Introduction to Chemical Reactions

In our daily lives, we see many changes happening around us.

For instance, milk turns sour when left out in the summer. An iron pan or nail develops rust in humid air. Grapes get fermented, food is cooked, and our bodies digest the food we eat.

Even the process of respiration is a change. In all these examples, the original substance changes its nature and identity.

We have learned about physical and chemical changes before. Whenever a chemical change happens, we can say that a chemical reaction has taken place.

But what exactly is a chemical reaction? And how can we tell if one has occurred?

To find the answers to these questions, let us perform some simple activities.

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Determining a Chemical Reaction

There are several key observations that help us determine if a chemical reaction has happened.

• Change in State: The physical state of substances might change, like a solid turning into a liquid or gas.
• Change in Colour: A visible change in colour is a common sign of a chemical reaction.
• Evolution of a Gas: The formation of bubbles or fumes shows that a gas is being produced.
• Change in Temperature: The reaction mixture can get hotter or colder, which means energy is being released or absorbed.

Activity 1: Burning of Magnesium Ribbon

First, a magnesium ribbon (about 3-4 cm long) is cleaned by rubbing it with sandpaper.

Using a pair of tongs, the ribbon is then burnt over a spirit lamp or burner. The ash that forms is collected in a watch-glass.

You will observe that the magnesium ribbon burns with a dazzling white flame.

It also changes into a white powder. This new substance is magnesium oxide.

It is formed because of the reaction between magnesium and the oxygen present in the air.

Observation: This activity clearly shows a change in state (from a solid metal to a powder) and a change in temperature (heat and light are produced).

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Activity 2: Reaction between Lead Nitrate and Potassium Iodide

In another activity, we take lead nitrate solution in a test tube.

To this, we add potassium iodide solution. A distinct change is observed immediately.

Observation: A yellow, insoluble substance is formed. This is known as a precipitate.

This result demonstrates a change in colour and a change in state, as two clear liquids react to form a yellow solid.

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Activity 3: Reaction between Zinc and Dilute Acid

For this activity, a few zinc granules are placed in a conical flask or a test tube.

Dilute hydrochloric acid or sulphuric acid is then carefully added.

You will notice something happening around the zinc granules. If you touch the flask, you will also feel a change.

Observations: Bubbles start to form around the zinc. This indicates the evolution of a gas, which is hydrogen gas.

The flask also becomes warm to the touch. This shows a change in temperature, as the reaction releases heat.

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Activity 4: Reaction of Quicklime with Water

Take a small amount of calcium oxide, also known as quicklime, in a beaker.

Slowly and carefully add water to the beaker.

Touch the outside of the beaker to feel for any change.

Observations: You will hear a hissing sound as the water is added.

The beaker will become very hot, indicating a significant change in temperature.

The solid quicklime reacts with water to form a milky suspension, showing a change in the substance's appearance.

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Activity 5: Reaction of Baking Soda with Vinegar

Take a spoonful of baking soda (sodium bicarbonate) and place it in a glass or a beaker.

Pour some vinegar (acetic acid) onto the baking soda.

Observe the reaction that takes place immediately.

Observations: The mixture will fizz and bubble vigorously. This is a clear sign of the evolution of a gas (carbon dioxide).

If you touch the beaker, you may notice it feels slightly cooler, indicating a change in temperature (this reaction absorbs a small amount of heat).

Writing Chemical Equations

Describing a chemical reaction in a sentence is often long and cumbersome. To represent chemical reactions in a more concise, informative, and efficient way, we use chemical equations.

Word Equations

The simplest method for describing a chemical reaction is to write it as a word equation.

In a word equation, the names of the substances that react are written on the left-hand side (LHS). These substances are called the reactants.

The names of the new substances formed are written on the right-hand side (RHS). These are called the products.

An arrow ($\rightarrow$) is placed between the reactants and products. The arrowhead points towards the products and indicates the direction of the reaction.

If there is more than one reactant or product, they are linked with a plus sign (+).

For the burning of magnesium ribbon (Activity 1), the word equation is:

$$\underset{\text{(Reactants)}}{\text{Magnesium + Oxygen}} \rightarrow \underset{\text{(Product)}}{\text{Magnesium Oxide}}$$

Similarly, for the reaction of zinc with acid (Activity 3), the word equation would be:

Zinc + Sulphuric acid $\rightarrow$ Zinc sulphate + Hydrogen

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Skeletal Chemical Equations

While word equations are a step up from sentences, they can be made even more concise and useful. This is done by using chemical formulae instead of words.

A chemical equation that uses symbols and formulae to represent a chemical reaction is the standard method used in chemistry.

Using formulae, the word equation for burning magnesium becomes:

$Mg + O_2 \rightarrow MgO$

This type of initial, unbalanced equation is known as a skeletal chemical equation.

It is called 'skeletal' or 'unbalanced' because the mass is not the same on both sides of the equation. If you count the atoms of each element, you will find they are not equal on the LHS and RHS.

For example, in the equation $Mg + O_2 \rightarrow MgO$:

• On the LHS (Reactants): There is 1 atom of Magnesium (Mg) and 2 atoms of Oxygen (O).
• On the RHS (Product): There is 1 atom of Magnesium (Mg) but only 1 atom of Oxygen (O).

Since the number of oxygen atoms is not the same on both sides, the equation is unbalanced. Such an equation violates the Law of Conservation of Mass, which is a fundamental principle in chemistry.

Balanced Chemical Equations

Recall the Law of Conservation of Mass that you studied in Class IX. It states that mass can neither be created nor destroyed in a chemical reaction.

This means that 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.

For this to be true, the number of atoms of each element must remain the same before and after a chemical reaction. Hence, we need to balance skeletal chemical equations to make them consistent with this law. A balanced equation is a true representation of a chemical reaction.

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Example of an Already Balanced Equation

Let's consider the reaction from Activity 1.3:

Word Equation: Zinc + Sulphuric acid $\rightarrow$ Zinc sulphate + Hydrogen

The skeletal equation is: $Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2$

Let us examine the number of atoms of different elements on both sides of the arrow.

Element	Number of atoms in reactants (LHS)	Number of atoms in products (RHS)
Zn	1	1
H	2	2
S	1	1
O	4	4

As the number of atoms of each element is the same on both sides of the arrow, this is a balanced chemical equation.

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Steps to Balance a Chemical Equation (Hit-and-Trial Method)

Let us learn to balance an unbalanced equation step-by-step. We will use the reaction of iron with steam as an example:

$Fe + H_2O \rightarrow Fe_3O_4 + H_2$

Step I: Draw boxes around each formula.

First, draw boxes around each formula. It is important not to change anything inside the boxes while balancing. Changing the formula (e.g., changing $H_2O$ to $H_2O_2$) would mean changing the substance itself.

$\boxed{Fe} + \boxed{H_2O} \rightarrow \boxed{Fe_3O_4} + \boxed{H_2}$

Step II: List the number of atoms of different elements.

List the number of atoms for each element present in the unbalanced equation.

Element	Number of atoms in reactants (LHS)	Number of atoms in products (RHS)
Fe	1	3
H	2	2
O	1	4

Step III: Start balancing with the most complex compound.

It is often convenient to start balancing with the compound that contains the maximum number of atoms. In that compound, select the element which has the maximum number of atoms.

In our example, we select $Fe_3O_4$. The element oxygen has 4 atoms on the RHS and only 1 on the LHS. To balance oxygen atoms, we place a coefficient '4' in front of $H_2O$ on the LHS.

The partly balanced equation now becomes:

$\boxed{Fe} + 4\boxed{H_2O} \rightarrow \boxed{Fe_3O_4} + \boxed{H_2}$

Step IV: Balance the other atoms.

Now, Fe and H atoms are still not balanced. Let us balance hydrogen atoms. The number of H atoms on the LHS is now $4 \times 2 = 8$. On the RHS, there are only 2. To balance, we place a coefficient '4' in front of $H_2$ on the RHS.

The equation becomes:

$\boxed{Fe} + 4\boxed{H_2O} \rightarrow \boxed{Fe_3O_4} + 4\boxed{H_2}$

Step V: Balance the last remaining element.

Examine the equation again and pick the element which is still unbalanced. Here, it is iron (Fe). There is 1 Fe atom on the LHS and 3 on the RHS. To balance, we place a coefficient '3' in front of $Fe$ on the LHS.

$3\boxed{Fe} + 4\boxed{H_2O} \rightarrow \boxed{Fe_3O_4} + 4\boxed{H_2}$

Step VI: Finally, check the correctness of the balanced equation.

We count the atoms of each element on both sides of the final equation to verify the balancing.

$3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$

• Iron (Fe): 3 atoms on LHS $\rightarrow$ 3 atoms on RHS. (Balanced)
• Hydrogen (H): $4 \times 2 = 8$ atoms on LHS $\rightarrow$ $4 \times 2 = 8$ atoms on RHS. (Balanced)
• Oxygen (O): $4 \times 1 = 4$ atoms on LHS $\rightarrow$ 4 atoms on RHS. (Balanced)

The numbers of atoms are equal on both sides. The equation is now balanced.

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More Examples of Balancing Equations

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Step VII: Writing Symbols of Physical States and Reaction Conditions

To make a chemical equation fully informative, we add physical states and reaction conditions.

The physical states are: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution.

The balanced equation for iron and steam becomes:

$3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$

Reaction conditions like temperature, pressure, or a catalyst are indicated above or below the arrow.

Examples of Reaction Conditions

1. Photosynthesis:

This process requires sunlight and the catalyst chlorophyll.

$6CO_2(aq) + 6H_2O(l) \xrightarrow[Chlorophyll]{Sunlight} C_6H_{12}O_6(aq) + 6O_2(g)$

2. Thermal Decomposition of Calcium Carbonate:

This reaction requires a high temperature to proceed. The delta symbol ($\Delta$) signifies heat.

$CaCO_3(s) \xrightarrow{\Delta} CaO(s) + CO_2(g)$

3. The Haber Process for Ammonia Production:

This industrial process requires specific conditions of temperature, pressure, and a catalyst.

$N_2(g) + 3H_2(g) \xrightarrow[200 \ atm, \ Fe \ catalyst]{450^\circ C} 2NH_3(g)$

4. Decomposition of Potassium Chlorate:

This reaction is sped up by a catalyst ($MnO_2$) and requires heat.

$2KClO_3(s) \xrightarrow[MnO_2]{\Delta} 2KCl(s) + 3O_2(g)$

5. Electrolysis of Water:

This reaction requires electrical energy to break down water.

$2H_2O(l) \xrightarrow{Electricity} 2H_2(g) + O_2(g)$

Types of Chemical Reactions

In Class IX, we learned that during a chemical reaction, atoms of one element do not change into those of another. Nor do atoms disappear from the mixture or appear from elsewhere. Actually, chemical reactions involve the breaking and making of bonds between atoms to produce new substances. Reactions are classified into different types based on the nature of this chemical change.

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1. Combination Reaction

A reaction in which a single product is formed from two or more reactants is known as a combination reaction. The reactants can be elements or compounds.

The general form of a combination reaction is: $A + B \rightarrow AB$

Activity 4: Reaction of Calcium Oxide with Water

Take a small amount of calcium oxide (quick lime) in a beaker. Slowly add water to this. When you touch the beaker, you will feel that it has become hot.

In this reaction, calcium oxide and water combine to form a single product, calcium hydroxide (slaked lime), releasing a large amount of heat.

$$\underset{\text{(Quick lime)}}{CaO(s)} + \underset{\text{(Water)}}{H_2O(l)} \rightarrow \underset{\text{(Slaked lime)}}{Ca(OH)_2(aq)} + Heat$$

Reactions in which heat is released along with the formation of products are called exothermic chemical reactions.

More Examples of Combination Reactions:

(i) Burning of coal: Carbon (coal) combines with oxygen to form a single product, carbon dioxide.

$C(s) + O_2(g) \rightarrow CO_2(g)$

(ii) Formation of water: Hydrogen gas combines with oxygen gas to form a single product, water.

$2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$

More Examples of Exothermic Reactions:

(i) Burning of natural gas (methane):

$CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g) + Heat$

(ii) Respiration: During digestion, carbohydrates (from rice, potatoes, bread) are broken down into simpler substances like glucose. This glucose combines with oxygen in our cells and provides energy. This special reaction is called respiration.

$$\underset{\text{(Glucose)}}{C_6H_{12}O_6(aq)} + 6O_2(aq) \rightarrow 6CO_2(aq) + 6H_2O(l) + Energy$$

(iii) Decomposition of vegetable matter into compost is also an example of an exothermic reaction.

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2. Decomposition Reaction

A reaction in which a single reactant breaks down to give simpler products is known as a decomposition reaction. It is the opposite of a combination reaction.

The general form is: $AB \rightarrow A + B$

Decomposition reactions require energy in the form of heat, light, or electricity for breaking down the reactants. Reactions that absorb energy are known as endothermic reactions.

Thermal Decomposition (using heat)

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

(i) Decomposition of Ferrous Sulphate (Activity 1.5): When green crystals of ferrous sulphate ($FeSO_4 \cdot 7H_2O$) are heated, they first lose water. Upon further heating, the anhydrous ferrous sulphate decomposes to form ferric oxide (a brown solid), sulphur dioxide, and sulphur trioxide gas, which have a characteristic odour of burning sulphur.

$$\underset{\text{(Ferrous sulphate)}}{2FeSO_4(s)} \xrightarrow{Heat} \underset{\text{(Ferric oxide)}}{Fe_2O_3(s)} + SO_2(g) + SO_3(g)$$

(ii) Decomposition of Lead Nitrate (Activity 1.6): When lead nitrate powder is heated in a boiling tube, you will observe the emission of brown fumes. These fumes are of nitrogen dioxide ($NO_2$).

$$\underset{\text{(Lead nitrate)}}{2Pb(NO_3)_2(s)} \xrightarrow{Heat} \underset{\text{(Lead oxide)}}{2PbO(s)} + \underset{\text{(Nitrogen dioxide)}}{4NO_2(g)} + \underset{\text{(Oxygen)}}{O_2(g)}$$

Electrolytic Decomposition (using electricity)

Electrolysis of Water (Activity 1.7): When an electric current is passed through acidulated water, it decomposes into hydrogen and oxygen gas. Hydrogen gas is collected at the negative electrode (cathode) and oxygen gas is collected at the positive electrode (anode). The volume of hydrogen produced is double the volume of oxygen.

$$2H_2O(l) \xrightarrow{Electricity} 2H_2(g) + O_2(g)$$

Photolytic Decomposition (using light)

Decomposition of Silver Chloride (Activity 1.8): Take some white silver chloride in a china dish and place it in sunlight. You will see that the white silver chloride turns grey. This is because light causes it to decompose into silver (grey) and chlorine gas.

$$2AgCl(s) \xrightarrow{Sunlight} 2Ag(s) + Cl_2(g)$$

Silver bromide also behaves in the same way. These reactions are used in black and white photography.

$$2AgBr(s) \xrightarrow{Sunlight} 2Ag(s) + Br_2(g)$$

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3. Displacement Reaction

A reaction in which a more reactive element displaces or removes another, less reactive element from its salt solution. This is based on the reactivity series of metals.

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

Activity 1.9: Reaction of Iron with Copper Sulphate

When an iron nail is dipped in a blue copper sulphate solution, the following changes are observed after about 20 minutes: the iron nail becomes brownish in colour, and the blue colour of the copper sulphate solution fades.

This happens because iron is more reactive than copper. It displaces copper from the copper sulphate solution. The brown coating on the nail is copper metal, and the solution turns pale green due to the formation of iron sulphate.

$$\underset{\text{(Iron)}}{Fe(s)} + \underset{\text{(Copper sulphate)}}{CuSO_4(aq)} \rightarrow \underset{\text{(Iron sulphate)}}{FeSO_4(aq)} + \underset{\text{(Copper)}}{Cu(s)}$$

More Examples of Displacement Reactions:

Zinc and lead are also more reactive than copper and will displace it from its compounds.

$$\underset{\text{(Zinc)}}{Zn(s)} + \underset{\text{(Copper sulphate)}}{CuSO_4(aq)} \rightarrow \underset{\text{(Zinc sulphate)}}{ZnSO_4(aq)} + \underset{\text{(Copper)}}{Cu(s)}$$

$$\underset{\text{(Lead)}}{Pb(s)} + \underset{\text{(Copper chloride)}}{CuCl_2(aq)} \rightarrow \underset{\text{(Lead chloride)}}{PbCl_2(aq)} + \underset{\text{(Copper)}}{Cu(s)}$$

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4. Double Displacement Reaction

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

The general form is: $AB + CD \rightarrow AD + CB$

Activity 1.10: Reaction of Sodium Sulphate with Barium Chloride

When a solution of sodium sulphate is mixed with a solution of barium chloride, a white, insoluble substance is formed immediately. This insoluble substance is known as a precipitate. Any reaction that produces a precipitate can be called a precipitation reaction.

The white precipitate of barium sulphate ($BaSO_4$) is formed by the reaction of $Ba^{2+}$ ions and $SO_4^{2-}$ ions. The other product, sodium chloride, remains dissolved in the solution.

$$\underset{\text{(Sodium sulphate)}}{Na_2SO_4(aq)} + \underset{\text{(Barium chloride)}}{BaCl_2(aq)} \rightarrow \underset{\text{(Barium sulphate)}}{BaSO_4(s)} + \underset{\text{(Sodium chloride)}}{2NaCl(aq)}$$

Another Example of a Double Displacement Reaction:

When lead(II) nitrate solution is mixed with potassium iodide solution, a yellow precipitate of lead(II) iodide is formed.

$$\underset{\text{(Lead(II) nitrate)}}{Pb(NO_3)_2(aq)} + \underset{\text{(Potassium iodide)}}{2KI(aq)} \rightarrow \underset{\text{(Lead(II) iodide)}}{PbI_2(s)} + \underset{\text{(Potassium nitrate)}}{2KNO_3(aq)}$$

Oxidation-Reduction and its Effects in Daily Life

Oxidation and Reduction (Redox Reactions)

Oxidation and reduction are fundamental chemical processes that are often linked. They can be defined in several ways, but a simple definition involves the transfer of oxygen or hydrogen.

• Oxidation is a process which involves the gain of oxygen or the loss of hydrogen by a substance.
• Reduction is a process which involves the loss of oxygen or the gain of hydrogen by a substance.

Activity 1.11: Oxidation of Copper

If you heat a china dish containing about 1 g of copper powder, you will observe that the surface of the copper powder becomes coated with a black substance. Why has this black substance formed?

This is because oxygen from the air is added to copper, and copper(II) oxide is formed. In this reaction, copper has gained oxygen, so it is said to be oxidised.

$$\underset{\text{(Copper)}}{2Cu} + \underset{\text{(Oxygen)}}{O_2} \xrightarrow{Heat} \underset{\text{(Copper(II) oxide)}}{2CuO}$$

Now, if hydrogen gas is passed over this heated black material (CuO), the black coating on the surface turns brown. This is because the reverse reaction takes place, and copper is obtained again.

$$\underset{\text{(Copper(II) oxide)}}{CuO} + \underset{\text{(Hydrogen)}}{H_2} \xrightarrow{Heat} \underset{\text{(Copper)}}{Cu} + \underset{\text{(Water)}}{H_2O}$$

In this reaction (1.29), two things are happening simultaneously:

• The copper(II) oxide (CuO) is losing oxygen and is being reduced to copper (Cu).
• The hydrogen ($H_2$) is gaining oxygen and is being oxidised to water ($H_2O$).

In other words, one reactant gets oxidised while the other gets reduced during a reaction. Such reactions are called oxidation-reduction reactions or, more commonly, redox reactions.

More Examples of Redox Reactions:

(i) $$\underset{\text{(Zinc oxide)}}{ZnO} + \underset{\text{(Carbon)}}{C} \rightarrow \underset{\text{(Zinc)}}{Zn} + \underset{\text{(Carbon monoxide)}}{CO}$$

In this reaction, Carbon (C) is oxidised to Carbon monoxide (CO), and Zinc oxide (ZnO) is reduced to Zinc (Zn).

(ii) $$\underset{\text{(Manganese dioxide)}}{MnO_2} + \underset{\text{(Hydrochloric acid)}}{4HCl} \rightarrow \underset{\text{(Manganese chloride)}}{MnCl_2} + \underset{\text{(Water)}}{2H_2O} + \underset{\text{(Chlorine)}}{Cl_2}$$

In this reaction, HCl is oxidised to $Cl_2$ (as it loses hydrogen), whereas $MnO_2$ is reduced to $MnCl_2$ (as it loses oxygen).

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Effects of Oxidation in Everyday Life

Oxidation reactions have a significant impact on our daily lives, and not all of them are beneficial. Two common damaging effects are corrosion and rancidity.

1. Corrosion

You must have observed that iron articles are shiny when new, but get coated with a reddish-brown powder when left for some time. This process is commonly known as the rusting of iron.

Corrosion is the process in which a metal is attacked by substances around it such as moisture, air, acids, etc. The metal is said to to corrode, and it slowly deteriorates.

• The black coating on silver and the green coating on copper are other examples of corrosion.

Corrosion causes damage to car bodies, bridges, iron railings, ships, and all objects made of metals, especially those of iron. Corrosion of iron is a serious problem. Every year, an enormous amount of money is spent to replace damaged iron structures.

2. Rancidity

Have you ever tasted or smelt food containing fat or oil that has been left for a long time? When fats and oils are oxidised, they become rancid, and their smell and taste change, making them unsuitable for consumption.

This process of oxidation of fats and oils is called rancidity. There are several ways to slow down or prevent this process:

• Using Antioxidants: Usually, substances which prevent oxidation (antioxidants) are added to foods containing fats and oil. These substances slow down the oxidation process.
• Airtight Containers: Keeping food in airtight containers helps to slow down oxidation because it reduces the food's exposure to oxygen in the air.
• Flushing with Nitrogen: Do you know that chip manufacturers usually flush bags of chips with an inert gas such as nitrogen? This is done to prevent the chips from coming into contact with oxygen and getting oxidised.

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