NCERT Class 10 Science Textbook made Screen Readable by Raunak Padhi from USA and Dr T K Bansal.

We have learnt in Class 9 that during a chemical reaction, atoms of one element do not change into those of another element. 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. You will study about types of bonds formed between atoms in Chapters 3 and 4.

1.2.1 - Combination Reaction

Activity 1.4:

● Take a small amount of calcium oxide or quicklime in a beaker.
● Slowly add water to this.
● Touch the beaker as shown in Figure 1.3.
● Do you feel any change in temperature?

Figure 1.3: Formation of slaked lime through the reaction of calcium oxide with water

Calcium oxide reacts vigorously with water to produce slaked lime, or calcium hydroxide, releasing a large amount of heat. The chemical equation for this is:

$$CaO (s) + H_2O (l) → Ca(OH)_2 (aq) + ΔH .. .. .. (1.13)$$

Where, CaO is quick lime, Ca(OH)₂ is slaked lime, and ΔH is the heat evolved.

In this reaction, calcium oxide and water combine to form a single product, calcium hydroxide. Such a reaction in which a single product is formed from two or more reactants is known as a combination reaction.

Do you know?

A solution of slaked lime produced by the reaction 1.13 is used for whitewashing walls. Calcium hydroxide reacts slowly with the carbon dioxide in air to form a thin layer of calcium carbonate on the walls. 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

$$Ca(OH)_2 (aq) + CO_2 (g) → CaCO_3 (s) + H_2O (l) .. .. ..(1.14)$$

Where, Ca(OH)₂ is calcium hydroxide, CaCO₃ is calcium carbonate.

Let us discuss some more examples of combination reactions.

(i) Burning of coal

$$C (s) + O_2 (g) → CO_2 (g) .. .. .. (1.15)$$

(ii) Formation of water from H₂ (g) and O₂ (g)

$$2 H_2 (g) + O_2 (g) → 2 H_2O (l) .. .. .. (1.16)$$

In simple language we can say that when two or more substances (elements or compounds) combine to form a single product, the reactions are called combination reactions.

In Activity 1.4, we observed that a large amount of heat is also evolved. This makes the reaction mixture warm. Reactions in which heat is released along with the formation of products are called exothermic chemical reactions.

Other examples of exothermic reactions are:

(i) Burning of natural gas

$$CH_4 (g) + 2 O_2 (g) → CO_2 (g) + 2 H_2O (g) .. .. .. (1.17)$$

(ii) Do you know that respiration is also an exothermic process?

We all know that we need energy to stay alive. We get this energy from the food we eat. During digestion, food is broken down into simpler substances. For example, rice, potatoes and bread contain carbohydrates. These carbohydrates are broken down to form glucose. This glucose combines with oxygen in the cells of our body and provides energy. The special name of this reaction is respiration, the process of which you will study in Chapter 6.

$$C_6H_{12}O_6 (aq) + 6O_2 (aq) → 6CO_2 (aq) + 6H_2O (l) + ΔH .. .. .. (1.18)$$

Where, C₆H₁₂O6 is glucose.

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

Identify the type of the reaction taking place in Activity 1.1, where heat is given out along with the formation of a single product.

1.2.2 Decomposition Reaction

Activity 1.5

● Take about 2 gram ferrous sulphate crystals in a dry boiling tube.
● Note the colour of the ferrous sulphate crystals.
● Heat the boiling tube over the flame of a burner or spirit lamp as shown in Fig. 1.4.
● Observe the colour of the crystals after heating.

Figure 1.4 Correct way of heating the boiling tube containing crystals of ferrous sulphate and of smelling the odour

Did you noticed that the green colour of the ferrous sulphate crystals has changed? You can also smell the characteristic odour of burning sulphur.

ferrous sulphate + Heat gives Ferric Oxide + sulphur dioxide + sulphur trioxide

Or in equation form it can be written as:

$$2 FeSO_4 (s) → Fe_2O_3 (s) + SO_2 (g) + SO_3 (g) .. .. .. .. (1.19)$$

(FeSO₄) is ferrous sulphate, (Fe₂O₃) is ferric oxide.

In this reaction you can observe that a single reactant breaks down to give simpler products. This is a decomposition reaction. Ferrous sulphate crystals (FeSO₄.7H₂O) lose water when heated and the colour of the crystals changes. It then decomposes to ferric oxide (Fe₂O₃), sulphur dioxide (SO₂) and sulphur trioxide (SO₃). Ferric oxide is a solid, while SO₂ and SO₃ are gases.

Decomposition of calcium carbonate to calcium oxide and carbon dioxide on heating is an important decomposition reaction used in various industries. Calcium oxide is called lime or quicklime. It has many uses - one is in the manufacture of cement. When a decomposition reaction is carried out by heating, it is called thermal decomposition.

calcium carbonate + heat (gives) calcium oxide + carbon dioxide

Writing in the equation form yields:

$$CaCO_3 (s) + ΔH → CaO (s) + CO_2 (g) .. .. .. (1.20)$$

Where, ΔH represents heat.

Another example of a thermal decomposition reaction is given in Activity 1.6.

Activity 1.6.

● Take about 2 gram lead nitrate powder in a boiling tube.
● Hold the boiling tube with a pair of tongs and heat it over a flame, as shown in Fig. 1.5.
● What do you observe? Note down the change, if any.

Figure 1.5 Heating of lead nitrate and emission of nitrogen dioxide

You will observe the emission of brown fumes. These fumes are of nitrogen dioxide (NO2). The reaction that takes place is:

$$2Pb(NO_3)_2 (s) + ΔH → 2PbO (s) + 4NO_2 (g) + O_2 (g) .. .. .. (1.21)$$

Pb(NO₃)₂ is lead nitrate, PbO is lead oxide, NO₂ is nitrogen dioxide

Let us perform some more decomposition reactions as given in Activities 1.7 and 1.8.

Activity 1.7

● Take a plastic mug. Drill two holes at its base and fit rubber stoppers in these holes. Insert carbon electrodes in these rubber stoppers as shown in Fig. 1.6.
● Connect these electrodes to a 6 volt battery.
● Fill the mug with water such that the electrodes are immersed. Add a few drops of dilute sulphuric acid to the water.
● Take two test tubes filled with water and invert them over the two carbon electrodes.
● Switch on the current and leave the apparatus undisturbed for some time.
● You will observe the formation of bubbles at both the electrodes. These bubbles displace water in the test tubes.
● Is the volume of the gas collected the same in both the test tubes?
● Once the test tubes are filled with the respective gases, remove them carefully.
● Test these gases one by one by bringing a burning candle close to the mouth of the test tubes.
● CAUTION: This step must be performed carefully by the teacher.
● What happens in each case?
● Which gas is present in each test tube?

Figure 1.6 Electrolysis of water

Activity 1.8.

● Take about 2 grams of silver chloride in a china dish.
● What is its colour? (it is white in color)
● Place this china dish in sunlight for some time (Fig. 1.7).
● Observe the colour of the silver chloride after some time. (It has turned grey.)

Figure 1.7 Silver chloride turns grey in sunlight to form silver metal

You will see that white silver chloride turns grey in sunlight. This is due to the decomposition of silver chloride into silver and chlorine by the action of light.

$$2 AgCl (s) + hν → 2Ag (s) + Cl_2 (g) .. .. .. (1.22)$$

Where, hν represents sunlight.

Silver bromide also behaves in the same way.

$$2AgBr (s) + hν → 2Ag (s) + Br_2 (g) .. .. .. (1.23)$$

The above reactions are used in black and white photography.

What form of energy is causing these decomposition reactions?

We have seen that the decomposition reactions require energy either in the form of heat, light or electricity for breaking down the reactants. Reactions in which energy is absorbed are known as endothermic reactions.

Carry out the following Activity:

Take about 2 grams of barium hydroxide in a test tube. Add 1 gram of ammonium chloride and mix with the help of a glass rod. Touch the bottom of the test tube with your palm. What do you feel? Is this an exothermic or endothermic reaction?

Questions:

Q1. A solution of a substance ‘X’ is used for whitewashing.
a. Name the substance ‘X’ and write its formula.
b. Write the reaction of the substance ‘X’ named in part a above with water.

Q2. Why is the amount of gas collected in one of the test tubes in Activity 1.7 double of the amount collected in the other? Name this gas.

1.2.3 Displacement Reaction

Activity 1.9

● Take three iron nails and clean them by rubbing with a piece of sandpaper.
● Take two test tubes marked as (A) and (B). In each test tube, take about 10 mL copper sulphate solution.
● Tie two iron nails with a thread and immerse them carefully in the copper sulphate solution in test tube B for about 20 minutes [Fig. 1.8 (a)]. Keep one iron nail aside for comparison.
● After 20 minutes, take out the iron nails from the copper sulphate solution.
● Compare the intensity of the blue colour of copper sulphate solutions in test tubes
● (A) and (B) [Fig. 1.8 (b)].
● Also, compare the colour of the iron nails dipped in the copper sulphate solution
● with the one kept aside [Fig. 1.8 (b)].

Figure 1.8 (a) Iron nails dipped in copper sulphate solution

Figure 1.8(b) Iron nails and copper sulphate solutions compared before and after the experiment

In the second part we compared the change in intensity of colour in test tube A and test tube B. The colour of solution in test tube B changes from blue to pale green.

Why does the iron nail become brownish in colour and the blue colour of copper sulphate solution fades?

The following chemical reaction takes place in this Activity:

$$Fe (s) + CuSO_4 (aq) → FeSO_4 (aq) + Cu (s) .. .. .. (1.24)$$

CuSO₄ is copper sulphate, FeSO₄ is iron sulphate

In this reaction, iron has displaced or removed another element, copper, from copper sulphate solution. This reaction is known as displacement reaction.

Other examples of displacement reactions are:

$$Zn (s) + CuSO_4 (aq) → ZnSO_4 (aq) + Cu (s) .. .. .. (1.25)$$

CuSO₄ is copper sulphate, ZnSO₄ is zinc sulphate

$$Pb (s) + CuCl_2 (aq) → PbCl_2 (aq) + Cu (s) .. .. .. (1.26)$$

Where, CuCl₂ is copper chloride, PbCl₂ is lead chloride

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

1.2.4 Double Displacement Reaction

Activity 1.10

● Take about 3 mL of sodium sulphate solution in a test tube.
● In another test tube, take about 3 mL of barium chloride solution.
● Mix the two solutions (Fig. 1.9).
● What do you observe?

Figure 1.9 Formation of barium sulphate and sodium chloride

$$Na_2SO_4 (aq) + BaCl_2 (aq) → BaSO_4 (s) + 2NaCl (aq) .. .. .. (1.27)$$

Where, Na₂SO₄ is sodium sulphate, BaCl₂ is barium chloride, BaSO₄ is barium sulphate, NaCl is sodium chloride

What causes this? The white precipitate of BaSO₄ is formed by the reaction of

(SO₄), which has a charge of negative 2, and Ba, which has a charge of positive 2. The other product formed is sodium chloride which remains in the solution. Such reactions in which there is an exchange of ions between the reactants are called double displacement reactions.

Recall Activity 1.2, where you have mixed the solutions of lead (II) nitrate and potassium iodide.
(a) What was the colour of the precipitate formed? Can you name the compound precipitated?
(b) Write the balanced chemical equation for this reaction.
(c) Is this also a double displacement reaction?

1.2.5 Oxidation and Reduction

Activity 1.11

● Heat a china dish containing about 1 gram of copper powder (Figure 1.10).
● What do you observe?

Figure 1.10 Oxidation of copper to copper oxide

 

The surface of copper powder becomes coated with black copper(II) oxide. Why has this black substance formed?

This is because oxygen is added to the copper and copper oxide is formed.

$$2 Cu + O_2 + ΔH → 2 CuO .. .. .. (1.28)$$

If hydrogen gas is passed over this heated material (CuO), the black coating on the surface turns brown as the reverse reaction takes place and copper is obtained.

$$CuO + H_2 + ΔH → Cu + H_2O .. .. .. (1.29)$$

If a substance gains oxygen during a reaction, it is said to be oxidised.

If a substance loses oxygen during a reaction, it is said to be reduced.

During this reaction (1.29), the copper(II) oxide is losing oxygen and is being reduced. The hydrogen is gaining oxygen and is being oxidised. In other words, one reactant gets oxidised while the other gets reduced during a reaction. Such reactions are called oxidation-reduction reactions or redox reactions.

Some other examples of redox reactions are:

$$ZnO + C → Zn + CO (1.31)$$

$$MnO_2 + 4 HCl → MnCl_2 + H_2O + Cl_2 .. .. .. (1.32)$$

In reaction (1.31) carbon is oxidised to CO and ZnO is reduced to Zn.

In reaction (1.32) HCl is oxidised to Cl₂ whereas MnO₂ is reduced to MnCl₂.

From the above examples we can say that if a substance gains oxygen or loses hydrogen during a reaction, it is said to be oxidised. If a substance loses oxygen or gains hydrogen during a reaction, it is said to be reduced.

Recall Activity 1.1, where a magnesium ribbon burns with a dazzling flame in air (oxygen) and changes into a white substance, magnesium oxide. Is magnesium being oxidised or reduced in this reaction?