Metals and Non-Metals
Metals are the solid materials which are typically hard, malleable, ductile and conduct heat and electricity, and also posses metallic lusture.
Example- Iron, Gold, Aluminium, Silver, Copper etc.
Nonmetals are chemical elements which lacks metallic properties. Non metals are either solids or gases except Bromine (Br2), which occurs as liquid. Non-metals vaporizes easily, insulator of heat and electricity. Non metals have high ionization energy and elctronegativity values.
Example- Hydrogen, Helium, Nitrogen, O2, F2, Ne, Cl, Ar, Kr, Xe, Rn, Br, C, P, S, Se, I.
Exception in metals and non-metals
· Mercury is liquid at room temperature while other metals are solid.
· Gallium and cesium have very low M.P. while other metals have very high M.P.
· Iodine is non-metal but it has lusture (shiny).
· Carbon (Non-Metal) exist in different forms, and these forms are known as allotrope.
· Graphite (allotrope of carbon) conduct electricity and Diamond (allotrope of carbon) having very high M.P. and B.P. is hardest natural substance known.
Chemical properties of metal
Burning of metals in air
Metals burn in air ( as oxygen present in air) to produce metal oxide.
Metal + Oxygen -------> Metal Oxide
Example- 2Cu (Copper) + O2 (air) ---> 2CuO (Copper Oxide)
4Al (Aluminium) + 3O2 -------> 2Al2O3 (Aluminium Oxide)
· Mostly metal oxide are basic in nature but some metal oxide are amphoteric in nature i.e. they show acidic as well as basic behavior like Aluminium Oxide, Zinc Oxide.
Al2O3 + 6HCl -------> 2AlCl3 + 3H2O
Al2O3 + 2NaOH ----> 2NaAlO2 (Sodium Aluminate) + H2O
· Mostly metal oxide do not dissolve in water but some dissolve to form alkali.
Na2O (s) + H2O (l) ---------> 2NaOH (aq)
K2O (s) + H2O (l) ---------> 2KOH (aq)
· Some metals prevent further corrosion by making protective oxide layer on itself like Al, Zn, Pb etc.
· Some metals like Na, K catches accidental fire so to prevent it, they are kept fully immersed in kerosene oil.
Reaction of metals with water
Most metal react with water to form metal oxide and hydrogen gas.
Metal + Water -----> Metal Oxide + Hydrogen Gas
2K (s) + H2O (l) -------> K2O (s) + H2 (g)
Not all, but some metal oxide react further with water to give metal hydroxide.
K2O (s) + H2O (l) -----> 2KOH (aq)
Some metal do not react with water like copper, lead, silver and gold.
Reaction of metals with acids
Most metal react with acids to produce salt and hydrogen gas.
Metal + Diluted Acid --------> Salt + Hydrogen Gas
2Al + 6HCl (dil) ---------> 2AlCl3 + 3H2
· As HNO3 is strong oxidizing agent so, hydrogen gas not evolve when reaction take place between metal and nitric acid Because HNO3 reduces itself to nitrogen oxides (N2O, NO, and NO2).
· Aqua-Regia (Royal Water) is freshly prepared mixture of concentrated Hydrochloric Acid (Conc. HCl) and concentrated Nitric Acid (Conc. HNO3) in the ratio of 3:1.
Aqua Regia have great dissolving power, highly corrosive, fuming liquids. Aqua Regia have ability to dissolve gold and platinum.
Reaction of metals with solution of other metal salts
More reactive metal have ability to displace less reactive metal from their compounds in molten or solution form.
For example, If metal A is more reactive then metal B then it displaces metal B from solution of metal B.
Metal A + Salt Solution of B ------> Salt Solution of A + Metal B
Fe (s) + CuSO4 (aq) -----> FeSO4 (aq) + Cu (s)
Reactivity Series of Metals
Reactivity series is the series in which metals are arranged as per order of decreasing activity. This reactivity series is developed after performing so many displacement experiments. This series is also called as activity series.
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How Metals react with Non-metals
As we know every element has tendency to achieve completely filled valence shell to get electronic configuration of nearby noble gas, for example a metal say sodium (Na) react with non-metal chlorine (Cl) to get complete valence shell so Na loose one electron from its outermost shell which is taken-up by chlorine (Cl), as a result they both have completely filled valence shell. In this process sodium become +vely charged and chlorine –vely charged and these charge attract each other to form NaCl.
Na (2,8,1) --------> Na+ (2,8) + e-
Na+ is sodium cation
Cl (2,8,7) + e- ----------> Cl-
Cl- is chlorine anion
Formation of Sodium Chloride
Na+ + Cl- -----------> NaCl
Properties of Ionic Compounds
Ionic Compounds also called Electrovalent Compounds and their properties are described below-
1. Physical Nature of Ionic Compounds
Ionic compounds are solid and hard due to strong force of attraction between +ve and –ve ions. Ionic compounds are also brittle.
2. Melting and Boiling Points of Ionic Compounds
Ionic Compounds have high M.P. and B.P. as large amount of energy is needed to breakup ionic bonds.
3. Solubility of Ionic Compounds
Ionic compounds are soluble in polar solvents like water and insoluble in non-polar solvents like ether, kerosene etc.
4. Electric Conductivity of Ionic Compounds
Occurrence of Metals
Maximum metals occur in earth’s crust and some metal occur in sea water. Metals and its compounds exist as minerals and if the percentages of metals in minerals are large then they are known as ores.
1. Extraction of Metals
Reactivity series is very helpful in metal extraction as metals present at the bottom of reactivity series are least reactive so found in Free State like gold, silver and platinum found in Free-State. Metal at top is most reactive and metals in the middle are also reactive so found in combined form. Metals generally found as oxides, sulphides and carbonates on earth’s crust.
Steps involved in extraction of metals from ores-
2. Enrichment of Ore
The process of removal of impurities or gangue from ore, before extraction of metal is known as enrichment of ore.
Gangue is terminology used for impurities like sand, soil etc. present in ore.
3. Extracting Metals Low in Reactivity Series (or Activity Series)
Metals present at bottom (or low position) in activity series are very unreactive and can be obtained in pure metallic form by just heating alone.
Example- Cinnabar (HgS), ore of Mercury (Hg)
2HgS (s) + 3O2 (g) + heat ---------> 2HgO (s) + 2SO2 (g)
2HgO (s) + Heat ------------> 2Hg (l) + O2 (g)
Cu2S, ore of copper (Cu)
2Cu2S + 3O2 (g) + Heat ---------> 2Cu2O (s) + 2SO2 (g)
2Cu2O + Cu2S + Heat ----------> 6Cu (s) + SO2 (g)
4. Extracting Metals in Middle of Activity Series
Metals in middle like iron, zinc, lead etc. are moderately reactive and present as sulphides or carbonates. Metals can be easily extracted from its oxides so sulphides and carbonates are reduced to oxides. Then these metal oxides are reduced to corresponding metal by using suitable reducing agent like carbon.
It is a process of converting sulphide ores into oxides by heating strongly in the presence of excess air.
2ZnS (s) + 3O2 (g) + Heat --------> 2ZnO (s) + 2SO2 (g)
It is a process of converting carbonates ores into oxides by heating strongly in the presence of limited air.
ZnCO3 (s) + Heat ----------> ZnO (s) +CO2 (g)
Oxides of ores are reduced to metal by using suitable reducing agent like carbon (Coke), or highly reactive metals.
ZnO (s) + C (s) ---------> Zn (s) + CO (s)
3MnO2 (s) + Al (s) -----------> 3Mn (l) + Al2O3 (s) + Heat
5. Extracting Metals at Top of Activity Series
Metals present at top in activity series are very reactive and they are not obtained by heating their compounds with carbon, for example Sodium, Calcium, Magnesium, Aluminium etc. cannot be obtained by reducing with carbon as these metals have more affinity for oxygen than carbon. So these metals are obtained by electrolytic reduction.
In electrolytic reduction, the metals get deposited at cathode (-ve electrode) and gas like chlorine get liberated at anode (+ve electrode)
Reaction for molten Sodium Chloride-
At Cathode :- Na+ + e- -------> Na
At Anode :- 2Cl- -------> Cl2 + 2e-
Refining of Metals
Refining of metals are done to obtain metals in very pure form by removing impurities present in it. Electrolytic refining is widely used method for this purpose.
Electrolytic refining is the method of obtaining very pure metals from impure metal. Metals like copper, zinc, nickel, silver, tin, gold etc. are refined electrolytically.
In electrolytic refining, anode (+ve) is made from impure metal and cathode (-ve) is made from thin strip of pure metal. Metal salt solution works as an electrolyte. When we applied electric current across the electrodes then current starts flow through electrolytic solution. Pure metal comes out from anode and dissolve in electrolyte and equivalent amount (i.e. to that comes from anode) of this pure metal from electrolyte solution get deposited on cathode.
“In simple way we can say that pure metal come from anode and get deposited on cathode by using electrolyte solution and electric current.”
Insoluble impurities settle down below anode at bottom and we say it as anode mud, while soluble impurities mix in electrolyte.
Natural process of conversion of refined metal to its high stable form like oxides or hydroxides of metals is known as corrosion. Corrosion is the process of gradual destruction of any material like metals by environment and chemical reaction.
Example - Rusting of Iron
Prevention of Corrosion
There are so many methods to prevent corrosion like-
1. Applied Coating
Applied coating is surface treatment method. Planting, enamel application and painting are applied coating method to prevent corrosion. These methods create barrier between metal and environment.
It is anode surface treatment process in which we made thicker oxide layer at metal surface.
Galvanization is the process of coating steel and iron with very thin layer of zinc to protect them from rusting.
Overall painting, greasing, oiling, chrome plating, galvanizing, alloy making and anodizing are some ways for the prevention of corrosion.