CHAPTER 3 METALS AND NON-METALS About 118 elements are known today. There are more than 90 metals, 22 non metals and a fewmetalloids. Sodium (Na), potassium (K),magnesium(Mg), aluminium(al), calcium(ca), Iron(Fe), Barium(Ba) are some metals. Oxygen(O), hydrogen(h),nitrogen(n), sulphur(s), phosphorus(p), fluorine(f),chlorine(cl), bromine(br), iodine(l) are some non-metals Physical properties of metals: Solid at room temperature except mercury Ductile (drawn into wires) Malleable (beaten into thin sheets) Sonorous(produce ringing sound) Lustrous(natural shine) Have high melting point. Cesium and gallium have very low melting point. Generally good conductor of heat and electricity, except lead and mercury which are comparatively poor conductors. Silver and copper are best conductors. Have high density. Sodium and potassium can be cut with knife, theyhave low density. Physical properties of non-metals: Occur as solid or gas. Bromine is liquid. Generally bad conductors of heat and electricity. Graphite a natural form of carbon is a good conductor. Non-sonorous, Non-lustrous, only iodine has lustre. Metals form basic oxides like Magnesium oxide(mgo), while non-metals form acidic oxides (as in acid rain). Chemical properties of metals: 1. Reaction with air Metals can burn in air, react or don't react with air. Metal oxygen Metal Oxide Some metals like Na and K are kept immersed in kerosene oil as they react vigorously with air and catch fire. Some metals like Mg, Al, Zn, Pb react slowly with air and form a protective layer. Mg can also burn in air with a white dazzling light to form its oxide Fe and Cu don't burn in air but combine with oxygen to form oxide. When heated iron filings burn when sprinkled over flame. Metals like silver, platinum and gold don't burn or react with air. 2Na O 2 Na 2 O 2Mg O 2 2MgO 2Cu O 2 2CuO 4Al 30 2 2Al 2 The oxides of metals are generally basic in nature they reacts with acids and give salt and
water e.g. MgO 2HCl MgCl 2 O CaO SO 4 CaSO 4 O Amphoteric Oxides : metal oxides which react with both acids as well as bases to form salt and water e.g. Al 2, ZnO. Al 2 HCl AlCl 3 O Al 2 NaOH NaAlO 2 O 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 2.Reaction with water : Metal water Metal hydroxide or Metal oxide gas e.g. Na O NaOH K O Ca O Ca(OH) 2 KOH Mg O Mg(OH) 2 In case of Ca and Mg, the metal starts floating due to bubbles of hydrogen gas sticking to its surface. Al O Al 2 Fe O Fe 3 O 4 3. Reaction with acids: Metal dilute acid Salt Hydrogen gas Metals react with dilute hydrochloric acid and dilute sulphuric acid to form salt and hydrogen gas. Fe 2HCl FeCl 2 Mg 2HCl MgCl 2 Zn 2HCl ZnCl 2 2Al 6HCl 2AlCl 3 3 Hydrogen gas is not evolved when a metal reacts with nitric acid. It is because HN is a strong oxidising agent. It oxidises the produced to water and itself gets reduced to any of the nitrogen oxides (N 2 O, NO,NO 2 ). But Mg and Mn, react with very dilute nitric acid to evolve hydrogen gas.
Mg 2HN Mg(N ) 2 Copper, mercury and silver generally don t react with dilute acids. They generally react with conc. SO 4 and conc.hn but don t liberate gas,thses liberates SO 2 or NO 2. e.g. 4.Reaction of metals with other metal salts : Metal A salt solution of metal B salt solution of metal A Metal B All metals are not equally reactive. Reactive metals can displace less reactive metals from their compounds in solution. This forms the basis of reactivity series of metals. e.g. Zn(s) CuSO 4 (aq) ZnSO 4 (aq) Cu (s) Cu (s) 2AgN (aq) Cu(N ) 2 (aq) 2Ag(s) Reactivity Series - The reactivity series is a list of metals arranged in the order of their decreasing activities. Li Lithium Most reactive K Potassium Na Sodium A metal can displace all the metals from its Ca Calcium compound which lie below it in activity Mg Magnesium series. Al Aluminium Zn Zinc Reactivity decreases r e a c t i v i t y s e r i e s Fe Iron Pb Lead H Hydrogen Cu Copper Hg Mercury Ag Silver Au Gold Pt Platinum Fe(s) CuSO 4 (aq) FeSO 4 (aq) Cu(s) Least reactive 5.Reaction between Metals and Non-Metals :Ionic compound formation Reactivity of elements can be understood as a tendency to attain a completely filled valence shell. We know that noble gases, which have a completely filled valence shell( or octet configuration), Show little chemical activity.by chemical combination the elements achieve noble gas electronic configuration Atom of metals can lose electrons from valence shells to form cations (ve ions). Atom of non-metals gain electrons in valence shell to form anions ( ve ions). Metals lying above hydrogen can displace hydrogen as gas from acid and from water. Metals lying below hydrogen can t displace hydrogen as gas from acid and from water. Oppositely charged ions attract each other and are held by strong electrostatic forces of attraction forming ionic compounds
Formation of NaCl Na(2,8,1) atoms lose one electron each to form Na ions, which contain only ten electrons, the same number as the preceding noble gas, neon(2,8). In contrast, Cl atoms(2,8,7) gain one electron each to form Cl - ions, which contain 18 electrons(2,8,8). This is the same number as the following noble gas, argon. These processes can be represented compactly as Formation of MgCl 2 Cl e Cl 2,8,7 2,8,8 (Chloride ion) Properties of Ionic Compounds : Are solid and mostly brittle. Don t exist as molecule but form ionic crystal lattice in which cations and anions are arranged in 3dimentional network in a regular pattern Have high melting and boiling points. More energy is required to break the strong inter-ionic attraction. Generally soluble in water and insoluble in kerosene,petrol. Ionic compounds are soluble in polar solvent like water, methanol etc but petrol kerosene benzene etc are non polar solvents hence these are insoluble. [LIKE DISSOLVES LIKE]. Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to the rigid structure of crystal lattice. But Conduct electricity in solution and in molten state. In both cases, free ions are formed and conduct electricity. Occurrence of Metals and its extraction Minerals : The elements or compounds, which occur naturally in the earth s crust, are known as minerals. The earth s crust is the major source of metals. Seawater also contains some soluble salts such as sodium chloride, magnesium chloride, etc. ORES : mineral from which metal can be profitably extracted is an ore. For example, sulphide ore, oxide ore, carbonate ore. Metals at the bottom of activity series like gold, platinum, silver, copper generally occur in free state. But copper and silver also occur in sulphide and oxide ores.
Metals of medium reactivity(zn, Fe, Pb etc.) occur mainly as oxides, sulphides or carbonates. Metals of high reactivity (K, Na, Ca, Mg and Al) are very reactive and thus found in combined state. The process of extraction of a metal from its ore is called as metallurgy. Some important ores of metals are as follow- Metals Ores metals Ores Iron Hematite: Fe 2 Magnetite: Fe 3 O 4 Lead Galena: PbS Copper Chalcopyrite: CuFeS 2 Malachite: Cu 2 C (OH) 2 Aluminium Bauxite Al 2 Mercury Cinnabar: HgS Silver Argentite Ag 2S Metallurgy : all the process involved in extraction of a metal from its ore is called as metallurgy. The main processes are -. *Enrichment of ore *Obtaining metal from enriched ore. *Refining of impure metal to obtain pure metal. A)Enrichment of ore Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue. The removal of these impurities from ore is called concentration of ore or enrichment of ore. 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, these are a. Hydraulic washing b. Magnetic separation (when ore or gangue any one is magnetic e.g. haematite and magnetite c. Froath-floation process ( used for sulphide ores CuFeS2, PbS, ZnS) B)Obtaining metal from enriched ore. This involve many processes depending upon the reactivity of metals. Generally ore is converted into oxide which is then reduced to metal by using coke (carbon) or other suitable reducing agent. Extracting Metals Low in the Activity Series : The oxides of these metals can be reduced to metals by heating alone at high temperature. *Mercury from cinnabar 2HgS 3O 2 He at 2HgO 2SO 2 2HgO He at 2Hg O 2 * Copper from copper sulphide Cu 2 S 3O 2 He at 2Cu 2 O _ 2SO 2 2Cu 2 O Cu 2 S He at 6Cu SO 2 Extracting Metals in the Middle of Activity Series : Conversion of ore into metal oxide -Metals are easier to obtain from oxide ores, thus, sulphide and carbonate ores are converted into oxides. Roasting It involve strongly heating of ore (specially sulphide ore) in excess of air so that sulphide ore is converted into oxide and volatile impurities are removed off. 2ZnS 3O 2 He at 2ZnO 2SO 2
Calcination - Metal ore (specially carbonate ore) is heated strongly in limited or no supply of air (Calcination) Reduction of Metal Oxide : ZnC He at ZnO CO 2 1. USING COKE: Coke as a reducing agent. ZnO C He at Zn CO 2. USING DISPLACEMENT REACTION : highly reactive metal like Na,Ca and Al are used to displace metals of lower reactivity from their compounds. MnO 2 4Al He at 3Mn 2Al 2 heat Fe 2 2Al He at 2Fe Al 2 heat 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 ) with aluminium is used to join railway tracks or cracked machine parts. This reaction is known as the thermit reaction Extracting Metals at the Top of Activity Series :. These metals have more affinity for oxygen than carbon. are obtained by electrolytic reduction. e.g. Sodium is obtained by electrolysis of its molten chloride NaCl Na Cl As electricity is passed through the solution metal gets deposited at cathode and nonmetal at anode. At cathode : At anode : Na e Na 2Cl Cl 2 2e Refining of Metals : Removal of impurities present in the metal obtained from ore is called refining. Impurities present in the metal can be removed by electrolytic refining. e.g. Copper is obtained using this method. Following are present inside the electrolytic tank. Anode slab of impure copper Cathode slab of pure copper Solution aqueous solution of copper sulphate with some dilute sulphuric acid From anode copper ions are released in the solution and equivalent amount of copper from solution is deposited at cathode. Impurities containing silver and gold gets deposited at the bottom of anode as anode mud.
Corrosion : When a metal is attacked by substances around it such as moisture,oxygen, Carbon dioxide & acids, etc., it is said to corrode and this process is called corrosion. e.g. Silver - it reacts with sulphur in air to form silver sulphide and articles become black. 2Ag S Ag 2 S Copper - reacts with moist carbon dioxide in air and gains a green coat of copper carbonate. 2Cu O 2 O CO 2 CuC.Cu(OH) 2 Iron-acquires a coating of a brown flaky substance called rust. Both air and moisture are necessary for rusting of iron. 2Fe O 2 x O Fe 2.x O Prevention of corrosion: Rusting of iron is prevented bypainting, oiling, greasing, galvanizing, chrome plating, anodising and making alloys. In galvanization, iron or steel is coated with a layer of zinc because zinc is preferably oxidized than iron. Alloys : These are mixture of metals with metals or non-metals Adding small amount of carbon makes iron hard and strong. Stainless steel is obtained by mixing iron with nickel and chromium. It is hard and doesn t rust. Alloy of Mercury with other metals is called as amalgam. Brass : alloy of copper and zinc. Bronze : alloy of copper and tin. In brass and bronze, melting point and electrical conductivity is lowerthan that of puremetal. Solder : alloy of lead and tin has low melting point and is used for welding electrical wires.