Metals And Non-Metals

Hello students, welcome to today's science class. I'm so happy to see you all here ready to learn something new and exciting. Today, we are going to study Chapter 3 of your Science textbook, and that is "Metals and Non-metals". This is a very important chapter, and I promise you that by the end of this lesson, you will have a complete understanding of everything in this chapter. So sit back, relax, and let's begin our journey into the fascinating world of metals and non-metals.

Now students, let me ask you something. Look around you. What do you see? You see your bench, maybe some metal parts of your desk, your pen, maybe your watch if you wear one. Think about your kitchen at home. What vessels does your mother use for cooking? She uses stainless steel pots, aluminum vessels, maybe some copper items too. And what about non-metals? The air you breathe contains oxygen, which is a non-metal. The water you drink contains hydrogen and oxygen. The pencil you use has graphite, which is a non-metal. So you see, metals and non-metals are all around us, and understanding their properties is very important.

In Class IX, you have already learnt that elements can be classified as metals or non-metals on the basis of their properties. But today, we are going to study this in much greater detail. We will learn about the physical properties of metals and non-metals, their chemical properties, how they react with each other, and how metals are extracted from their ores. We will also learn about a very important phenomenon called corrosion and how to prevent it. So let's get started.

Students, the first thing we need to understand is the physical properties of metals. How can we identify a metal? What are the characteristics that make a metal different from a non-metal? Let us perform some activities to find out.

For performing the first few activities, we need samples of metals like iron, copper, aluminium, magnesium, sodium, lead, zinc and any other metal that is easily available. Ask your teacher to arrange these samples for you.

Now, let's look at Activity 3.1. Take samples of iron, copper, aluminium and magnesium. First, note the appearance of each sample. What do you see? You will notice that these metals have a shining surface. Now, clean the surface of each sample by rubbing them with sandpaper and note their appearance again. You will see that after cleaning, they still have a shining surface. This property is called metallic lustre. So students, whenever you see something shining like a metal, you can say it has metallic lustre. Gold and silver are the most lustrous metals, which is why they are used for making jewellery.

Now, let's move to Activity 3.2. Take small pieces of iron, copper, aluminium, and magnesium. Try to cut these metals with a sharp knife and note your observations. You will find that metals are generally hard. But the hardness varies from metal to metal. Some metals like sodium and potassium are so soft that they can be cut with a knife. Now, let me tell you something interesting about sodium. Sodium metal is stored in kerosene because it reacts vigorously with air and moisture. So always handle sodium metal with care. If you want to cut sodium, first dry it by pressing between the folds of a filter paper, then put it on a watch-glass and try to cut it with a knife. What do you observe? You will see that even though sodium is soft, it still shows some resistance to cutting. So remember, most metals are hard, but there are exceptions like sodium and potassium.

Moving on to Activity 3.3. Take pieces of iron, zinc, lead and copper. Place any one metal on a block of iron and strike it four or five times with a hammer. What do you observe? You will see that the metal flattens or changes its shape. Repeat with other metals and record the change in the shape. You will find that some metals can be beaten into thin sheets. This property is called malleability. Did you know that gold and silver are the most malleable metals? This means you can beat gold into very thin sheets called gold leaf, which is used for decorative purposes. Isn't that fascinating?

Now, let's discuss Activity 3.4. List the metals whose wires you have seen in daily life. You must have seen wires made of copper and aluminium. These wires are used for electrical wiring in homes. The ability of metals to be drawn into thin wires is called ductility. Gold is the most ductile metal. You will be surprised to know that a wire of about 2 km length can be drawn from just one gram of gold! That is incredible, isn't it? It is because of their malleability and ductility that metals can be given different shapes according to our needs. Think about the various shapes of metal objects you see around you - from thin wires to flat sheets to different utensils.

Now students, let me ask you a question. Which metals are used for making cooking vessels? Do you know why these metals are used for making vessels? Let us do Activity 3.5 to find out the answer.

Take an aluminium or copper wire. Clamp this wire on a stand. Fix a pin to the free end of the wire using wax. Now, heat the wire with a spirit lamp, candle or a burner near the place where it is clamped. What do you observe after some time? You will see that the wax melts and the pin falls off. But does the metal wire melt? No, it doesn't. This shows that metals have high melting points. The above activity demonstrates that metals are good conductors of heat and have high melting points. The best conductors of heat are silver and copper. That is why copper is used for making cooking vessels and electrical wires. Lead and mercury are comparatively poor conductors of heat.

Now, do metals also conduct electricity? Let us find out through Activity 3.6. Set up an electric circuit as shown in Figure 3.2 in your textbook. Place the metal to be tested in the circuit between terminals A and B. Does the bulb glow? If it does, it means the metal conducts electricity. You must have seen that the wires that carry current in your homes have a coating of polyvinylchloride (PVC) or a rubber-like material. Why are electric wires coated with such substances? It is to prevent electric shock, because these coatings are insulators and do not conduct electricity. So students, remember that metals are good conductors of electricity.

Now, what happens when metals strike a hard surface? Do they produce a sound? The metals that produce a sound on striking a hard surface are said to be sonorous. Can you now say why school bells are made of metals? Yes, because metals are sonorous. When you strike a metal bell, it produces a loud, ringing sound.

So students, let me quickly recap what we have learned about the physical properties of metals. Metals have metallic lustre, they are generally hard, they are malleable (can be beaten into sheets), they are ductile (can be drawn into wires), they are good conductors of heat and electricity, and they are sonorous (produce sound when struck). Now, let's learn about non-metals.

In the previous Class, you have learnt that there are very few non-metals as compared to metals. Some examples of non-metals are carbon, sulphur, iodine, oxygen, hydrogen, etc. The non-metals are either solids or gases except bromine, which is a liquid. So students, remember that most non-metals are solids or gases, and bromine is the only liquid non-metal at room temperature.

Now, do non-metals also have physical properties similar to that of metals? Let us find out through Activity 3.7. Collect samples of carbon (coal or graphite), sulphur and iodine. Carry out the Activities 3.1 to 3.4 and 3.6 with these non-metals and record your observations.

After performing these activities, you will notice that non-metals do not have metallic lustre. They are not malleable or ductile - in fact, they are brittle. They are poor conductors of heat and electricity. Graphite is an exception - it conducts electricity. Iodine is a non-metal but it is lustrous, which is an exception to the general property.

Now students, compile your observations regarding metals and non-metals in Table 3.1 as given in your textbook. On the basis of these observations, we can say that metals and non-metals have different physical properties. However, we cannot group elements according to their physical properties alone, as there are many exceptions. Let me tell you about some important exceptions.

First, all metals except mercury exist as solids at room temperature. Mercury is the only liquid metal at room temperature. Second, metals have high melting points, but gallium and caesium have very low melting points. These two metals will melt if you keep them on your palm! Third, iodine is a non-metal but it is lustrous. Fourth, carbon is a non-metal that can exist in different forms. Each form is called an allotrope. Diamond, an allotrope of carbon, is the hardest natural substance known and has a very high melting and boiling point. Graphite, another allotrope of carbon, is a conductor of electricity. Fifth, alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.

So students, you see that there are many exceptions to the general properties. This is why elements can be more clearly classified as metals and non-metals on the basis of their chemical properties, which we will study now.

Now, let's perform Activity 3.8 to understand the chemical properties of metals and non-metals. Take a magnesium ribbon and some sulphur powder. Burn the magnesium ribbon. Collect the ashes formed and dissolve them in water. Test the resultant solution with both red and blue litmus paper. What do you observe? The blue litmus paper turns red, which means the product is acidic. Now burn sulphur powder. Place a test tube over the burning sulphur to collect the fumes produced. Add some water to the above test tube and shake. Test this solution with blue and red litmus paper. What do you observe? The red litmus paper turns blue, which means the product is basic.

So students, what did we learn from this? We learned that when magnesium burns, it produces a basic oxide. When sulphur burns, it produces an acidic oxide. Most non-metals produce acidic oxides when dissolved in water. On the other hand, most metals give rise to basic oxides. You will be learning more about these metal oxides in the next section.

Now students, let's study the chemical properties of metals in detail. We will learn about what happens when metals are burnt in air, when they react with water, when they react with acids, and when they react with solutions of other metal salts.

Let's start with Section 3.2.1: What happens when Metals are burnt in Air?

You have seen in Activity 3.8 that magnesium burns in air with a dazzling white flame. Do all metals react in the same manner? Let us check by performing Activity 3.9.

CAUTION: The following activity needs the teacher's assistance. It would be better if students wear eye protection.

Hold any of the metal samples with a pair of tongs and try burning over a flame. Repeat with the other metal samples. Collect the product if formed. Let the products and the metal surface cool down. Which metals burn easily? What flame colour did you observe when the metal burnt? How does the metal surface appear after burning? Arrange the metals in the decreasing order of their reactivity towards oxygen. Are the products soluble in water?

From this activity, you will observe that almost all metals combine with oxygen to form metal oxides. The general reaction is: Metal + Oxygen → Metal oxide.

For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide. The reaction is: 2Cu + O₂ → 2CuO. Similarly, aluminium forms aluminium oxide: 4Al + 3O₂ → 2Al₂O₃.

Now students, recall from Chapter 2, how copper oxide reacts with hydrochloric acid. We have learnt that metal oxides are basic in nature. But some metal oxides, such as aluminium oxide and zinc oxide, show both acidic as well as basic behaviour. Such metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides. Aluminium oxide reacts in the following manner with acids and bases:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O

Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (Sodium aluminate)

Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium oxide and potassium oxide dissolve in water to produce alkalis as follows:

Na₂O(s) + H₂O(l) → 2NaOH(aq)

K₂O(s) + H₂O(l) → 2KOH(aq)

Now students, we have observed in Activity 3.9 that all metals do not react with oxygen at the same rate. Different metals show different reactivities towards oxygen. Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil. At ordinary temperature, the surfaces of metals such as magnesium, aluminium, zinc, lead, etc., are covered with a thin layer of oxide. The protective oxide layer prevents the metal from further oxidation. Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner. Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide. Silver and gold do not react with oxygen even at high temperatures.

Now, let me tell you about a process called anodising. Anodising is a process of forming a thick oxide layer of aluminium. Aluminium develops a thin oxide layer when exposed to air. This aluminium oxide coat makes it resistant to further corrosion. The resistance can be improved further by making the oxide layer thicker. During anodising, a clean aluminium article is made the anode and is electrolysed with dilute sulphuric acid. The oxygen gas evolved at the anode reacts with aluminium to make a thicker protective oxide layer. This oxide layer can be dyed easily to give aluminium articles an attractive finish. This is why you see coloured aluminium windows and frames nowadays.

After performing Activity 3.9, you must have observed that sodium is the most reactive of the samples of metals taken here. The reaction of magnesium is less vigorous implying that it is not as reactive as sodium. But burning in oxygen does not help us to decide about the reactivity of zinc, iron, copper or lead. Let us see some more reactions to arrive at a conclusion about the order of reactivity of these metals.

Now, let's move to Section 3.2.2: What happens when Metals react with Water?

Perform Activity 3.10. CAUTION: This Activity needs the teacher's assistance.

Collect the samples of the same metals as in Activity 3.9. Put small pieces of the samples separately in beakers half-filled with cold water. Which metals reacted with cold water? Arrange them in the increasing order of their reactivity with cold water. Did any metal produce fire on water? Does any metal start floating after some time? Put the metals that did not react with cold water in beakers half-filled with hot water. For the metals that did not react with hot water, arrange the apparatus as shown in Figure 3.3 and observe their reaction with steam. Which metals did not react even with steam? Arrange the metals in the decreasing order of reactivity with water.

From this activity, you will learn that some metals react with cold water, some react with hot water, some react with steam, and some don't react with water at all. For example, sodium and potassium react violently with cold water, producing hydrogen gas and enough heat to cause the hydrogen to catch fire. Calcium reacts less vigorously with cold water. Magnesium reacts with hot water. Aluminium, zinc and iron react with steam. Copper does not react with water at all.

Now, let's move to Section 3.2.3: What happens when Metals react with Acids?

You have already learnt that metals react with acids to give a salt and hydrogen gas. The general reaction is: Metal + Dilute acid → Salt + Hydrogen

But do all metals react in the same manner? Let us find out through Activity 3.11.

Collect all the metal samples except sodium and potassium again. If the samples are tarnished, rub them clean with sand paper. CAUTION: Do not take sodium and potassium as they react vigorously even with cold water.

Put the samples separately in test tubes containing dilute hydrochloric acid. Suspend thermometers in the test tubes, so that their bulbs are dipped in the acid. Observe the rate of formation of bubbles carefully. Which metals reacted vigorously with dilute hydrochloric acid? With which metal did you record the highest temperature? Arrange the metals in the decreasing order of reactivity with dilute acids.

Write equations for the reactions of magnesium, aluminium, zinc and iron with dilute hydrochloric acid.

The reactions are:

Mg + 2HCl → MgCl₂ + H₂

2Al + 6HCl → 2AlCl₃ + 3H₂

Zn + 2HCl → ZnCl₂ + H₂

Fe + 2HCl → FeCl₂ + H₂

Now students, hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HNO₃ is a strong oxidising agent. It oxidises the H₂ produced to water and itself gets reduced to any of the nitrogen oxides (N₂O, NO, NO₂). But magnesium (Mg) and manganese (Mn) react with very dilute HNO₃ to evolve H₂ gas.

You must have observed in Activity 3.11 that the rate of formation of bubbles was the fastest in the case of magnesium. The reaction was also the most exothermic in this case. The reactivity decreases in the order Mg > Al > Zn > Fe. In the case of copper, no bubbles were seen and the temperature also remained unchanged. This shows that copper does not react with dilute HCl.

Now, let me tell you about aqua regia. Aqua regia (Latin for 'royal water') is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1. It can dissolve gold, even though neither of these acids can do so alone. Aqua regia is a highly corrosive, fuming liquid. It is one of the few reagents that is able to dissolve gold and platinum. This is very important in the refining of gold.

Now, let's move to Section 3.2.4: How do Metals react with Solutions of other Metal Salts?

Perform Activity 3.12. Take a clean wire of copper and an iron nail. Put the copper wire in a solution of iron sulphate and the iron nail in a solution of copper sulphate taken in test tubes. Record your observations after 20 minutes.

In which test tube did you find that a reaction has occurred? On what basis can you say that a reaction has actually taken place? Can you correlate your observations for the Activities 3.9, 3.10 and 3.11? Write a balanced chemical equation for the reaction that has taken place. Name the type of reaction.

From this activity, you will observe that when iron nail is placed in copper sulphate solution, the iron nail gets coated with a reddish-brown deposit of copper, and the blue colour of copper sulphate fades. This is because iron is more reactive than copper and displaces copper from its salt solution. The reaction is: Fe + CuSO₄ → FeSO₄ + Cu

When copper wire is placed in iron sulphate solution, no change occurs because copper is less reactive than iron.

So students, the key point here is: Reactive metals can displace less reactive metals from their compounds in solution or molten form. This is known as a displacement reaction.

We have seen in the previous sections that all metals are not equally reactive. We checked the reactivity of various metals with oxygen, water and acids. But all metals do not react with these reagents. So we were not able to put all the metal samples we had collected in decreasing order of their reactivity. Displacement reactions give better evidence about the reactivity of metals. It is simple and easy - if metal A displaces metal B from its solution, it is more reactive than B.

The general equation is: Metal A + Salt solution of B → Salt solution of A + Metal B

Now students, let's talk about the Reactivity Series. The reactivity series is a list of metals arranged in the order of their decreasing activities. After performing displacement experiments, the following series, known as the reactivity or activity series, has been developed. This is given in Table 3.2 in your textbook.

From most reactive to least reactive: Potassium (K) is the most reactive, then Sodium (Na), Calcium (Ca), Magnesium (Mg), Aluminium (Al), Zinc (Zn), Iron (Fe), Lead (Pb), then comes Hydrogen (H), then Copper (Cu), Mercury (Hg), Silver (Ag), and Gold (Au) is the least reactive.

Students, notice that hydrogen is included in the reactivity series. Hydrogen is not a metal, but it is included because it is used as a reference to determine the reactivity of metals. Metals above hydrogen can displace hydrogen from dilute acids, while metals below hydrogen cannot.

Now, let's answer the questions given after Section 3.2.

Question 1: Why is sodium kept immersed in kerosene oil?

Answer: Sodium is a very reactive metal. It reacts vigorously with oxygen and moisture present in the air. If kept in the open, it can catch fire or get covered with a layer of oxide, making it useless. Therefore, sodium is kept immersed in kerosene oil to prevent its contact with air and moisture.

Question 2: Write equations for the reactions of (i) iron with steam, (ii) calcium and potassium with water.

Answer: (i) Iron reacts with steam to form iron(II,III) oxide and hydrogen gas. The reaction is: 3Fe + 4H₂O → Fe₃O₄ + 4H₂

(ii) Calcium reacts with water to form calcium hydroxide and hydrogen gas: Ca + 2H₂O → Ca(OH)₂ + H₂

Potassium reacts violently with water to form potassium hydroxide and hydrogen gas: 2K + 2H₂O → 2KOH + H₂

Question 3: Samples of four metals A, B, C and D were taken and added to the following solution one by one. The results obtained have been tabulated as follows.

Metal | Iron(II) sulphate | Copper(II) sulphate | Zinc sulphate | Silver nitrate A | No reaction | Displacement | | B | Displacement | | No reaction | C | No reaction | No reaction | No reaction | Displacement D | No reaction | No reaction | No reaction | No reaction

Use the Table above to answer the following questions about metals A, B, C and D.

(i) Which is the most reactive metal?

Answer: Metal B is the most reactive because it displaces iron from iron(II) sulphate.

(ii) What would you observe if B is added to a solution of Copper(II) sulphate?

Answer: Since B is more reactive than copper, it will displace copper from copper(II) sulphate. We would observe a reddish-brown deposit of copper on B, and the blue colour of copper sulphate solution would fade.

(iii) Arrange the metals A, B, C and D in the order of decreasing reactivity.

Answer: From the table, we can see that: - B displaces iron from iron(II) sulphate, so B is more reactive than iron. - A displaces copper from copper(II) sulphate, so A is more reactive than copper. - C displaces silver from silver nitrate, so C is more reactive than silver. - D shows no reaction with any of the salts, so D is the least reactive.

We know that iron is more reactive than copper, and copper is more reactive than silver. So the order of decreasing reactivity is: B > A > C > D

Question 4: Which gas is produced when dilute hydrochloric acid is added to a reactive metal? Write the chemical reaction when iron reacts with dilute H₂SO₄.

Answer: Hydrogen gas is produced when dilute hydrochloric acid is added to a reactive metal. When iron reacts with dilute H₂SO₄, the reaction is: Fe + H₂SO₄ → FeSO₄ + H₂

Question 5: What would you observe when zinc is added to a solution of iron(II) sulphate? Write the chemical reaction that takes place.

Answer: Zinc is more reactive than iron. So when zinc is added to a solution of iron(II) sulphate, zinc will displace iron from its salt solution. We would observe that the pale green colour of iron(II) sulphate solution would become colourless as zinc sulphate is colourless, and a greyish-brown deposit of iron would form on zinc. The reaction is: Zn + FeSO₄ → ZnSO₄ + Fe

Now students, let's move to Section 3.3: How do Metals and Non-metals React?

In the above activities, you saw the reactions of metals with a number of reagents. Why do metals react in this manner? Let us recall what we learnt about the electronic configuration of elements in Class IX. We learnt that noble gases, which have a completely filled valence shell, show little chemical activity. We, therefore, explain the reactivity of elements as a tendency to attain a completely filled valence shell.

Let us have a look at the electronic configuration of noble gases and some metals and non-metals. This is given in Table 3.3 in your textbook.

We can see that: - Helium (He) has 2 electrons in its outermost shell (K shell). - Neon (Ne) has 2, 8 electrons in K and L shells. - Argon (Ar) has 2, 8, 8 electrons in K, L and M shells.

These noble gases have a completely filled valence shell, which makes them stable and unreactive.

Now, let's look at metals: - Sodium (Na) has atomic number 11, so its electronic configuration is 2, 8, 1. It has one electron in its outermost shell. - Magnesium (Mg) has atomic number 12, so its electronic configuration is 2, 8, 2. It has two electrons in its outermost shell. - Aluminium (Al) has atomic number 13, so its electronic configuration is 2, 8, 3. It has three electrons in its outermost shell. - Potassium (K) has atomic number 19, so its electronic configuration is 2, 8, 8, 1. It has one electron in its outermost shell. - Calcium (Ca) has atomic number 20, so its electronic configuration is 2, 8, 8, 2. It has two electrons in its outermost shell.

Now, let's look at non-metals: - Nitrogen (N) has atomic number 7, so its electronic configuration is 2, 5. It has five electrons in its outermost shell. - Oxygen (O) has atomic number 8, so its electronic configuration is 2, 6. It has six electrons in its outermost shell. - Fluorine (F) has atomic number 9, so its electronic configuration is 2, 7. It has seven electrons in its outermost shell. - Phosphorus (P) has atomic number 15, so its electronic configuration is 2, 8, 5. It has five electrons in its outermost shell. - Sulphur (S) has atomic number 16, so its electronic configuration is 2, 8, 6. It has six electrons in its outermost shell. - Chlorine (Cl) has atomic number 17, so its electronic configuration is 2, 8, 7. It has seven electrons in its outermost shell.

Now students, let's understand how sodium and chlorine react. We can see from the table that a sodium atom has one electron in its outermost shell. If it loses the electron from its M shell then its L shell now becomes the outermost shell and that has a stable octet. The nucleus of this atom still has 11 protons but the number of electrons has become 10, so there is a net positive charge giving us a sodium cation Na⁺.

On the other hand, chlorine has seven electrons in its outermost shell and it requires one more electron to complete its octet. If sodium and chlorine were to react, the electron lost by sodium could be taken up by chlorine. After gaining an electron, the chlorine atom gets a unit negative charge, because its nucleus has 17 protons and there are 18 electrons in its K, L and M shells. This gives us a chloride anion Cl⁻.

So both these elements can have a give-and-take relation between them. The reactions are:

Na → Na⁺ + e⁻ (Sodium loses an electron and becomes a positive ion) 2,8,1 → 2,8

Cl + e⁻ → Cl⁻ (Chlorine gains an electron and becomes a negative ion) 2,8,7 → 2,8,8

Na + Cl → (Na⁺)(Cl⁻) (Sodium chloride is formed)

Sodium and chloride ions, being oppositely charged, attract each other and are held by strong electrostatic forces of attraction to exist as sodium chloride (NaCl). It should be noted that sodium chloride does not exist as molecules but aggregates of oppositely charged ions.

Let us see the formation of one more ionic compound, magnesium chloride.

Mg → Mg²⁺ + 2e⁻ (Magnesium loses two electrons) 2,8,2 → 2,8

Cl + e⁻ → Cl⁻ (Chlorine gains one electron) 2,8,7 → 2,8,8

Mg + 2Cl → (Mg²⁺)(Cl⁻)₂ (Magnesium chloride is formed)

The compounds formed in this manner by the transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds. In MgCl₂, the cation is Mg²⁺ and the anion is Cl⁻.

Now students, let's learn about the properties of ionic compounds through Activity 3.13.

Take samples of sodium chloride, potassium iodide, barium chloride or any other salt from the science laboratory. What is the physical state of these salts? You will find that they are all solids. Take a small amount of a sample on a metal spatula and heat directly on the flame. Repeat with other samples. What did you observe? Did the samples impart any colour to the flame? Do these compounds melt? You will observe that these compounds melt only at high temperatures. Try to dissolve the samples in water, petrol and kerosene. Are they soluble? You will find that they are soluble in water but insoluble in petrol and kerosene. Make a circuit as shown in Figure 3.8 and insert the electrodes into a solution of one salt. What did you observe? You will see that the bulb glows, indicating that the solution conducts electricity. Test the other salt samples too in this manner. What is your inference about the nature of these compounds?

From this activity, we can conclude that ionic compounds have specific properties. Let me summarize the properties of ionic compounds:

First, physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between the positive and negative ions. These compounds are generally brittle and break into pieces when pressure is applied.

Second, melting and boiling points: Ionic compounds have high melting and boiling points. This is because a considerable amount of energy is required to break the strong inter-ionic attraction. For example, NaCl has a melting point of 1074 K and a boiling point of 1686 K.

Third, solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.

Fourth, conduction of electricity: The conduction of electricity through a solution involves the movement of charged particles. A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure. But ionic compounds conduct electricity in the molten state. This is possible in the molten state since the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat. Thus, the ions move freely and conduct electricity.

Now students, let's answer the questions given after Section 3.3.

Question 1: (i) Write the electron-dot structures for sodium, oxygen and magnesium. (ii) Show the formation of Na₂O and MgO by the transfer of electrons. (iii) What are the ions present in these compounds?

Answer: (i) Sodium has one electron in its outermost shell, so its electron-dot structure is Na•. Oxygen has six electrons in its outermost shell, so its electron-dot structure is :Ö:. Magnesium has two electrons in its outermost shell, so its electron-dot structure is Mg:.

(ii) Sodium has one electron in its outermost shell. Oxygen needs two electrons to complete its octet. So two sodium atoms each give one electron to one oxygen atom. The reactions are:

Na → Na⁺ + e⁻ :Ö: + 2e⁻ → O²⁻ 2Na + :Ö: → 2Na⁺ + O²⁻ → Na₂O

Similarly, magnesium has two electrons in its outermost shell. Oxygen needs two electrons to complete its octet. So one magnesium atom gives two electrons to one oxygen atom. The reactions are:

Mg → Mg²⁺ + 2e⁻ :Ö: + 2e⁻ → O²⁻ Mg + :Ö: → Mg²⁺ + O²⁻ → MgO

(iii) In Na₂O, the ions present are Na⁺ (sodium ion) and O²⁻ (oxide ion). In MgO, the ions present are Mg²⁺ (magnesium ion) and O²⁻ (oxide ion).

Question 2: Why do ionic compounds have high melting points?

Answer: Ionic compounds have high melting points because of the strong electrostatic forces of attraction between the positive and negative ions. A considerable amount of energy is required to break these forces. That is why ionic compounds have high melting and boiling points.

Now students, let's move to Section 3.4: Occurrence of Metals.

The earth's crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride, magnesium chloride, etc. The elements or compounds, which occur naturally in the earth's crust, are known as minerals. At some places, minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.

So students, remember the difference: A mineral is a naturally occurring substance from which a metal can be extracted profitably. The impurities present in the ore are called gangue.

Now, let's learn about the extraction of metals. You have learnt about the reactivity series of metals. Having this knowledge, you can easily understand how a metal is extracted from its ore. Some metals are found in the earth's crust in the free state. Some are found in the form of their compounds. The metals at the bottom of the activity series are the least reactive. They are often found in a free state. For example, gold, silver, platinum and copper are found in the free state. Copper and silver are also found in the combined state as their sulphide or oxide ores.

The metals at the top of the activity series (K, Na, Ca, Mg and Al) are so reactive that they are never found in nature as free elements. The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive. They are found in the earth's crust mainly as oxides, sulphides or carbonates. You will find that the ores of many metals are oxides. This is because oxygen is a very reactive element and is very abundant on the earth.

Thus, on the basis of reactivity, we can group the metals into the following three categories: (i) Metals of low reactivity; (ii) Metals of medium reactivity; (iii) Metals of high reactivity. Different techniques are to be used for obtaining the metals falling in each category.

Several steps are involved in the extraction of pure metal from ores. First, the ore is enriched to remove gangue. Then, the metal is extracted from the enriched ore. Finally, the metal is refined to obtain pure metal.

Now, let's discuss enrichment of ores. Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue. The impurities must be removed from the ore prior to the extraction of the metal. The processes used for removing the gangue from the ore are based on the differences between the physical or chemical properties of the gangue and the ore. Different separation techniques are accordingly employed.

Now, let's discuss extracting metals low in the activity series. Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone. For example, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO). Mercuric oxide is then reduced to mercury on further heating.

2HgS(s) + 3O₂(g) → 2HgO(s) + 2SO₂(g)

2HgO(s) → 2Hg(l) + O₂(g)

Similarly, copper which is found as Cu₂S in nature can be obtained from its ore by just heating in air.

2Cu₂S + 3O₂(g) → 2Cu₂O(s) + 2SO₂(g)

2Cu₂O + Cu₂S → 6Cu(s) + SO₂(g)

Now, let's discuss extracting metals in the middle of the activity series. The metals in the middle of the activity series such as iron, zinc, lead, copper, are moderately reactive. These are usually present as sulphides or carbonates in nature. It is easier to obtain a metal from its oxide, as compared to its sulphides and carbonates. Therefore, prior to reduction, the metal sulphides and carbonates must be converted into metal oxides.

The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This process is known as roasting. The carbonate ores are changed into oxides by heating strongly in limited air. This process is known as calcination.

The chemical reaction that takes place during roasting and calcination of zinc ores can be shown as follows:

Roasting: 2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)

Calcination: ZnCO₃(s) → ZnO(s) + CO₂(g)

The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc.

ZnO(s) + C(s) → Zn(s) + CO(g)

You are already familiar with the process of oxidation and reduction explained in the first Chapter. Obtaining metals from their compounds is also a reduction process.

Besides using carbon (coke) to reduce metal oxides to metals, sometimes displacement reactions can also be used. The highly reactive metals such as sodium, calcium, aluminium, etc., are used as reducing agents because they can displace metals of lower reactivity from their compounds. For example, when manganese dioxide is heated with aluminium powder, the following reaction takes place:

3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + Heat

Can you identify the substances that are getting oxidised and reduced? In this reaction, manganese dioxide is being reduced to manganese, and aluminium is being oxidised to aluminium oxide.

These displacement reactions are highly exothermic. The amount of heat evolved is so large that the metals are produced in the molten state. In fact, the reaction of iron(III) oxide (Fe₂O₃) with aluminium is used to join railway tracks or cracked machine parts. This reaction is known as the thermit reaction.

Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat

Now, let's discuss extracting metals towards the top of the activity series. The metals high up in the reactivity series are very reactive. They cannot be obtained from their compounds by heating with carbon. For example, carbon cannot reduce the oxides of sodium, magnesium, calcium, aluminium, etc., to the respective metals. This is because these metals have more affinity for oxygen than carbon. These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode (the negatively charged electrode), whereas chlorine is liberated at the anode (the positively charged electrode). The reactions are:

At cathode: Na⁺ + e⁻ → Na

At anode: 2Cl⁻ → Cl₂ + 2e⁻

Similarly, aluminium is obtained by the electrolytic reduction of aluminium oxide.

Now, let's discuss refining of metals. The metals produced by various reduction processes described above are not very pure. They contain impurities, which must be removed to obtain pure metals. The most widely used method for refining impure metals is electrolytic refining.

Electrolytic Refining: Many metals, such as copper, zinc, tin, nickel, silver, gold, etc., are refined electrolytically. In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. A solution of the metal salt is used as an electrolyte. On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte is deposited on the cathode. The soluble impurities go into the solution, whereas the insoluble impurities settle down at the bottom of the anode and are known as anode mud.

Now students, let's answer the questions given after Section 3.4.

Question 1: Define the following terms: (i) Mineral, (ii) Ore, (iii) Gangue.

Answer: (i) Mineral: The elements or compounds, which occur naturally in the earth's crust, are known as minerals.

(ii) Ore: At some places, minerals contain a very high percentage of a particular metal and the metal can be profitably extracted from it. These minerals are called ores.

(iii) Gangue: The impurities present in the ore such as soil, sand, etc., are called gangue.

Question 2: Name two metals which are found in nature in the free state.

Answer: Gold and platinum are two metals which are found in nature in the free state. Silver and copper can also be found in the free state sometimes.

Question 3: What chemical process is used for obtaining a metal from its oxide?

Answer: Reduction is the chemical process used for obtaining a metal from its oxide. For example, when zinc oxide is heated with carbon, it is reduced to zinc metal.

Now students, let's move to Section 3.5: Corrosion.

You have learnt the following about corrosion in Chapter 1: - Silver articles become black after some time when exposed to air. This is because it reacts with sulphur in the air to form a coating of silver sulphide. - Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is basic copper carbonate. - Iron when exposed to moist air for a long time acquires a coating of a brown flaky substance called rust.

Let us find out the conditions under which iron rusts through Activity 3.14.

Take three test tubes and place clean iron nails in each of them. Label these test tubes A, B and C. Pour some water in test tube A and cork it. Pour boiled distilled water in test tube B, add about 1 mL of oil and cork it. The oil will float on water and prevent the air from dissolving in the water. Put some anhydrous calcium chloride in test tube C and cork it. Anhydrous calcium chloride will absorb the moisture, if any, from the air. Leave these test tubes for a few days and then observe.

You will observe that iron nails rust in test tube A, but they do not rust in test tubes B and C. In the test tube A, the nails are exposed to both air and water. In the test tube B, the nails are exposed to only water, and the nails in test tube C are exposed to dry air. What does this tell us about the conditions under which iron articles rust?

From this activity, we can conclude that iron rusts when it is exposed to both air and water. If either air or water is absent, rusting does not occur. So, the presence of both moisture and oxygen is necessary for rusting.

Now, let's discuss prevention of corrosion. The rusting of iron can be prevented by painting, oiling, greasing, galvanising, chrome plating, anodising or making alloys.

Galvanisation is a method of protecting steel and iron from rusting by coating them with a thin layer of zinc. The galvanised article is protected against rusting even if the zinc coating is broken. This is because zinc is more reactive than iron, so it reacts with oxygen and moisture first, sacrificing itself to protect the iron. This is called sacrificial protection.

Alloying is a very good method of improving the properties of a metal. We can get the desired properties by this method. For example, iron is the most widely used metal. But it is never used in its pure state. This is because pure iron is very soft and stretches easily when hot. But if it is mixed with a small amount of carbon (about 0.05%), it becomes hard and strong. When iron is mixed with nickel and chromium, we get stainless steel, which is hard and does not rust. Thus, if iron is mixed with some other substance, its properties change. In fact, the properties of any metal can be changed if it is mixed with some other substance. The substance added may be a metal or a non-metal. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by first melting the primary metal, and then dissolving the other elements in it in definite proportions. It is then cooled to room temperature.

Now, let me tell you about gold. Pure gold, known as 24 carat gold, is very soft. It is, therefore, not suitable for making jewellery. It is alloyed with either silver or copper to make it hard. Generally, in India, 22 carat gold is used for making ornaments. It means that 22 parts of pure gold is alloyed with 2 parts of either copper or silver.

If one of the metals is mercury, then the alloy is known as an amalgam. The electrical conductivity and melting point of an alloy is less than that of pure metals. For example, brass, an alloy of copper and zinc (Cu and Zn), and bronze, an alloy of copper and tin (Cu and Sn), are not good conductors of electricity whereas copper is used for making electrical circuits. Solder, an alloy of lead and tin (Pb and Sn), has a low melting point and is used for welding electrical wires together.

Now, let me tell you about the wonder of ancient Indian metallurgy. The iron pillar near the Qutub Minar in Delhi was built more than 1600 years ago by the iron workers of India. They had developed a process which prevented iron from rusting. For its quality of rust resistance it has been examined by scientists from all parts of the world. The iron pillar is 8 m high and weighs 6 tonnes (6000 kg). This shows that ancient Indians had advanced knowledge of metallurgy!

Now students, let's answer the questions given after Section 3.5.

Question 1: Metallic oxides of zinc, magnesium and copper were heated with the following metals:

Metal | Zinc oxide | Magnesium oxide | Copper oxide Zinc | | | Magnesium | | | Copper | | |

In which cases will you find displacement reactions taking place?

Answer: A displacement reaction takes place when a more reactive metal displaces a less reactive metal from its oxide. From the reactivity series, we know that magnesium is more reactive than zinc, which is more reactive than copper. So: - When zinc is heated with magnesium oxide: No reaction (Zn is less reactive than Mg) - When magnesium is heated with zinc oxide: Displacement reaction (Mg displaces Zn) - Mg + ZnO → MgO + Zn - When copper is heated with zinc oxide: No reaction (Cu is less reactive than Zn) - When zinc is heated with copper oxide: Displacement reaction (Zn displaces Cu) - Zn + CuO → ZnO + Cu - When magnesium is heated with copper oxide: Displacement reaction (Mg displaces Cu) - Mg + CuO → MgO + Cu - When copper is heated with magnesium oxide: No reaction (Cu is less reactive than Mg)

So the displacement reactions take place when: magnesium is heated with zinc oxide, zinc is heated with copper oxide, and magnesium is heated with copper oxide.

Question 2: Which metals do not corrode easily?

Answer: Metals like gold, platinum, silver and copper do not corrode easily. They are called noble metals because they are very less reactive.

Question 3: What are alloys?

Answer: An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by first melting the primary metal, and then dissolving the other elements in it in definite proportions. Examples of alloys are brass (copper and zinc), bronze (copper and tin), stainless steel (iron, chromium and nickel), etc.

Now students, we have completed the entire chapter. Now let's solve the exercises at the end of the chapter.

Exercise Question 1: Which of the following pairs will give displacement reactions?

(a) NaCl solution and copper metal (b) MgCl₂ solution and aluminium metal (c) FeSO₄ solution and silver metal (d) AgNO₃ solution and copper metal.

Answer: Let's check each option:

(a) NaCl solution and copper metal: Sodium is more reactive than copper, but copper is less reactive than sodium. So copper cannot displace sodium from NaCl. No reaction.

(b) MgCl₂ solution and aluminium metal: Magnesium is more reactive than aluminium. So aluminium cannot displace magnesium from MgCl₂. No reaction.

(c) FeSO₄ solution and silver metal: Iron is more reactive than silver. So silver cannot displace iron from FeSO₄. No reaction.

(d) AgNO₃ solution and copper metal: Copper is more reactive than silver. So copper will displace silver from AgNO₃. Yes, displacement reaction takes place. The reaction is: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag

So the answer is (d).

Exercise Question 2: Which of the following methods is suitable for preventing an iron frying pan from rusting?

(a) Applying grease (b) Applying paint (c) Applying a coating of zinc (d) All of the above.

Answer: All of the above methods can prevent rusting. Applying grease or oil creates a protective layer that prevents contact with air and water. Applying paint also creates a protective layer. Applying a coating of zinc is called galvanisation, which protects iron from rusting. So the answer is (d) All of the above.

Exercise Question 3: An element reacts with oxygen to give a compound with a high melting point. This compound is also soluble in water. The element is likely to be

(a) calcium (b) carbon (c) silicon (d) iron.

Answer: Let's analyse each option:

(a) Calcium: Calcium reacts with oxygen to form calcium oxide (CaO). Calcium oxide has a high melting point (about 2600°C) and is soluble in water, forming calcium hydroxide.

(b) Carbon: Carbon reacts with oxygen to form carbon dioxide (CO₂), which is a gas, not a compound with a high melting point.

(c) Silicon: Silicon reacts with oxygen to form silicon dioxide (SiO₂), which has a high melting point but is not soluble in water.

(d) Iron: Iron reacts with oxygen to form iron oxide (Fe₂O₃), which has a high melting point but is not very soluble in water.

So the answer is (a) calcium.

Exercise Question 4: Food cans are coated with tin and not with zinc because

(a) zinc is costlier than tin. (b) zinc has a higher melting point than tin. (c) zinc is more reactive than tin. (d) zinc is less reactive than tin.

Answer: Zinc is more reactive than tin. If food cans were coated with zinc, the zinc would react with the food and dissolve, making the food toxic. Tin is less reactive and does not react with food. So the answer is (c) zinc is more reactive than tin.

Exercise Question 5: You are given a hammer, a battery, a bulb, wires and a switch.

(a) How could you use them to distinguish between samples of metals and non-metals?

(b) Assess the usefulness of these tests in distinguishing between metals and non-metals.

Answer: (a) We can set up an electric circuit using the battery, bulb, wires and switch. Place the sample in the circuit between the terminals. If the bulb glows, the sample is a metal (good conductor of electricity). If the bulb does not glow, the sample is a non-metal (poor conductor of electricity), except for graphite which is a non-metal but conducts electricity.

We can also use the hammer to test hardness and malleability. Metals are generally hard and malleable, while non-metals are generally brittle.

(b) The electrical conductivity test is useful but not foolproof because graphite (a non-metal) conducts electricity. The hammer test is also useful but not definitive because there are exceptions like sodium (a metal that is soft) and iodine (a non-metal that is lustrous). So these tests give us an indication but we need to consider multiple properties to correctly identify metals and non-metals.

Exercise Question 6: What are amphoteric oxides? Give two examples of amphoteric oxides.

Answer: Amphoteric oxides are metal oxides that react with both acids as well as bases to produce salts and water. They show both acidic as well as basic behaviour. Two examples of amphoteric oxides are aluminium oxide (Al₂O₃) and zinc oxide (ZnO).

Exercise Question 7: Name two metals which will displace hydrogen from dilute acids, and two metals which will not.

Answer: Metals that are more reactive than hydrogen (i.e., metals above hydrogen in the reactivity series) will displace hydrogen from dilute acids. Examples: Magnesium and Zinc.

Metals that are less reactive than hydrogen (i.e., metals below hydrogen in the reactivity series) will not displace hydrogen from dilute acids. Examples: Copper and Silver.

Exercise Question 8: In the electrolytic refining of a metal M, what would you take as the anode, the cathode and the electrolyte?

Answer: In electrolytic refining: - The anode is made of the impure metal M. - The cathode is made of a thin strip of pure metal M. - The electrolyte is a solution of a salt of metal M (e.g., MSO₄).

Exercise Question 9: Pratyush took sulphur powder on a spatula and heated it. He collected the gas evolved by inverting a test tube over it, as shown in figure below.

(a) What will be the action of gas on (i) dry litmus paper? (ii) moist litmus paper?

(b) Write a balanced chemical equation for the reaction taking place.

Answer: When sulphur is burned in air, it produces sulphur dioxide gas.

(a) (i) Dry litmus paper: Sulphur dioxide gas does not have any direct effect on dry litmus paper because it needs moisture to show its acidic properties.

(ii) Moist litmus paper: Sulphur dioxide dissolves in the moisture present on the moist litmus paper to form sulphurous acid, which turns blue litmus paper red.

(b) The balanced chemical equation is: S + O₂ → SO₂

Exercise Question 10: State two ways to prevent the rusting of iron.

Answer: Two ways to prevent rusting of iron are: 1. Painting or applying a protective coating to the iron surface. 2. Galvanisation, i.e., coating iron with a layer of zinc.

Other methods include oiling, greasing, and making alloys like stainless steel.

Exercise Question 11: What type of oxides are formed when non-metals combine with oxygen?

Answer: When non-metals combine with oxygen, they form acidic oxides or neutral oxides. For example, carbon forms carbon dioxide (CO₂), which is acidic; sulphur forms sulphur dioxide (SO₂), which is acidic; nitrogen forms nitrous oxide (N₂O), which is neutral.

Exercise Question 12: Give reasons

(a) Platinum, gold and silver are used to make jewellery.

Answer: Platinum, gold and silver are used to make jewellery because they are lustrous, malleable, ductile, and do not corrode easily. They are also precious and have high value.

(b) Sodium, potassium and lithium are stored under oil.

Answer: Sodium, potassium and lithium are very reactive metals. They react vigorously with oxygen and moisture in the air. If stored in the open, they can catch fire or get covered with a layer of oxide. Therefore, they are stored under oil to prevent their contact with air and moisture.

(c) Aluminium is a highly reactive metal, yet it is used to make utensils for cooking.

Answer: Aluminium is a highly reactive metal, but it forms a thin layer of aluminium oxide on its surface when exposed to air. This oxide layer is protective and prevents further reaction of aluminium with food. Moreover, aluminium is a good conductor of heat, lightweight, and inexpensive. That is why it is used for making utensils for cooking.

(d) Carbonate and sulphide ores are usually converted into oxides during the process of extraction.

Answer: It is easier to obtain a metal from its oxide as compared to its carbonate or sulphide. Therefore, carbonate and sulphide ores are converted into oxides by the processes of calcination and roasting respectively, before extracting the metal.

Exercise Question 13: You must have seen tarnished copper vessels being cleaned with lemon or tamarind juice. Explain why these sour substances are effective in cleaning the vessels.

Answer: Tarnished copper vessels have a coating of basic copper carbonate (green) and copper sulphide (black) on their surface. Lemon juice and tamarind juice contain acids (citric acid and tartaric acid respectively). These acids react with the basic copper carbonate and copper sulphide, dissolving them and cleaning the vessel. This is a classic example of how acids react with basic substances.

Exercise Question 14: Differentiate between metal and non-metal on the basis of their chemical properties.

Answer: The differences between metals and non-metals on the basis of their chemical properties are:

1. Reaction with oxygen: Metals form basic oxides (or amphoteric oxides), while non-metals form acidic oxides (or neutral oxides).

2. Reaction with water: Metals react with water (except some like gold, silver, copper), but non-metals generally do not react with water.

3. Reaction with acids: Metals react with dilute acids to produce hydrogen gas, while non-metals do not react with dilute acids.

4. Reaction with salt solutions: More reactive metals can displace less reactive metals from their salt solutions, while non-metals do not show this property.

5. Formation of ions: Metals form positive ions by losing electrons, while non-metals form negative ions by gaining electrons.

6. Nature of compounds: Metals form ionic compounds, while non-metals form covalent compounds.

Exercise Question 15: A man went door to door posing as a goldsmith. He promised to bring back the glitter of old and dull gold ornaments. An unsuspecting lady gave a set of gold bangles to him which he dipped in a particular solution. The bangles sparkled like new but their weight was reduced drastically. The lady was upset but after a futile argument the man beat a hasty retreat. Can you play the detective to find out the nature of the solution he had used?

Answer: The solution used was aqua regia, which is a mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3:1. Aqua regia can dissolve gold, which is a noble metal. When the gold bangles were dipped in aqua regia, the gold dissolved (reacted with the acid), reducing the weight of the bangles. The glitter was probably due to the remaining gold or some other effect. This is a classic example of how one should be careful of such fraudsters!

Exercise Question 16: Give reasons why copper is used to make hot water tanks and not steel (an alloy of iron).

Answer: Copper is used to make hot water tanks because: 1. Copper is a better conductor of heat than steel. 2. Copper does not rust (corrode) easily, while steel (iron) can rust when exposed to water and air. 3. Copper can withstand high temperatures without getting damaged.

Steel, on the other hand, is an alloy of iron and can rust when exposed to hot water and air. Therefore, copper is preferred over steel for making hot water tanks.

Now students, we have completed all the exercises. Let me give you a quick summary of everything we have learned in this chapter.

What you have learnt:

- Elements can be classified as metals and non-metals. - Metals are lustrous, malleable, ductile and are good conductors of heat and electricity. They are solids at room temperature, except mercury which is a liquid. - Metals can form positive ions by losing electrons to non-metals. - Metals combine with oxygen to form basic oxides. Aluminium oxide and zinc oxide show the properties of both basic as well as acidic oxides. These oxides are known as amphoteric oxides. - Different metals have different reactivities with water and dilute acids. - A list of common metals arranged in order of their decreasing reactivity is known as an activity series. - Metals above hydrogen in the Activity series can displace hydrogen from dilute acids. - A more reactive metal displaces a less reactive metal from its salt solution. - Metals occur in nature as free elements or in the form of their compounds. - The extraction of metals from their ores and then refining them for use is known as metallurgy. - An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. - The surface of some metals, such as iron, is corroded when they are exposed to moist air for a long period of time. This phenomenon is known as corrosion. - Non-metals have properties opposite to that of metals. They are neither malleable nor ductile. They are bad conductors of heat and electricity, except for graphite, which conducts electricity. - Non-metals form negatively charged ions by gaining electrons when reacting with metals. - Non-metals form oxides which are either acidic or neutral. - Non-metals do not displace hydrogen from dilute acids. They react with hydrogen to form hydrides.

Students, I hope you have understood all the concepts in this chapter. This is a very important chapter, and you will find these concepts useful in your daily life as well as in higher classes. Remember to revise this chapter regularly and practice the exercises. If you have any doubts, please ask your teacher.

Thank you for listening attentively. Have a great day and happy learning!