NCERT Class 10 Science Chapter 1: Chemical Reactions and Equations - A Detailed Explanation

Introduction to Chemical Reactions and Equations

Welcome, students! Have you ever wondered what happens when milk is left at room temperature during summer, an iron pan is left exposed to humid air, or when you digest food? In each of these scenarios, the nature and the identity of the initial substance have changed. This transformation is what we call a chemical change. Science describes these changes through a process known as a chemical reaction. In essence, a chemical reaction is a process that leads to the chemical transformation of one set of chemical substances to another. This chapter, 'Chemical Reactions and Equations', is your first step into the fascinating world of chemistry in Class 10. We will learn how to represent these reactions concisely using chemical equations, how to balance them, and explore the various types of reactions that occur all around us, from a simple fire to complex biological processes. Understanding these fundamentals is crucial as it lays the groundwork for all subsequent chemistry chapters.

A chemical reaction involves the breaking of old chemical bonds and the formation of new ones, resulting in new substances with entirely different properties. For instance, when hydrogen gas (flammable) reacts with oxygen gas (supports combustion), they form water, a substance that \textinguishes fire. This dramatic change in properties is a hallmark of a chemical reaction. Throughout this chapter, we will delve deep into identifying, representing, and classifying these incredible transformations.

Chemical Equations

How do we describe a chemical reaction? We could write a long sentence, but that's inefficient. Chemists use a shorthand notation called a chemical equation. It's a symbolic representation of a chemical reaction in the form of symbols and formulae, where the reactant entities are given on the left-hand side and the product entities on the right-hand side.

Writing a Chemical Equation

Let's consider an activity from your textbook: burning a magnesium ribbon in the air. When you burn a magnesium ribbon, it burns with a dazzling white flame and changes into a white powder. This powder is magnesium oxide. The sentence description is quite long. We can write this in a shorter form called a word equation:

Magnesium + Oxygen → Magnesium oxide

In this equation, the substances that undergo change, Magnesium and Oxygen, are called reactants. The new substance formed, Magnesium oxide, is called the product. An arrow points from the reactants to the products, indicating the direction of the reaction.

To make the equation even more concise and universally understood, we use chemical formulae instead of words. The chemical equation for the above reaction is:

Mg + O₂ → MgO

This is known as a skeletal chemical equation because it's unbalanced. If you count the atoms of each element on both sides, you'll notice they are not equal. There is one magnesium atom on both sides, but two oxygen atoms on the left (in O₂) and only one on the right (in MgO). This brings us to a fundamental concept in chemistry.

Balanced Chemical Equations

According to the Law of Conservation of Mass, which you studied in Class 9, mass can neither be created nor destroyed in a chemical reaction. This means the total mass of the elements present in the products of a chemical reaction must be equal to the total mass of the elements present in the reactants. In simpler terms, the number of atoms of each element must remain the same before and after a chemical reaction. Therefore, we must balance the skeletal chemical equation.

Let's learn how to balance an equation using the 'hit-and-trial' method with an example: the reaction of iron with steam.

Word Equation: Iron + Steam → Iron(II,III) oxide + Hydrogen

Skeletal Equation: Fe + H₂O → Fe₃O₄ + H₂

Step-by-step Balancing:

• Step 1: Draw boxes around each formula. Do not change anything inside the boxes while balancing. Fe + H₂O → Fe₃O₄ + H₂ becomes [Fe] + [H₂O] → [Fe₃O₄] + [H₂].
• Step 2: List the number of atoms of different elements on both sides of the equation.

• Step 3: Start balancing with the compound that contains the maximum number of atoms. It can be a reactant or a product. In that compound, select the element which has the maximum number of atoms. In our example, this is Fe₃O₄, and the element is Oxygen. There are 4 oxygen atoms on the RHS and only 1 on the LHS. To balance the oxygen atoms, we place a coefficient '4' before the H₂O molecule on the LHS.

Now the equation is: [Fe] + 4[H₂O] → [Fe₃O₄] + [H₂]

Let's check the atoms again: Oxygen is balanced (4 on both sides), but now Hydrogen has 4 x 2 = 8 atoms on the LHS and only 2 on the RHS.

• Step 4: Balance the hydrogen atoms. To make the hydrogen atoms 8 on the RHS, we place a coefficient '4' before the H₂ molecule.

Now the equation is: [Fe] + 4[H₂O] → [Fe₃O₄] + 4[H₂]

Hydrogen is now balanced (8 on both sides).

• Step 5: Balance the iron atoms. There is 1 iron atom on the LHS and 3 on the RHS. To balance, we place a coefficient '3' before Fe on the LHS.

Now the equation is: 3[Fe] + 4[H₂O] → [Fe₃O₄] + 4[H₂]

• Step 6: Finally, check the correctness of the balanced equation by counting the atoms of each element on both sides one last time.

The number of atoms of each element is equal on both sides. The equation is now balanced!

Writing Symbols of Physical States

A chemical equation can be made more informative by mentioning the physical states of the reactants and products. The gaseous, liquid, aqueous, and solid states are represented by the notations (g), (l), (aq), and (s), respectively. The word 'aqueous' (aq) is used when a reactant or product is present as a solution in water.

Our balanced equation becomes:

3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Note that we use (g) for water because in this reaction, it is used in the form of steam.

Sometimes, reaction conditions such as temperature, pressure, or the presence of a catalyst are also indicated above or below the arrow. For example:

• Photosynthesis: 6CO₂(aq) + 6H₂O(l) --(Sunlight / Chlorophyll)--> C₆H₁₂O₆(aq) + 6O₂(aq)

Types of Chemical Reactions

Chemical reactions are classified into different types based on the nature of the chemical change occurring. Let's explore the major types of reactions.

Combination Reaction

As the name suggests, a combination reaction is a reaction in which two or more reactants combine to form a single product.

The general form is: A + B → C

Examples:

• Burning of coal: Solid carbon burns in the presence of oxygen to form carbon dioxide gas.

  C(s) + O₂(g) → CO₂(g)

• Formation of water: Hydrogen gas combines with oxygen gas to form liquid water.

  2H₂(g) + O₂(g) → 2H₂O(l)

• Formation of slaked lime: Calcium oxide (quick lime) reacts vigorously with water to produce calcium hydroxide (slaked lime), releasing a large amount of heat.

  CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

  Reactions in which heat is released along with the formation of products are called exothermic chemical reactions. The reaction above is a prime example. Respiration is another example of an exothermic reaction where glucose combines with oxygen to produce energy.

Decomposition Reaction

A decomposition reaction is the opposite of a combination reaction. In this reaction, a single compound breaks down to give two or more simpler substances.

The general form is: C → A + B

These reactions require energy in the form of heat, light, or electricity to break the bonds of the reactant. Reactions that absorb energy are known as endothermic reactions. Decomposition reactions are typically endothermic.

Thermal Decomposition

When a decomposition reaction is carried out by heating, it is called thermal decomposition.

Examples:

• Decomposition of ferrous sulphate: When green crystals of ferrous sulphate (FeSO₄·7H₂O) are heated, they first lose water. On further heating, the anhydrous ferrous sulphate (FeSO₄) decomposes to form ferric oxide (Fe₂O₃), sulphur dioxide (SO₂), and sulphur trioxide (SO₃). The green color changes to reddish-brown, and a characteristic smell of burning sulphur is observed.

  2FeSO₄(s) --(Heat)--> Fe₂O₃(s) + SO₂(g) + SO₃(g)

• Decomposition of calcium carbonate: On heating, calcium carbonate (limestone) decomposes to form calcium oxide (quick lime) and carbon dioxide. This is a very important reaction used in various industries, including the manufacture of cement.

  CaCO₃(s) --(Heat)--> CaO(s) + CO₂(g)

• Decomposition of lead nitrate: When white lead nitrate powder is heated, it decomposes to form yellow lead oxide (PbO), brown fumes of nitrogen dioxide (NO₂), and oxygen gas (O₂).

  2Pb(NO₃)₂(s) --(Heat)--> 2PbO(s) + 4NO₂(g) + O₂(g)

Electrolytic Decomposition (Electrolysis)

When decomposition is carried out by passing electricity, it is known as electrolytic decomposition or electrolysis.

Example:

• Electrolysis of water: When an electric current is passed through acidulated water, it decomposes into hydrogen gas and oxygen gas. Hydrogen gas is collected at the cathode (negative electrode) and oxygen gas is collected at the anode (positive electrode). The volume of hydrogen gas collected is double the volume of oxygen gas.

  2H₂O(l) --(Electricity)--> 2H₂(g) + O₂(g)

Photolytic Decomposition (Photolysis)

When decomposition is carried out in the presence of sunlight, it is called photolytic decomposition or photolysis.

Examples:

• Decomposition of silver chloride: When white silver chloride is exposed to sunlight, it decomposes into grey silver and chlorine gas.

  2AgCl(s) --(Sunlight)--> 2Ag(s) + Cl₂(g)

• Decomposition of silver bromide: Similarly, silver bromide also decomposes in sunlight.

  2AgBr(s) --(Sunlight)--> 2Ag(s) + Br₂(g)

  This reaction is the principle behind black-and-white photography.

Displacement Reaction

A displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its compound (usually in its salt solution).

The general form is: A + BC → AC + B (where A is more reactive than B)

Examples:

• Iron and Copper Sulphate: When an iron nail is dipped in a blue solution of copper sulphate, the iron, being more reactive than copper, displaces copper from the solution. The blue color of the copper sulphate solution fades and turns light green due to the formation of iron(II) sulphate, and a reddish-brown coating of copper metal is deposited on the iron nail.

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

• Zinc and Copper Sulphate: Zinc is also more reactive than copper. When a zinc strip is placed in a copper sulphate solution, it displaces copper.

  Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

• Lead and Copper Chloride: Lead is more reactive than copper. It will displace copper from its salt solution.

  Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)

The ability of a metal to displace another depends on its position in the reactivity series.

Double Displacement Reaction

A double displacement reaction is a reaction in which there is an exchange of ions between the reactants. Two compounds react to form two new compounds.

The general form is: AB + CD → AD + CB

These reactions often occur in aqueous solutions and one of the products formed is usually a precipitate or a gas.

Example: Precipitation Reaction

• When a solution of sodium sulphate is mixed with a solution of barium chloride, a white, insoluble substance called a precipitate of barium sulphate is formed, along with a solution of sodium chloride.

  Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)

  Any reaction that produces a precipitate is called a precipitation reaction.

Oxidation and Reduction (Redox Reactions)

Oxidation and reduction are two simultaneous processes that occur together in a single reaction, known as a redox reaction.

Oxidation

Oxidation is a process which involves:

• The gain of oxygen by a substance.
• The loss of hydrogen from a substance.

For example, when copper powder is heated in a china dish, the surface of the copper powder becomes coated with a black substance, which is copper(II) oxide. Here, copper has gained oxygen and is said to be oxidized.

2Cu(s) + O₂(g) --(Heat)--> 2CuO(s)

Reduction

Reduction is a process which involves:

• The loss of oxygen from a substance.
• The gain of hydrogen by a substance.

For example, if hydrogen gas is passed over the heated black copper(II) oxide (CuO), the black coating on the surface turns brown as the reverse reaction takes place and copper is obtained. Here, copper oxide is losing oxygen and is said to be reduced.

CuO(s) + H₂(g) --(Heat)--> Cu(s) + H₂O(l)

Redox Reactions

In the reaction above (CuO + H₂), copper oxide is losing oxygen (reduction), while hydrogen is gaining oxygen (oxidation). Such reactions, where one reactant gets oxidized while the other gets reduced, are called oxidation-reduction reactions or redox reactions.

CuO + H₂ → Cu + H₂O

• Substance Oxidized: H₂ (it gains oxygen)
• Substance Reduced: CuO (it loses oxygen)
• Oxidizing Agent: The substance that gives oxygen or causes oxidation. Here, it is CuO.
• Reducing Agent: The substance that takes oxygen or causes reduction. Here, it is H₂.

Another example:

ZnO + C → Zn + CO

• Zinc oxide (ZnO) is reduced to Zinc (Zn).
• Carbon (C) is oxidized to Carbon monoxide (CO).
• ZnO is the oxidizing agent.
• C is the reducing agent.

Effects of Oxidation in Everyday Life

Oxidation has damaging effects on metals and food. Let's discuss two common phenomena: corrosion and rancidity.

Corrosion

Corrosion is the process of slow deterioration of metals when they are exposed to substances such as moisture, air, water, acids, etc. The most common example of corrosion is the rusting of iron. When an iron article is exposed to moist air for a long time, it acquires a coating of a brown, flaky substance called rust. Rust is hydrated iron(III) oxide (Fe₂O₃·xH₂O).

Corrosion causes significant damage to car bodies, bridges, iron railings, ships, and all objects made of metal, especially iron. It is a wasteful process.

Other examples of corrosion include:

• The formation of a black coating on silver articles (silver sulfide).
• The formation of a green coating on copper articles (basic copper carbonate).

Prevention of Corrosion: Corrosion can be prevented by shielding the metal surface from air and moisture. Methods include:

• Painting: Applying a coat of paint.
• Oiling or Greasing: Applying a layer of oil or grease.
• Galvanization: Coating iron with a layer of a more reactive metal like zinc.
• Electroplating: Coating with less reactive metals like chromium or tin.
• Alloying: Mixing the metal with other metals or non-metals to form an alloy, which is more resistant to corrosion (e.g., stainless steel).

Rancidity

Have you ever tasted or smelled old chips or nuts? They have an unpleasant taste and smell. This is due to rancidity. When foods containing fats and oils are left exposed to air, they get oxidized. This oxidation process makes them rancid, resulting in a change in their smell and taste, rendering them unfit for consumption.

Prevention of Rancidity:

• Adding Antioxidants: These are substances that prevent oxidation. They are commonly added to fatty foods. BHA (Butylated hydroxyanisole) and BHT (Butylated hydroxytoluene) are common antioxidants.
• Vacuum Packing / Using Inert Gas: To prevent oxidation, manufacturers of potato chips flush bags of chips with an inert gas like nitrogen. This prevents the chips from coming into contact with oxygen.
• Refrigeration: Keeping food at low temperatures slows down the oxidation process.
• Airtight Containers: Storing food in airtight containers reduces its exposure to oxygen from the air.

Important Questions and Answers

Here are some solved questions from the chapter to help you test your understanding.

Question 1: Why should a magnesium ribbon be cleaned before burning in air?

Answer: A magnesium ribbon should be cleaned before burning in air to remove the protective layer of magnesium oxide (MgO) from its surface. Magnesium is a reactive metal and it readily reacts with oxygen from the air to form a thin, dull layer of magnesium oxide. This layer is quite stable and can prevent or slow down the burning of the magnesium ribbon. Cleaning the ribbon with sandpaper removes this layer, exposing the pure magnesium metal, which can then readily ignite and burn with a dazzling white flame as intended.

Question 2: Write the balanced equation for the following chemical reactions.

(i) Hydrogen + Chlorine → Hydrogen chloride
(ii) Barium chloride + Aluminium sulphate → Barium sulphate + Aluminium chloride
(iii) Sodium + Water → Sodium hydroxide + Hydrogen

Answer:

(i) Hydrogen + Chlorine → Hydrogen chloride

• Skeletal Equation: H₂ + Cl₂ → HCl
• Balancing: There are 2 H atoms and 2 Cl atoms on the LHS, but only 1 of each on the RHS. We place a coefficient '2' before HCl.
• Balanced Equation: H₂(g) + Cl₂(g) → 2HCl(g)

(ii) Barium chloride + Aluminium sulphate → Barium sulphate + Aluminium chloride

• Skeletal Equation: BaCl₂ + Al₂(SO₄)₃ → BaSO₄ + AlCl₃
• Balancing: Let's balance the SO₄ ions first. There are 3 on the LHS and 1 on the RHS. Place a '3' before BaSO₄: BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + AlCl₃
• Now, Ba is unbalanced (1 on LHS, 3 on RHS). Place a '3' before BaCl₂: 3BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + AlCl₃
• Now, Cl is unbalanced (3x2=6 on LHS, 3 on RHS). Place a '2' before AlCl₃: 3BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + 2AlCl₃
• Finally, check Al atoms. There are 2 on the LHS and 2 on the RHS. The equation is now balanced.
• Balanced Equation: 3BaCl₂(aq) + Al₂(SO₄)₃(aq) → 3BaSO₄(s) + 2AlCl₃(aq)

(iii) Sodium + Water → Sodium hydroxide + Hydrogen

• Skeletal Equation: Na + H₂O → NaOH + H₂
• Balancing: H is unbalanced (2 on LHS, 3 on RHS). Let's put a '2' before H₂O: Na + 2H₂O → NaOH + H₂. Now H is 4 on LHS, 3 on RHS.
• Let's try putting a '2' before NaOH: Na + H₂O → 2NaOH + H₂. Now Na is unbalanced. Let's put a '2' before Na: 2Na + H₂O → 2NaOH + H₂. Now H is unbalanced (2 on LHS, 4 on RHS).
• Let's restart with the original skeletal equation and use a different approach. The odd number of H atoms on the RHS is the problem. Let's make it even by placing a '2' before NaOH: Na + H₂O → 2NaOH + H₂.
• Now Na is 1 on LHS and 2 on RHS. Place a '2' before Na: 2Na + H₂O → 2NaOH + H₂.
• Now H is 2 on LHS and (2x1 + 2) = 4 on RHS. Place a '2' before H₂O: 2Na + 2H₂O → 2NaOH + H₂.
• Let's check all atoms: Na (2 LHS, 2 RHS), H (4 LHS, 4 RHS), O (2 LHS, 2 RHS). It is balanced.
• Balanced Equation: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

Question 3: What is a balanced chemical equation? Why should chemical equations be balanced?

Answer: A balanced chemical equation is a chemical equation in which the number of atoms of each element is equal on both the reactant side and the product side. The coefficients (stoichiometric coefficients) are adjusted to ensure this equality. For example, 2H₂ + O₂ → 2H₂O is a balanced equation.

Chemical equations should be balanced to satisfy the Law of Conservation of Mass. This fundamental law of chemistry states that mass can neither be created nor destroyed during a chemical reaction. This implies that the total mass of the reactants must be equal to the total mass of the products. Since mass is directly related to the number of atoms of each element, balancing the number of atoms on both sides of the equation ensures that the law is upheld and the equation accurately represents the chemical transformation.

Question 4: A shiny brown coloured element ‘X’ on heating in air becomes black in colour. Name the element ‘X’ and the black coloured compound formed.

Answer:

• The shiny brown coloured element ‘X’ is Copper (Cu).
• When copper is heated in the presence of air (which contains oxygen), it undergoes oxidation. It reacts with oxygen to form a black coloured compound.
• This black coloured compound is Copper(II) oxide (CuO).
• The chemical reaction involved is:

  2Cu(s) + O₂(g) --(Heat)--> 2CuO(s)

  (Brown)                        (Black)

Chapter Summary

Here is a quick summary of the key concepts from this chapter for your revision:

• A chemical reaction is a process that involves the rearrangement of the structure of molecules or ions, resulting in the formation of new substances.
• Chemical reactions are represented by chemical equations, which use symbols and formulae for reactants and products.
• A chemical equation must be balanced to adhere to the Law of Conservation of Mass, ensuring an equal number of atoms of each element on both sides.
• Combination Reaction: Two or more substances combine to form a single product (A + B → C).
• Decomposition Reaction: A single compound breaks down into two or more simpler substances (C → A + B). It requires energy and can be thermal, electrolytic, or photolytic.
• Displacement Reaction: A more reactive element displaces a less reactive element from its salt solution (A + BC → AC + B).
• Double Displacement Reaction: An exchange of ions occurs between two reacting compounds (AB + CD → AD + CB). Precipitation reactions are a common type.
• Redox Reactions: Reactions involving simultaneous oxidation (gain of oxygen/loss of hydrogen) and reduction (loss of oxygen/gain of hydrogen).
• Corrosion: The gradual deterioration of metals due to their reaction with air, moisture, or chemicals in their surroundings. Rusting of iron is a prime example.
• Rancidity: The oxidation of fats and oils in food items, leading to an unpleasant smell and taste. It can be prevented by using antioxidants, nitrogen gas, or refrigeration.