Chemical Reactions And Equations

Hello my dear students! Welcome to today's science class. I am so happy to see all of you here. Today, we are going to start a very interesting and important chapter from your Science textbook. We will be studying Chapter 1: Chemical Reactions and Equations. This chapter is going to be extremely useful for you, not just for your exams but also for understanding the world around you. So, sit back, relax, and let me take you through this fascinating journey of chemical reactions.

Now, before we begin, let me ask you something. Have you ever thought about what happens in our daily lives? Think about these situations. What happens when milk is left at room temperature during summers? It becomes sour, isn't it? What happens when an iron tawa or a pan or a nail is left exposed to a humid atmosphere? It gets rusted. What happens when grapes get fermented? They turn into wine or alcohol. What happens when we cook food? The raw ingredients change into cooked food. And what happens when food gets digested in our body? It gets broken down into simpler substances. And what happens when we respire? We breathe in oxygen and breathe out carbon dioxide.

In all these situations, the nature and the identity of the initial substance have somewhat changed. We have already learnt about physical and chemical changes of matter in our previous classes. Whenever a chemical change occurs, we can say that a chemical reaction has taken place.

So, the big question is: what exactly is a chemical reaction? How do we come to know that a chemical reaction has taken place? Let us perform some activities together to find the answer to these questions.

Let's start with Activity 1.1. Now, this activity needs careful handling, so please pay attention. We are going to take a magnesium ribbon, about 3 to 4 centimeters long. First, we clean it by rubbing it with sandpaper to remove any oxide layer that might have formed on its surface. Then, we hold it with a pair of tongs and burn it using a spirit lamp or burner. We collect the ash formed in a watch-glass. Now, very important: we burn the magnesium ribbon keeping it away from our eyes as far as possible because the flame is very bright.

What do we observe? You must have observed that the magnesium ribbon burns with a dazzling white flame and changes into a white powder. This white powder is magnesium oxide. It is formed due to the reaction between magnesium and oxygen present in the air. This is a chemical reaction!

Now, let's look at Activity 1.2. We take some lead nitrate solution in a test tube. Then we add potassium iodide solution to this. What do we observe? We see a yellow precipitate forming. This is because lead nitrate reacts with potassium iodide to form lead iodide, which is a yellow solid.

Now, Activity 1.3. We take a few zinc granules in a conical flask or a test tube. We add dilute hydrochloric acid or sulphuric acid to this. Now, handle the acid with care! Do we observe anything happening around the zinc granules? Yes, we see bubbles forming. These bubbles are of hydrogen gas. And if we touch the conical flask or test tube, we feel that it has become warm. So, there is a change in temperature as well.

From these three activities, we can say that any of the following observations help us determine whether a chemical reaction has taken place. First, a change in state. For example, magnesium solid changes to magnesium oxide solid. Second, a change in colour. For example, the colour change when lead nitrate reacts with potassium iodide. Third, the evolution of a gas. For example, the bubbles of hydrogen gas when zinc reacts with acid. And fourth, a change in temperature. For example, the flask becoming warm when zinc reacts with acid.

As we observe the changes around us, we can see that there is a large variety of chemical reactions taking place around us. We will study about the various types of chemical reactions and their symbolic representation in this chapter.

Now, let's move on to section 1.1: Chemical Equations.

Activity 1.1 can be described as: when a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide. This description of a chemical reaction in a sentence form is quite long. It can be written in a shorter form. The simplest way to do this is to write it in the form of a word-equation.

The word-equation for the above reaction would be: Magnesium plus Oxygen arrow Magnesium oxide. The substances that undergo chemical change in the reaction, magnesium and oxygen, are the reactants. The new substance, magnesium oxide, formed during the reaction, is the product.

A word-equation shows change of reactants to products through an arrow placed between them. The reactants are written on the left-hand side with a plus sign between them. Similarly, products are written on the right-hand side with a plus sign between them. The arrowhead points towards the products and shows the direction of the reaction.

Now, is there any other shorter way for representing chemical equations? Yes, there is! Chemical equations can be made more concise and useful if we use chemical formulae instead of words. A chemical equation represents a chemical reaction.

If you recall the formulae of magnesium, oxygen, and magnesium oxide, the above word-equation can be written as: Mg plus O2 arrow MgO.

Now, let's count and compare the number of atoms of each element on the left-hand side and the right-hand side of the arrow. Is the number of atoms of each element the same on both sides? Let me check. On the left-hand side, we have one magnesium atom and two oxygen atoms. On the right-hand side, we have one magnesium atom and one oxygen atom. So, the number of oxygen atoms is not the same on both sides. Therefore, the equation is unbalanced because the mass is not the same on both sides of the equation. Such a chemical equation is called a skeletal chemical equation for a reaction. Equation Mg plus O2 arrow MgO is a skeletal chemical equation for the burning of magnesium in air.

Now, let's understand why we need to balance chemical equations. Recall the law of conservation of mass that you studied in Class IX. Mass can neither be created nor destroyed in a chemical reaction. That is, the total mass of the elements present in the products of a chemical reaction has to be equal to the total mass of the elements present in the reactants. In other words, the number of atoms of each element remains the same, before and after a chemical reaction. Hence, we need to balance a skeletal chemical equation.

Is the chemical equation Mg plus O2 arrow MgO balanced? No, it is not. Let us learn about balancing a chemical equation step by step.

But first, let me give you another example. The word-equation for Activity 1.3 may be represented as: Zinc plus Sulphuric acid arrow Zinc sulphate plus Hydrogen.

The above word-equation may be represented by the following chemical equation: Zn plus H2SO4 arrow ZnSO4 plus H2.

Let us examine the number of atoms of different elements on both sides of the arrow. We have Zn: 1 atom on left, 1 atom on right. H: 2 atoms on left, 2 atoms on right. S: 1 atom on left, 1 atom on right. O: 4 atoms on left, 4 atoms on right. As the number of atoms of each element is the same on both sides of the arrow, this equation is a balanced chemical equation.

Now, let's learn how to balance an unbalanced equation. Let us try to balance the following chemical equation: Fe plus H2O arrow Fe3O4 plus H2.

Step I: To balance a chemical equation, first draw boxes around each formula. Do not change anything inside the boxes while balancing the equation. So we have: Fe box plus H2O box arrow Fe3O4 box plus H2 box.

Step II: List the number of atoms of different elements present in the unbalanced equation. We have: Fe: 1 atom on left, 3 atoms on right. H: 2 atoms on left, 2 atoms on right. O: 1 atom on left, 4 atoms on right.

Step III: It is often convenient to start balancing with the compound that contains the maximum number of atoms. It may be a reactant or a product. In that compound, select the element which has the maximum number of atoms. Using these criteria, we select Fe3O4 and the element oxygen in it. There are four oxygen atoms on the right-hand side and only one on the left-hand side.

To balance the oxygen atoms, we need to make them equal. Initially, we have 1 oxygen atom in H2O on the left and 4 oxygen atoms in Fe3O4 on the right. To equalize, we multiply H2O by 4. So we write: 4 H2O. Now we have 4 oxygen atoms on both sides.

Step IV: Fe and H atoms are still not balanced. Pick any of these elements to proceed further. Let us balance hydrogen atoms in the partly balanced equation. We now have: Fe plus 4 H2O arrow Fe3O4 plus H2.

Initially, we have 8 hydrogen atoms in 4 H2O on the left and 2 hydrogen atoms in H2 on the right. To balance, we multiply H2 by 4. So we write: 4 H2. Now we have 8 hydrogen atoms on both sides.

Step V: Examine the above equation and pick up the third element which is not balanced. You find that only one element is left to be balanced, that is, iron. Initially, we have 1 iron atom on the left and 3 iron atoms in Fe3O4 on the right. To balance, we multiply Fe by 3. So we write: 3 Fe.

Step VI: Finally, to check the correctness of the balanced equation, we count atoms of each element on both sides of the equation. We have: 3 Fe plus 4 H2O arrow Fe3O4 plus 4 H2. Let's count: Fe: 3 on left, 3 on right. H: 8 on left, 8 on right. O: 4 on left, 4 on right. The numbers of atoms of elements on both sides are equal. This equation is now balanced. This method of balancing chemical equations is called the hit-and-trial method as we make trials to balance the equation by using the smallest whole number coefficient.

Step VII: Writing Symbols of Physical States. Carefully examine the above balanced equation. Does this equation tell us anything about the physical state of each reactant and product? No information has been given in this equation about their physical states.

To make a chemical equation more informative, the physical states of the reactants and products are mentioned along with their chemical formulae. The gaseous, liquid, aqueous, and solid states of reactants and products are represented by the notations (g), (l), (aq), and (s), respectively. The word aqueous (aq) is written if the reactant or product is present as a solution in water.

The balanced equation becomes: 3 Fe(s) plus 4 H2O(g) arrow Fe3O4(s) plus 4 H2(g).

Note that the symbol (g) is used with H2O to indicate that in this reaction water is used in the form of steam. Usually, physical states are not included in a chemical equation unless it is necessary to specify them.

Sometimes the reaction conditions, such as temperature, pressure, catalyst, etc., for the reaction are indicated above and/or below the arrow in the equation. For example, CO(g) plus 2 H2(g) arrow CH3OH(l), and the condition 340 atm is written above the arrow. Another example is: 6 CO2(aq) plus 12 H2O(l) arrow C6H12O6(aq) plus 6 O2(aq) plus 6 H2O(l), with sunlight and chlorophyll written above the arrow.

Now, let me ask you a question. Why should a magnesium ribbon be cleaned before burning in air? This is because magnesium ribbon often has a layer of magnesium oxide on its surface due to exposure to air. This layer prevents the magnesium from burning properly. So, we clean it with sandpaper to remove this layer.

Now, let's solve some questions. Question 2: Write the balanced equation for the following chemical reactions.

(i) Hydrogen plus Chlorine arrow Hydrogen chloride. The balanced equation is: H2 plus Cl2 arrow 2 HCl.

(ii) Barium chloride plus Aluminium sulphate arrow Barium sulphate plus Aluminium chloride. The balanced equation is: 3 BaCl2 plus Al2(SO4)3 arrow 3 BaSO4 plus 2 AlCl3.

(iii) Sodium plus Water arrow Sodium hydroxide plus Hydrogen. The balanced equation is: 2 Na plus 2 H2O arrow 2 NaOH plus H2.

Question 3: Write a balanced chemical equation with state symbols for the following reactions.

(i) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride. The balanced equation is: BaCl2(aq) plus Na2SO4(aq) arrow BaSO4(s) plus 2 NaCl(aq).

(ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water. The balanced equation is: NaOH(aq) plus HCl(aq) arrow NaCl(aq) plus H2O(l).

Now, let's move on to section 1.2: Types of Chemical Reactions.

We have learnt in Class IX 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.

Let's start with 1.2.1: Combination Reaction.

Let's perform Activity 1.4. Take a small amount of calcium oxide or quick lime in a beaker. Slowly add water to this. Touch the beaker. Do you feel any change in temperature? Yes, the beaker becomes warm. In fact, it becomes quite hot!

Calcium oxide reacts vigorously with water to produce slaked lime (calcium hydroxide) releasing a large amount of heat. The equation is: CaO(s) plus H2O(l) arrow Ca(OH)2(aq) plus Heat.

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 that a solution of slaked lime produced by this reaction 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. The equation is: Ca(OH)2(aq) plus CO2(g) arrow CaCO3(s) plus H2O(l).

Let us discuss some more examples of combination reactions.

(i) Burning of coal: C(s) plus O2(g) arrow CO2(g).

(ii) Formation of water from H2(g) and O2(g): 2 H2(g) plus O2(g) arrow 2 H2O(l).

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 also observed that a large amount of heat is 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: CH4(g) plus 2 O2(g) arrow CO2(g) plus 2 H2O(g).

(ii) Do you know that respiration is 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 equation is: C6H12O6(aq) plus 6 O2(aq) arrow 6 CO2(aq) plus 6 H2O(l) plus energy.

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

Now, let me ask you: 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. This is a combination reaction, and it is also exothermic because heat is given out.

Now, let's move on to 1.2.2: Decomposition Reaction.

Let's perform Activity 1.5. Take about 2 grams of ferrous sulphate crystals in a dry boiling tube. Note the colour of the ferrous sulphate crystals. It is green. Heat the boiling tube over the flame of a burner or spirit lamp. Observe the colour of the crystals after heating. Have you noticed that the green colour of the ferrous sulphate crystals has changed? It becomes reddish-brown. You can also smell the characteristic odour of burning sulphur.

The reaction is: 2 FeSO4(s) arrow Fe2O3(s) plus SO2(g) plus SO3(g).

In this reaction, you can observe that a single reactant breaks down to give simpler products. This is a decomposition reaction. Ferrous sulphate crystals 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 quick lime. 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. The equation is: CaCO3(s) arrow CaO(s) plus CO2(g).

Another example of a thermal decomposition reaction is given in Activity 1.6. Take about 2 grams of lead nitrate powder in a boiling tube. Hold the boiling tube with a pair of tongs and heat it over a flame. What do you observe? You will observe the emission of brown fumes. These fumes are of nitrogen dioxide (NO2). The reaction that takes place is: 2 Pb(NO3)2(s) arrow 2 PbO(s) plus 4 NO2(g) plus O2(g).

Let us perform some more decomposition reactions. Activity 1.7 involves electrolysis of water. We take a plastic mug, drill two holes at its base and fit rubber stoppers. Insert carbon electrodes in these rubber stoppers. 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? No, the volume of gas collected in one test tube is double the volume collected in the other. 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. What happens in each case? The gas in one test tube burns with a pop sound - this is hydrogen. The gas in the other test tube rekindles a glowing splinter - this is oxygen. So, water decomposes to form hydrogen and oxygen. The equation is: 2 H2O(l) arrow 2 H2(g) plus O2(g). Since twice as much hydrogen as oxygen is produced, the volume of hydrogen gas collected is double the volume of oxygen.

Now, Activity 1.8. Take about 2 grams of silver chloride in a china dish. What is its colour? It is white. Place this china dish in sunlight for some time. Observe the colour of the silver chloride after some time. 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 light. The equation is: 2 AgCl(s) arrow 2 Ag(s) plus Cl2(g). Silver bromide also behaves in the same way: 2 AgBr(s) arrow 2 Ag(s) plus Br2(g). The above reactions are used in black and white photography.

What form of energy is causing these decomposition reactions? In the case of ferrous sulphate, it is heat. In the case of lead nitrate, it is also heat. In the case of water, it is electricity. And in the case of silver chloride and silver bromide, it is light.

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.

Now, let me ask you: 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? You will feel that the test tube is cold. So, this is an endothermic reaction because heat is absorbed.

Question 1: A solution of a substance 'X' is used for whitewashing. (i) Name the substance 'X' and write its formula. The substance X is calcium hydroxide, also called slaked lime, with formula Ca(OH)2. (ii) Write the reaction of the substance 'X' named in (i) above with water. The reaction is: CaO(s) plus H2O(l) arrow Ca(OH)2(aq).

Question 2: 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. The gas is hydrogen. The reason is that water decomposes to form two molecules of hydrogen and one molecule of oxygen. So, the volume of hydrogen gas collected is double the volume of oxygen.

Now, let's move on to 1.2.3: Displacement Reaction.

Let's perform Activity 1.9. Take three iron nails and clean them by rubbing with sand paper. 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. 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. Also, compare the colour of the iron nails dipped in the copper sulphate solution with the one kept aside.

Why does the iron nail become brownish in colour and the blue colour of copper sulphate solution fades? This is because iron is more reactive than copper. So, iron displaces copper from copper sulphate solution. The following chemical reaction takes place: Fe(s) plus CuSO4(aq) arrow FeSO4(aq) plus Cu(s).

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

Other examples of displacement reactions are: Zn(s) plus CuSO4(aq) arrow ZnSO4(aq) plus Cu(s). And Pb(s) plus CuCl2(aq) arrow PbCl2(aq) plus Cu(s). Zinc and lead are more reactive elements than copper. They displace copper from its compounds.

Now, let's move on to 1.2.4: Double Displacement Reaction.

Let's perform 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. What do you observe? You will observe that a white substance, which is insoluble in water, is formed. This insoluble substance formed is known as a precipitate. Any reaction that produces a precipitate can be called a precipitation reaction.

The equation is: Na2SO4(aq) plus BaCl2(aq) arrow BaSO4(s) plus 2 NaCl(aq).

What causes this? The white precipitate of BaSO4 is formed by the reaction of SO4²⁻ and Ba²⁺. 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.

Now, recall Activity 1.2, where we mixed solutions of lead(II) nitrate and potassium iodide. (i) What was the colour of the precipitate formed? Can you name the compound precipitated? The precipitate is yellow lead iodide, PbI2. (ii) Write the balanced chemical equation for this reaction. The equation is: Pb(NO3)2(aq) plus 2 KI(aq) arrow PbI2(s) plus 2 KNO3(aq). (iii) Is this also a double displacement reaction? Yes, it is.

Now, let's move on to 1.2.5: Oxidation and Reduction.

Let's perform Activity 1.11. Heat a china dish containing about 1 gram copper powder. What do you observe? 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 copper and copper oxide is formed. The equation is: 2 Cu plus O2 arrow 2 CuO.

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. The equation is: CuO plus H2 arrow Cu plus H2O.

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, 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 plus C arrow Zn plus CO. And MnO2 plus 4 HCl arrow MnCl2 plus 2 H2O plus Cl2.

In the first reaction, carbon is oxidised to CO and ZnO is reduced to Zn. In the second reaction, 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 oxidised. If a substance loses oxygen or gains hydrogen during a reaction, it is reduced.

Now, 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? Magnesium is being oxidised because it gains oxygen.

Now, let's move on to section 1.3: Have you observed the effects of oxidation reactions in everyday life?

First, let's talk about 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 rusting of iron. Some other metals also get tarnished in this manner. Have you noticed the colour of the coating formed on copper and silver? When a metal is attacked by substances around it such as moisture, acids, etc., it is said to corrode and this process is called corrosion. 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 to all objects made of metals, specially those of iron. Corrosion of iron is a serious problem. Every year an enormous amount of money is spent to replace damaged iron.

Now, let's talk about rancidity. Have you ever tasted or smelt the fat or oil containing food materials left for a long time? When fats and oils are oxidised, they become rancid and their smell and taste change. Usually, substances which prevent oxidation (antioxidants) are added to foods containing fats and oil. Keeping food in air tight containers helps to slow down oxidation. Do you know that chips manufacturers usually flush bags of chips with gas such as nitrogen to prevent the chips from getting oxidised?

Now, let's answer some questions. Question 1: Why does the colour of copper sulphate solution change when an iron nail is dipped in it? This is because iron displaces copper from copper sulphate solution. The blue colour of copper sulphate fades as copper is displaced and iron sulphate, which is greenish, is formed.

Question 2: Give an example of a double displacement reaction other than the one given in Activity 1.10. An example is: Pb(NO3)2(aq) plus 2 KI(aq) arrow PbI2(s) plus 2 KNO3(aq).

Question 3: Identify the substances that are oxidised and the substances that are reduced in the following reactions.

(i) 4 Na(s) plus O2(g) arrow 2 Na2O(s). Here, sodium is gaining oxygen, so it is being oxidised. Oxygen is being reduced because it is the oxidising agent that provides oxygen to sodium.

(ii) CuO(s) plus H2(g) arrow Cu(s) plus H2O(l). Here, copper oxide is losing oxygen, so it is being reduced. Hydrogen is gaining oxygen, so it is being oxidised.

Now, let's look at the summary of what we have learnt:

A complete chemical equation represents the reactants, products, and their physical states symbolically.

A chemical equation is balanced so that the numbers of atoms of each type involved in a chemical reaction are the same on the reactant and product sides of the equation. Equations must always be balanced.

In a combination reaction, two or more substances combine to form a new single substance.

Decomposition reactions are opposite to combination reactions. In a decomposition reaction, a single substance decomposes to give two or more substances.

Reactions in which heat is given out along with the products are called exothermic reactions.

Reactions in which energy is absorbed are known as endothermic reactions.

When an element displaces another element from its compound, a displacement reaction occurs.

Two different atoms or groups of atoms (ions) are exchanged in double displacement reactions.

Precipitation reactions produce insoluble salts.

Reactions also involve the gain or loss of oxygen or hydrogen by substances. Oxidation is the gain of oxygen or loss of hydrogen. Reduction is the loss of oxygen or gain of hydrogen.

Now, let's solve the exercises at the end of the chapter.

Exercise 1: Which of the statements about the reaction below are incorrect? 2 PbO(s) plus C(s) arrow 2 Pb(s) plus CO2(g).

(a) Lead is getting reduced. This is incorrect because it is lead oxide (PbO) that is getting reduced to lead (Pb), not lead itself getting reduced.

(b) Carbon dioxide is getting oxidised. This is incorrect because carbon dioxide is a product, not a reactant. Carbon is getting oxidised to carbon dioxide.

(c) Carbon is getting oxidised. This is correct because carbon is gaining oxygen to form carbon dioxide.

(d) Lead oxide is getting reduced. This is correct because lead oxide is losing oxygen.

So, the incorrect statements are (a) and (b). The answer is (i).

Exercise 2: Fe2O3 plus 2 Al arrow Al2O3 plus 2 Fe. The above reaction is an example of a (a) combination reaction, (b) double displacement reaction, (c) decomposition reaction, (d) displacement reaction.

This is a displacement reaction because aluminium displaces iron from iron oxide. So, the answer is (d).

Exercise 3: What happens when dilute hydrochloric acid is added to iron fillings? Tick the correct answer. (a) Hydrogen gas and iron chloride are produced. (b) Chlorine gas and iron hydroxide are produced. (c) No reaction takes place. (d) Iron salt and water are produced.

When dilute hydrochloric acid is added to iron fillings, iron chloride and hydrogen gas are produced. The equation is: Fe(s) plus 2 HCl(aq) arrow FeCl2(aq) plus H2(g). So, the answer is (a).

Exercise 4: What is a balanced chemical equation? Why should chemical equations be balanced?

A balanced chemical equation is one in which the number of atoms of each element is the same on both the reactant and product sides. Chemical equations should be balanced because of the law of conservation of mass, which states that mass can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of reactants must equal the total mass of products.

Exercise 5: Translate the following statements into chemical equations and then balance them.

(a) Hydrogen gas combines with nitrogen to form ammonia. The equation is: N2 plus 3 H2 arrow 2 NH3.

(b) Hydrogen sulphide gas burns in air to give water and sulphur dioxide. The equation is: 2 H2S plus 3 O2 arrow 2 H2O plus 2 SO2.

(c) Barium chloride reacts with aluminium sulphate to give aluminium chloride and a precipitate of barium sulphate. The equation is: 3 BaCl2 plus Al2(SO4)3 arrow 2 AlCl3 plus 3 BaSO4.

(d) Potassium metal reacts with water to give potassium hydroxide and hydrogen gas. The equation is: 2 K plus 2 H2O arrow 2 KOH plus H2.

Exercise 6: Balance the following chemical equations.

(a) HNO3 plus Ca(OH)2 arrow Ca(NO3)2 plus H2O. Balanced: 2 HNO3 plus Ca(OH)2 arrow Ca(NO3)2 plus 2 H2O.

(b) NaOH plus H2SO4 arrow Na2SO4 plus H2O. Balanced: 2 NaOH plus H2SO4 arrow Na2SO4 plus 2 H2O.

(c) NaCl plus AgNO3 arrow AgCl plus NaNO3. This is already balanced: NaCl plus AgNO3 arrow AgCl plus NaNO3.

(d) BaCl2 plus H2SO4 arrow BaSO4 plus HCl. Balanced: BaCl2 plus H2SO4 arrow BaSO4 plus 2 HCl.

Exercise 7: Write the balanced chemical equations for the following reactions.

(a) Calcium hydroxide plus Carbon dioxide arrow Calcium carbonate plus Water. The equation is: Ca(OH)2 plus CO2 arrow CaCO3 plus H2O.

(b) Zinc plus Silver nitrate arrow Zinc nitrate plus Silver. The equation is: Zn plus 2 AgNO3 arrow Zn(NO3)2 plus 2 Ag.

(c) Aluminium plus Copper chloride arrow Aluminium chloride plus Copper. The equation is: 2 Al plus 3 CuCl2 arrow 2 AlCl3 plus 3 Cu.

(d) Barium chloride plus Potassium sulphate arrow Barium sulphate plus Potassium chloride. The equation is: BaCl2 plus K2SO4 arrow BaSO4 plus 2 KCl.

Exercise 8: Write the balanced chemical equation for the following and identify the type of reaction in each case.

(a) Potassium bromide(aq) plus Barium iodide(aq) arrow Potassium iodide(aq) plus Barium bromide(s). The balanced equation is: 2 KBr(aq) plus BaI2(aq) arrow 2 KI(aq) plus BaBr2(s). This is a double displacement reaction.

(b) Zinc carbonate(s) arrow Zinc oxide(s) plus Carbon dioxide(g). The balanced equation is: ZnCO3(s) arrow ZnO(s) plus CO2(g). This is a decomposition reaction.

(c) Hydrogen(g) plus Chlorine(g) arrow Hydrogen chloride(g). The balanced equation is: H2 plus Cl2 arrow 2 HCl. This is a combination reaction.

(d) Magnesium(s) plus Hydrochloric acid(aq) arrow Magnesium chloride(aq) plus Hydrogen(g). The balanced equation is: Mg plus 2 HCl arrow MgCl2 plus H2. This is a displacement reaction.

Exercise 9: What does one mean by exothermic and endothermic reactions? Give examples.

Exothermic reactions are those in which heat is released along with the formation of products. Examples: burning of natural gas, respiration, combustion of coal.

Endothermic reactions are those in which heat is absorbed. Examples: decomposition of calcium carbonate, decomposition of silver chloride (by light), reaction between barium hydroxide and ammonium chloride.

Exercise 10: Why is respiration considered an exothermic reaction? Explain.

Respiration is considered an exothermic reaction because during respiration, glucose combines with oxygen in the cells of our body to produce carbon dioxide, water, and energy. This energy is released in the form of heat, which keeps our body warm and provides us with the energy to carry out various life processes. The equation is: C6H12O6 plus 6 O2 arrow 6 CO2 plus 6 H2O plus energy.

Exercise 11: Why are decomposition reactions called the opposite of combination reactions? Write equations for these reactions.

Decomposition reactions are called the opposite of combination reactions because in a combination reaction, two or more substances combine to form a single product, whereas in a decomposition reaction, a single substance breaks down to give two or more substances.

Example of combination reaction: 2 H2 plus O2 arrow 2 H2O. Example of decomposition reaction: 2 H2O arrow 2 H2 plus O2.

Exercise 12: Write one equation each for decomposition reactions where energy is supplied in the form of heat, light, or electricity.

Heat: CaCO3(s) arrow CaO(s) plus CO2(g). Light: 2 AgCl(s) arrow 2 Ag(s) plus Cl2(g). Electricity: 2 H2O(l) arrow 2 H2(g) plus O2(g).

Exercise 13: What is the difference between displacement and double displacement reactions? Write equations for these reactions.

In a displacement reaction, one element displaces another element from its compound. Example: Fe(s) plus CuSO4(aq) arrow FeSO4(aq) plus Cu(s).

In a double displacement reaction, two different atoms or groups of atoms (ions) are exchanged between the reactants. Example: Na2SO4(aq) plus BaCl2(aq) arrow BaSO4(s) plus 2 NaCl(aq).

Exercise 14: In the refining of silver, the recovery of silver from silver nitrate solution involved displacement by copper metal. Write down the reaction involved.

The reaction is: 2 AgNO3(aq) plus Cu(s) arrow Cu(NO3)2(aq) plus 2 Ag(s).

Exercise 15: What do you mean by a precipitation reaction? Explain by giving examples.

A precipitation reaction is a reaction in which an insoluble substance (precipitate) is formed when two solutions are mixed. Example 1: When sodium sulphate solution is mixed with barium chloride solution, a white precipitate of barium sulphate is formed. The equation is: Na2SO4(aq) plus BaCl2(aq) arrow BaSO4(s) plus 2 NaCl(aq). Example 2: When lead nitrate solution is mixed with potassium iodide solution, a yellow precipitate of lead iodide is formed. The equation is: Pb(NO3)2(aq) plus 2 KI(aq) arrow PbI2(s) plus 2 KNO3(aq).

Exercise 16: Explain the following in terms of gain or loss of oxygen with two examples each.

(a) Oxidation: Oxidation is the gain of oxygen or loss of hydrogen. Examples: (i) 4 Na plus O2 arrow 2 Na2O (sodium gains oxygen). (ii) CuO plus H2 arrow Cu plus H2O (hydrogen gains oxygen).

(b) Reduction: Reduction is the loss of oxygen or gain of hydrogen. Examples: (i) CuO plus H2 arrow Cu plus H2O (copper oxide loses oxygen). (ii) Fe2O3 plus 3 C arrow 2 Fe plus 3 CO (iron oxide loses oxygen).

Exercise 17: A shiny brown coloured element 'X' on heating in air becomes black in colour. Name the element 'X' and the black coloured compound formed.

The element X is copper. The black coloured compound formed is copper(II) oxide, CuO. The equation is: 2 Cu plus O2 arrow 2 CuO.

Exercise 18: Why do we apply paint on iron articles?

We apply paint on iron articles to prevent corrosion (rusting). Paint acts as a protective coating that prevents iron from coming in contact with moisture and oxygen in the air, which are the main causes of rusting.

Exercise 19: Oil and fat containing food items are flushed with nitrogen. Why?

Oil and fat containing food items are flushed with nitrogen to prevent rancidity. Nitrogen is an inert gas that does not react with the oils and fats, thus preventing their oxidation. This helps to keep the food fresh for a longer time.

Exercise 20: Explain the following terms with one example each.

(a) Corrosion: Corrosion is the process by which a metal is attacked by substances around it such as moisture, acids, etc., leading to its deterioration. Example: Rusting of iron, where iron reacts with oxygen and moisture to form rust (iron oxide), which is reddish-brown in colour.

(b) Rancidity: Rancidity is the process by which fats and oils become rancid (oxidised) when exposed to air, resulting in a change in their smell and taste. Example: When chips are kept open for a long time, they become stale and taste bad due to oxidation of the oils used in them.

Now, we have covered the entire chapter. Let me give you a brief summary of what we have learnt today.

In this chapter, we learnt about chemical reactions and equations. We started by understanding what a chemical reaction is and how to identify it through changes in state, colour, evolution of gas, and change in temperature. We then learnt about chemical equations, how to write them, and how to balance them using the hit-and-trial method. We also learnt about the importance of balancing chemical equations according to the law of conservation of mass.

We then studied different types of chemical reactions. Combination reactions are those in which two or more substances combine to form a single product. Decomposition reactions are those in which a single substance breaks down into two or more substances. Displacement reactions are those in which one element displaces another from its compound. Double displacement reactions are those in which ions are exchanged between reactants. We also learnt about oxidation and reduction reactions, where substances either gain or lose oxygen.

Finally, we discussed the effects of oxidation reactions in everyday life, including corrosion and rancidity, and how to prevent them.

This is a very important chapter that will help you understand many chemical processes in your daily life and also form the basis for many topics you will study in higher classes.

Thank you for listening attentively. I hope you enjoyed this lesson. Keep studying and stay curious. See you in the next class!