Chemical Reactions and Equations
Class 10 Science — Chapter 1 Complete Chapter Notes
In our daily lives, physical changes involve alterations in appearance or state without creating new substances, like melting ice. On the other hand, chemical changes result in the formation of entirely new substances, such as cooking food or rusting iron. Recognizing these changes helps us understand when a chemical reaction occurs.
1. What is a Chemical Reaction?
A chemical reaction is a process in which one or more substances, the Reactants, are converted to one or more different substances, the Products.
2. Chemical Equations
A chemical equation is a symbolic representation of a chemical reaction using the symbols and formulas of the substances involved.
The substances that undergo chemical change in the reaction, Magnesium and Oxygen, are the Reactants. When magnesium burns in air, it combines with oxygen to form magnesium oxide.
Chemical equations can be made more concise and useful if we use chemical formulae instead of words. Example: The chemical formula for water is H₂O, which indicates that it is made up of two hydrogen atoms and one oxygen atom.
1. Writing a Chemical Equation
Representation of a chemical reaction in terms of symbols and chemical formulae of the reactants and products is known as a chemical equation.
Physical state notation:
For solids, the symbol is "(s)".
For liquids, it is "(l)".
For gases, it is "(g)".
For aqueous solutions, it is "(aq)".
For gas produced in the reaction, it is represented by "(↑)".
For precipitate formed in the reaction, it is represented by "(↓)".
The reactants are on the left (LHS) of the arrow, while the products are on the right (RHS). A plus sign (+) links the different reactants and products together.
2. Balanced Chemical Equations
A balanced account of a chemical reaction is a complete chemical equation, which symbolically depicts the reactants, products, and their physical states.
The Law of Conservation of Mass states that in a chemical reaction, atoms can't be created or destroyed. This means that the total number of atoms for each element in the starting materials (reactants) must be the same as in the end products, keeping the overall mass the same.
Balanced chemical equation definition: A chemical equation is considered balanced when the number of atoms for each element on the reactant side is identical to the product side.
Steps for Balancing Chemical Equations
• To balance chemical equations, coefficients (numeric values preceding chemical symbols or formulas) are utilized. These coefficients represent the number of atoms or molecules involved. Adjust the coefficients as needed to make sure that the number of each type of atom is the same on both sides of the equation.
• For instance, in the equation Zn + HCl → ZnCl₂ + H₂, balancing involves adding coefficients to achieve equilibrium: Zn + 2HCl → ZnCl₂ + H₂. This balancing process often involves trial and error adjustments to ensure the equality of atoms on both sides.
Detailed Step-by-Step Example: Fe + H₂O → Fe₃O₄ + H₂
Informative Notation and Reaction Conditions:
• Including physical states makes chemical equations more informative: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution.
• Reaction conditions (temperature, pressure, catalyst, sunlight, chlorophyll) are indicated above and/or below the arrow in the chemical equation.
• Example (Photosynthesis):
6CO₂(g) + 6H₂O(l) —[Sunlight / Chlorophyll]→ C₆H₁₂O₆(aq) + 6O₂(g)
3. Types of Chemical Reactions
Chemical reactions are classified into several categories based on various factors:
1. Combination Reaction
A reaction in which a single product is formed from two or more reactants is known as a combination reaction.
Examples:
(i) Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide), releasing a large amount of heat:
CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat
(ii) White-washing of walls: A solution of slaked lime, Ca(OH)₂,reacts slowly with carbon dioxide in air to form a thin layer of calcium carbonate (CaCO₃) on walls, giving a shiny finish after 2–3 days. (Note: Marble also has the chemical formula CaCO₃):
Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
Exothermic Chemical Reactions: Reactions that release heat along with the formation of products are called exothermic chemical reactions.
Other Examples of Exothermic Reactions:
• Burning of Coal: C(s) + O₂(g) → CO₂(g)
• Formation of Water: 2H₂(g) + O₂(g) → 2H₂O(l)
• Burning of Natural Gas: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)
• Respiration: Digestion breaks carbohydrates (rice, potatoes, bread) into glucose. Glucose combines with oxygen in cells to yield energy:
C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(g) + 6H₂O(l) + Energy
2. Decomposition Reaction
A reaction in which a single reactant breaks down when exposed to heat, light, or electricity to give two or more simpler products is called a decomposition reaction.
There are three types of decomposition reactions:
(a) Thermal Decomposition
A decomposition reaction carried out by heating is called thermal decomposition.
Example 1: Heating of Ferrous Sulphate
Green ferrous sulphate crystals (FeSO₄·7H₂O) lose water when heated and turn white. On further heating, it decomposes into ferric oxide (Fe₂O₃ - solid), sulphur dioxide (SO₂ - gas), and sulphur trioxide (SO₃ - gas). The characteristic smell of burning sulphur is observed.
2FeSO₄(s) —[Heat]→ Fe₂O₃(s) + SO₂(g) + SO₃(g)
Example 2: Decomposition of Calcium Carbonate
Heating calcium carbonate breaks it down into calcium oxide (quick lime) and carbon dioxide gas. Calcium oxide is widely used in cement manufacturing.
CaCO₃(s) —[Heat]→ CaO(s) + CO₂(g)
Example 3: Heating of Lead Nitrate Powder
Heating lead nitrate powder produces brown fumes of nitrogen dioxide (NO₂) along with lead oxide and oxygen gas.
2Pb(NO₃)₂(s) —[Heat]→ 2PbO(s) + 4NO₂(g) + O₂(g)
(b) Photolytic Decomposition
A decomposition reaction carried out in the presence of light energy.
Example 1: Decomposition of Hydrogen Peroxide
Hydrogen peroxide (H₂O₂) breaks down into water and oxygen gas when exposed to sunlight.
Example 2: Effect of Sunlight on Silver Chloride and Silver Bromide
White silver chloride turns grey in sunlight due to decomposition into silver metal and chlorine gas:
2AgCl(s) —[Sunlight]→ 2Ag(s) + Cl₂(g)
Similarly, silver bromide decomposes:
2AgBr(s) —[Sunlight]→ 2Ag(s) + Br₂(g)
Note: These reactions of silver halides are used in black and white photography.
(c) Electrolytic Decomposition
A decomposition reaction carried out by passing an electric current through an aqueous solution or molten state of a compound.
Example: Electrolysis of Water — Decomposition of water into hydrogen gas and oxygen gas by passing electric current through acidulated water:
2H₂O(l) —[Electric Current]→ 2H₂(g) + O₂(g)
Endothermic Reactions: Decomposition reactions require energy in the form of heat, light, or electricity to break bonds. Reactions in which energy is absorbed are called endothermic reactions.
3. Displacement Reaction
A displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its salt solution. Both metals and non-metals take part in displacement reactions.
Important Reaction: Iron Nails in Copper Sulphate Solution
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
Observation: The iron nail becomes brownish due to copper deposition, and the blue colour of the copper sulphate solution fades to pale green (FeSO₄).
Other Examples:
• Zinc displaces copper: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
• Lead displaces copper: Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)
Zinc and lead are more reactive elements than copper.
4. Double Displacement Reaction
A double displacement reaction is a chemical reaction where two compounds react by exchanging their ions (cations and anions) to form two new compounds.
Example: Reaction of Sodium Sulphate and Barium Chloride
Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq)
Observation: A white insoluble precipitate of barium sulphate (BaSO₄) is formed.
Precipitation&Neutralisation Reactions:
• Any reaction that produces an insoluble precipitate is called a precipitation reaction.
• Neutralisation reactions (Acid + Base → Salt + Water) are also examples of double displacement reactions.
5. Endothermic and Exothermic Reactions Summary
(i) Exothermic Reaction: Energy is released into the surroundings, usually as heat. Example: C(s) + O₂(g) → CO₂(g) + Heat
(ii) Endothermic Reaction: Energy is absorbed from the surroundings in the form of heat, light, or electricity. Example (Photosynthesis): 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
6. Oxidation and Reduction (Redox Reactions)
Oxidation and Reduction occur simultaneously in chemical reactions, known as redox reactions.
Definitions:
• Oxidation: Gain of oxygen OR loss of hydrogen / loss of electrons (increase in oxidation state).
• Reduction: Loss of oxygen OR gain of hydrogen / gain of electrons (decrease in oxidation state).
Oxidation of Copper to Copper Oxide:
When copper powder is heated in air, it turns black due to the formation of copper(II) oxide:
2Cu + O₂ —[Heat]→ 2CuO
If hydrogen gas is passed over heated CuO, the black coating turns brown as the reverse reaction occurs:
CuO + H₂ —[Heat]→ Cu + H₂O
Here, CuO loses oxygen (reduced) and H₂ gains oxygen (oxidized).
Other Examples of Redox Reactions:
• ZnO + C → Zn + CO (ZnO is reduced to Zn; C is oxidized to CO)
• MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂ (MnO₂ is reduced; HCl is oxidized to Cl₂)
Oxidizing Agent and Reducing Agent:
• Reducing Agent: The substance that gets oxidized (causes reduction of the other substance by donating electrons or adding hydrogen/removing oxygen).
• Oxidizing Agent: The substance that gets reduced (causes oxidation of the other substance by accepting electrons or supplying oxygen/removing hydrogen).
• Example: In Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g), Zinc is oxidized (reducing agent) and HCl is reduced (oxidizing agent).
4. Effects of Oxidation Reaction in Everyday Life
1. Corrosion
When a metal is attacked by substances around it such as moisture, air, oxygen, and acids, it is said to corrode, and this process is called corrosion. Corrosion causes flaking, structural weakening, and damage to bridges, iron railings, ships, and automobile bodies.
Chemical Equations of Corrosion:
• Rusting of Iron: 4Fe(s) + 3O₂(from air) + xH₂O(moisture) → 2Fe₂O₃·xH₂O (reddish-brown rust)
• Corrosion of Copper: Cu(s) + H₂O(moisture) + CO₂(air) → CuCO₃·Cu(OH)₂ (green coating)
• Corrosion of Silver: 2Ag(s) + H₂S(air) → Ag₂S(black coating) + H₂(g)
2. Rancidity
When fats and oils present in food materials are left exposed to air for a long time, they undergo oxidation, resulting in a change in taste and odour. This phenomenon is called rancidity.
Methods to Prevent Rancidity:
Seal food materials in air-tight containers.
Flush food packaging with inert gases like nitrogen (e.g., chips bags).
Keep food in refrigerators to slow down oxidation.
Add antioxidants or food preservatives.
Chapter Revision Questions&Solutions
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