Chemical Reactions and Equations Class 10 Notes Science Chapter 1

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.

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.  

Chemical Equations

Chemical Equation is representation of chemical reaction using symbols and formulae of the substances.  A chemical equation represents a chemical reaction. 

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.

• 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.

A balanced account of a chemical transaction is a complete chemical equation, which symbolically depicts the reactants, products, and their physical states.

2. Balanced Chemical Equations

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.

(ii) Balanced chemical equation

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.

Here are the steps to write a balanced chemical equation: Fe + H₂O → Fe₃O₄ + H₂

Step 1:  Write the number of atoms of elements present in reactants and in products in a table as shown here.

Step 2:  Balance the atom which is maximum in number on either side of a chemical equation.
Start by balancing the number of oxygen atoms, which has the maximum count on the right-hand side (RHS).
To balance oxygen, multiply the number of oxygen atoms on the left-hand side (LHS) by 4.
Fe + 4 × H₂O → Fe₃O₄ + H₂

Step 3: Balancing oxygen created an imbalance in the number of hydrogen atoms. There are now eight hydrogen atoms on the LHS and two on the RHS.
To balance hydrogen, multiply the number of hydrogen atoms on the RHS by 4.
Fe + 4 × H₂O → Fe₃O₄ + 4 × H₂

Step 4: There is only one iron atom on the LHS and three on the RHS.
To balance iron, multiply the number of iron atoms on the LHS by 3.
3 × Fe + 4 × H₂O → Fe3O4 + 4 × H2

The equation is now balanced, and the same number of atoms of each element is present on both sides.

After balancing, the above equation can be written as follows:
3Fe + 4H₂O → Fe₃O₄ + 4H₂ 

• To make a chemical equation more informative, it is important to include the physical state of the substances involved.
• The symbol (g) is used to represent a gaseous state, (l) for liquid state, (s) for solid state, and (aq) for an aqueous solution.
• The conditions under which the reaction takes place are also important to include in the chemical equation.
• These conditions can be written above and/or below the arrow in the chemical equation.
• Including this information can help provide a clearer understanding of the chemical reaction taking place.
• Example : 
• In this equation, (g) indicates that carbon dioxide is a gas, (l) indicates that water is a liquid, and (aq) indicates that glucose is dissolved in water. Sunlight and chlorophyll are specifically mentioned as the sources of energy used in the reaction.  This highlights the importance of these specific components in the photosynthesis process.

Types of Chemical Reactions

Considering various factors, chemical reactions are classified into several categories.

1. Combination Reaction  

This is the equation of a combination reaction. 

Example:   

(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)  A solution of slaked lime produced by the reaction is used for white-washing walls. Calcium hydroxide reacts with carbon dioxide in the air to form calcium carbonate. Calcium carbonate is formed after two to three days of whitewashing and gives a shiny finish to the walls. It is interesting to note that the chemical formula for marble is also CaCO₃.

Note:  Reactions that release heat when products are formed are known as exothermic chemical reactions.

Other Examples of Exothermic Reactions 

(a) Burning of Coal: When carbon is burnt in oxygen (air), carbon dioxide is formed. In this reaction, carbon is combined with oxygen.
C(s) + O₂ (g) → CO₂ (g) 

(b) Formation of H_{2 and O2}

2H₂ (g) + O₂ (g) → 2H₂O (l)

(c) Burning of natural gas 

CH₄ (g) + 2O₂ (g) → CO₂ (g) + 2H₂O (g)

Note: Do you know that respiration is an exothermic process?

We need energy to live, and we get it from the food we eat. When we digest food, it's broken down into simpler forms. For instance, rice, potatoes, and bread have carbohydrates, which are turned into glucose. This glucose then mixes with oxygen in our cells to produce energy.

C₆H₁₂O₆ (aq) + 6O₂ (aq) → 6CO₂ (aq) + 6H₂O (l) + energy

2. 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

Ferrous sulphate crystals (FeSO₄. 7H₂O) lose water when heated and the green colour of ferrous sulphate crystals fades. It then decomposes to ferric oxide (Fe₂0₃), sulphur dioxide (SO₂) and sulphur trioxide (SO₃). Ferric oxide is a solid, while S0₂ and SO₃ are gases.

Example 2: Decomposition of Calcium Carbonate

Heating calcium carbonate breaks it down into calcium oxide (lime) and carbon dioxide. This reaction is important in industries, with calcium oxide being used in cement production.

Example 3: Heating of Lead Nitrate Powder

(b) Photolytic decomposition

where light is needed for the reaction. An example is the photolytic decomposition of H₂O₂.
Photolytic Decomposition

Example: Effect of sunlight on Silver Chloride

White silver chloride turns grey in sunlight. This occurs due to the decomposition of silver chloride into silver and chlorine when exposed to light.
Silver bromide also behaves in the same way.
The above reactions are used in black and white photography.

(c) Electrolytic decomposition

where electricity is necessary for the reaction. An example is the electrolytic decomposition of H₂O. Electrolysis of water is the decomposition into oxygen and hydrogen gas due to an electric current passed through the water.

Note: What form of energy is causing these decomposition reactions?

Decomposition reactions need energy to break down the substances involved. This energy can come from heat, light, or electricity. These types of reactions, where energy is taken in rather than given off, are called endothermic reactions.

3. Displacement Reaction

Some Important Reactions are:
Reaction of iron nails with copper sulphate solution.

Fe (s) + CuSO₄(aq) →  FeSO₄(g) + Cu (s)

The iron nail becomes brownish in colour and the blue colour of copper sulphate solution fades. In this reaction, iron has displaced or removed another element, copper from the copper sulphate solution.

Other examples of displacement reactions are:

Zinc and lead are more reactive elements than copper. They displacecopper from its compounds.

4. Double Displacement Reaction  

Example: When the solution of barium chloride reacts with the solution of sodium sulphate, white precipitate of barium sulphate is formed along with sodium chloride.
Na₂SO₄ (aq) + BaCl₂ (aq) →  BaSO₄ (s) + NaCl (aq)

Note: Double Displacement Reaction, in which precipitate is formed, is also known as precipitation reaction. Neutralisation reactions are also examples of double displacement reactions

5. Endothermic and Exothermic Reaction

(i) Exothermic Reaction:  An exothermic reaction is a type of chemical reaction in which energy is released from the reaction system into the surroundings, usually in the form of heat, light, or sound. 

Example: Formation of Carbon dioxide
The chemical reaction can be depicted as:  
C(s) + O₂(g) → CO₂(g) + Heat

(ii) Endothermic Reaction: An endothermic reaction is a type of chemical reaction in which energy is absorbed from the surroundings in the form of heat, light, or electricity.  

Example: Photosynthesis
Plants absorb heat energy from sunlight to convert carbon dioxide and water into glucose and oxygen. 

Reaction as:   6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂

6. Oxidation and Reduction

(i) Oxidation is the process in which an atom, ion, or molecule loses one or more electrons. During oxidation, the oxidation state of the atom or molecule increases, since it becomes more positive or less negative.   

1. Corrosion

• Numerous metals, prone to chemical activity, easily react with moisture, oxygen, and acids. Iron, for example, starts as shiny but eventually rusts, forming a reddish-brown powder due to oxidation.

• Metals, like iron, face corrosion as their oxides don't firmly attach, causing flaking, structural weakening, and breakdown. The process of metal deterioration due to environmental substances is called corrosion.
• Copper and silver items tarnish when exposed to air and water. Copper develops a green oxide layer, while silver acquires a black oxide covering. These oxide formations act as barriers, limiting further exposure and reducing corrosion.
• Rusting:

  4Fe(s) + 3O₂(from air) + xH₂O(moisture) → 2Fe₂O₃.xH₂O(rust)
  Corrosion of copper:
  Cu(s) + H₂O(moisture) + CO₂(from air) → CuCO₃.Cu(OH)₂(green)

  Corrosion of silver:
  Ag(s) + H₂S (from air) → Ag₂S(black) + H₂(g)
  Rusting

2. Rancidity

• The taste and odour of food materials containing fat and oil change when they are left exposed to air for a long time. This is called rancidity. It is caused by the oxidation of fats and oils present in food materials.
• Prevention:
  (i) Seal food in air-tight containers.
  (ii) Employ nitrogen packaging.
  (iii) Store in refrigeration.
  (iv) Add antioxidants or preservatives.