Welcome dear students! Today we are going to learn about Metals and Non-metals from Class 10 Science.
In Class Nine you have learnt about various elements. You have seen that elements can be classified as metals or non-metals on the basis of their properties. Think of some uses of metals and non-metals in your daily life. What properties did you think of while categorising elements as metals or non-metals? How are these properties related to the uses of these elements? Let us look at some of these properties in detail.
We begin with the physical properties of metals. The easiest way to start grouping substances is by comparing their physical properties. Let us study this with the help of activities. For performing Activities three point one to three point six, collect samples of iron, copper, aluminium, magnesium, sodium, lead, zinc and any other metal that is easily available.
Activity three point one instructs you to take samples of iron, copper, aluminium and magnesium. Note the appearance of each sample. Clean the surface of each sample by rubbing them with sand paper and note their appearance again. You will observe that metals, in their pure state, have a shining surface. This property is called metallic lustre.
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Activity three point two asks you to take small pieces of iron, copper, aluminium, and magnesium. Try to cut these metals with a sharp knife and note your observations. Next, hold a piece of sodium metal with a pair of tongs. Always handle sodium metal with care. Dry it by pressing between the folds of a filter paper. Put it on a watch-glass and try to cut it with a knife. What do you observe? You will notice that sodium is soft and can be cut easily with a knife.
Activity three point three requires you to 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. 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? Repeat the hammering with other metals and record the change in shape. You will also find that metals are generally hard, though the hardness varies from metal to metal.
Activity three point four asks you to list the metals whose wires you have seen in daily life. 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 two kilometres length can be drawn from one gram of gold. It is because of their malleability and ductility that metals can be given different shapes according to our needs. Can you name some metals that are used for making cooking vessels? Do you know why these metals are used for making vessels? Let us do the following activity to find out the answer.
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Activity three point five instructs you to take an aluminium or copper wire. Clamp this wire on a stand, as shown in Figure three point one. In this diagram, you can see a metal wire clamped horizontally on a laboratory stand. Fix a pin to the free end of the wire using wax. Heat the wire with a spirit lamp, candle or a burner near the place where it is clamped. After some time, you will observe that the wax melts and the pin falls off. The metal wire does not melt. This activity shows that metals are good conductors of heat and have high melting points. The best conductors of heat are silver and copper. Lead and mercury are comparatively poor conductors of heat.
Do metals also conduct electricity? Let us find out. You must have seen that the wires that carry current in your homes have a coating of polyvinylchloride or PVC or a rubber-like material. Why are electric wires coated with such substances? This coating provides essential electrical insulation for safety. 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. This is why school bells are made of metals.
Now let us move on to 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, and so on. The non-metals are either solids or gases except bromine which is a liquid. Do non-metals also have physical properties similar to that of metals? Let us find out.
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Activity three point six asks you to set up an electric circuit as shown in Figure three point two. In this figure, a battery, a bulb, a switch, and two terminals labelled A and B are connected in series. Place the metal to be tested in the circuit between terminals A and B. The bulb glows, indicating that metals are good conductors of electricity.
Activity three point seven instructs you to collect samples of carbon, specifically coal or graphite, sulphur and iodine. Carry out Activities three point one to three point four and three point six with these non-metals and record your observations. You will compile these in a table. Based on these observations, you will conclude that we cannot group elements according to their physical properties alone, as there are many exceptions. For instance, all metals except mercury exist as solids at room temperature. In Activity three point five, you observed that 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. Iodine is a non-metal but it is lustrous. 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. Alkali metals, which are lithium, sodium, and potassium, are so soft that they can be cut with a knife. They have low densities and low melting points. Elements can be more clearly classified as metals and non-metals on the basis of their chemical properties.
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Activity three point eight asks you to 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. The product formed on burning magnesium is basic. 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. The product formed on burning sulphur is acidic. You can write equations for these reactions.
Let us address the questions following this section. Question one asks for an example of a metal which is a liquid at room temperature. The answer is mercury. Which can be easily cut with a knife? The answer is sodium. Which is the best conductor of heat? The answer is silver. Which is a poor conductor of heat? The answer is lead or mercury. Question two asks to explain the meanings of malleable and ductile. Malleable means the ability of a substance to be beaten into thin sheets. Ductile means the ability of a substance to be drawn into thin wires.
Now we move to the chemical properties of metals. We will learn about these in the following sections. Most non-metals produce acidic oxides when dissolved in water. On the other hand, most metals give rise to basic oxides.
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Section three point two point one asks what happens when metals are burnt in air. You have seen in Activity three point eight that magnesium burns in air with a dazzling white flame. Do all metals react in the same manner? Let us check by performing Activity three point nine. This 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. You will observe which metals burn easily, the flame colour, and how the metal surface appears after burning. Arrange the metals in decreasing order of their reactivity towards oxygen. Check if the products are soluble in water. Almost all metals combine with oxygen to form metal oxides. The reaction is Metal plus Oxygen yields Metal oxide. For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide. The equation is 2Cu + O₂ → 2CuO. Similarly, aluminium forms aluminium oxide. The equation is 4Al + 3O₂ → 2Al₂O₃. Recall from Chapter two 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, which is sodium aluminate. Most metal oxides are insoluble in water but some 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).
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We have observed in Activity three point nine 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, and so on, 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.
Do You Know? 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.
After performing Activity three point nine, you must have observed that sodium is the most reactive of the samples taken. 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.
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Section three point two point two asks what happens when metals react with water. Activity three point ten needs the teacher’s assistance. Collect the same metal samples as in Activity three point nine. Put small pieces of the samples separately in beakers half-filled with cold water. Observe which metals reacted with cold water and arrange them in increasing order of reactivity with cold water. Check if any metal produced fire on water or started floating after some time. Put the metals that did not react with cold water in beakers half-filled with hot water. For metals that did not react with hot water, arrange the apparatus as shown in Figure three point three and observe their reaction with steam. In this figure, a metal sample is placed in a boiling tube, which is connected to a delivery tube passing through a water-filled container to collect gas, while the metal is heated. Observe which metals did not react even with steam and arrange the metals in decreasing order of reactivity with water. Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to further form metal hydroxide. But all metals do not react with water. The reactions are Metal plus Water yields Metal oxide plus Hydrogen, and Metal oxide plus Water yields Metal hydroxide. Metals like potassium and sodium react violently with cold water. In case of sodium and potassium, the reaction is so violent and exothermic that the evolved hydrogen immediately catches fire. The equations are 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g) + heat energy, and 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + heat energy. The reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to catch fire. The equation is Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). Calcium starts floating because the bubbles of hydrogen gas formed stick to the surface of the metal. Magnesium does not react with cold water. It reacts with hot water to form magnesium hydroxide and hydrogen. It also starts floating due to the bubbles of hydrogen gas sticking to its surface. Metals like aluminium, iron and zinc do not react either with cold or hot water. But they react with steam to form the metal oxide and hydrogen. The equations are 2Al(s) + 3H₂O(g) → Al₂O₃(s) + 3H₂(g), and 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g). Metals such as lead, copper, silver and gold do not react with water at all.
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Section three point two point three asks what happens when metals react with acids. You have already learnt that metals react with acids to give a salt and hydrogen gas. The reaction is Metal plus Dilute acid yields Salt plus Hydrogen. But do all metals react in the same manner? Let us find out with Activity three point eleven. Collect all the metal samples except sodium and potassium. If tarnished, rub them clean with sand paper. 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. Note which metals reacted vigorously and with which metal you recorded the highest temperature. Arrange the metals in decreasing order of reactivity with dilute acids. Write equations for the reactions of magnesium, aluminium, zinc and iron with dilute hydrochloric acid. 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, which are N₂O, NO, and NO₂. But magnesium and manganese react with very dilute HNO₃ to evolve H₂ gas. You must have observed in Activity three point eleven 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.
Do You Know? Aqua regia, Latin for royal water, is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of three to one. 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.
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Section three point two point four asks how metals react with solutions of other metal salts. Activity three point twelve instructs you to 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, as shown in Figure three point four. In this figure, two test tubes are shown side by side. One contains a copper wire dipped in a greenish iron sulphate solution. The other contains an iron nail dipped in a blue copper sulphate solution. Record your observations after twenty minutes. You will find that a reaction has occurred in the test tube with the iron nail and copper sulphate. The blue colour fades and a brown coating of copper appears on the iron nail. You can say a reaction has taken place because of this colour change and deposition. You can correlate this with Activities three point nine, three point ten, and three point eleven. The balanced chemical equation is Fe + CuSO₄ → FeSO₄ + Cu. This is a displacement reaction. Reactive metals can displace less reactive metals from their compounds in solution or molten form. We have seen that all metals are not equally reactive. We checked reactivity with oxygen, water and acids. But all metals do not react with these reagents. So we were not able to put all metal samples in decreasing order of reactivity. Displacement reactions studied in Chapter one give better evidence. It is simple and easy if metal A displaces metal B from its solution, it is more reactive than B. The reaction is Metal A plus Salt solution of B yields Salt solution of A plus Metal B. According to Activity three point twelve, iron is more reactive than copper.
Section three point two point five introduces 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. Potassium is the most reactive. Next is Sodium, then Calcium, Magnesium, Aluminium, Zinc, Iron, Lead, Hydrogen, Copper, Mercury, Silver, and Gold is the least reactive.
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Let us answer the questions following this section. Question one asks why sodium is kept immersed in kerosene oil. Sodium reacts so vigorously with oxygen and moisture in air that it catches fire. Kerosene oil prevents this contact. Question two asks for equations. Iron with steam yields Fe₃O₄ plus H₂. The equation is 3Fe + 4H₂O → Fe₃O₄ + 4H₂. Calcium with water yields Ca(OH)₂ plus H₂. The equation is Ca + 2H₂O → Ca(OH)₂ + H₂. Potassium with water yields KOH plus H₂ plus heat. The equation is 2K + 2H₂O → 2KOH + H₂ + heat energy. Question three gives a table of metals A, B, C, D reacting with iron(II) sulphate, copper(II) sulphate, zinc sulphate, and silver nitrate. A shows no reaction with iron(II) sulphate, displacement with copper(II) sulphate, no reaction with zinc sulphate, and no reaction with silver nitrate. B shows displacement with iron(II) sulphate, no reaction with copper(II) sulphate, no reaction with zinc sulphate, and no reaction with silver nitrate. C shows no reaction with iron(II) sulphate, no reaction with copper(II) sulphate, no reaction with zinc sulphate, and displacement with silver nitrate. D shows no reaction with any. Which is the most reactive metal? B is the most reactive as it displaces iron. What would you observe if B is added to copper(II) sulphate? B would displace copper, so the blue colour would fade and copper metal would deposit. Arrange A, B, C, D in decreasing reactivity: B, A, C, D. Question four asks which gas is produced when dilute hydrochloric acid is added to a reactive metal. Hydrogen gas is produced. Write the reaction when iron reacts with dilute H₂SO₄. The equation is Fe + H₂SO₄ → FeSO₄ + H₂. Question five asks what you would observe when zinc is added to iron(II) sulphate. Zinc is more reactive, so it will displace iron. The green colour of iron(II) sulphate will fade, and a greyish deposit of iron will form. The equation is Zn + FeSO₄ → ZnSO₄ + Fe.
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Now we move to section three point three, how do metals and non-metals react? In the above activities, you saw the reactions of metals with reagents. Why do metals react this way? Recall the electronic configuration from Class Nine. Noble gases have a completely filled valence shell and show little chemical activity. We explain reactivity as a tendency to attain a completely filled valence shell. Look at Table three point three. A sodium atom has one electron in its outermost shell. If it loses the electron from its M shell, its L shell becomes the outermost shell with a stable octet. The nucleus still has eleven protons but electrons become ten, giving a net positive charge, a sodium cation Na⁺. Chlorine has seven electrons in its outermost shell and requires one more to complete its octet. If sodium and chlorine react, the electron lost by sodium is taken by chlorine. After gaining an electron, chlorine gets a unit negative charge, giving a chloride anion Cl⁻. So both elements have a give-and-take relation. Sodium loses one electron to become Na⁺. Chlorine gains one electron to become Cl⁻. Sodium and chloride ions attract each other and are held by strong electrostatic forces to exist as sodium chloride, NaCl. Sodium chloride does not exist as molecules but as aggregates of oppositely charged ions. Let us see the formation of magnesium chloride. Magnesium loses two electrons to become Mg²⁺. Two chlorine atoms each gain one electron to become Cl⁻. The compounds formed by transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds.
Section three point three point one covers properties of ionic compounds. Perform Activity three point thirteen. Take samples of sodium chloride, potassium iodide, barium chloride or any other salt. Note their physical state; they are solids. Take a small amount on a metal spatula and heat directly on a flame, as shown in Figure three point seven. In this figure, a metal spatula holding a white salt is held over a Bunsen burner flame. Observe if they impart colour to the flame or melt. Try to dissolve them in water, petrol, and kerosene. They are soluble in water but insoluble in petrol and kerosene. Make a circuit as shown in Figure three point eight. In this figure, a battery, bulb, switch, and two electrodes are connected in series, with the electrodes dipped in a beaker of salt solution. Insert electrodes into a salt solution. The bulb glows, indicating conductivity. Your inference is that these are ionic compounds.
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General properties for ionic compounds are as follows. Physical nature: Ionic compounds are solids and are somewhat hard because of the strong force of attraction between positive and negative ions. They are generally brittle and break into pieces when pressure is applied. Melting and boiling points: Ionic compounds have high melting and boiling points. Table three point four shows NaCl melts at 1074 Kelvin and boils at 1686 Kelvin. LiCl melts at 887 Kelvin and boils at 1600 Kelvin. CaCl₂ melts at 1045 Kelvin and boils at 1900 Kelvin. CaO melts at 2850 Kelvin and boils at 3120 Kelvin. MgCl₂ melts at 981 Kelvin and boils at 1685 Kelvin. This is because considerable energy is required to break strong inter-ionic attraction. Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene and petrol. Conduction of electricity: Conduction through a solution involves movement of charged particles. An ionic compound solution contains ions that move to opposite electrodes when electricity is passed. Ionic compounds in solid state do not conduct electricity because ion movement is not possible due to rigid structure. But they conduct electricity in molten state because electrostatic forces are overcome by heat, allowing ions to move freely.
Let us answer the questions. Question one part one asks for electron-dot structures of sodium, oxygen, and magnesium. Sodium has one dot. Oxygen has six dots. Magnesium has two dots. Part two asks to show formation of Na₂O and MgO by electron transfer. Two sodium atoms each lose one electron to one oxygen atom, forming Na⁺ and O²⁻ ions. One magnesium atom loses two electrons to one oxygen atom, forming Mg²⁺ and O²⁻ ions. Part three asks for ions present. Na₂O contains Na⁺ and O²⁻. MgO contains Mg²⁺ and O²⁻. Question two asks why ionic compounds have high melting points. It is because a considerable amount of energy is required to break the strong inter-ionic attraction between the positive and negative ions.
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Now we cover section three point four, occurrence of metals. The earth’s crust is the major source of metals. Seawater also contains soluble salts like sodium chloride and magnesium chloride. Elements or compounds occurring naturally in the earth’s crust are known as minerals. At some places, minerals contain a high percentage of a particular metal that can be profitably extracted. These minerals are called ores.
Section three point four point one covers extraction of metals. You have learnt the reactivity series. Having this knowledge, you can understand metal extraction. Some metals are found free in the crust. Some are in compounds. Metals at the bottom of the activity series are least reactive and often found free, like gold, silver, platinum, and copper. Copper and silver are also found combined as sulphide or oxide ores. Metals at the top, K, Na, Ca, Mg, Al, are so reactive they are never free. Metals in the middle, Zn, Fe, Pb, are moderately reactive and found as oxides, sulphides, or carbonates. Many ores are oxides because oxygen is reactive and abundant. Based on reactivity, metals are grouped into metals of low reactivity, medium reactivity, and high reactivity. Figure three point nine shows this grouping. K, Na, Ca, Mg, Al require electrolysis. Zn, Fe, Pb, Cu require reduction using carbon. Ag, Au are found in native state. Figure three point ten shows steps for extraction: enrichment of ores, extraction of metal from enriched ore, and refining of metal.
Section three point four point two covers enrichment of ores. Ores are contaminated with impurities like soil and sand, called gangue. Impurities must be removed before extraction. Processes used are based on physical or chemical property differences.
Section three point four point three covers extracting metals low in the activity series. These are very unreactive. Their oxides can be reduced by heating alone. Cinnabar, HgS, is an ore of mercury. Heated in air, it converts to mercuric oxide, HgO. Mercuric oxide reduces to mercury on further heating. The equations are 2HgS(s) + 3O₂(g) → 2HgO(s) + 2SO₂(g), followed by 2HgO(s) → 2Hg(l) + O₂(g). Copper found as Cu₂S is obtained by heating in air. 2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂, followed by 2Cu₂O + Cu₂S → 6Cu(s) + SO₂(g).
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Section three point four point four covers extracting metals in the middle of the series, like iron, zinc, lead, copper. They are moderately reactive, usually as sulphides or carbonates. It is easier to obtain metal from oxide. Prior to reduction, sulphides and carbonates must become oxides. Sulphide ores are converted to oxides by heating strongly in excess air, known as roasting. Carbonate ores are changed to oxides by heating strongly in limited air, known as calcination. For zinc ores: Roasting is 2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g). Calcination is ZnCO₃(s) → ZnO(s) + CO₂(g). Metal oxides are reduced using carbon. For example, ZnO(s) + C(s) → Zn(s) + CO(g). Obtaining metals from compounds is a reduction process. Besides carbon, displacement reactions use highly reactive metals like sodium, calcium, aluminium as reducing agents. For example, 3MnO₂(s) + 4Al(s) → 3Mn(l) + 2Al₂O₃(s) + Heat. These are highly exothermic, producing molten metals. The reaction of iron(III) oxide with aluminium joins railway tracks. This is the thermit reaction. Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat. Figure three point eleven shows the thermit process for joining railway tracks, where a crucible containing the mixture is placed over the track joint and ignited, producing molten iron that fills the gap.
Section three point four point five covers extracting metals at the top of the series. They are very reactive and cannot be reduced by carbon. They are obtained by electrolytic reduction. Sodium, magnesium, calcium are obtained by electrolysis of molten chlorides. Metals deposit at the cathode, chlorine liberates at the anode. Reactions: At cathode, Na⁺ + e⁻ → Na. At anode, 2Cl⁻ → Cl₂ + 2e⁻. Aluminium is obtained by electrolytic reduction of aluminium oxide.
Section three point four point six covers refining of metals. Metals produced are impure. The most widely used method is electrolytic refining. Many metals like copper, zinc, tin, nickel, silver, gold are refined electrolytically. The impure metal is the anode. A thin strip of pure metal is the cathode. A solution of the metal salt is the electrolyte. Figure three point twelve shows electrolytic refining of copper. The electrolyte is acidified copper sulphate. The anode is impure copper, the cathode is pure copper. Passing current dissolves pure metal from the anode into the electrolyte. An equivalent amount deposits on the cathode. Soluble impurities go into solution. Insoluble impurities settle at the anode bottom as anode mud.
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Let us answer the questions. Question one defines mineral, ore, gangue. Mineral is an element or compound occurring naturally in the earth’s crust. Ore is a mineral from which a metal can be profitably extracted. Gangue is the impurity like soil or sand mixed with the ore. Question two asks for two metals found free in nature. Gold and silver. Question three asks for the chemical process to obtain metal from its oxide. Reduction.
Now we perform Activity three point fourteen. Take three test tubes and place clean iron nails in each. Label them A, B, C. Pour water in A and cork it. Pour boiled distilled water in B, add one millilitre of oil, and cork it. The oil prevents air from dissolving. Put anhydrous calcium chloride in C and cork it. It absorbs moisture from air. Leave for a few days and observe, as shown in Figure three point thirteen. In this figure, tube A shows rusted nails in water. Tube B shows unrusted nails in water with an oil layer. Tube C shows unrusted nails in dry air. You will observe iron nails rust in A, but not in B and C. In A, nails are exposed to both air and water. In B, only water. In C, only dry air. This tells us that iron rusts only in the presence of both air and water.
Section three point five covers corrosion. You have learnt that silver articles become black due to silver sulphide coating. Copper reacts with moist carbon dioxide to form a green coat of basic copper carbonate. Iron exposed to moist air forms a brown flaky substance called rust.
Do You Know? Pure gold, twenty four carat gold, is very soft and not suitable for jewellery. It is alloyed with silver or copper to make it hard. In India, twenty two carat gold is used, meaning twenty two parts pure gold alloyed with two parts copper or silver. If one metal is mercury, the alloy is an amalgam. Alloy electrical conductivity and melting point are less than pure metals. Brass, copper and zinc, and bronze, copper and tin, are poor conductors. Solder, lead and tin, has a low melting point and is used for welding electrical wires.
Section three point five point one covers prevention of corrosion. Rusting of iron can be prevented by painting, oiling, greasing, galvanising, chrome plating, anodising, or making alloys. Galvanisation protects steel and iron from rusting by coating with zinc. It protects even if the zinc coating is broken because zinc is more reactive and oxidises first. Alloying improves properties. Iron mixed with zero point zero five percent carbon becomes hard and strong. Iron mixed with nickel and chromium gives stainless steel, which is hard and does not rust. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal. It is prepared by melting the primary metal, dissolving other elements in definite proportions, and cooling.
The Iron pillar at Delhi is a wonder of ancient Indian metallurgy. Built over one thousand six hundred years ago, it is eight metres high and weighs six tonnes. It resists rusting due to an ancient process.
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Let us answer the questions. Question one asks in which cases displacement reactions take place. They take place when a more reactive metal is added to a salt solution of a less reactive metal. Question two asks which metals do not corrode easily. Gold, silver, and platinum. Question three asks what are alloys. An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal.
Now we cover what you have learnt. Elements can be classified as metals and non-metals. Metals are lustrous, malleable, ductile, good conductors of heat and electricity, solids at room temperature except mercury. Metals form positive ions by losing electrons. Metals combine with oxygen to form basic oxides. Aluminium oxide and zinc oxide are amphoteric oxides. Different metals have different reactivities with water and dilute acids. The activity series lists metals in decreasing reactivity. Metals above hydrogen displace hydrogen from dilute acids. A more reactive metal displaces a less reactive metal from its salt solution. Metals occur as free elements or compounds. Metallurgy is extraction and refining. An alloy is a homogeneous mixture. Corrosion occurs when metals like iron are exposed to moist air. Non-metals have opposite properties. They are not malleable or ductile. They are bad conductors except graphite. Non-metals form negative ions by gaining electrons. Non-metals form acidic or neutral oxides. Non-metals do not displace hydrogen from dilute acids but react with hydrogen to form hydrides.
Now we will solve the exercises completely.
Exercise one asks which pair gives displacement reactions. Option a is NaCl and copper. Copper is less reactive than sodium, so no reaction. Option b is MgCl₂ and aluminium. Aluminium is less reactive than magnesium, so no reaction. Option c is FeSO₄ and silver. Silver is less reactive than iron, so no reaction. Option d is AgNO₃ and copper. Copper is more reactive than silver, so it displaces silver. The correct answer is d.
Exercise two asks which method prevents an iron frying pan from rusting. Applying grease, paint, or zinc coating all work. The correct answer is d, all of the above.
Exercise three describes an element reacting with oxygen to give a high melting point compound soluble in water. Calcium oxide has a high melting point and dissolves in water to form calcium hydroxide. The correct answer is a, calcium.
Exercise four asks why food cans are coated with tin and not zinc. Zinc is more reactive than tin, so it would react with food. The correct answer is c, zinc is more reactive than tin.
Exercise five gives a hammer, battery, bulb, wires, switch. Part a asks how to distinguish metals from non-metals. Strike the sample with a hammer. If it flattens into a sheet, it is malleable and likely a metal. If it shatters, it is brittle and likely a non-metal. Next, connect the sample in the circuit with the battery, bulb, and switch. If the bulb glows, it conducts electricity and is a metal. If not, it is a non-metal. Part b asks to assess usefulness. These tests are useful but not absolute. Graphite conducts electricity but is a non-metal. Mercury is a metal but is liquid and cannot be hammered into sheets.
Exercise six asks what amphoteric oxides are and for two examples. Amphoteric oxides are metal oxides which react with both acids as well as bases to produce salts and water. Examples are aluminium oxide and zinc oxide.
Exercise seven asks for two metals that displace hydrogen from dilute acids, and two that do not. Metals above hydrogen in the activity series displace hydrogen, like zinc and iron. Metals below hydrogen do not, like copper and silver.
Exercise eight asks what to take as anode, cathode, and electrolyte in electrolytic refining of metal M. The anode is impure metal M. The cathode is a thin strip of pure metal M. The electrolyte is a solution of a salt of metal M.
Exercise nine describes Pratyush heating sulphur powder and collecting gas. Part a asks for action on dry and moist litmus paper. The gas is sulphur dioxide. On dry litmus paper, it has no action. On moist litmus paper, it turns blue litmus red because it forms sulphurous acid. Part b asks for the balanced equation. S + O₂ → SO₂.
Exercise ten asks for two ways to prevent rusting of iron. Painting or oiling the surface, and galvanising with a zinc coating.
Exercise eleven asks what type of oxides are formed when non-metals combine with oxygen. Acidic or neutral oxides are formed.
Exercise twelve asks for reasons. Part a: Platinum, gold, silver are used for jewellery because they are highly lustrous, do not corrode easily, and are malleable. Part b: Sodium, potassium, lithium are stored under oil because they are highly reactive with air and moisture and would catch fire otherwise. Part c: Aluminium is highly reactive but forms a protective oxide layer when exposed to air, which prevents further reaction, making it safe for cooking utensils. Part d: Carbonate and sulphide ores are converted to oxides because it is easier to reduce oxides to metals using carbon or other reducing agents.
Exercise thirteen asks why tarnished copper vessels are cleaned with lemon or tamarind juice. The tarnish is basic copper carbonate. Lemon and tamarind juice contain acids which react with the basic copper carbonate to form soluble salts, removing the green layer and restoring shine.
Exercise fourteen asks to differentiate metal and non-metal based on chemical properties. Metals form basic oxides, displace hydrogen from dilute acids, and form positive ions by losing electrons. Non-metals form acidic or neutral oxides, do not displace hydrogen from dilute acids, and form negative ions by gaining electrons.
Exercise fifteen describes a man dipping gold bangles in a solution, making them sparkle but reducing weight drastically. The solution used is aqua regia, a mixture of concentrated hydrochloric acid and concentrated nitric acid in a three to one ratio. It dissolves gold, removing the tarnished outer layer and reducing weight.
Exercise sixteen asks why copper is used for hot water tanks and not steel. Copper does not react with water or steam, whereas iron in steel reacts with steam to form iron oxide and hydrogen, causing corrosion. Copper is also a better conductor of heat and does not rust.
[CHECKPOINT]
Thank you for listening! Keep revising and practicing. Goodbye! [CHAPTER_COMPLETE]