CH 01: Chemical Reactions and Equations of Class 10th Science

Chemical Reactions and Equations
By Pratap Sanjay Sir

Chemical Reaction:

➥ It is the process in which new substances with new properties are formed.

Reactants (A + B) ➔ (AB)Products

Reactants: The substances which take part in a chemical reaction are called reactants.

Products: The new substances produced as a result of a chemical reaction are called products.

Example: Rusting of iron, Burning of wood, Respiration… etc.

➥ We use (g) for gas, (l) for liquid, (s) for solid, and (aq) for aqueous.

Chemical Equation:

➥ The symbolic representation of a chemical reaction is said to be a chemical equation.

Example: Zinc metal reacts with dilute sulphuric acid to form zinc sulfate and hydrogen gas.

Zinc + Sulphuric acid → Zinc sulfate + Hydrogen

Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

Magnesium + Oxygen → Magnesium oxide

(Reactants) → (Product)

Characteristics of Chemical Reactions:

1. Change in colour: Fe + CuSO₄ (blue) → FeSO₄ (blue-green) + Cu
2. Change in temperature: CaO + H₂O → Ca(OH)₂ + Heat
3. Change in state: H₂(g) + O₂(g) → H₂O(l)
4. Evolution of gas: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
5. Formation of precipitate: Pb(NO₃)₂(aq) + KI(aq) → PbI₂(s) + KNO₃(aq)
6. Endothermic reaction: CaCO₃ + Heat → CaO + CO₂
7. Exothermic reaction: CaO + H₂O → Ca(OH)₂ + Heat

Classification of Chemical Reactions Based on Energy Change:

1. Exothermic Reactions: Reaction in which heat is released along with the formation of products.
   Example: CH₄ + 2O₂ ➔ CO₂ + 2H₂O + Heat


2. Endothermic Reactions: The reactions which require energy in the form of heat, light, or electricity to break reactants are called endothermic reactions.
   Example: 2Pb(NO₃)₂ + Heat ➔ 2PbO + 4NO₂ + O₂

Balanced Chemical Equation:

➥ Number of atoms of each element in reactants = number of atoms of each element in products.

➥ Law of Conservation of Mass: Mass of reactants = Mass of products.

➥ Mass is neither created nor destroyed in a chemical reaction.

Example: Iron and Water Reaction

Unbalanced Equation: Fe + H₂O → Fe₃O₄ + H₂

Balanced Equation: 3Fe + 4H₂O → Fe₃O₄ + 4H₂

Element	Reactants	Products	Balanced
Fe	1	3	✔
H	4	4	✔
O	4	4	✔

Unbalanced Equation: CH₄ + O₂ → CO₂ + H₂O

Balanced Equation: CH₄ + 2O₂ → CO₂ + 2H₂O

Element	Reactants	Products	Balanced
C	1	1	✔
H	4	4	✔
O	4	4	✔

Types of Chemical Reactions:

1. Combination Reaction

➥ The reaction in which two or more substances combine to form a single new substance is called a combination reaction.

Examples:
1. CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat
(Quick lime reacts with water to form slaked lime.)

2. N₂(g) + 3H₂(g) → 2NH₃(g)
(Nitrogen and hydrogen combine to form ammonia.)

3. C(s) + O₂(g) → CO₂(g)
(Carbon reacts with oxygen to form carbon dioxide.)

2. Decomposition Reaction

➥The reaction in which a single substance breaks down into two or more simpler substances is called a decomposition reaction.

Examples:
1. 2H₂O(l) → 2H₂(g) + O₂(g)
(Electrolysis of water decomposes it into hydrogen and oxygen gases.)

2. CaCO₃(s) → CaO(s) + CO₂(g)
(Calcium carbonate decomposes to form calcium oxide and carbon dioxide.)

3. 2FeSO₄(s) → Fe₂O₃(s) + SO₂(g) + SO₃(g)
(Ferrous sulfate decomposes to ferric oxide, sulfur dioxide, and sulfur trioxide.)

Decomposition reactions can be of three types:

1. Thermal Decomposition or Thermolysis:

   ➥Thermal decomposition is a chemical reaction in which a substance breaks down into simpler substances when heated.

   Examples:

   1. CaCO₃ (s) → CaO (s) + CO₂ (g)
      (Calcium carbonate → Quicklime + Carbon dioxide)
   2. 2Pb(NO₃)₂ (s) → 2PbO (s) + 4NO₂ (g) + O₂ (g)
      (Lead nitrate → Lead oxide + Nitrogen dioxide + Oxygen)
   3. 2KClO₃ (s) → 2KCl (s) + 3O₂ (g)
      (Potassium chlorate → Potassium chloride + Oxygen)
2. Electrolytic Decomposition or Electrolysis:

   ➥Electrolytic decomposition occurs when an electric current is passed through an electrolyte, causing it to decompose.

   Examples:

   1. 2H₂O (l) → 2H₂ (g) + O₂ (g)
      (Water decomposes into hydrogen and oxygen gases.)
   2. 2NaCl (aq) → 2Na (s) + Cl₂ (g)
      (Electrolysis of brine produces sodium metal and chlorine gas.)
   3. Al₂O₃ → 4Al (s) + 3O₂ (g)
      (Electrolysis of alumina produces aluminum metal and oxygen gas.)

   Applications: Electrolysis is widely used in industries for metal extraction and chemical production.

3. Photochemical decomposition:

   ➥Photochemical decomposition occurs when a substance breaks down in the presence of light.

   Examples:

   1. 2AgCl (s) → 2Ag (s) + Cl₂ (g)
      (Silver chloride decomposes into silver and chlorine gas in sunlight.)
   2. 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
      (Photosynthesis: Carbon dioxide and water form glucose and oxygen.)
   3. 2HgO (s) → 2Hg (l) + O₂ (g)
      (Mercuric oxide decomposes into mercury and oxygen under sunlight.)

   Applications: Photochemical reactions are critical in natural processes like photosynthesis and are used in photographic films.

3. Displacement Reaction

➥A reaction in which a more reactive element displaces a less reactive element from its compound is called a displacement reaction.

Examples:
1. Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
(Zinc displaces copper from copper sulfate.)

2. Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)
(Iron displaces copper from copper sulfate.)

3. Pb(s) + 2AgNO₃(aq) → Pb(NO₃)₂(aq) + 2Ag(s)
(Lead displaces silver from silver nitrate.)

4. Double Displacement Reaction

➥A reaction in which two compounds react by exchanging ions to form two new compounds is called a double displacement reaction.

Examples:
1. NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl(s)
(Sodium chloride reacts with silver nitrate to form sodium nitrate and a white precipitate of silver chloride.)

2. BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
(Barium chloride reacts with sodium sulfate to form barium sulfate and sodium chloride.)

3. Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
(Lead nitrate reacts with potassium iodide to form lead iodide and potassium nitrate.)

5. Precipitation Reaction

➥A reaction in which an insoluble substance (precipitate) is formed during the reaction is called a precipitation reaction.

Examples:
1. BaCl₂(aq) + H₂SO₄(aq) → BaSO₄(s) + 2HCl(aq)
(Barium sulfate is formed as a precipitate.)

2. Na₂CO₃(aq) + CaCl₂(aq) → CaCO₃(s) + 2NaCl(aq)
(Calcium carbonate is formed as a precipitate.)

3. Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
(Lead iodide is formed as a yellow precipitate.)

Reactivity Series:

Symbol	Element	Reactivity
K	Potassium	Most Reactive
Na	Sodium
Ca	Calcium
Mg	Magnesium
Al	Aluminium
Zn	Zinc
Fe	Iron	Reactivity decreases
Pb	Lead
H	Hydrogen
Cu	Copper
Hg	Mercury
Ag	Silver
Au	Gold	Least Reactive

Trick to remember the Reactivity Series:

Please Stop Calling Me A Zebra, I Like Her Cute Smart Goat.

Oxidation: +Oxygen or -Hydrogen

Reduction: -Oxygen or +Hydrogen

Oxidation:

➥Oxidation is the process where a substance gains oxygen or loses hydrogen. This process involves the loss of electrons.

Example 1: 2Cu + O₂ → 2CuO

Here, copper (Cu) gains oxygen to form copper oxide (CuO), so copper is oxidized.

Example 2: 2Mg + O₂ → 2MgO

Magnesium (Mg) reacts with oxygen to form magnesium oxide (MgO), indicating oxidation.

Example 3: 4Fe + 3O₂ → 2Fe₂O₃

Iron (Fe) oxidizes in the presence of oxygen to form iron(III) oxide.

Reduction:

➥Reduction is the process where a substance loses oxygen or gains hydrogen. This involves the gain of electrons.

Example 1: CuO + H₂ → Cu + H₂O

Here, copper(II) oxide (CuO) is reduced by hydrogen gas (H₂) to form copper metal (Cu) and water.

Example 2: Fe₂O₃ + 3CO → 2Fe + 3CO₂

Iron(III) oxide (Fe₂O₃) is reduced by carbon monoxide (CO) to form iron (Fe) and carbon dioxide (CO₂).

Example 3: CuO + H₂ → Cu + H₂O

Copper(II) oxide (CuO) is reduced by hydrogen gas to form copper metal (Cu) and water.

Oxidizing Agent:

An oxidizing agent is a substance that causes oxidation by accepting electrons, and as a result, it gets reduced.

Example 1: 2Cu + O₂ → 2CuO

Oxygen (O₂) is the oxidizing agent because it accepts electrons and causes copper to oxidize.

Reducing Agent:

➥A reducing agent is a substance that causes reduction by losing electrons, and thus, it gets oxidized.

Example 1: CuO + H₂ → Cu + H₂O

Hydrogen (H₂) is the reducing agent because it donates electrons, reducing copper oxide to copper.

Redox Reaction:

➥A Redox reaction is a reaction where one reactant is oxidized and another reactant is reduced.

Example 1: ZnO + C → Zn + CO

In this reaction, zinc oxide (ZnO) is reduced to zinc (Zn), while carbon (C) is oxidized to carbon monoxide (CO).

Example 2: MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂

Manganese dioxide (MnO₂) is reduced to manganese chloride (MnCl₂), while chlorine gas (Cl₂) is released, showing oxidation.

Example 3: H₂ + Cl₂ → 2HCl

Hydrogen (H₂) is oxidized (loses electrons), and chlorine (Cl₂) is reduced (gains electrons) to form hydrochloric acid (HCl).

The Effects of Oxidation Reactions in Everyday Life

1. Corrosion

Definition: Corrosion is the gradual degradation of metals due to their reaction with atmospheric oxygen, moisture, or other chemicals.

Effect of Oxidation: Oxidation of metals, particularly iron, leads to the formation of metal oxides, commonly known as rust.

Impact: Corrosion weakens the metal structure, affecting its strength and durability.

Prevention:

1. By Painting
2. By Greasing and Oiling
3. By Galvanisation (coating of zinc layer over the surface of iron metal)

Example 1: Silver

Silver develops a black coating after some time:

Ag₂ + H₂S → Ag₂S + H₂

Example 2: Copper

Copper develops a green coating after some time due to oxidation in air and moisture:

2Cu + CO₂ + H₂O + O₂ → CuCO₃·Cu(OH)₂

Result: Copper Carbonate x Copper Hydroxide (Greenish colour)

Example 3: Iron

Iron rusts when exposed to oxygen:

Fe + O₂ → Fe₂O₃

Rusting continues with the addition of moisture:

4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃

2. Rancidity

Definition: Rancidity is the development of undesirable odors and flavors in fats and oils due to their exposure to oxygen.

Effect of Oxidation: Oxidation of the unsaturated fatty acids in fats and oils leads to the formation of rancid compounds.

Impact: Rancidity imparts unpleasant tastes and smells to food products, making them unpalatable and reducing their shelf life.

Prevention:

1. Refrigerating the foodstuff
2. Replacing air with nitrogen gas
3. Using airtight containers
4. Adding antioxidants

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