxygen-free) water. What difference will he observe and why?
Q25
Give reason: the burning of magnesium is a combination reaction, but the heating of calcium carbonate is not.
Q26
Explain using the redox concept why zinc cannot reduce copper(II) oxide to copper at room temperature by being physically mixed, but does displace copper from CuSO₁₀ solution.
Q27
A teacher says "rusting is a slow oxidation reaction." Support this statement with the products formed.
22.6 Case based questions (Q28–Q30)
Q28
Case: A factory fills chips packets with nitrogen gas. (i) What is the purpose? (ii) Name the problem this avoids.
Q29
Case: Silver jewelry turns black in a polluted city. (i) Name the compound (ii) Write its formula.
Q30
Case: On heating lead nitrate, a brown gas is evolved. (i) Write the equation (ii) Name the gas.
23. Answers to Self Practice Questions
23.1 Conceptual: Q1–Q6
Q1. A process in which one or more substances change into new substances is a chemical reaction. Example: burning of magnesium in air, combustion of LPG, rusting of iron.
Q2. A balanced chemical equation has equal numbers of atoms of each element on both sides. It follows the Law of Conservation of Mass.
Q3. (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).
Q4. Exothermic releases heat (e.g. CaO + H₂O → Ca(OH)₂, burning of coal). Endothermic absorbs heat (e.g. thermal decomposition of CaCO₃).
Q5. A combination reaction joins simpler substances into one product (A + B → AB), while a decomposition reaction splits one compound into simpler products (AB → A + B). They are reverse of each other.
Q6. Rancidity is the change in taste/smell of oily food due to oxidation of fats. Prevention: antioxidants, airtight containers, refrigeration, nitrogen flushing.
23.2 Balancing: Q7–Q13
Q7. 2H₂ + O₂ → 2H₂O.
Q8. N₂ + 3H₂ → 2NH₃.
Q9. CH₁₀ + 2O₂ → CO₂ + 2H₂O.
Q10. 4Al + 3O₂ → 2Al₂O₃.
Q11. CaCO₃ + 2HCl → CaCl₂ + CO₂ + H₂O.
Q12. 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂.
Q13. 3Fe + 4H₂O → Fe₃O₁₀ + 4H₂.
23.3 Identification: Q14–Q18
Q14. Decomposition (electrolytic decomposition of water).
Q15. Displacement reaction.
Q16. Double displacement (precipitation) reaction.
Q17. Combination (exothermic) reaction.
Q18. Photochemical decomposition.
23.4 Application: Q19–Q23
Q19. Paint stops air/moisture reaching iron (prevents rusting). Connecting protective metal (sacrificial protection) keeps the iron safe.
Q20. Carbon dioxide (CO₂), which gives milky precipitate CaCO₃ with lime water.
Q21. Smell of burning sulphur comes from sulphur dioxide (SO₂) gas.
Q22. Low temperature slows oxidation of ghee, delaying rancidity.
Q23. Zn + H₂SO₁₀ → ZnSO₁₀ + H₂.
23.5 Competency: Q24–Q27
Q24. The nail in tap water rusts; the nail in boiled water does not. Rusting needs both oxygen and moisture; boiled water has lost dissolved oxygen, so no rusting.
Q25. Burning of magnesium combines Mg and O₂ into a single product (combination). Heating CaCO₃ splits one compound into CaO and CO₂ (decomposition), not a combination.
Q26. In solution, the ions exist freely, so zinc (more reactive) displaces copper from CuSO₁₀ solution: Zn + CuSO₁₀ → ZnSO₁₀ + Cu. Simply mixing zinc and copper(II) oxide powders in the solid state gives no such ion exchange, so no displacement happens at room temperature.
Q27. Rusting is the slow oxidation of iron to hydrated iron(III) oxide Fe₂O₃ · xH₂O, using the oxygen of air. Being slow, it is called slow oxidation.
23.6 Case based: Q28–Q30
Q28. (i) Nitrogen is flushed in to avoid oxidation of fats/oils and to keep chips fresh and crisp. (ii) The problem avoided is rancidity.
Q29. (i) Silver sulphide (Ag₂S), formed by reaction with hydrogen sulphide gas in the air. (ii) Ag₂S.
Q30. (i) 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂. (ii) The brown gas is nitrogen dioxide (NO₂).