• About 118 elements are known today. There are more than 90 metals, 22 non metals and a few metalloids.
• 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 sound).
• Lustrous(natural shine).
• Have high melting point. Cesiumand galliumhave 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, they have 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-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 do not 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 do not burn or react with air.
2Na + O2 ⇒ Na2O
2Mg + O2 ⇒ 2MgO
2Cu + O2 ⇒ 2CuO
4Al + 302 ⇒ 2Al2O3
⇒ Amphoteric Oxides
• Metal oxides which react with both acids as well as bases to form salt and water e.g. Al2O3, ZnO.
Al2O3 + HCl – AlCl3 + H2O
Al2O3 + NaOH – NaAlO2 + H2O
2. Reaction with water
Na + H2O ⇒ NaOH + H2
K + H2O ⇒ KOH + H2
Ca + H2O ⇒ Ca(OH)2 + H2
Mg + H2O ⇒ Mg(OH)2 + H2
In case of Ca and Mg, the metal starts floating due to dubbles of hydrogen gas sticking to its surface.
Al + H2O ⇒ Al2O3 + H2
Fe + H2O ⇒ Fe3O4 + H2
Try Balancing these Chemical equations yourself
3. Reaction with dilute 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 – FeCl2 + H2
Mg + 2HCl – MgCl2 + H2
Zn + 2HCl – ZnCl2 + H2
2Al + 6HCl – 2AlCl3 + 3H2
• Copper, mercury and silver don’t react with dilute acids.
• Hydrogen gas produced is oxidised to water when metals react with nitric acid. But Mg and Mn, react with very dilute nitric acid to evolve hydrogen gas.
Mg + 2HNO3 ⇒ Mg(NO3)2 + H2
4. Reaction of metals with other metal salts
Metal A + salt solution B – salt solution 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.
• Reactivity series is a list of metals arranged in order of their decreasing activities.
K > NA > Ca > Mg > Al > Zn > Fe > Pb > Cu > Hg > Ag > Au
Fe + CuSO4 ⇒ FeSO4 + Cu
Zn + CuSO4 ⇒ ZnSO4 + Cu
5. Reaction between Metals and Non-Metals
• Reactivity of elements can be understood as a tendency to attain a completely filled valence shell.
• 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).
• Oppositely charged ions attract each other and are held by strong electrostatic forces of attraction forming ionic compounds.
• Formation of MgCl2
Mg – Mg2+ + 2e–
Cl2 + 2e– – 2Cl–
⇒ Properties of Ionic Compounds
• Are solid and mostly brittle.
• 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.
• Conduct electricity in solution and in molten state. In both cases, free ions are formed and conduct electricity.
⇒ Occurance of Metals
• Minerals : Elements of compounds occuring naturally are minerals.
• 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.
• Gangue : Ores are naturally found mixed impurities like soil, sand, etc. called gangue. The gangue is removed from the ore.
• Metallurgy : Step-wise process of obtaining metal from its ore.
⇒ Enrichment of ore
• Obtaining metal from enriched ore.
• Refining of impure metal to obtain pure metal.
⇒ Extracting Metals Low in the Activity Series
1. By heating the ores in air at high temperature.
• Mercury from cinnabar
2HgS + 3O2 ⇒ 2HgO + 2SO2
2HgO – 2Hg + O2
• Copper from copper sulphide
Cu2S + 3O2 ⇒ 2Cu2O _ 2SO2
2Cu2O + Cu2S ⇒ 6Cu + SO2
⇒ Extracting Metals in the Middle of Activity Series
• Metals are easier to obtain from oxide ores, thus, sulphide and carbonate ores are converted into oxides.
• Metal ore heated strongly in excess of air (Roasting)
2ZnS + 3O2 ⇒ 2ZnO + 2SO2
• Metal ore heated strongly in limited or no supply of air (Calcination)
ZnCO3 ⇒ ZnO + CO2
⇒ Reduction of Metal Oxide
1. Using coke: Coke as a reducing agent.
ZnO + C ⇒ Zn + CO\n
2. Using displacement reaction : Highly reactive metal like Na, Ca and Al are used to displace metals of lower reactivity from their compounds.
MnO2 + 4Al – 3Mn + 2Al2O3 + heat
Fe2O3 + 2Al – 2Fe + Al2O3 + heat
• In the above reaction molten iron is formed and is used to join railway tracks.This is called thermit reaction.
Extracting Metals at the Top of Activity Series
• have more affinity for oxygen than carbon.
• are obtained by electrolytic reduction. 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 non-metal at anode.
• At cathode :
Na+ + e– – Na
• At anode :
2Cl– – Cl2 + 2e–
⇒ Refining of Metals
• Impurities present in the obtained metal can be removed by electrolytic refining.
• 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.
• Metals are attacked by substances in surroundings like moisture and acids.
• Silver – it reacts with sulphur in air to form silver sulphide and articles become black.
• Copper – reacts with moist carbon dioxide in air and gains a green coat of copper carbonate.
• Iron-acquires a coating of a brown flaky substance called rust. Both air and moisture are necessary for rusting of iron.
⇒ Prevention of corrosion
• Rusting of iron is prevented by painting, 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 bymixing iron with nickel and chromium. It is hard and doesn’t rust.
• Mercury is added to other metals to make amalgam.
⇒ Brass : Alloy of copper and zinc.
⇒ Bronze : Alloy of copper and tin.
• In brass and bronze, melting point and electrical conductivity is lower than that of pure metal.
⇒ Solder : Alloy of lead and tin has low melting point and is used for soldering electrical wires.