Metals and Non-metals, Study notes of Chemistry

Download Metals and Non-metals and more Study notes Chemistry in PDF only on Docsity! Metals and Non-metals 3CHAPTER In Class IX 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. 3.1 PHYSICAL PROPERTIES 3.1.1 Metals The easiest way to start grouping substances is by comparing their physical properties. Let us study this with the help of the following activities. For Activities 3.1 to 3.6, collect the samples of following metals – iron, copper, aluminium, magnesium, sodium, lead, zinc and any other metal that is easily available. Activity 3.1 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. Metals, in their pure state, have a shining surface. This property is called metallic lustre. 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. Hold a piece of sodium metal with a pair of tongs. CAUTION: 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? Science38 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? 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? Repeat with other metals. Record the change in the shape of these metals. You will find that metals are generally hard. The hardness varies from metal to metal. Activity 3.4 Consider some metals such as iron, copper, aluminium, lead, etc. Which of the above metals are also available in the form of wires? 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 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. Activity 3.5 Take an aluminium or copper wire. Clamp this wire on a stand, as shown in Fig. 3.1. 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. What do you observe after some time? Note your observations. Does the metal wire melt? The above 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. Figure 3.1 Metals are good conductors of heat. Metals and Non-metals 41 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 the following Activity. 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 samples taken above 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? Almost all metals combine with oxygen to form metal oxides. Metal + Oxygen → Metal oxide For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide. 2Cu + O 2 → 2CuO (Copper) (Copper(II) oxide) Similarly, aluminium forms aluminium oxide. 4Al + 3O2 → 2Al2O3 (Aluminium) (Aluminium oxide) 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, zinc oxide, etc., 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 Al 2 O 3 + 2NaOH → 2NaAlO 2 + H 2 O (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) K 2 O(s) + H 2 O(l) → 2KOH(aq) Science42 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. D o Y o u K no w ? 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 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. 3.2.2 What happens when Metals react with Water? 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 Fig. 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. Metals and Non-metals 43 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. Metal + Water → Metal oxide + Hydrogen Metal oxide + Water → 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. 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) + heat energy 2Na(s) + 2H 2 O(l) → 2NaOH(aq) + H 2 (g) + heat energy The reaction of calcium with water is less violent. The heat evolved is not sufficient for the hydrogen to catch fire. Ca(s) + 2H 2 O(l) → Ca(OH) 2 (aq) + H 2 (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. 2Al(s) + 3H2O(g) → Al2O3(s) + 3H2(g) 3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g) Metals such as lead, copper, silver and gold do not react with water at all. 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. Figure 3.3 Action of steam on a metal Science46 ? Q U E S T I O N S 1. Why is sodium kept immersed in kerosene oil? 2. Write equations for the reactions of (i) iron with steam (ii) calcium and potassium with water 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. 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. We can see from Table 3.3 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 Use the Table above to answer the following questions about metals A, B, C and D. (i) Which is the most reactive metal? (ii) What would you observe if B is added to a solution of Copper(II) sulphate? (iii) Arrange the metals A, B, C and D in the order of decreasing reactivity. 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. 5. What would you observe when zinc is added to a solution of iron(II) sulphate? Write the chemical reaction that takes place. 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 Metals and Non-metals 47 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 C1–. So both these elements can have a give-and-take relation between them as follows (Fig. 3.5). Na Na + e 2,8,1 2,8 + (Sodium cation) → – Cl +e Cl 2,8,7 2,8,8 (Chloride anion) – –→ Figure 3.5 Formation of sodium chloride 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 (Fig. 3.6). Table 3.3 Electronic configuration of some elements Type of Element Atomic Number of element number electrons in shells K L M N Noble gases Helium (He) 2 2 Neon (Ne) 10 2 8 Argon (Ar) 18 2 8 8 Metals Sodium (Na) 11 2 8 1 Magnesium (Mg) 12 2 8 2 Aluminium (Al) 13 2 8 3 Potassium (K) 19 2 8 8 1 Calcium (Ca) 20 2 8 8 2 Non-metals Nitrogen (N) 7 2 5 Oxygen (O) 8 2 6 Fluorine (F) 9 2 7 Phosphorus (P) 15 2 8 5 Sulphur (S) 16 2 8 6 Chlorine (Cl) 17 2 8 7 Science48 Mg Mg e2+ (Magnesium cation) ⎯ →⎯ + 2 2 8 2 2 8 – , , , Cl +e Cl 2,8,7 2,8,8 (Chloride anion) – –⎯ →⎯ Figure 3.6 Formation of magnesium chloride 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. Can you name the cation and anion present in MgCl2? 3.3.1 Properties of Ionic Compounds To learn about the properties of ionic compounds, let us perform the following Activity: 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? Take a small amount of a sample on a metal spatula and heat directly on the flame (Fig. 3.7). Repeat with other samples. What did you observe? Did the samples impart any colour to the flame? Do these compounds melt? Try to dissolve the samples in water, petrol and kerosene. Are they soluble? Make a circuit as shown in Fig. 3.8 and insert the electrodes into a solution of one salt. What did you observe? Test the other salt samples too in this manner. What is your inference about the nature of these compounds? Figure 3.7 Heating a salt sample on a spatula Figure 3.8 Testing the conductivity of a salt solution Table 3.4 Melting and boiling points of some ionic compounds Ionic Melting point Boiling point compound (K) (K) NaCl 1074 1686 LiCl 887 1600 CaCl2 1045 1900 CaO 2850 3120 MgCl2 981 1685 Metals and Non-metals 51 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. 3.4.3 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) + 3O (g) 2HgO(s) + 2SO (g)2 2 Heat⎯ →⎯⎯⎯ 2HgO(s) 2Hg(l) + O (g)2 Heat⎯ →⎯⎯⎯ Similarly, copper which is found as Cu 2 S in nature can be obtained from its ore by just heating in air. 2Cu S + 3O (g) 2Cu O(s) + 2SO (g) 2Cu O + Cu S 2 2 2 2 2 2 Heat Heat ⎯ →⎯⎯⎯ ⎯ →⎯⎯⎯ 6Cu(s) + SO (g)2 3.4.4 Extracting Metals in the Middle of the Activity Series The metals in the middle of the activity series such as iron, zinc, lead, copper, etc., 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) + 3O (g) 2ZnO(s) + 2SO (g)2 2 Heat⎯ →⎯⎯⎯ Calcination ZnCO (s) ZnO(s) + CO (g)3 2 Heat⎯ →⎯⎯⎯ 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 Science52 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 – 3MnO 2 (s) + 4Al(s) → 3Mn(l) + 2Al 2 O 3 (s) + Heat Can you identify the substances that are getting oxidised and reduced? 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 (Fe 2 O 3 ) 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 3.4.5 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. 3.4.6 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. The apparatus is set up as shown in Fig. 3.12. On passing the current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure Figure 3.11 Thermit process for joining railway tracks Figure 3.12 Electrolytic refining of copper. The electrolyte is a solution of acidified copper sulphate. The anode is impure copper, whereas, the cathode is a strip of pure copper. On passing electric current, pure copper is deposited on the cathode. Metals and Non-metals 53 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. 1. Define the following terms. (i) Mineral (ii) Ore (iii) Gangue 2. Name two metals which are found in nature in the free state. 3. What chemical process is used for obtaining a metal from its oxide? 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 (Fig. 3.13). Figure 3.13 Investigating the conditions under which iron rusts. In tube A, both air and water are present. In tube B, there is no air dissolved in the water. In tube C, the air is dry. 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? Q U E S T I O N S ? 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 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. A B C Science56 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. E X E R C I S E S 1. Which of the following pairs will give displacement reactions? (a) NaCl solution and copper metal (b) MgCl2 solution and aluminium metal (c) FeSO 4 solution and silver metal (d) AgNO3 solution and copper metal. 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. 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. 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. 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. 6. What are amphoteric oxides? Give two examples of amphoteric oxides. 7. Name two metals which will displace hydrogen from dilute acids, and two metals which will not. Metals and Non-metals 57 8. In the electrolytic refining of a metal M, what would you take as the anode, the cathode and the electrolyte? 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. 10. State two ways to prevent the rusting of iron. 11. What type of oxides are formed when non-metals combine with oxygen? 12. Give reasons (a) Platinum, gold and silver are used to make jewellery. (b) Sodium, potassium and lithium are stored under oil. (c) Aluminium is a highly reactive metal, yet it is used to make utensils for cooking. (d) Carbonate and sulphide ores are usually converted into oxides during the process of extraction. 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. 14. Differentiate between metal and non-metal on the basis of their chemical properties. 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? 16. Give reasons why copper is used to make hot water tanks and not steel (an alloy of iron). Collection of gas