Chemical Reactions and Equations - Notes

Chemical Reactions and Equations

Introduction

Chemistry is a branch of science that deals with the composition and properties of matter and the changes that matter undergoes through different interactions.

Changes around us

Chemistry is the study of change. All changes that happen around us can be classified into two categories.

1. Physical Change

A change in which only the physical properties of a substance change is called a physical change.

In a physical change, no new substance is formed.

Examples: Melting of ice, folding of paper, breaking of glass, melting of candle wax and dissolving sugar in water.

2. Chemical Change

A change in which the chemical composition of substances changes is called a chemical change.

In a chemical change, one or more new substances are formed.

Examples: Rusting of iron and burning of wood.

A chemical compound is formed as a result of a chemical change. In this process, different types of energies are either absorbed or evolved.

The total mass of substances remains the same throughout the chemical change.

Chemical Reactions

The process by which a compound undergoes a chemical change to form a new compound is known as a chemical reaction.

This chemical change is shown with the help of a chemical equation.

The substance which undergoes the change in a chemical reaction is called the reactant.

The new substance formed in a chemical reaction is called the product.

Characteristics of a Chemical Reaction

A chemical reaction can be identified by the following changes:

1. Change in State

2. Change in Colour

3. Evolution of Gas

4. Change in Temperature

5. Formation of Precipitate

1. Change in State

Some chemical reactions are shown by a change in state, such as solid, liquid or gas.

Example:

2H₂ (g) + O₂ (g) → 2H₂O (l)

In this reaction, hydrogen and oxygen are gases, but they combine to form water, which is liquid in state.

2. Change in Colour

Some chemical reactions are shown by a change in the colour of the reacting compound.

Example:

2Pb(NO₃)₂ (s) → 2PbO (s) + 4NO₂ (g) + O₂ (g)

Lead nitrate is white in colour, but on heating it forms lead oxide, which is yellow in colour.

3. Evolution of Gases

In some chemical reactions, gas is released during the reaction. This is called evolution of gas.

Example:

Zn (s) + H₂SO₄ (dil.) → ZnSO₄ (aq) + H₂ (g)

In this reaction, hydrogen gas is evolved.

4. Change in Temperature or Energy Change

All chemical reactions take place either with the absorption or release of energy.

(a) Chemical Reaction with Release of Energy

Example:

2Mg (s) + O₂ (g) → 2MgO (s) + Heat and Light

When magnesium wire is heated from its tip in a Bunsen flame, it catches fire and burns with a dazzling white flame. It releases heat and light energy.

(b) Chemical Reaction with Absorption of Energy

Example:

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

Light energy is essential for biochemical reactions such as photosynthesis, in which green plants prepare their food from carbon dioxide and water.

5. Formation of Precipitate

Some chemical reactions result in the formation of an insoluble substance called a precipitate.

When an aqueous solution of a chemical compound is mixed with another aqueous solution, a precipitate may be formed.

Example:

AgNO₃ (aq) + NaCl (aq) → NaNO₃ (aq) + AgCl (s)

In this reaction, silver chloride is formed as a white precipitate.

Chemical Equations

All chemical reactions are represented by chemical equations.

A chemical equation is a shorthand representation of a chemical reaction using symbols and formulae of the substances involved in the chemical reaction.

The symbols and formulae of the elements or compounds are written to show the reactants and products of a chemical reaction.

Types of Chemical Equations

There are two types of chemical equations:

1. Word Equation

A word equation is an equation which links together the names of the reactants and products.

Example: When sodium metal reacts with water, sodium hydroxide and hydrogen gas are formed. The word equation is written as:

Sodium + Water → Sodium hydroxide + Hydrogen gas

Rules for Writing a Word Equation

There are some simple rules for writing a word equation.

• On the left side, write the names of the reactants with a plus sign (+) between them.
• On the right side, write the names of the products with a plus sign (+) between them.
• The direction of the arrow shows the direction of the reaction.
• An arrow (→) is placed between the reactants and products.

2. Symbol Equation

A symbol equation gives the chemical reaction in terms of formulae and symbols of the elements and compounds involved in the reaction.

In a symbol equation, symbols and formulae are written instead of their word names.

Example: Sodium metal reacts with water to give sodium hydroxide and hydrogen gas.

Na + H₂O → NaOH + H₂

Unbalanced and Balanced Equations

In an unbalanced chemical equation, the number of atoms of different elements on both sides of the equation are not equal.

For example, in the equation given below, the number of chlorine atoms and hydrogen atoms on both sides is not equal. So, it is called an unbalanced equation.

Mg + HCl → MgCl₂ + H₂

An unbalanced equation is also called a skeletal equation.

In a balanced chemical equation, the number of atoms of different elements on both sides of the equation are equal.

The balanced equation for magnesium reacting with hydrochloric acid is:

Mg + 2HCl → MgCl₂ + H₂

Importance of Balanced Chemical Equation

Balancing of a chemical equation is necessary to follow the Law of Conservation of Mass.

According to this law, mass can neither be created nor destroyed in a chemical reaction.

So, the total mass of reactants must be equal to the total mass of products.

Question. Why should chemical reactions be balanced?

Answer: Chemical reactions should be balanced because the total number of atoms of each element must remain the same on both sides of the equation. This follows the Law of Conservation of Mass.

Balancing of Chemical Equations

Chemical equations are balanced by making the number of atoms of different elements equal on both sides of the equation.

Balancing is usually done by the hit and trial method.

In this method, coefficients are placed before symbols or formulae of reactants and products.

The number of atoms of each element on both sides of the arrow are made equal.

This balancing is also called mass balancing, because the atoms of elements on both sides are equal and their masses are also equal.

Steps Involved in Balancing a Chemical Equation

• Write the chemical equation in the form of word equation.
• Convert the word equation into symbol equation.
• Make the atoms of different elements equal on both sides by adding suitable coefficients.
• Try to make the equation more informative, if possible.

Example 1: Zinc reacts with dilute sulphuric acid

Zinc reacts with dilute sulphuric acid to give zinc sulphate and hydrogen gas.

The word equation is:

Zinc + Sulphuric acid → Zinc sulphate + Hydrogen

The symbol equation is:

Zn + H₂SO₄ → ZnSO₄ + H₂

Now, count the number of atoms of all elements on both sides.

Element	No. of atoms of reactants (L.H.S.)	No. of atoms of products (R.H.S.)
Zn	1	1
S	1	1
H	2	2
O	4	4

In this equation, the number of atoms of all elements are equal on both sides. So, the equation is already balanced and no further balancing is needed.

Example 2: Iron reacts with water steam

Iron reacts with water steam to form iron (III) oxide and hydrogen gas.

The word equation is:

Iron + Water → Iron (III) oxide + Hydrogen gas

The symbol equation is:

Fe + H₂O → Fe₂O₃ + H₂

Element	No. of atoms of reactants (L.H.S.)	No. of atoms of products (R.H.S.)
Fe	1	2
O	1	3
H	2	2

Here, hydrogen atoms are equal on both sides, but iron and oxygen atoms are not equal.

First, balance oxygen atoms. The product side has 3 oxygen atoms, so we put coefficient 3 before H₂O.

Fe + 3H₂O → Fe₂O₃ + H₂

Now oxygen atoms are balanced, but hydrogen atoms are not balanced.

To balance hydrogen atoms, put coefficient 3 before H₂ on the product side.

Fe + 3H₂O → Fe₂O₃ + 3H₂

Now hydrogen and oxygen atoms are balanced, but iron atoms are still not balanced.

To balance iron atoms, put coefficient 2 before Fe on the reactant side.

2Fe + 3H₂O → Fe₂O₃ + 3H₂

Now all atoms are equal on both sides. So, the equation is balanced.

Question. Balance the following equations.

1. Fe (s) + H₂O (l) → Fe₃O₄ (s) + H₂ (g)

2. NH₄Cl (s) → NH₃ (g) + HCl (g)

3. Pb(NO₃)₂ (s) → PbO (s) + NO₂ (g) + O₂ (g)

4. C₃H₈ (g) + O₂ (g) → CO₂ (g) + H₂O (g)

5. C₂H₆ (g) + O₂ (g) → CO₂ (g) + H₂O (g)

Answers:

1. 3Fe + 4H₂O → Fe₃O₄ + 4H₂

2. NH₄Cl → NH₃ + HCl

3. 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

4. C₃H₈ + 5O₂ → 3CO₂ + 4H₂O

5. 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O

Writing State Symbols

A chemical equation becomes more informative when the physical states of reactants and products are also written.

These symbols are called state symbols.

• (s) is used for solid state.
• (l) is used for liquid state.
• (g) is used for gaseous state.
• (aq) is used for aqueous solution, which means the substance is dissolved in water.

A gas evolved in a reaction is shown by the symbol (↑).

A precipitate formed in a reaction is shown by the symbol (↓).

Short form ppt is also used to show precipitate formation.

Examples:

ZnCO₃ (s) → ZnO (s) + CO₂ ↑

AgNO₃ (aq) + KCl (aq) → AgCl ↓ + KNO₃ (aq)

Importance of State Symbols

State symbols are very important for chemical reactions that involve the release or absorption of heat energy.

Example:

2H₂ (g) + O₂ (g) → 2H₂O (l) + 572 kJ

2H₂ (g) + O₂ (g) → 2H₂O (g) + 484 kJ

Both reactions show the formation of water, but the heat released is different because water is formed in different physical states.

When water is liquid, it is written as H₂O (l). When water is gas, it is written as H₂O (g).

Importance of Chemical Equation

1. The weight of reactants or products can be found from the chemical equation.
2. The chemical equation follows the Law of Conservation of Mass. So, the total weight of reactants is equal to the total weight of products.
3. It gives information about the substances taking part and formed in the reaction.
4. The information about the number of molecules of reactants and products is revealed from the chemical equation.

Some Limitations of Chemical Equations

1. It does not give information about the physical state of reactants and products. For example, solid, liquid or gas.
2. It does not give information about the concentration of reactants and products.
3. It does not give information about the speed of reaction.
4. It does not tell about favourable conditions such as temperature, catalyst and pressure.
5. It does not show whether heat is absorbed or released during the reaction.
6. It does not show the percentage of reactants converted into products.
7. It does not show whether the reaction is reversible or irreversible.

Types of Chemical Reactions

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1. Combination Reaction

Two or more reactants give single product.

2. Decomposition Reaction

Single reactant gives multiple products

3. Displacement Reaction

Strong element replacing the weak element from its compound.

4. Double Displacement Reaction

Anions and cations of two substances are exchanged.

5. Redox reactions

Transfer of electrons, hydrogen atoms or oxygen atoms takes place

Types of Chemical Reactions

1. Combination / Addition Reaction

A Combination (Addition) Reaction is a chemical reaction in which two or more substances or elements combine to form a single new substance (compound).

Such reactions can occur by applying heat, light, pressure or electricity.

Example:

H₂(g) + Cl₂(g) → 2HCl(g)

Here, hydrogen and chlorine, which are two elements, combine to form a new compound called hydrogen chloride (HCl).

🧪 Activity – Combination Reaction

🎯 Aim

To study the combination reaction between calcium oxide (quick lime) and water.

📝 Procedure

• Take a small amount of calcium oxide (quick lime) in a beaker.
• Slowly add water to it.
• Carefully touch the outside of the beaker.

👀 Observation

A vigorous reaction takes place and the beaker becomes very hot.

💡 Conclusion

Calcium oxide combines with water to form calcium hydroxide, also called slaked lime.

This reaction releases a large amount of heat. Therefore, it is a highly exothermic reaction.

CaO(s) + H₂O(l) → Ca(OH)₂(aq.) + Heat

Quick lime + Water → Slaked lime + Heat

2. Decomposition Reaction

A Decomposition Reaction is a chemical reaction in which one compound breaks down into two or more simpler substances.

Decomposition reactions take place by applying heat, electricity or light.

(i) Thermal Decomposition Reaction

A decomposition reaction that takes place by the application of heat is called a thermal decomposition reaction.

Examples:

CaCO₃ Δ ⟶ CaO + CO₂

Calcium carbonate decomposes on heating to form calcium oxide and carbon dioxide.

CaCO₃ Δ ⟶ CaO + CO₂

During this reaction:

• Lead nitrate is white in colour.
• Lead oxide (PbO) is yellow.
• Nitrogen dioxide (NO₂) is a brown coloured gas.

💡 More to Know

2NaN₃ → 2Na + 3N₂

• This reaction is used in airbags of automobiles.
• On decomposition, sodium azide (NaN₃) rapidly releases nitrogen gas (N₂), which inflates the airbag.
• A decomposition reaction is the opposite of a combination reaction.

(ii) Electrolytic (Electrical) Decomposition Reaction

A decomposition reaction that takes place by the application of electric current is called an electrolytic decomposition reaction or electrolysis.

Examples:

2NaCl Electricity ⟶ 2Na + Cl₂

2Al₂O₃ Electricity ⟶ 4Al + 3O₂

🧪 Activity – Electrolysis of Water

🎯 Aim

To study the electrolysis of water.

📝 Procedure

• Take a plastic mug and make two holes at its base.
• Insert rubber stoppers and fix carbon electrodes through them.
• Connect the electrodes to a 6 V battery.
• Fill the mug with water so that the electrodes are immersed.
• Add a few drops of dilute sulphuric acid to the water.
• Fill two test tubes with water and invert them over the two electrodes.
• Switch on the electric current and leave the setup undisturbed for some time.
• Bubbles start forming at both electrodes and gradually collect inside the test tubes.
• After the test tubes are filled with gases, remove them carefully.
• Test the gases by bringing a burning candle near the mouth of each test tube.

👀 Observation

• Bubbles are produced at both electrodes.
• The gas collected at the negative electrode is twice the volume of the gas collected at the positive electrode.
• The gas collected at the negative electrode burns with a pop sound. This gas is hydrogen.
• The gas collected at the positive electrode makes the candle burn more brightly. This gas is oxygen.

💡 Conclusion

Electrolysis of water produces 2 volumes of hydrogen and 1 volume of oxygen.

Therefore, water is a compound made up of hydrogen and oxygen in the ratio of 2 : 1 by volume.

2H₂O(l) Electricity → 2H₂(g) + O₂(g)

Water → Hydrogen + Oxygen

(iii) Photolytic (Photo) Decomposition Reaction

A decomposition reaction that takes place in the presence of light is called a photolytic (photo) decomposition reaction.

Examples:

2AgBr Light ⟶ 2Ag + Br₂

Silver bromide is yellow in colour, while silver formed is grey.

2AgCl Light ⟶ 2Ag + Cl₂

Silver chloride is white in colour, while silver formed is grey.

📸 Note

These reactions are used in black and white photography.

🧪 Activity – Decomposition of Ferrous Sulphate

🎯 Aim

To study the decomposition of ferrous sulphate on heating.

📝 Procedure

• Take about 2 g of ferrous sulphate crystals in a dry boiling tube.
• Observe and note the colour of the crystals.
• Heat the boiling tube over the flame of a burner or spirit lamp.
• Observe the colour of the substance after heating.

👀 Observation

• Green coloured ferrous sulphate crystals (FeSO₄.7H₂O) first lose water on heating.
• They form anhydrous ferrous sulphate (FeSO₄).
• On further heating, anhydrous ferrous sulphate decomposes.
• A reddish-brown residue is left behind.
• Sulphur dioxide (SO₂) and sulphur trioxide (SO₃) gases are evolved.

💡 Conclusion

The reddish-brown residue is ferric oxide (Fe₂O₃).

The decomposition takes place in two steps.

Step 1: Removal of water

FeSO₄.7H₂O Δ ⟶ FeSO₄ + 7H₂O

Ferrous sulphate crystals → Anhydrous ferrous sulphate + Water

Step 2: Decomposition of anhydrous ferrous sulphate

2FeSO₄ Δ ⟶ Fe₂O₃ + SO₂ + SO₃

Anhydrous ferrous sulphate → Ferric oxide + Sulphur dioxide + Sulphur trioxide

3. Displacement Reaction

A Displacement Reaction is a chemical reaction in which a more reactive element replaces a less reactive element from its compound.

Displacement reactions take place according to the reactivity series of metals.

Examples:

1. Zinc displaces hydrogen from sulphuric acid.

Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂ ↑

Here, zinc is more reactive than hydrogen, so it displaces hydrogen from sulphuric acid.

2. Iron displaces copper from copper sulphate solution.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Here, iron is more reactive than copper, so it displaces copper from copper sulphate solution.

💡 Note

Only a stronger (more reactive) metal can replace a weaker (less reactive) metal from its compound.

The decreasing order of the strength of metals is called the Reactivity Series.

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au

From Potassium (K) to Gold (Au), the ability of metals to displace other metals gradually decreases.

4. Double Displacement Reaction

A Double Displacement Reaction is a chemical reaction in which the positive and negative ions of two compounds exchange with each other to form two new compounds.

(i) Double Displacement with Precipitate Formation

In some double displacement reactions, an insoluble solid called a precipitate is formed.

Examples:

NaCl(aq) + AgNO₃(aq) → AgCl(s) ↓ + NaNO₃(aq)

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

Here, AgCl and BaSO₄ are insoluble precipitates.

💡 Note

Any reaction that produces an insoluble substance (precipitate) is called a precipitation reaction.

(ii) Acid–Base Neutralization Reaction

An acid reacts with a base to form salt and water. This reaction is called an acid–base neutralization reaction.

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Here, hydrochloric acid reacts with sodium hydroxide to produce sodium chloride and water.

(iii) Evolution of Gas

Some double displacement reactions also produce a gas.

Examples:

ZnS + 2HCl → ZnCl₂ + H₂S ↑

In this reaction, hydrogen sulphide (H₂S) gas is evolved.

Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂ ↑

In this reaction, carbon dioxide (CO₂) gas is evolved.

💡 Note

All precipitation reactions are double displacement reactions, but all double displacement reactions are not precipitation reactions.

Try It Yourself

Questions:

1. Sodium chlorate is a colourless solid. It decomposes on heating to form sodium chloride and oxygen gas. Balance the chemical equation.

2. Classify the following reactions as decomposition, combination, single-displacement or double-displacement reactions.

(a) Na₂S(aq) + 2HCl(aq) → H₂S(g) + 2NaCl(aq)

(b) H₂(g) + CuO(s) → Cu(s) + H₂O(g)

(c) 2Fe(s) + 3Cl₂(g) → 2FeCl₃(s)

Answers:

1.

2NaClO₃ Δ ⟶ 2NaCl + 3O₂

2.

(a) Double-displacement reaction with gas evolution.

(b) Single-displacement reaction / Redox reaction.

(c) Combination reaction.

5. Oxidation and Reduction

Oxidation

Oxidation is a chemical reaction in which there is:

• Gain of oxygen, or
• Gain of any electronegative atom, or
• Loss of hydrogen, or
• Loss of any electropositive atom.

(i) Gain of Oxygen Atom

Examples:

S(s) + O₂(g) → SO₂(g)

Here, sulphur gains oxygen. So, sulphur is oxidised.

2SO₂(g) + O₂(g) → 2SO₃(g)

Here, sulphur dioxide gains oxygen. So, it is oxidised to sulphur trioxide.

(ii) Gain of Electronegative Atom

Example:

Mg(s) + Cl₂(g) → MgCl₂(s)

Here, magnesium gains chlorine, which is an electronegative element. So, magnesium is oxidised.

(iii) Removal of Hydrogen Atom

Example:

2HCl(g) → Cl₂(g) + H₂(g)

Here, hydrogen is removed from hydrogen chloride. So, hydrogen chloride is oxidised.

(iv) Loss of Electropositive Element

Example:

2KI(aq) + H₂O₂(aq) → 2KOH(aq) + I₂(s)

Here, potassium iodide loses potassium, an electropositive element. So, iodide is oxidised to iodine.

Reduction

Reduction is a chemical reaction in which there is:

• Gain of hydrogen, or
• Gain of any electropositive atom, or
• Loss of oxygen, or
• Loss of any electronegative atom.

(i) Gain of Hydrogen

Example:

Cl₂(g) + H₂S(g) → 2HCl(g) + S(s)

Here, chlorine gains hydrogen. So, chlorine is reduced.

(ii) Gain of Electropositive Element

Example:

2Na(s) + Cl₂(g) → 2NaCl(s)

Here, chlorine gains sodium, an electropositive element. So, chlorine is reduced.

(iii) Removal of Oxygen Atom

Example:

ZnO(s) + C(s) → Zn(s) + CO(g)

Here, zinc oxide loses oxygen. So, zinc oxide is reduced to zinc.

(iv) Loss of Electronegative Element

Example:

FeS(s) → Fe(s) + S(s)

Here, iron sulphide loses sulphur, an electronegative element. So, iron sulphide is reduced.

Redox Reactions

Reactions in which oxidation and reduction take place at the same time are called redox reactions.

CuO(s) + H₂(g) → Cu(s) + H₂O(g)

When hydrogen gas is passed over hot cupric oxide, hydrogen is oxidised to water.

At the same time, cupric oxide is reduced to metallic copper by losing oxygen.

In this reaction, H₂ helps in reducing cupric oxide to copper. So, H₂ acts as a reducing agent.

CuO helps in oxidising hydrogen to water. So, CuO acts as an oxidising agent.

💡 Important Terms

• The substance which itself gets oxidised acts as a reducing agent or reductant.
• The substance which itself gets reduced acts as an oxidising agent or oxidant.

Effects of Oxidation Reactions in Life

1. Corrosion

Corrosion is the slow deterioration of metals when they are exposed to air and moisture for a long time.

During corrosion, metals are converted into compounds such as oxides or sulphides.

Examples:

• A black coating forms on silver.
• A green coating forms on copper.
• Corrosion of iron is called rusting.

Rust is a brown coloured substance formed due to the action of moist air (oxygen and water) on iron.

The chemical formula of rust is Fe₂O₃.xH₂O.

Rusting is a slow process.

2Fe(s) + 3O₂(g) + xH₂O(l) → Fe₂O₃.xH₂O(s)

Both corrosion and rusting damage buildings, railway tracks, vehicles, bridges and many other metal objects.

Examples of Corrosion

1. Corrosion of Copper

When copper is exposed to air and moisture, it develops a green coating of basic copper carbonate.

Cu + H₂O + CO₂ + O₂ → CuCO₃.Cu(OH)₂

2. Corrosion of Silver

Silver reacts with hydrogen sulphide present in air to form a black coating of silver sulphide.

4Ag + 2H₂S + O₂ → 2Ag₂S + 2H₂O

Prevention of Corrosion

Corrosion can be prevented by the following methods:

Prevention of Corrosion

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🎨 Painting

🛢️ Painting protects the objects from oxygen and moisture. It is generally done to objects placed outside.

🛢️ Oiling & Greasing

Metal parts should be oiled and greased regularly. It is used for motion machines and mechanical tools

⚡ Electroplating

A thin layer of metal resistant to corrosion is layered around metals like iron to prevent it from corrosion. For example, chromeplating (using Chromium (Cr) metal) and tin plating (using Tin (Sn) metal)

🛡️ Galvanisation

A thin layer of zinc metal is attached around iron body to prevent it from corrosion.

📘 Note

These methods will be explained in detail in Chapter 3 – Metals and Non-metals.

2. Rancidity

Rancidity is the process in which fats and oils get oxidised when exposed to air for a long time.

As a result, the taste and smell of food change and become unpleasant.

Prevention of Rancidity

• Manufacturers add antioxidants such as BHA and BHT to food to prevent oxidation.
• Food should be stored in air-tight containers.
• Food should be kept in a refrigerator because low temperature slows down oxidation.
• Packets of chips are filled with nitrogen gas. Nitrogen is an unreactive gas, so it prevents oxidation and keeps the chips fresh.