1. What is a Chemical Reaction?
A chemical reaction is a process in which one or more substances undergo chemical change to form new substances with different properties.
Examples from daily life:
- Milk turning sour
- Iron rusting
- Food being cooked
- Digestion of food
- Respiration
- Fermentation of grapes
In a chemical reaction, the identity and nature of the original substances change.
How can we identify a chemical reaction?
A chemical reaction may be indicated by:
- Change in state
- Change in colour
- Evolution of a gas
- Change in temperature
These observations are useful signs, but the key idea is the formation of new substances.
2. Chemical Equations
A chemical reaction can first be represented using words.
Example: Burning of magnesium
Magnesium + Oxygen → Magnesium oxide
- Substances on the left side are called reactants.
- Substances on the right side are called products.
- The arrow → indicates the direction of the reaction.
Magnesium burns with a dazzling white flame and produces a white powder of magnesium oxide.
Chemical equation
Instead of writing names, we can use chemical formulae:
Mg + O₂ → MgO
This is called a skeletal chemical equation because it is not balanced.
3. Balanced Chemical Equation
According to the law of conservation of mass, matter cannot be created or destroyed during a chemical reaction.
Therefore:
The number of atoms of every element must be equal on both sides of a chemical equation.
For example:
Mg + O₂ → MgO ❌
Oxygen atoms are unequal.
The balanced equation is:
2Mg + O₂ → 2MgO ✅
Important rule
While balancing an equation:
Never change the chemical formula of a substance.
Change only the coefficients placed before formulae.
For example, to balance H₂O, use 2H₂O, not H₂O₂ or (H₂O)₂.
4. Hit-and-Trial Method of Balancing
The chapter demonstrates balancing using:
Fe + H₂O → Fe₃O₄ + H₂
A systematic approach is:
- Write the skeletal equation.
- Count atoms of each element on both sides.
- Begin with an appropriate compound/element, often the one with the larger number of atoms.
- Put suitable whole-number coefficients before formulae.
- Balance the remaining elements.
- Check the final atom count on both sides.
- Use the smallest whole-number coefficients possible.
Final equation:
3Fe + 4H₂O → Fe₃O₄ + 4H₂
This method is called the hit-and-trial method.
Remember
Subscripts cannot be changed.
Coefficients can be changed.
This is one of the most important rules for balancing equations.
5. Physical States in Chemical Equations
A chemical equation can be made more informative by showing the physical state of each substance.
| Symbol | Meaning |
|---|---|
| (s) | Solid |
| (l) | Liquid |
| (g) | Gas |
| (aq) | Aqueous solution |
Example:
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)
Here, H₂O(g) means water is present as steam.
Reaction conditions such as heat, sunlight, pressure or catalyst may also be written near the arrow.
6. Types of Chemical Reactions
The major reaction types in this chapter are:
- Combination reaction
- Decomposition reaction
- Displacement reaction
- Double displacement reaction
- Oxidation and reduction (redox reactions)
7. Combination Reaction
Definition
A reaction in which two or more reactants combine to form a single product is called a combination reaction.
General pattern
A + B → AB
Example 1
CaO + H₂O → Ca(OH)₂ + Heat
Calcium oxide and water combine to form calcium hydroxide.
Example 2
C + O₂ → CO₂
Carbon combines with oxygen to form carbon dioxide.
Example 3
2H₂ + O₂ → 2H₂O
Two substances combine to form one product.
Quick identification
Many reactants → ONE product = Combination
8. Exothermic Reaction
A reaction in which heat is released along with the products is called an exothermic reaction.
Examples
CaO + H₂O → Ca(OH)₂ + Heat
CH₄ + 2O₂ → CO₂ + 2H₂O + Energy
Respiration is exothermic
During respiration, glucose reacts with oxygen and releases energy:
C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy
Thus, respiration is an exothermic reaction because energy is released.
9. Decomposition Reaction
Definition
A reaction in which one reactant breaks down into two or more simpler substances is called a decomposition reaction.
General pattern
AB → A + B
It is essentially the opposite of a combination reaction.
Types of decomposition reactions
A. Thermal decomposition
Decomposition caused by heat.
Example:
CaCO₃(s) —Heat→ CaO(s) + CO₂(g)
Calcium carbonate decomposes into calcium oxide and carbon dioxide.
This reaction is important industrially because calcium oxide (quick lime) has several uses, including cement manufacture.
B. Decomposition by electricity
Water can be decomposed using electricity.
2H₂O → 2H₂ + O₂
The volume of hydrogen produced is twice that of oxygen because of the ratio of atoms in water.
C. Photochemical decomposition
Decomposition caused by light.
2AgCl —Sunlight→ 2Ag + Cl₂
Silver chloride changes from white to grey because silver is formed.
Similar behaviour occurs with silver bromide:
2AgBr —Sunlight→ 2Ag + Br₂
These reactions are associated with black-and-white photography.
10. Endothermic Reaction
A reaction in which energy is absorbed is called an endothermic reaction.
Decomposition reactions commonly require energy in the form of:
- Heat
- Light
- Electricity
Therefore, such decomposition processes can be endothermic.
Easy memory trick
Exothermic → energy exits/released
Endothermic → energy enters/absorbed
11. Displacement Reaction
Definition
A reaction in which a more reactive element displaces a less reactive element from its compound is called a displacement reaction.
Example
Fe + CuSO₄ → FeSO₄ + Cu
Iron displaces copper from copper sulphate.
What do we observe?
- Copper sulphate solution is initially blue.
- Iron is more reactive than copper.
- Iron replaces copper.
- Copper is deposited.
- The solution changes because iron sulphate is formed.
- The iron nail develops a brownish coating of copper.
Other examples:
Zn + CuSO₄ → ZnSO₄ + Cu
Pb + CuCl₂ → PbCl₂ + Cu
Identification
Element + compound → new compound + displaced element
12. Double Displacement Reaction
Definition
A reaction in which the ions of two compounds exchange partners to form two new compounds is called a double displacement reaction.
Example
Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl
Here, ions exchange partners:
- Na⁺ combines with Cl⁻
- Ba²⁺ combines with SO₄²⁻
A white insoluble substance, BaSO₄, is formed.
General pattern
AB + CD → AD + CB
13. Precipitation Reaction
A reaction in which an insoluble solid substance is formed from solutions is called a precipitation reaction.
The insoluble solid is called a precipitate.
Example:
Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)
BaSO₄ is the white precipitate.
Key connection
Precipitation reaction → insoluble precipitate formed
14. Oxidation
In terms used in this chapter:
A substance is oxidised when it:
- Gains oxygen, or
- Loses hydrogen
Example
2Cu + O₂ → 2CuO
Copper gains oxygen, so copper is oxidised.
15. Reduction
A substance is reduced when it:
- Loses oxygen, or
- Gains hydrogen
Example
CuO + H₂ → Cu + H₂O
CuO loses oxygen → CuO is reduced.
H₂ gains oxygen → H₂ is oxidised.
16. Redox Reaction
When oxidation and reduction occur together in the same reaction, the reaction is called an oxidation-reduction or redox reaction.
Example:
CuO + H₂ → Cu + H₂O
| Substance | Change | Process |
|---|---|---|
| CuO | Loses oxygen | Reduction |
| H₂ | Gains oxygen | Oxidation |
Memory rule
OIL RIG
- Oxidation Is Loss
- Reduction Is Gain
For this chapter, remember the loss/gain specifically in terms of oxygen/hydrogen.
17. Corrosion
Definition
When metals are attacked by substances in their surroundings, such as moisture or acids, the process is called corrosion.
Examples
- Iron develops a reddish-brown coating → rusting
- Silver develops a black coating
- Copper develops a green coating
Corrosion damages:
- Iron railings
- Bridges
- Ships
- Car bodies
- Other metal structures
Why is corrosion harmful?
It gradually damages metal objects and reduces their useful life.
18. Rancidity
Definition
When fats and oils undergo oxidation, they become rancid, causing undesirable changes in their smell and taste.
Prevention of rancidity
Oxidation can be slowed by:
- Adding antioxidants
- Storing food in airtight containers
- Flushing packets of oily food with nitrogen
Nitrogen reduces the contact of food with oxygen and therefore slows oxidation.
⭐ Most Important Reactions to Learn
| Reaction | Equation | Type |
|---|---|---|
| Magnesium burning | 2Mg + O₂ → 2MgO | Combination / oxidation |
| Quick lime + water | CaO + H₂O → Ca(OH)₂ + Heat | Combination / exothermic |
| Carbon burning | C + O₂ → CO₂ | Combination |
| Hydrogen + oxygen | 2H₂ + O₂ → 2H₂O | Combination |
| Calcium carbonate heating | CaCO₃ → CaO + CO₂ | Decomposition |
| Lead nitrate heating | 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂ | Decomposition |
| Silver chloride in sunlight | 2AgCl → 2Ag + Cl₂ | Photochemical decomposition |
| Iron + copper sulphate | Fe + CuSO₄ → FeSO₄ + Cu | Displacement |
| Sodium sulphate + barium chloride | Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl | Double displacement / precipitation |
| Copper oxidation | 2Cu + O₂ → 2CuO | Oxidation |
| Copper oxide + hydrogen | CuO + H₂ → Cu + H₂O | Redox |
Conditions such as heat/sunlight should be written above the arrow where required.
🔥 High-Value Differences
Combination vs Decomposition
| Combination | Decomposition |
|---|---|
| Two or more reactants combine | One reactant breaks down |
| Usually gives one product | Gives two or more products |
| A + B → AB | AB → A + B |
| Example: 2H₂ + O₂ → 2H₂O | Example: CaCO₃ → CaO + CO₂ |
Displacement vs Double Displacement
| Displacement | Double displacement |
|---|---|
| One element replaces another | Ions/groups exchange partners |
| Usually element + compound | Usually two compounds react |
| A + BC → AC + B | AB + CD → AD + CB |
| Fe + CuSO₄ → FeSO₄ + Cu | Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl |
Exothermic vs Endothermic
| Exothermic | Endothermic |
|---|---|
| Energy/heat is released | Energy is absorbed |
| Surroundings may become warmer | Reaction may require energy input |
| Respiration is an example | Decomposition requiring energy is an example |
🧠 Exam-Focused Concepts
Why should magnesium ribbon be cleaned before burning?
Magnesium ribbon may have a surface coating that interferes with its reaction with oxygen. Cleaning exposes the magnesium surface so it can burn properly.
Why does the colour of CuSO₄ solution change when iron is dipped into it?
Iron is more reactive than copper and displaces copper from copper sulphate:
Fe + CuSO₄ → FeSO₄ + Cu
Thus, a new iron salt is formed and copper is deposited.
Why are chemical equations balanced?
Because of the law of conservation of mass. The number of atoms of each element must remain the same before and after a reaction.
Why is respiration exothermic?
Glucose reacts with oxygen during respiration and energy is released, which the body uses.
Why are decomposition reactions opposite to combination reactions?
A combination reaction forms one product from two or more substances, whereas a decomposition reaction breaks one substance into two or more products.
Why are chips packets flushed with nitrogen?
To reduce contact with oxygen and thereby slow oxidation of fats/oils, preventing or delaying rancidity.
Why is paint applied to iron articles?
Painting provides a protective layer that helps prevent iron from coming into contact with substances such as moisture and oxygen, thereby reducing corrosion.
⚡ One-Page Revision Map
Chemical Reaction
→ New substances formed
→ Signs: state / colour / gas / temperature change
Chemical Equation
→ Reactants → Products
→ Must be balanced
→ Physical states: (s), (l), (g), (aq)
Types
Combination
→ Many → One
Decomposition
→ One → Many
→ Heat / Light / Electricity
Displacement
→ More reactive element replaces less reactive element
Double Displacement
→ Exchange of ions
Precipitation
→ Insoluble solid formed
Oxidation
→ Gain O / Loss H
Reduction
→ Loss O / Gain H
Redox
→ Oxidation + Reduction together
Everyday effects
→ Corrosion
→ Rancidity
🎯 What to prepare especially well for exams
- Definition of chemical reaction
- Four signs of a chemical reaction
- Reactants and products
- Balanced vs skeletal equations
- Rules for balancing equations
- Physical-state symbols
- Combination and decomposition reactions
- Exothermic and endothermic reactions
- Displacement and double displacement
- Precipitation reaction
- Oxidation, reduction and redox
- Corrosion and rancidity
- Important equations and their reaction types
- Reason-based questions, especially on magnesium, iron/CuSO₄, respiration, corrosion and rancidity.
Introduction
- Chemical reaction: Process where reactants transform into products with new chemical properties.
- Examples:
- Combustion: C + O₂ → CO₂
- Rusting: Fe + O₂ + H₂O → Fe₂O₃·xH₂O
Types of Chemical Reactions
- Combination Reaction (Synthesis)
- Two or more substances combine to form a single product
- Example: 2H₂ + O₂ → 2H₂O
- Decomposition Reaction
- A single compound breaks into two or more simpler substances
- Example: 2HgO → 2Hg + O₂
- Displacement Reaction
- More reactive element displaces a less reactive element from its compound
- Example: Zn + CuSO₄ → ZnSO₄ + Cu
- Double Displacement Reaction
- Exchange of ions between two compounds
- Example: AgNO₃ + NaCl → AgCl + NaNO₃
- Redox Reaction (Oxidation–Reduction)
- Oxidation: Loss of electrons
- Reduction: Gain of electrons
- Example: Zn + Cu²⁺ → Zn²⁺ + Cu
- Neutralization Reaction
- Acid reacts with base to form salt and water
- Example: HCl + NaOH → NaCl + H₂O
- Exothermic and Endothermic Reactions
- Exothermic: Release energy (e.g., combustion)
- Endothermic: Absorb energy (e.g., photosynthesis)
Chemical Equations
- Chemical equation: Representation of a chemical reaction using symbols and formulas.
- Balanced chemical equation: Number of atoms of each element is same on both sides.
Steps to Balance Chemical Equations
- Write reactants and products.
- Count atoms of each element.
- Use coefficients to balance atoms.
- Check that all atoms are balanced.
Example:
Unbalanced: H₂ + O₂ → H₂O
Balanced: 2H₂ + O₂ → 2H₂O
Important Concepts
- Reactants and Products:
- Reactants → Substances initially present
- Products → New substances formed
- Indicators of Chemical Reactions:
- Gas evolution, color change, temperature change, precipitate formation
- Conservation of Mass:
- Mass of reactants = Mass of products