Welcome dear students! Today we are going to learn about Chemical Reactions and Equations from Class 10 Science.
Consider the following situations of daily life and think what happens when milk is left at room temperature during summers, an iron tawa or pan or nail is left exposed to humid atmosphere, grapes get fermented, food is cooked, food gets digested in our body, and we respire. In all the above 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. You may perhaps be wondering as to what is actually meant by a chemical reaction. How do we come to know that a chemical reaction has taken place? Let us perform some activities to find the answer to these questions.
Activity 1.1. CAUTION: This Activity needs the teacher’s assistance. It would be better if students wear suitable eyeglasses. Clean a magnesium ribbon about three to four centimetres long by rubbing it with sandpaper. Hold it with a pair of tongs. Burn it using a spirit lamp or burner and collect the ash so formed in a watch glass. Burn the magnesium ribbon keeping it away as far as possible from your eyes. What do you observe? You must have observed that magnesium ribbon burns with a dazzling white flame and changes into a white powder. This powder is magnesium oxide. It is formed due to the reaction between magnesium and oxygen present in the air.
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Activity 1.2. Take lead nitrate solution in a test tube. Add potassium iodide solution to this. What do you observe? Activity 1.3. Take a few zinc granules in a conical flask or a test tube. Add dilute hydrochloric acid or sulphuric acid to this. CAUTION: Handle the acid with care. Do you observe anything happening around the zinc granules? Touch the conical flask or test tube. Is there any change in its temperature? From the above three activities, we can say that any of the following observations helps us to determine whether a chemical reaction has taken place. Change in state. Change in colour. Evolution of a gas. Change in temperature. 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.
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 yields Magnesium oxide. The substances that undergo chemical change in this reaction, magnesium and oxygen, are the reactants. The new substance is magnesium oxide, formed during the reaction, as a 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.
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Section 1.1.1, Writing a Chemical Equation. Is there any other shorter way for representing chemical equations? 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 formulae of magnesium, oxygen and magnesium oxide, the above word equation can be written as Mg plus O₂ yields MgO. Count and compare the number of atoms of each element on the left hand side and right hand side of the arrow. Is the number of atoms of each element the same on both the sides? If yes, then the equation is balanced. If not, then the equation is unbalanced because the mass is not the same on both sides of the equation. Such a chemical equation is a skeletal chemical equation for a reaction. This equation is a skeletal chemical equation for the burning of magnesium in air.
Section 1.1.2, Balanced Chemical Equations. Recall the law of conservation of mass that you studied in Class Nine. 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. Let us learn about balancing a chemical equation step by step. The word equation for Activity 1.3 may be represented as Zinc plus Sulphuric acid yields Zinc sulphate plus Hydrogen. The above word equation may be represented by the following chemical equation, Zn plus H₂SO₄ yields ZnSO₄ plus H₂. Let us examine the number of atoms of different elements on both sides of the arrow. For Zinc, there is one atom in reactants and one in products. For Hydrogen, there are two atoms in reactants and two in products. For Sulphur, there is one atom in reactants and one in products. For Oxygen, there are four atoms in reactants and four in products. 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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Let us try to balance the following chemical equation, Fe plus H₂O yields Fe₃O₄ plus H₂. Step one. To balance a chemical equation, first draw boxes around each formula. Do not change anything inside the boxes while balancing the equation. Step two. List the number of atoms of different elements present in the unbalanced equation. For Iron, there is one on the left and three on the right. For Hydrogen, there are two on the left and two on the right. For Oxygen, there is one on the left and four on the right. Step three. 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, initially we have one in water and four in iron oxide. To balance, we multiply water by four. To equalise the number of atoms, it must be remembered that we cannot alter the formulae of the compounds or elements involved in the reactions. For example, to balance oxygen atoms we can put coefficient four as four H₂O and not H₂O₄ or H₂O in brackets four. Now the partly balanced equation becomes Fe plus four H₂O yields Fe₃O₄ plus H₂.
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Step four. Iron and hydrogen atoms are still not balanced. Pick any of these elements to proceed further. Let us balance hydrogen atoms in the partly balanced equation. To equalise the number of hydrogen atoms, make the number of molecules of hydrogen as four on the right hand side. Initially we have eight hydrogen atoms in four water molecules and two in hydrogen gas. To balance, we multiply hydrogen gas by four. The equation would be Fe plus four H₂O yields Fe₃O₄ plus four H₂. Step five. 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 one iron atom in iron and three in iron oxide. To balance, we multiply iron by three. To equalise iron, we take three atoms of iron on the left hand side. Three Fe plus four H₂O yields Fe₃O₄ plus four H₂. Step six. Finally, to check the correctness of the balanced equation, we count atoms of each element on both sides of the equation. The numbers of atoms of elements on both sides are equal. This equation is now balanced. This method of balancing chemical equations is called hit and trial method as we make trials to balance the equation by using the smallest whole number coefficient. Step seven. 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 is written if the reactant or product is present as a solution in water. The balanced equation becomes three Fe(s) plus four H₂O(g) yields Fe₃O₄(s) plus four H₂(g). Note that the symbol (g) is used with water 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, and so on, for the reaction are indicated above and or below the arrow in the equation. For example, CO(g) plus two H₂(g) yields CH₃OH(l) at three hundred forty atmospheres. Six CO₂(aq) plus twelve H₂O(l) yields C₆H₁₂O₆(aq) plus six O₂(aq) plus six H₂O(l) under sunlight and chlorophyll. Using these steps, can you balance the magnesium burning equation given earlier?
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Now let us answer the questions from this section. Question one. Why should a magnesium ribbon be cleaned before burning in air? Answer. Magnesium ribbon is cleaned before burning to remove the protective layer of magnesium oxide or magnesium carbonate that forms on its surface when exposed to air. This layer prevents the magnesium metal from reacting with oxygen efficiently, so cleaning ensures a proper and vigorous reaction. Question two. Write the balanced equation for the following chemical reactions. Part one. Hydrogen plus Chlorine yields Hydrogen chloride. Balanced equation is H₂(g) plus Cl₂(g) yields two HCl(g). Part two. Barium chloride plus Aluminium sulphate yields Barium sulphate plus Aluminium chloride. Balanced equation is three BaCl₂(aq) plus Al₂(SO₄)₃(aq) yields three BaSO₄(s) plus two AlCl₃(aq). Part three. Sodium plus Water yields Sodium hydroxide plus Hydrogen. Balanced equation is two Na(s) plus two H₂O(l) yields two NaOH(aq) plus H₂(g). Question three. Write a balanced chemical equation with state symbols for the following reactions. Part one. Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride. Balanced equation is BaCl₂(aq) plus Na₂SO₄(aq) yields BaSO₄(s) plus two NaCl(aq). Part two. Sodium hydroxide solution in water reacts with hydrochloric acid solution in water to produce sodium chloride solution and water. Balanced equation is NaOH(aq) plus HCl(aq) yields NaCl(aq) plus H₂O(l).
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Section 1.2, Types of Chemical Reactions. We have learnt in Class Nine 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 three and four. Section 1.2.1, Combination Reaction. Let us look at 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? Calcium oxide reacts vigorously with water to produce slaked lime, which is calcium hydroxide, releasing a large amount of heat. The equation is CaO(s) plus H₂O(l) yields Ca(OH)₂(aq) plus Heat. Quick lime reacts to form slaked lime. 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 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 reaction is Ca(OH)₂(aq) plus CO₂(g) yields CaCO₃(s) plus H₂O(l). Let us discuss some more examples of combination reactions. First, burning of coal. C(s) plus O₂(g) yields CO₂(g). Second, formation of water from hydrogen gas and oxygen gas. Two H₂(g) plus O₂(g) yields two H₂O(l). In simple language we can say that when two or more substances, whether 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, first, burning of natural gas. CH₄(g) plus two O₂(g) yields CO₂(g) plus two H₂O(g). Second, 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 process of which you will study in Chapter six. The equation is C₆H₁₂O₆(aq) plus six O₂(aq) yields six CO₂(aq) plus six H₂O(l) plus energy. Third, 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. It is a combination reaction.
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Section 1.2.2, Decomposition Reaction. Let us perform Activity 1.5. Take about two grams of 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. Observe the colour of the crystals after heating. Have you noticed that the green colour of the ferrous sulphate crystals has changed? You can also smell the characteristic odour of burning sulphur. The equation is two FeSO₄(s) yields 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, sulphur dioxide and sulphur trioxide. 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) yields CaO(s) plus CO₂(g).
Another example of a thermal decomposition reaction is given in Activity 1.6. Take about two 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? Note down the change, if any. You will observe the emission of brown fumes. These fumes are of nitrogen dioxide. The reaction that takes place is two Pb(NO₃)₂(s) yields two PbO(s) plus four NO₂(g) plus O₂(g). 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. Connect these electrodes to a six 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? The volume of gas collected at the cathode is double that at the anode. The gas at the cathode is hydrogen, which burns with a pop sound. The gas at the anode is oxygen, which supports combustion and makes the candle burn brighter.
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Activity 1.8. Take about two grams of silver chloride in a china dish. What is its colour? 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 two AgCl(s) yields two Ag(s) plus Cl₂(g). Silver bromide also behaves in the same way. Two AgBr(s) yields two Ag(s) plus Br₂(g) in the presence of sunlight. 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 two grams of barium hydroxide in a test tube. Add one 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? You will feel the test tube becoming cold, indicating that heat is absorbed from the surroundings. This is an endothermic reaction.
Let us answer the questions for this section. Question one. A solution of a substance X is used for whitewashing. Part one. Name the substance X and write its formula. Answer. The substance X is calcium oxide, also known as quick lime. Its chemical formula is CaO. Part two. Write the reaction of the substance X named above with water. Answer. CaO(s) plus H₂O(l) yields Ca(OH)₂(aq) plus Heat. Question two. 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. Answer. In the electrolysis of water, water decomposes into hydrogen and oxygen gases in a two to one ratio by volume. The chemical formula of water is H₂O, meaning it contains two hydrogen atoms for every one oxygen atom. Therefore, the volume of hydrogen gas collected is double that of oxygen gas. The gas collected in larger volume is hydrogen.
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Section 1.2.3, Displacement Reaction. Let us look at 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 ten millilitres of 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 twenty minutes. Keep one iron nail aside for comparison. After twenty 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 fade? The following chemical reaction takes place in this activity. Fe(s) plus CuSO₄(aq) yields 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 displacement reaction. Other examples of displacement reactions are Zn(s) plus CuSO₄(aq) yields ZnSO₄(aq) plus Cu(s), and Pb(s) plus CuCl₂(aq) yields PbCl₂(aq) plus Cu(s). Zinc and lead are more reactive elements than copper. They displace copper from its compounds.
Section 1.2.4, Double Displacement Reaction. Activity 1.10. Take about three millilitres of sodium sulphate solution in a test tube. In another test tube, take about three millilitres 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) yields BaSO₄(s) plus two NaCl(aq). What causes this? The white precipitate of barium sulphate is formed by the reaction of sulphate ions and barium ions. 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 two nitrate and potassium iodide. Part one. What was the colour of the precipitate formed? Can you name the compound precipitated? Answer. A yellow precipitate of lead iodide is formed. Part two. Write the balanced chemical equation for this reaction. Answer. Pb(NO₃)₂(aq) plus two KI(aq) yields PbI₂(s) plus two KNO₃(aq). Part three. Is this also a double displacement reaction? Answer. Yes, it is a double displacement reaction because the lead and potassium ions exchange their partners.
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Section 1.2.5, Oxidation and Reduction. Activity 1.11. Heat a china dish containing about one gram of copper powder. What do you observe? The surface of copper powder becomes coated with black copper two oxide. Why has this black substance formed? This is because oxygen is added to copper and copper oxide is formed. The equation is two Cu plus O₂ yields two CuO. If hydrogen gas is passed over this heated material, the black coating on the surface turns brown as the reverse reaction takes place and copper is obtained. The equation is CuO plus H₂ yields 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 two 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 yields Zn plus CO, and MnO₂ plus four HCl yields MnCl₂ plus two H₂O plus Cl₂. In the first reaction, carbon is oxidised to carbon monoxide and zinc oxide is reduced to zinc. In the second reaction, hydrochloric acid is oxidised to chlorine gas whereas manganese dioxide is reduced to manganese chloride. 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. Recall Activity 1.1, where a magnesium ribbon burns with a dazzling flame in air and changes into a white substance, magnesium oxide. Is magnesium being oxidised or reduced in this reaction? Magnesium gains oxygen to form magnesium oxide, so magnesium is being oxidised.
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Section 1.3, Have you observed the effects of oxidation reactions in everyday life? Section 1.3.1, 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 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, and so on, 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. You will learn more about corrosion in Chapter three. Section 1.3.2, 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, known as 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?
Let us answer the questions for this section. Question one. Why does the colour of copper sulphate solution change when an iron nail is dipped in it? Answer. Iron is more reactive than copper. When an iron nail is dipped in copper sulphate solution, iron displaces copper from the solution to form iron sulphate, which is greenish in colour, and copper metal gets deposited. Hence, the blue colour of copper sulphate fades to green. Question two. Give an example of a double displacement reaction other than the one given in Activity 1.10. Answer. Silver nitrate reacts with sodium chloride to form silver chloride precipitate and sodium nitrate. The equation is AgNO₃(aq) plus NaCl(aq) yields AgCl(s) plus NaNO₃(aq). Question three. Identify the substances that are oxidised and the substances that are reduced in the following reactions. Part one. Four Na(s) plus O₂(g) yields two Na₂O(s). Answer. Sodium gains oxygen, so sodium is oxidised. Oxygen loses oxygen equivalent, so oxygen is reduced. Part two. CuO(s) plus H₂(g) yields Cu(s) plus H₂O(l). Answer. Copper oxide loses oxygen, so copper oxide is reduced. Hydrogen gains oxygen, so hydrogen is oxidised.
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What you 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, which are 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 us proceed to the Exercises. Question one. Which of the statements about the reaction below are incorrect? Two PbO(s) plus C(s) yields two Pb(s) plus CO₂(g). Options are a Lead is getting reduced. b Carbon dioxide is getting oxidised. c Carbon is getting oxidised. d Lead oxide is getting reduced. Answer. In this reaction, lead oxide loses oxygen to become lead, so it is reduced. Carbon gains oxygen to become carbon dioxide, so it is oxidised. Therefore, statements a and b are incorrect. The correct option is i, a and b. Question two. Fe₂O₃ plus two Al yields Al₂O₃ plus two Fe. The above reaction is an example of a combination reaction, double displacement reaction, decomposition reaction, or displacement reaction. Answer. Aluminium displaces iron from iron oxide. Therefore, it is a displacement reaction. The correct option is d. Question three. What happens when dilute hydrochloric acid is added to iron fillings? Tick the correct answer. Options are 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. Answer. Iron reacts with dilute hydrochloric acid to form iron chloride and hydrogen gas. The correct option is a.
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Question four. What is a balanced chemical equation? Why should chemical equations be balanced? Answer. 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 must be balanced to obey the law of conservation of mass, which states that mass can neither be created nor destroyed in a chemical reaction. Question five. Translate the following statements into chemical equations and then balance them. Part a. Hydrogen gas combines with nitrogen to form ammonia. Answer. Three H₂(g) plus N₂(g) yields two NH₃(g). Part b. Hydrogen sulphide gas burns in air to give water and sulphur dioxide. Answer. Two H₂S(g) plus three O₂(g) yields two H₂O(l) plus two SO₂(g). Part c. Barium chloride reacts with aluminium sulphate to give aluminium chloride and a precipitate of barium sulphate. Answer. Three BaCl₂(aq) plus Al₂(SO₄)₃(aq) yields two AlCl₃(aq) plus three BaSO₄(s). Part d. Potassium metal reacts with water to give potassium hydroxide and hydrogen gas. Answer. Two K(s) plus two H₂O(l) yields two KOH(aq) plus H₂(g). Question six. Balance the following chemical equations. Part a. HNO₃ plus Ca(OH)₂ yields Ca(NO₃)₂ plus H₂O. Balanced: two HNO₃ plus Ca(OH)₂ yields Ca(NO₃)₂ plus two H₂O. Part b. NaOH plus H₂SO₄ yields Na₂SO₄ plus H₂O. Balanced: two NaOH plus H₂SO₄ yields Na₂SO₄ plus two H₂O. Part c. NaCl plus AgNO₃ yields AgCl plus NaNO₃. Balanced: NaCl plus AgNO₃ yields AgCl plus NaNO₃. It is already balanced. Part d. BaCl₂ plus H₂SO₄ yields BaSO₄ plus HCl. Balanced: BaCl₂ plus H₂SO₄ yields BaSO₄ plus two HCl.
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Question seven. Write the balanced chemical equations for the following reactions. Part a. Calcium hydroxide plus Carbon dioxide yields Calcium carbonate plus Water. Answer. Ca(OH)₂(aq) plus CO₂(g) yields CaCO₃(s) plus H₂O(l). Part b. Zinc plus Silver nitrate yields Zinc nitrate plus Silver. Answer. Zn(s) plus two AgNO₃(aq) yields Zn(NO₃)₂(aq) plus two Ag(s). Part c. Aluminium plus Copper chloride yields Aluminium chloride plus Copper. Answer. Two Al(s) plus three CuCl₂(aq) yields two AlCl₃(aq) plus three Cu(s). Part d. Barium chloride plus Potassium sulphate yields Barium sulphate plus Potassium chloride. Answer. BaCl₂(aq) plus K₂SO₄(aq) yields BaSO₄(s) plus two KCl(aq). Question eight. Write the balanced chemical equation for the following and identify the type of reaction in each case. Part a. Potassium bromide aqueous plus Barium iodide aqueous yields Potassium iodide aqueous plus Barium bromide solid. Answer. Two KBr(aq) plus BaI₂(aq) yields two KI(aq) plus BaBr₂(s). This is a double displacement reaction. Part b. Zinc carbonate solid yields Zinc oxide solid plus Carbon dioxide gas. Answer. ZnCO₃(s) yields ZnO(s) plus CO₂(g). This is a decomposition reaction. Part c. Hydrogen gas plus Chlorine gas yields Hydrogen chloride gas. Answer. H₂(g) plus Cl₂(g) yields two HCl(g). This is a combination reaction. Part d. Magnesium solid plus Hydrochloric acid aqueous yields Magnesium chloride aqueous plus Hydrogen gas. Answer. Mg(s) plus two HCl(aq) yields MgCl₂(aq) plus H₂(g). This is a displacement reaction.
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Question nine. What does one mean by exothermic and endothermic reactions? Give examples. Answer. Exothermic reactions are those in which heat is released along with the formation of products. An example is the combustion of methane: CH₄(g) plus two O₂(g) yields CO₂(g) plus two H₂O(g) plus heat. Endothermic reactions are those in which energy is absorbed from the surroundings. An example is the decomposition of calcium carbonate: CaCO₃(s) yields CaO(s) plus CO₂(g), which requires heat. Question ten. Why is respiration considered an exothermic reaction? Explain. Answer. Respiration is considered an exothermic reaction because during this process, glucose combines with oxygen in our cells to produce carbon dioxide, water, and a large amount of energy. Since energy is released in the form of heat and ATP, it is classified as exothermic. Question eleven. Why are decomposition reactions called the opposite of combination reactions? Write equations for these reactions. Answer. In a combination reaction, two or more reactants combine to form a single product. In a decomposition reaction, a single reactant breaks down into two or more products. Hence, they are opposite processes. Example of combination: C(s) plus O₂(g) yields CO₂(g). Example of decomposition: CaCO₃(s) yields CaO(s) plus CO₂(g). Question twelve. Write one equation each for decomposition reactions where energy is supplied in the form of heat, light or electricity. Answer. Heat: CaCO₃(s) yields CaO(s) plus CO₂(g). Light: two AgCl(s) yields two Ag(s) plus Cl₂(g) in sunlight. Electricity: two H₂O(l) yields two H₂(g) plus O₂(g).
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Question thirteen. What is the difference between displacement and double displacement reactions? Write equations for these reactions. Answer. In a displacement reaction, a more reactive element displaces a less reactive element from its compound. In a double displacement reaction, two compounds exchange their ions to form two new compounds. Displacement example: Fe(s) plus CuSO₄(aq) yields FeSO₄(aq) plus Cu(s). Double displacement example: Na₂SO₄(aq) plus BaCl₂(aq) yields BaSO₄(s) plus two NaCl(aq). Question fourteen. In the refining of silver, the recovery of silver from silver nitrate solution involved displacement by copper metal. Write down the reaction involved. Answer. Cu(s) plus two AgNO₃(aq) yields Cu(NO₃)₂(aq) plus two Ag(s). Question fifteen. What do you mean by a precipitation reaction? Explain by giving examples. Answer. A precipitation reaction is a chemical reaction in which two aqueous solutions react to form an insoluble solid called a precipitate. Example: AgNO₃(aq) plus NaCl(aq) yields AgCl(s) plus NaNO₃(aq). Here, silver chloride is the white precipitate. Question sixteen. Explain the following in terms of gain or loss of oxygen with two examples each. Part a. Oxidation. Answer. Oxidation is the gain of oxygen or loss of hydrogen during a reaction. Examples: two Mg(s) plus O₂(g) yields two MgO(s). Magnesium gains oxygen. C(s) plus O₂(g) yields CO₂(g). Carbon gains oxygen. Part b. Reduction. Answer. Reduction is the loss of oxygen or gain of hydrogen during a reaction. Examples: CuO(s) plus H₂(g) yields Cu(s) plus H₂O(l). Copper oxide loses oxygen. Fe₂O₃(s) plus three CO(g) yields two Fe(s) plus three CO₂(g). Iron oxide loses oxygen.
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Question seventeen. A shiny brown coloured element X on heating in air becomes black in colour. Name the element X and the black coloured compound formed. Answer. The element X is copper. The black coloured compound formed is copper oxide, CuO. Question eighteen. Why do we apply paint on iron articles? Answer. We apply paint on iron articles to prevent them from coming into contact with moisture and oxygen in the air. This protective coating stops the process of rusting, which is the corrosion of iron. Question nineteen. Oil and fat containing food items are flushed with nitrogen. Why? Answer. Nitrogen is an inert gas that does not react easily with fats and oils. Flushing food items with nitrogen removes oxygen from the packaging, which prevents the oxidation of fats and oils, thereby avoiding rancidity and preserving freshness. Question twenty. Explain the following terms with one example each. Part a. Corrosion. Answer. Corrosion is the gradual destruction of metals by chemical reaction with their environment, usually moisture and oxygen. Example: Rusting of iron, where iron reacts with oxygen and water to form hydrated iron oxide. Part b. Rancidity. Answer. Rancidity is the oxidation of fats and oils in food, leading to unpleasant smell and taste. Example: Potato chips turning stale and smelling bad when left open to air for a long time.
Group Activity. Perform the following activity. Take four beakers and label them as A, B, C and D. Put twenty five millilitres of water in A, B and C beakers and copper sulphate solution in beaker D. Measure and record the temperature of each liquid contained in the beakers above. Add two spatulas of potassium sulphate, ammonium nitrate, anhydrous copper sulphate and fine iron fillings to beakers A, B, C and D respectively and stir. Finally measure and record the temperature of each of the mixture above. Find out which reactions are exothermic and which ones are endothermic in nature. Answer. In this activity, adding anhydrous copper sulphate to water and adding iron fillings to copper sulphate solution will release heat, making the beaker warm. These are exothermic processes. Adding ammonium nitrate to water will absorb heat, making the beaker cold. This is an endothermic process. Adding potassium sulphate to water shows little to no temperature change, indicating a physical dissolution process. By comparing initial and final temperatures, you can classify the reactions accordingly.
Thank you for listening! Keep revising and practicing. Goodbye! [CHAPTER_COMPLETE]