Metals and Non-metals PDF

3 Metals and Non-metals ◼ Introduction Everything around us is made up of different elements. These elements can be classified into metals or non-metals on the basis of their properties. A few elements have properties common to both metals and non-metals. Thes...

3 Metals and Non-metals ◼ Introduction Everything around us is made up of different elements. These elements can be classified into metals or non-metals on the basis of their properties. A few elements have properties common to both metals and non-metals. These are called semi-metals or metalloids. Metal are elements which are hard, lustrous, malleable, ductile and possess good electrical and thermal conductivity. Whereas, non-metals do not possess lustre, are bad conductors of heat and electricity, are not malleable and ductlie but are brittle. Some uses of metals (i) In construction of buildings and bridges. (ii) In making of coins. (iii) In making of utensils and kitchen wares. (iv) In making of jewellery. (v) In making of machine parts, automobiles, etc. Metal currency Some uses of non-metals (i) Oxygen is used for respiration by living things. (ii) Nitrogen is the main constituent of fertilizers. (iii) The food we eat consist of many non-metals like carbon, hydrogen, oxygen, nitrogen, sulphur, etc. (iv) Sulphur is used as a fungicide and used for making gun powder. (v) Carbon, in the form of graphite, is used as electrodes in electrolytic cells and dry cells. In the body : Chromium assists in the metabolism of sugar, cobalt is present in vitamin B 12, iodine is necessary for the proper functioning of the thyroid gland, manganese plays a role in maintaining proper calcium levels in the bones, and copper is involved in the production of red blood cells. Found. JEE-NEET : Class 8 ◼ Physical properties (I) Metals (1) Lustre of metals Most of the metals, in their pure state, have a shining surface. This property is called metallic lustre. For example, gold is shining yellow, copper is brown, iron, aluminium and zinc are lustrous grey. 1 Why do metals possess lustre? Explanation When light falls on the surface of a metal, the atoms absorb photons as energy. They get excited and start vibrating. These vibrating electrons release energy in the form of light. Therefore, metal surface shines and metals possess lustre. (2) Hardness of metals Most of the metals are hard, but all metals are not equally hard. The hardness of metal varies from metal to metal. (3) Malleability of metals The property according to which metals can be beaten with a hammer into very thin sheets without breaking is called malleability. Gold and silver are the most malleable metals. Aluminium and copper are also highly malleable metals. All of these metals can be beaten with a hammer to form Being highly malleable, silver foil is used for very thin sheets, called foil. wrapping sweets and (4) Ductility of metals chocolates. Ductility is also an important property of metals. The ability of metals to be drawn (stretched) into thin wires is called ductility. Generally, wires are made up of iron, copper and aluminium. For example, 100 mg of silver can be drawn into a thin wire of about 200 metres length. Copper and aluminium are also very ductile, and therefore, they can be drawn into thin wires which are used in electrical wiring. (5) Thermal conductivity of metals The process in which a metal allows the flow of heat through it is called its thermal conductivity. Most of the metals are good conductors of heat, such as silver, gold, iron, copper and aluminium.   Chemistry 2 How metals conduct heat? Explanation Metals conduct heat by the process of conduction. When we heat a piece of metal, metal atoms gain energy and start vibrating. During vibration, they transfer their energy to neighbouring atom and the neighbouring atom repeats the same and pass on its energy to the next atom. In this way, transfer of heat energy takes place. (6) Electrical conductivity of metals The property in which metal facilitates the flow of electric current through it is called electrical conductivity. All metals are good conductors of electricity because they contain free or mobile electrons. These free electrons conduct electric current. (7) Sonorous The property of metals in which metals produce sound when they strike a hard object or other surface is called sonority or sonorisity. Some metals like copper, silver, gold, aluminium give musical sound when they are struck by themselves or any other object. Memory map Sonorous Hard Lusture Thermal Electrical METALS conductance conductance Ductile Malleable High M.P. & B.P. Note: Difference in properties of different materials is based on their internal structures. (II) Non-metals Among the total known elements, there are only 22 non-metals, out of which 11 are gases like oxygen, nitrogen, hydrogen, one is a liquid (bromine) and the rest 10 are solids such as sulphur, phosphorus, iodine and the allotropes of carbon (diamond and graphite). Do non-metals also have physical properties similar to that of metals? The important physical properties of non-metals are given below : (1) Non-metals may be solids (such as sulphur, phosphorus and diamond), liquid (bromine), or gases (such as oxygen, nitrogen, hydrogen, neon, argon, etc.) at room temperature. (2) Non-metals are usually brittle and cannot be used to make sheets and wires. (3) Non-metals are non-lustrous and cannot be polished. (Exception : Graphite and iodine are lustrous non-metals). Found. JEE-NEET : Class 8 (4) Non-metals are generally bad conductor of heat and electricity. Exception : Graphite is a good conductor of electricity. Non-metals do not conduct the electric current due to absence of mobile electrons. (5) Non-metals can be easily broken due to their low tensile strength. (6) Non-metals are generally light and have low densities. (7) Unlike metals, non-metals do not produce any ringing sound when struck with an object. (8) Non-metals are soft (Exception : Diamond) (9) Non-metals have low melting and boiling points. (Exception : Graphite has very high melting point (3730°C)) On the basis of the above discussion of the physical properties of metals and non-metals, we have concluded that elements cannot be grouped according to the physical properties alone, as there are many exceptions. 1 Q.1 Name one metal and one non-metal that exist in liquid state at room temperature. Also name two metals having melting point less than 310 K (37°C). Ans. Metal mercury (Hg) and non-metal bromine (Br) exist in liquid state at Gallium metal has room temperature. The two metals with melting point less than 310 K such a low melting are cesium (Cs) and gallium (Ga). point (30°C) that it Q.2 A non-metal A is an important constituent of our food and forms melts with the heat of a hand. two oxides B and C. Oxide B is neutral whereas C causes global warming. Identify A, B and C. Ans. The non-metal A is carbon. B is carbon monoxide (CO), C is carbon dioxide (CO2) Some exceptions (i) All metals except mercury are solid at room temperature. We know that metals have very high melting points but gallium (Ga) and cesium (Cs) have very low melting points. These two metals will melt if we keep them on our palm. (ii) Iodine is a non-metal but it is lustrous. (iii) Alkali metals such as lithium, sodium and potassium are so soft, that they can be easily cut with a knife i.e. they have low densities and low melting points. (iv) 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, which has very high melting and boiling point. Graphite is another allotrope of carbon which is good conductor of electricity.   Chemistry Memory map Low M.P. & B.P. Brittle Soft Low density Non-Metals Non-lustrous Non-conductor Non sonorous 2 Q.1 A non-metal A is an important constituent of our food and forms two oxides B and C. Oxide B is neutral whereas C causes global warming. Identify A, B and C. Ans. The non-metal A is carbon. B is carbon monoxide (CO), C is carbon dioxide (CO2) Elements can be more clearly classified as metals and non-metals on the basis of their chemical properties. ◼ Chemical properties of metals We have studied the physical characteristics of metals. Now let us focus our attention on their chemical properties. Metals in general have tendency to lose one or more electrons present in the valence shells of their atoms to form positive ions. Metals are therefore, regarded as electropositive elements. M ⎯→ Mn+ + ne– (Metal atom) The chemical properties of the metals are mostly linked with the electron releasing tendency of their atoms. Greater the tendency, more will be the reactivity of the metal. (1) Reaction of metal with oxygen Almost all metals combine with oxygen to form metal oxides. But they possess different reactivity towards oxygen. Almost all metals combine with oxygen to form metal oxides. Metal + Oxygen ⎯→ Metal oxide Sodium and potassium react vigorously with oxygen. 4Na(s) + O2(g) ⎯→ 2Na2O(s) 4K(s) + O2(g) ⎯→ 2K2O(s) Sodium and potassium burns with a golden yellow and lilac colour flame respectively to form sodium and potassium oxides, which dissolve with water to form alkali called sodium hydroxide and potassium hydroxide. Na2O(s) + H2O() ⎯→ 2NaOH(aq) K2O(s) + H2O() ⎯→ 2KOH(aq) Found. JEE-NEET : Class 8 Magnesium also burns easily, to form magnesium oxide 2Mg(s) + O2(g) ⎯→ 2MgO(s) Copper and Aluminium do not burn but on heating in air form black copper (II) oxide and white aluminium oxide (Al2O3) respectively. 2Cu(s) + O2(g) ⎯→ 2CuO(s) Copper Copper (II) oxide (Black) 4Al(s) + 3O2(g) ⎯→ 2Al2O3(s) Aluminium Aluminium oxide (White) These metal oxides are found to be insoluble in water. The order of reactivity with oxygen is : K > Na > Mg > Al > Cu At ordinary temperature, the surfaces of metals such as magnesium, zinc and lead, etc. are covered with a thin layer of the oxide. The protective layer of the oxide prevents the metal from further oxidation. Nature of metallic oxide Generally, metallic oxides are basic in nature except aluminium and zinc oxides which are amphoteric in nature. That means these oxides (Al2O3, ZnO) react with base as well as acid. The basic oxide of metals react with acid to give salt. For example, CuO(s) + H2SO4() ⎯→ CuSO4(aq) + H2O() Copper(II) Sulphuric Copper(II) Water oxide acid sulphate Some oxides of metals dissolve in water and form alkalis. For example, Na2O(s) + H2O() ⎯→ 2NaOH(aq) Sodium oxide Water Sodium hydroxide K2O(s) + H2O() ⎯→ 2KOH (aq) Potassium oxide Water Potassium hydroxide Reactions showing amphoteric nature of Al2O3 and ZnO. Al2O3(s) + 6HCl(aq) ⎯→ 2AlCl3(aq) + 3H2O() Aluminium Hydrochloric Aluminium Water oxide acid chloride Al2O3(s) + 2NaOH(aq) ⎯→ 2NaAlO2(aq) + H2O() Aluminium Sodium hydroxide Sodium meta Water oxide (base) aluminate   Chemistry Similarly, ZnO(s) + 2HCl(aq) ⎯→ ZnCl2(aq) + H2O() Zinc Hydrochloric Zinc Water oxide acid chloride ZnO(s) + 2NaOH(aq) ⎯→ Na2ZnO2(aq) + H2O() Zinc oxide Sodium hydroxide Sodium zincate Water Anodising of aluminium Anodising is a process of forming a thick oxide layer of aluminium. Aluminium develops a thin oxide layer when it exposed to air. This oxide coat of aluminium (Al) makes it resistant to further corrosion. During anodising, the resistance can be improved further by making the oxide layer thicker. In this process, a clean Al article is made the anode and dilute sulphuric acid (H2SO4) is used for electrolyte. The oxygen gas evolved at the anode react with Al to make a thicker protective oxide layer. This oxide layer can be dyed easily to give Al articles an attractive finishing. (2) Reaction of metals with water Metal reacts with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolve in it to form metal hydroxide. But all metals do not react with water. Metal + Water ⎯⎯→ Metal oxide + Hydrogen gas Metal oxide + Water ⎯⎯→ Metal hydroxide (i) Na and K metals react vigorously with cold water to form NaOH and KOH respectively and H2 gas is liberated. 2Na(s) + 2H2O() ⎯⎯→ 2NaOH(aq) + H2(g) Sodium Cold water Sodium hydroxide Hydrogen gas (ii) 2K(s) + 2H2O() ⎯⎯→ 2KOH(aq) + H2(g) Potassium Cold water Potassium hydroxide Hydrogen gas These reactions are so violent and exothermic that the H2 gas evolved, catches fire. (iii) Calcium reacts with cold water to form Ca(OH)2 and H2 gas. It is less violent. Ca(s) + 2H2O() ⎯⎯→ Ca(OH)2(aq) + H2(g) Calcium Cold water Calcium hydroxide Hydrogen gas (iv) Magnesium reacts with hot boiling water to form Mg(OH)2 and H2 gas. Mg(s) + H2O() ⎯⎯→ Mg(OH)2(aq) + H2(g) Magnesium Boiling water Magnesium hydroxide Hydrogen gas (v) Aluminium does not react either with cold or hot water. But it reacts only with steam to form aluminium oxide and hydrogen gas. 2Al(s) + 3H2O(g) ⎯⎯→ Al2O3(s) + 3H2(g) Aluminium Steam Aluminium oxide Hydrogen gas Found. JEE-NEET : Class 8 (vi) Similarly, zinc reacts with steam to form zinc oxide and H2 gas. Zn(s) + H2O(g) ⎯⎯→ ZnO(s) + H2(g) Zinc Steam Zinc oxide Hydrogen gas (vii) Copper do not react with water even under strong conditions. The above reactions indicate that sodium and potassium are the most reactive metals while copper is less reactive. The reactivity order of these metals with water are K > Na > Ca > Mg > Al > Zn > Fe > Cu Reactivity with water decreases ⎯→ (a) (b) (c) The reaction (a) Potassium metal (stored in mineral oil to prevent oxidation) and (b) water. (c) The reaction of potassium with water. The flame occurs because of the produced hydrogen gas. (H2(g) burns in air which reacts with O2(g), at the high temperature) (3) Reaction of metals with acids The highly reactive metals react with dilute acid to displace hydrogen from acid and give a salt. Metal + Dilute acid ⎯⎯→ Salt + Hydrogen In the test tube which contains Mg, the hydrogen bubbles appear very rapidly. Mg(s) + 2HCl(aq) ⎯⎯→ MgCl2(aq) + H2(g) Magnesium Hydrochloric acid Magnesium chloride Hydrogen In the test tubes containing Al and Zn reaction with acid is fast. 2Al(s) + 6HCl(aq) ⎯⎯→ 2AlCl3(aq) + 3H2(g) Aluminium Hydrochloric acid Aluminium chloride Hydrogen The reaction between Fe and acid is slow. Fe(s) + 2HCl(aq) ⎯⎯→ FeCl2(aq) + H2(g) Iron Hydrochloric acid Ferrous chloride Hydrogen No reaction is observed in the test tube which contain Cu and dil HCl. Cu(s) + HCl(aq) ⎯⎯→ No reaction Copper Hydrochloric acid Temperature was found to rise in case of all the metals that reacted with dilute acid showing that reaction is exothermic. The rise in temperature is maximum in case of magnesium.   Chemistry Important information Hydrogen gas is not evolved when metals such as Zn, Fe, Cu and Al react with nitric acid. Because HNO3 is strong oxidising agent. It oxidises, H2 gas to water and is itself reduced to oxides of nitrogen (NO, N2O and NO2). 3Fe(s) + 8HNO3(aq) ⎯→ 3Fe(NO3)2(aq) + 4H2O() + 2NO(g) Iron Nitric acid (dil) Iron(II) nitrate Water Nitric oxide 3Cu(s) + 8HNO3(aq) ⎯→ 3Cu(NO3)2(aq) + 4H2O() + 2NO(g) Copper Nitric acid Copper nitrate Water Nitric oxide But copper reacts with hot concentrated sulphuric acid (H2SO4) to produce copper sulphate, sulphur dioxide and water. Cu(s) + 2H2SO4(aq) ⎯→ CuSO4(aq) + SO2(g) + 2H2O() Copper Sulphuric acid Copper sulphate Sulphur dioxide Water Mg reacts with very dilute HNO3 to evolve H2 gas. Mg(s) + 2HNO3(aq) ⎯→ Mg(NO3)2(aq) + H2(g) Magnesium Nitric acid (dil) Magnesium nitrate Hydrogen Fe reacts with dil H2SO4 to evolve H2. Fe(s) + dil H2SO4 ⎯→ FeSO4(aq) + H2(g) Iron Sulphuric acid Ferrous sulphate Hydrogen Aqua Regia (Royal water): Aqua regia is a Latin word it means "royal water". It is a freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid in the ratio of 3 : 1. It is a highly corrosive, fuming liquid and is used to dissolve gold and platinum. (4) Reaction of metal with solutions of other metal salts When a more reactive metal is placed in a salt solution of less reactive metal, then the more reactive metal displaces the less reactive metal from its salt solution. This reaction is also known as displacement reaction. For example, Reaction is found to occur in the tube containing iron nail dipped in copper sulphate solution. This is because in this tube, blue colour of copper sulphate solution fades and light green colour due to formation of iron (II) sulphate appears. Moreover, a brown deposit of copper takes place on iron nail. Thus, the following reaction takes place. Fe(s) + CuSO4(aq) ⎯→ FeSO4(aq) + Cu(s) Iron Copper sulphate Iron (II) sulphate Copper (Blue) (Green) (Brown) Found. JEE-NEET : Class 8 Iron is more reactive than copper and displaces copper from copper sulphate solution. In general, a more reactive metal displaces a less reactive metal from its salt in the solution. Let us discuss one more example of displacement reaction. Reaction of copper with silver nitrate solution When a strip of copper metal is placed in a solution of AgNO3, the solution becomes gradually blue due to the formation of copper nitrate and a shining coating of silver metal gets deposited on the copper strip. The reaction may be written as: 2AgNO3(aq) + Cu(s) ⎯⎯→ Cu(NO3)2(aq) + 2Ag(s) Silver nitrate Copper Copper nitrate Silver (colourless solution) (blue colour) However, if we place silver wire in a copper sulphate solution no reaction occurs. This means copper can displace silver from its salt solution, but silver cannot displace copper from its solution. i.e. copper is more reactive metal than silver. The reactivity series The arrangement of metals in order of decreasing reactivities is called reactivity series or activity series of metals. After performing displacement experiments the following series has been developed. Reactivity series of metals Potassium K Sodium Na Calcium Ca Magnesium Mg Reactivity decreases Aluminium Al Reactivity increases Zinc Zn Iron Fe Lead Pb Hydrogen H Copper Cu Mercury Hg Silver Ag Gold Au   Chemistry ◼ Reaction of metals with non-metals Atoms of elements combine to form stable molecules. The combining power of an atom is expressed as valency. Each atom has a tendency to attain a completely filled valence shell. The noble gases, which have a completely filled valence shell or outermost shell, are very stable. The electronic configuration of noble gases and some metals and non-metals are given in the following table. Electronic configuration of some elements Types of Element Atomic Number of electrons in element number 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 It is clear from the above table that except helium, all other noble gases have 8 electrons (octet) in their outermost shell, which represent a highly stable electronic configuration. Due to this stable configuration, the noble gases have no tendency to lose or gain electrons. So they exist in monoatomic form. However, metals and non-metals which do not have complete octet will try to attain stability either by gaining or loosing electrons. Lets discuss formation of sodium chloride (NaCl). Sodium atom has one electron in its outermost shell. If it loses the electron from its M shell then its L shell becomes the outermost shell, which has stable octet like noble gases. The nucleus of this atom still has 11 protons but the number of electrons becomes 10. Therefore, it becomes positively charged sodium ion or cation (Na+). Found. JEE-NEET : Class 8 lose 1 electron Na ⎯⎯⎯⎯⎯⎯⎯→ Na+ + e– 2, 8, 1 2, 8 Sodium ion On the other hand, chlorine has seven electrons in its outer most shell and it require one more electron to complete its octet. The nucleus of chlorine atom has 17 protons and the number of electrons become 18. Therefore, it becomes negatively charged chloride ion (Cl–) or anion. gain 1 electron Cl + e– ⎯⎯⎯⎯⎯⎯⎯→ Cl– 2, 8, 7 2, 8, 8 Chloride ion So, Na+ and Cl– ions being oppositely charged attract to each other and are held by strong electrostatic forces of attraction to exist as NaCl. In other words, Na+ and Cl– ions are held together by electrovalent or ionic bond. Na + Cl ⎯→[Na+] [ Cl ] – or NaCl Na Cl Na Cl Sodium atom Chlorine atom Sodium atom Chlorine atom (a cation) (an anion) Sodium chloride (NaCl) Formation of sodium chloride. Electrovalent bond or ionic bond may be defined as the electrostatic force of attraction which holds the oppositely charged ions together or it may also be defined as a chemical bond formed between two atoms by complete transfer of electrons from one atom to another so as to complete their octet and hence acquire the stable nearest noble gas configuration. The number of electrons lost or gained by the atom is called its electro valency.   Chemistry Let us discuss the formation of one more ionic compound magnesium chloride (MgCl2). The electronic configuration of magnesium (Mg) and chlorine (Cl) atoms are Mg12 : 2, 8, 2 Cl17 : 2, 8, 7 Magnesium atom has two electrons in its valence shell, so has a tendency to lose both of its electrons to attain the nearest noble gas configuration i.e. Ne. Mg → Mg2+ 2, 8, 2 2,8. On the other hand, chlorine has only one electron less than the nearest noble gas (i.e. Ar) configuration. Magnesium transfers both its valence electrons to two chlorine atoms, each of which needs one electron to form Cl– ion. –2e Mg → Mg2+ – 2, 8, 2 2, 8 Mg + Cl ⎯→ [Mg2+] [ Cl ] 2 +2e 2Cl → 2Cl– or Cl MgCl2 2, 8, 7 2, 8, 8 Magnesium chloride Some common ionic compounds are 1. Magnesium Mg + O ⎯⎯→ Mg2+[O]2– or MgO oxide 2, 8, 2 2, 8, 6 2. Magnesium Mg + 2F ⎯⎯→ Mg2+2[F]– or MgF2 fluoride 2, 8, 2 2, 7 3. Calcium Ca + O ⎯⎯→ Ca2+[O]2– or CaO oxide 2, 8, 8, 2 2, 6 4. Aluminium 2Al + 3O ⎯⎯→ 2Al3+3[O]2– or Al2O3 oxide 2, 8, 3 2, 6 5. Magnesium Mg + 2Cl ⎯⎯→ Mg2+2[Cl]– or MgCl2 chloride 2, 8, 2 2, 8, 7 6. Aluminium Al + N ⎯⎯→ Al3+N3– or AlN nitride 2, 8, 3 2, 5 ◼ Properties of ionic compounds To learn about the properties of ionic compounds, let us perform the following activity. Following are the general properties of ionic compounds. (a) Physical state Ionic compounds are solids and relatively hard because of the strong force of attraction between the positive and negative ions. This force of attraction is also known as strong electrostatic force of attraction. These compounds are generally brittle and break into pieces when pressure is applied. (b) Solubility Electrovalent compounds are generally soluble in water (because of their polar nature) and insoluble in solvents such as kerosene, petrol, etc. (c) Melting and boiling points Ionic compounds have high melting and boiling points, due to the strong electrostatic force of attraction between the oppositely charged ions. Therefore, large amount of energy is needed to break these bonds. Found. JEE-NEET : Class 8 Ionic Compound Melting Point (K) Boiling Point (K) NaCl 1074 1738 LiCl 878 > 1655 KBr 1007 1708 KI 953 1600 CaCl2 1055 1870 CaO 2845 3123 MgCl2 987 1685 (d) Conduction of electricity Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid state is not possible due to their rigid structure. But they can conduct electricity in molten or aqueous state. (e) Colour to the flame Most of the salts when brought into the flame, impart characteristic colour to the flame. Flame colours are produced from the movement of the electrons in the metal ions present in the compounds. Metal Sodium Barium salts Potassium Colour of flame Yellow Green Lilac/violet Properties Metals Non-metals Physical properties State Metals are solids at ordinary Non-metals exist in all the three temperature. states i.e. solid, liquid and gas. Exception: Mercury is a liquid metal. Lustre They possess lustre or shine. They possess no lustre. Exceptions are Iodine and graphite. Malleability Metals are generally Non-metals are neither malleable nor and Ductility malleable and ductile, ductile. exceptions are alkali metals. Hardness Metals are generally hard. Non-metals possess varying Alkali metals are exceptions. hardness. Diamond is an exception. It is the hardest substance known to occur in nature. Density They have high densities. They generally possess low densities. Conductivity Metals are good conductors of Non-metals are poor conductors of heat and electricity. heat and electricity. The only exception is graphite which is a good conductor of electricity. Melting and They usually have high Their melting and boiling points are boiling points melting and boiling points. usually low. The only exceptions are boron, carbon and silicon.   Chemistry Distinction between Metals and Non-metals Chemical Properties Action with Metals generally react with Non-metals do not displace mineral acids dilute mineral acids to liberate hydrogen on reaction with H2 gas. dilute minerals acids. Nature of They form basic oxides. For Non-metals form acidic or oxides example, Na2O, MgO, etc. These neutral oxides. For example, oxides are ionic in nature. SO2, CO2, P2O5, etc. are acidic whereas CO, N2O, etc. are neutral. These oxides are covalent in nature. Electrochemical Metals are electropositive in Non-metals are electronegative behaviour character. in character. Oxidising or Metals behave as reducing Non-metals generally behave reducing agents. This is because of their as oxidising agents since they behaviour tendency to lose electrons. have the tendency to gain Na → Na+ + e– electrons. 1/2 Cl2 + e– → Cl– 3 Why sodium chloride has high melting point? Explanation Sodium chloride consist of Na+ and Cl– ions. These oppositely charged ions are strongly attracted towards each other. To break these strong forces of attraction, a large amount of energy is needed and hence sodium chloride has a high melting point. ◼ Occurrence of metals The earth's crust is the major source of metals. They are present in nature in the free state as well as in combined state. Oxygen & silicon are the main elements present in earth's crust. Both are non-metals. Seawater also contains some soluble salts such as NaCl, MgCl2, etc. Native and combined states of metals Metals occur in the crust of earth in two states : native state and combined state. Found. JEE-NEET : Class 8 (1) Native state A metal is said to occur in native or free state when a metal is found in nature in the elementary or metallic state. The metals at the bottom of the activity series are least reactive. They are often found in free state. For example, Gold, silver, copper and platinum are found in free state because they are very unreactive metals. So, they have no tendency to react with oxygen and they do not react with moisture, CO2 of air or any other non-metal. (2) Combined state The metals at the top of reactivity series are not expected to occur in free state due to their reactive nature. They exist in combination with other elements as oxides, carbonates, halides, sulphates, sulphides, etc. For example, Sodium, potassium, calcium, aluminium, magnesium, etc. are very reactive metals. All of these are lying at the top of activity series. These are never found in the free state. The metals in the middle of the activity series such as zinc, iron, lead, etc. are moderately reactive. They are found in the earth crust mainly as oxides, sulphide or carbonates. On the basis of reactivity, metals are divided into the following three categories. K Na Highly reactive metals Ca (Top of the activity series) Mg Never found in the free state. Al Metals above hydrogen Zn Fe Moderately reactive metals Ni (Middle of the activity series) Sn Found in combined state. Pb Cu Less reactive metals Ag (Just below hydrogen) Metals below hydrogen Found in the free state as well as combined state. Pt Least reactive metals Au (Bottom of the activity series). Found in the free state. Occurrences of metals in the activity series   Chemistry Minerals and ores The elementary state of the compounds in the form of which the metals occur in nature are called minerals. The minerals from where metals can be conveniently and profitably extracted are called ores. For example, Copper occurs in nature in the form of several mineral like copper pyrites(CuFeS2), copper glance (Cu2S) and cuprite (Cu2O). We obtain copper metal profitably from copper pyrites mineral, so it is called ore of copper. Note: All ores are minerals but all minerals are not ores. ◼ Extraction of metals : Metallurgy The process through which a pure metal is extracted from its ores is known as extraction of metals. The series of various processes involved in the extraction of metals from their ores, followed by refining of the metal is known as metallurgy. Various steps involved in the extraction of metals or metallurgical process: 1. Crushing and grinding of the ore. 2. Concentration of the ore or enrichment of the ore. 3. Extraction of metal from the concentrated ore. 4. Refining or purification of the impure metal. 1. Crushing and grinding of the ore Most of the ores in nature occur as big rocks. They are broken into small pieces with the help of crushers. These pieces are then reduced to fine powder with the help of a ball mill or a stamp mill. This process is known as pulverization of the ore. Large pieces of ore Crushed ore Crushed ore Finely divided powder (A) Crusher (B) Stamp mill Pulverization of ore Found. JEE-NEET : Class 8 2. Enrichment of ore or concentration of ore The ores mined from the earth's crust contain a number of impurities, such as soil, sand, etc. called gangue or matrix. The process of removal of impurities (gangue) from the ore is called enrichment of ore or concentration of ore. 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. (a) Levigation or gravity separation or hydraulic washing This method is based upon the difference in the densities of the ore particles and impurities (gangue). Generally metal ores are heavier than the gangue associated with them. For example, Haemetite ore of iron. Powdered ore gangue ore suspension Water Concentrated ore Gravity separation (b) Froth floatation This method is based on the difference in the wetting properties of the ore and gangue particles with water and oil. It is used for enrichment of sulphide ores. For example, ZnS, HgS.   Chemistry Froth bubbles carrying sulphide Compressed ore particles air Sulphide ore particles Water containing pine oil Gangue Froth floatation process for the concentration of sulphide ores Froth floatation (c) Liquation This method is based on difference in melting point of ore and gangue particles. For example, ore of tin and zinc. Crude metal Heat Pure metal Liquation (d) Magnetic separation This method is based on difference in the magnetic properties of the ore and gangue. For example, Magnetite (Fe3O4) ore of iron. Powered ore Magnetic roller Non-magnetic Leather belt impurities Magnetic ore Magnetic separation (e) Chemical separation When none of the physical property makes the difference, then we use chemical properties as the basis for enrichment. For example, Bayer's process for alumina enrichment. Found. JEE-NEET : Class 8 3. Extraction of metal from the enriched ore The method used for extraction of the metal from the concentrated ore depends upon the nature of metal. Based on the reactivity, the metals have been grouped into the following three categories : (I) Metals of low reactivity. (Low in the activity series) (II) Metals of medium reactivity. (In the middle of the activity series) (III) Metals of high reactivity. (At the top of the activity series) (I) Extraction of metals low in the activity series (Cu, Hg, Ag, Pt, Au) As these metals are very less reactive, they are either found in native state or in the form of sulphide ores. For example, Cinnabar (HgS) These sulphide ores can be converted to oxide ores on heating in the presence of excess of air called Roasting. 2HgS(s) + Heat→ 2HgO(s) 3O2(g) ⎯⎯⎯ + 2SO2(g) Mercuric sulphide Oxygen Mercuric Sulphur dioxide (Cinnabar) oxide This oxide can be reduced to metal by further heating. 2HgO Heat→ 2Hg () + O (g) ⎯⎯⎯ 2 Mercuric oxide Mercury Oxygen Similarly, when copper glance (Cu2S) an ore of copper, is subjected to roasting, it directly gives copper. 2Cu2S(s) + 3O2(g) Heat→ 2Cu O(s) ⎯⎯⎯ + 2SO2(g) 2 Copper glance Oxygen Cuprous oxide Sulphur dioxide 2Cu2O(s) + Cu2S(s) Heat→ 6Cu(s) ⎯⎯⎯ + SO2(g) Copper oxide Copper glance Copper Sulphur dioxide (II) Extraction of metals in the middle of the activity series (Fe, Zn, Pb, etc.) These metals are found in the form of their oxides, sulphides and carbonates. For easy extraction, sulphide and carbonate ores are first converted into the oxide. (a) Conversion into metal oxide (i) Calcination : for the conversion of carbonate ores into oxides. It is the process of heating the ore strongly in the absence of air. For example, Heat ZnCO3(s) ⎯⎯⎯⎯⎯⎯⎯⎯⎯ → ZnO(s) + CO2(g) (Absence of air) Zinc carbonate Zinc oxide Carbon dioxide (Calamine - ore of Zn)   Chemistry (ii) Roasting : for the conversion of sulphide ores into oxides. It is the process of heating the ore strongly in the presence of excess of air. 2ZnS(s) + 3O2(g) Heat ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ 2ZnO(s) + 2SO2(g) Presence of excess of air Zinc sulphide Zinc Sulphur (Zinc blende-ore of Zn) oxide dioxide (b) Reduction of the metal oxide to metal For reduction suitable reducing agents are used, like carbon, carbon monoxide, aluminium, sodium or calcium. (i) Reduction by heating with carbon (coke) When zinc oxide is heated with carbon, zinc metal is produced. ZnO(s) + C(s) ⎯⎯→ Zn(s) + CO(g) Zinc oxide Carbon Zinc metal Carbon monoxide (Reducing agent) Reduction by carbon is also known as smelting. Similarly, iron and lead are also obtained from their oxides by heating with carbon. Fe2O3(s) + 3C(s) Heat→ ⎯⎯⎯ 2Fe(s) + 3CO(g) PbO(s) + C(s) Heat→ ⎯⎯⎯ Pb(s) + CO(g) (ii) Reduction with CO Iron is obtained from ferric oxide by heating with CO. Fe2O3(s) Heat→ 2Fe(s) + + 3CO(g) ⎯⎯⎯ 3CO2(g) (iii) Reduction by aluminium 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 The thermite reaction gives off so much heat that compounds. the iron formed is molten Certain metal oxides are reduced by aluminium to metals. This method is known as aluminothermy or thermite process. Found. JEE-NEET : Class 8 For example, Chromium, manganese, vanadium metals are obtained by the reduction of their oxides with Al powder. The following reaction takes place 3MnO (s) + 4Al(s) ⎯⎯⎯Heat→ 3Mn() + 2Al O (s) + Heat 2 2 3 Cr2O3(s) + 2Al(s) ⎯⎯⎯Heat→ 2Cr() + Al O (s) + Heat 2 3 These displacement reactions are highly exothermic, so a large amount of heat is evolved and metals are produced in the molten state. In fact the reaction of iron(III) oxide (Fe2O3) with Thermite process for aluminium, is used to weld railway tracks or cracked joining railway tracks machine parts. This reaction is known as thermite reaction. The mixture of iron oxide and aluminium powder is called thermite. Fe O (s) + 2Al(s) ⎯⎯⎯Heat→ 2Fe() + Al O (s) + Heat 2 3 2 3 3 Q.1 Name two metals other than silver and gold which are not attacked even by steam. Ans. Lead and copper. Q.2 Name two metals which react with very dilute HNO3 to produce hydrogen gas. Ans. Magnesium and manganese. (III) Extraction of metals high up in the activity series (K, Ca, Na, Mg and Al) The highly reactive metals such as K, Na, Mg have strong affinity for oxygen, so they can not be reduced with the help of carbon. Hence these metals are obtained by electrolysis of their molten or fused oxides or chlorides, this method is called electrolytic reduction. On electrolysis, metal ions, being positive, are liberated at the cathode (negative electrode) where they gain electrons and convert in the metal atoms. For examples, (i) Sodium metal is obtained by electrolysis of molten sodium chloride. Heat to → Na+() + Cl–() NaCl(s) ⎯⎯⎯⎯⎯ melt At Cathode: Na+() + e– ⎯⎯→ Na(s) (Reduction) Sodium ion electron Sodium metal At Anode : Cl–() ⎯⎯→ Cl(g) + e– (Oxidation) Chloride ion Chlorine atom Cl(g) + Cl(g) ⎯⎯→ Cl2(g) Chlorine atoms Chlorine gas Thus, sodium metal is obtained at cathode whereas chlorine gas is obtained at the anode.   Chemistry (ii) Aluminium oxide is reduced to aluminium by the electrolysis of molten aluminium oxide. Heat → 2Al3+() + 3O2–() Al2O3(s) ⎯⎯⎯⎯⎯ to melt At Cathode : Al3+() + 3e– ⎯⎯→ Al(s) (Reduction) Aluminium ion Aluminium metal At Anode : O2–() ⎯⎯→ O(g) + 2e– (Oxidation) Oxide ions Oxygen atom O(g) + O(g) ⎯⎯→ O2(g) Oxygen atoms Oxygen gas (4) Refining of impure 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 of impure metals is electrolytic refining. Process (i) In this process, the impure metal is made the anode and a thin strip of pure metal is made the cathode. (ii) A solution of the metal salt is used as an electrolyte. On passing the electric current through the electrolyte, the pure metal from the anode dissolves into the electrolyte. (iii) An equivalent amount of pure metal from the electrolyte gets deposited on the cathode. The soluble impurities go into the solution, leaving the insoluble impurities which settle down at the bottom of the anode. At Anode : M(s) ⎯⎯→ Mn+(aq) + ne– Metal atom Metal ion (from anode) At cathode: Mn+(aq) + ne– ⎯⎯→ M(s) Metal ion Metal atom (from solution) Anode mud / anode sludge The soluble impurities present in the Key impure metal pass into solution whereas – – + –e insoluble impurities fall below the anode e – + as anode mud. Cathode Anode For example, Electrolytic refining of copper. Electrolytic refining of copper. The electrolyte is a solution of acidified Tank Acidified copper sulphate. The anode is impure copper copper, whereas the cathode is a strip of Impurities sulphate solution pure copper. On passing electric current, (anode mud) pure copper is deposited on the cathode. Found. JEE-NEET : Class 8 4 Explain why carbon can reduce copper oxide to copper but not sodium oxide to sodium? Explanation Carbon is a strong reducing agent. Hence, it can reduce copper oxide to copper as follows. CuO(s) + C(s) ⎯→ Cu(s) + CO(g) Sodium is much more reactive than copper. It has greater affinity for oxygen than the affinity for carbon. Moreover, at high temperature, sodium can combine with carbon to form sodium carbide. Process of metallurgy Ore (1) Crushing and grinding Powdered ore (2) Concentration of ore (3) Extraction of metal (4) Refining Removal of impurities Electrolytic refining Metals Metals Metals of high of medium of low reactivity reactivity reactivity Electrolysis of Sulphide molten ore Carbonate Sulphide ores ore ore Roasting Pure Calcination Roasting metal heating Metal Oxide of metal Carbon / Aluminium Reduction to metal ◼ Corrosion When the surface of a metal is attacked by the gases and water vapour present in the air, it is said to corrode and the phenomenon is called corrosion. Thus, corrosion may be defined as follows: The process of slowly eating up of metals due to their conversion into oxides, carbonates, sulphide, sulphates, etc. by the action of atmospheric gases and moisture is called corrosion. In case iron is the metal involved in the chemical process, then corrosion is called rusting.   Chemistry Factors which promote corrosion (a) Position of metal in the reactivity series: Active metals placed above hydrogen are easily corroded as compared to metals which are placed below hydrogen. (b) Presence of water vapours and gases like CO2, SO2, etc. in the air. (c) Presence of salts or electrolyte in water promotes corrosion. e.g. Rusting of iron is faster in sea water than in ordinary or distilled water. Example of corrosion (i) When iron is exposed to moist air for a long time, its surface acquires a brown flaky substance called rust and the process is known as rusting. Rust is a mixture of Fe2O3 and Fe(OH)3. (ii) Copper reacts with CO2 in the air and slowly loses its shiny brown surface and acquires a green coating of basic copper carbonate in moist air. 2Cu(s) + CO2(g) + O2(g) + H2O() ⎯→ CuCO3.Cu(OH)2 Copper Basic copper carbonate from moist air Green (iii) Silver articles becomes black after sometime when exposed to air. [Due to reaction with H2S in the air to form a black coating of silver sulphide(Ag2S).] (iv) Lead or stainless steel lose their lusture due to corrosion. Prevention of corrosion (i) By painting : The corrosion of a metal can be prevented simply by painting the metal surface by grease or varnish that forms a protective layer on the surface of the metal which protect the metal from moisture and air. (ii) Self prevention : Some metals form protective layers. For example, When zinc is left exposed to the atmosphere, it combines with the oxygen of air to form a layer of zinc oxide over its surface. The oxides layer does not allow air to go inside the metal. Thus, zinc is protected from corrosion by its own protective layer. Similarly, aluminium combines with oxygen to form a dull layer of aluminium oxide on its surface which protects aluminium from further corrosion. (iii) Cathodic protection : In this method, the more reactive metal which is more corrosion- prone is connected to a bar of another metal which is less reactive and to be protected. In this process, electron flows from more reactive metal to the less reactive metal. The metal to be protected becomes the cathode and the more reactive metal becomes the anode. In this way, the two metals form an electrochemical cell and oxidation of the metal is prevented. Found. JEE-NEET : Class 8 For example, The pipelines (iron) under the surface of the earth are protected from corrosion by connecting them to a more reactive metal (magnesium or Zn) which is buried in the earth and connected to the pipelines by a wire. e– Moist Soil (–) (+) Mg Fe Sacrificial Pipeline anode (cathode) (oxidation) (reduction) Cathodic protection (iv) Oiling and greasing : Both protect the surface of metal against moisture and chemicals, etc. In addition the oil and grease prevent the surface from getting scratched. (v) Electroplating : It is a very common and effective method to check corrosion. The surface of metal is coated with chromium, nickel or aluminium, etc. by electrolysis also called electroplating. They are quite resistant to the attack by both air and water and check corrosion. If the surface of metal is electroplated by zinc, it is known as galvanisation and in case tin metal is used, then the process is called tinning. (vi) By alloying : It is a very good method of improving the properties of a metal. An alloy is a homogeneous mixture of two or more metals or non-metal. It can be prepared by first melting the metal and then dissolving the other elements (metal or non-metal) in proper proportions. The physical properties of an alloy are different from the constituent metals (from which it is made). Some of the common alloys are (i) Steel : 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 to form stainless steel which is hard and does not rust.   Chemistry (ii) Amalgam : An alloy of mercury and one or more other metals is known as an amalgam. It may be solid or liquid. A solution of sodium metal in liquid mercury metal is called sodium amalgam, which is used as a reducing agent. Amalgam of silver, tin and zinc is used by dentists for filling in teeth. (iii) Brass : Brass is an alloy of copper (Cu) and Zn. It contains 80% copper and 20% zinc. It is more malleable and more stronger than pure copper. Brass is used for making cooking utensils, condenser sheets, pipe, screws, bolts, wire, scientific instruments, ornaments, etc. Brass, a solid solution of copper and zinc, is used to make musical instruments and many other objects. (iv) Bronze : It is also the alloy of copper. It contain 90% of copper and 10% tin. It is highly resistant to corrosion and used for making utensils, statues, coins, hardware, etc. (v) Solder : It is an alloy of lead (50%) and tin (50%). It is used for soldering (or welding) electrical wires together as it melts at a low temperature. (vi) Alloys of Gold : The purity of gold is expressed is terms of 'carats'. Pure gold is known as 24 carats gold. It is very soft due to which, it is not suitable for making jewellery. It is alloyed with either silver or copper to make it hard and more suitable for making ornaments. In India, gold ornaments are usually made of 22 carats gold. It is an alloy of gold with silver or copper. Found. JEE-NEET : Class 8 24 karat gold 18 karat gold 14 karat gold 100% 56% 20% 24% 36% 25% 37% Gold Silver Copper carat gold is an element. It contains only gold atoms. 14-carat and 18-carat gold are alloys. They contain a mixture of different atoms. The wonder of ancient Indian metallurgy Ashoka pillar near Qutub-Minar in Delhi The iron pillar near the Qutub-Minar in Delhi was made around 400 BC by the iron workers of India. They had developed a process which prevented the wrought iron pillar (a type of iron) from rusting even after thousands of years. This is likely because of formation of a thin film of magnetic oxide (Fe3O4) on the surface as a result of finishing treatment given to the pillar, painting it with a mixture of different salts then heating and quenching (rapid cooling). The iron pillar is 8 metres high and 6000 kg (6 tones) in weight. This tells us that ancient Indians had good knowledge of metals and their alloys. 