{"id":5985,"date":"2020-10-13T12:22:51","date_gmt":"2020-10-13T06:52:51","guid":{"rendered":"http:\/\/astan.lk\/al_virtualclassroom\/?p=5985"},"modified":"2020-10-13T12:22:45","modified_gmt":"2020-10-13T06:52:45","slug":"p-block-elements","status":"publish","type":"post","link":"https:\/\/astan.lk\/al_virtualclassroom\/p-block-elements\/","title":{"rendered":"p block elements"},"content":{"rendered":"

Aluminium<\/span><\/h3>\n

Aluminum (also called Aluminium) is the third most abundant element in the earth’s crust. \u00a0It is commonly used in the household as aluminum foil, in crafts such as dyeing and pottery, and also in construction to make alloys. In its purest form the metal is bluish-white and very ductile. It is an excellent conductor of heat and electricity and finds use in some wiring. When pure it is too soft for construction purposes but addition of small amounts of silicon and iron hardens it significantly.<\/p>\n

The electron configuration for Aluminum is\u00a01s2<\/sup>2s2<\/sup>2p6<\/sup>3s2<\/sup>3p1<\/sup>.<\/p>\n

Aluminum has three oxidation states. \u00a0The most common one is +3. \u00a0The other two are +1 and +2.<\/p>\n

Reaction of aluminium with air<\/strong><\/span><\/p>\n

Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react with air. If the oxide layer is damaged, the aluminium metal is exposed to attack. Aluminium will burn in oxygen with a brilliant white flame to form the trioxide alumnium(III) oxide, Al2<\/sub>O3<\/sub>.<\/p>\n

4Al(s) + 3O2<\/sub>(l) \u2192 2Al2<\/sub>O3<\/sub>(s)<\/p>\n

Reaction of aluminium with water<\/strong><\/span><\/p>\n

Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react with air. If the oxide layer is damaged, the aluminium metal is exposed to attack, even by water.<\/p>\n

Reaction of aluminium with the halogens<\/span><\/strong><\/p>\n

Aluminium metal reacts vigorously with all the halogens to form aluminium halides. So, it reacts with chlorine, Cl2<\/sub>, bromine, I2<\/sub>, and iodine, I2<\/sub>, to form respectively aluminium(III) chloride, AlCl3<\/sub>, aluminium(III) bromide, AlBr3<\/sub>, and aluminium(III) iodide, AlI3<\/sub>.<\/p>\n

2Al(s) + 3Cl2<\/sub>(l) \u2192 2AlCl3<\/sub>(s)\u00a0\u2192 Al2<\/sub>Cl6<\/sub>(s)<\/p>\n

AlCl3 is an electron defficient compound and also highly covalent when anhydrous. AlCl3 exists as a dimer thus attaining an octet of electrons.<\/p>\n

2Al(s) + 3Br2<\/sub>(l) \u2192 Al2<\/sub>Br6<\/sub>(s)<\/p>\n

2Al(s) + 3I2<\/sub>(l) \u2192 Al2<\/sub>I6<\/sub>(s)<\/p>\n

Reaction of aluminium with acids<\/strong><\/span><\/p>\n

Aluminium metal dissolves readily in dilute sulphuric acid to form solutions containing the aquated Al(III) ion together with hydrogen gas, H2<\/sub>. The corresponding reactions with dilute hydrochloric acid also give the aquated Al(III) ion. Concentrated nitric acid passivates aluminium metal.<\/p>\n

2Al(s) + 3H2<\/sub>SO4<\/sub>(aq) \u2192 2Al3+<\/sup>(aq) + 2SO4<\/sub>2-<\/sup>(aq) + 3H2<\/sub>(g)<\/p>\n

2Al(s) + 6HCl(aq) \u2192 2Al3+<\/sup>(aq) + 6Cl–<\/sup>(aq) + 3H2<\/sub>(g)<\/p>\n

Reaction of aluminium with bases<\/strong><\/span><\/p>\n

Aluminium dissolves in sodium hydroxide with the evolution of hydrogen gas, H2<\/sub>, and the formation of aluminates of the type [Al(OH)4<\/sub>]–<\/sup>.<\/p>\n

2Al(s) + 2NaOH(aq) + 6H2<\/sub>O \u2192 2Na+<\/sup>(aq) + 2[Al(OH)4<\/sub>]–<\/sup> + 3H2<\/sub>(g)<\/p>\n

\u00a0Carbon<\/span><\/h3>\n

Carbon is a Group 14 element and is distributed very widely in nature. It is found in abundance in the sun, stars, comets, and atmospheres of most planets. Carbon is present as carbon dioxide in the atmosphere and dissolved in all natural waters. It is a component of rocks as carbonates of calcium (limestone), magnesium, and iron.<\/p>\n

Coal, petroleum, and natural gas are chiefly hydrocarbons. Carbon is unique among the elements in the vast number of variety of compounds it can form.<\/p>\n

Carbon is found free in nature in three allotropic forms: amorphous, graphite, and diamond. Graphite is one of the softest known materials while diamond is one of the hardest<\/p>\n

More recently, another form of carbon, buckminsterfullerene, C60<\/sub>, was discovered. This form of carbon is the subject of great interest in research laboratories today.<\/p>\n

Pure carbon is available in a number of different forms (allotropes). The most common form of pure carbon is \u03b1-graphite. This is also the thermodynamically most stable form. Diamond is a second form of carbon but is much less common. Other forms of carbon include the fullerenes. Whereas diamond and graphite are infinite lattices, fullerenes such as buckminsterfullerene, C60<\/sub>, is a discrete molecular species. Amorphous<\/em> forms of carbon such as soot<\/em> and lampblack<\/em> are materials consisting of very small particles of graphite.<\/p>\n

\"graphite
\nAtom arrangements in the most common allotrops of carbon: \u03b1-graphite.<\/p>\n

As diamond has a slightly more compact structure its density is greater than that of graphite. The appearance of diamond is well known and it is also one of the hardest materials known. Like graphite, it is relatively unreactive but does burn in air at 600-800\u00b0C. Each carbon atom is bound to four neighbours at a distance of 154.45 pm in a tetrahedral fashion and so each diamond crystal is a single giant lattice structure. In principle (and in practice!) graphite may be converted into diamond by the application of heat and pressure.<\/p>\n

\"diamond
\nCrystal strucutres of diamond.<\/p>\n

Recently another allotrope of carbon was characterized. Whereas diamond and graphite are infinite lattices, buckminsterfullerene, C60<\/sub>, is a discrete molecular species. The buckminsterfullerene molecule is a net of 12 pentagons and 20 hexagons folded into a sphere. The effect is very similar to the patchwork of 12 pentagonal and 20 hexagonal pieces of leather that sewn together make up an association football (soccer ball). The name buckminsterfullerene<\/strong> (or buckyball<\/strong> was coined because of the relationship between the structure of C60<\/sub> and R. Buckminster Fuller’s geodesic dome designs. Buckminsterfullerene is now commercially available and has also been identified in interstellar space and soot.<\/p>\n

\"buckminsterfullerene\"
\nC60<\/sub>, Buckminsterfullerene.<\/p>\n

Nanotubes<\/strong> are related to fullerenes. They are tubes giving the appearance of rolled graphite, although they are made from graphite. They are open ended while fullerenes are closed structures.<\/p>\n

Oxides of carbon<\/span><\/strong><\/p>\n

CO is a colourless, neutral, and poisonous gas. CO is used as an industrial fuel. CO is a Lewis acid.<\/p>\n

CO2 is a colourless acidic gas and a non polar molecule. Solid CO2 (dry ice) contains dispersion interactions. Dry ice is used as a freezing agent in food industry, in causing artificial rain, etc.<\/p>\n

Carbonic acid<\/span><\/strong><\/p>\n

Carbonic acid is weak diprotic acid. There are two salts derived from H2CO3 .<\/p>\n

CO2<\/sub> + H2<\/sub>0\u00a0\u2192 H2<\/sub>CO3<\/sub><\/p>\n

H2<\/sub>CO3<\/sub>\u00a0\u2192 H+<\/sup> + HCO3<\/sub>–<\/sup><\/p>\n

HCO3<\/sub>–\u00a0\u2192 H+<\/sup> + CO3<\/sub>2-<\/sup><\/p>\n

Reaction of carbon with air<\/span><\/strong><\/p>\n

Carbon, as graphite, burns to form gaseous carbon (IV) oxide (carbon dioxide), CO2<\/sub>. Diamond is a form of carbon and also burns in air when heated to 600-800\u00b0C – an expensive way to make carbon dioxide!<\/p>\n

C(s) + O2<\/sub>(g) \u2192 CO2<\/sub>(g)<\/p>\n

When the air or oxygen supply is restricted, incomplete combustion to carbon monoxide, CO, occurs.<\/p>\n

2C(s) + O2<\/sub>(g) \u2192 2CO(g)<\/p>\n

This reaction is important. In industry, air is blown through hot coke. The resulting gas is called producer gas and is a mixture of carbon monoxide (25%), carbon dioxide (4%), nitrogen (70%), and traces of hydrogen (H2<\/sub>), methane (CH4<\/sub>), and oxygen (O2<\/sub>).<\/p>\n

Reaction of carbon with water<\/span><\/strong><\/p>\n

Carbon, either as graphite or diamond does not react with water under normal conditions. Under more forsing conditions, the reaction becomes important. In industry, water is blown through hot coke. The resulting gas is called water gas and is a mixture of hydrogen (H2<\/sub>, 50%), carbon monoxide (CO, 40%), carbon dioxide (CO2<\/sub>, 5%), nitrogen and methane (N2<\/sub> + CH4<\/sub>, 5%). It is an important feedstock gas for the chemical industry.<\/p>\n

C + H2<\/sub>O \u2192 CO + H2<\/sub><\/p>\n

This reaction is endothermic (\u0394H\u00b0 = +131.3 kJ mol-1<\/sup>; \u0394S\u00b0 = +133.7 J K-1<\/sup> mol-1<\/sup>) which means that the coke cools down during the reaction. To counteract this, the steam flow is replaced by air to reheat the coke allowing further reaction.<\/p>\n

Reaction of carbon with the halogens<\/span><\/strong><\/p>\n

Graphite reacts with fluorine, F2<\/sub>, at high temperatures to make a mixture of carbon tetrafluoride, CF4<\/sub>, together with some C2<\/sub>F6<\/sub> and C5<\/sub>F12<\/sub>.<\/p>\n

C(s) + excess F2<\/sub>(g) \u2192 CF4<\/sub>(g) + C2<\/sub>F6<\/sub> + C5<\/sub>F12<\/sub><\/p>\n

At room temperature, the reaction with fluorine is complex.<\/p>\n

The other halogens appear to not react with graphite.<\/p>\n

Applications<\/h3>\n

Carbon has a very high melting and boiling point and rapidly combines with oxygen at elevated temperatures. In small amounts it is an excellent hardener for iron, yielding the various steel alloys upon which so much of modern construction depends. An important (but rare) radioactive isotope of carbon, C-14, is used to date ancient objects of organic origin.<\/p>\n

Nitrogen<\/span><\/h3>\n

N2<\/sub>\u00a0is a colourless gas. It contains a triple bond, , with a short bond length of 1.09 \u00b0A. It has high dissociation bond energy 946 kJ mol-1 and it is an inert gas.<\/p>\n

\u2022 Liquid \u00a0N2<\/sub>\u00a0(boiling point -196 \u00b0C) is used as a coolant. \u00a0N2<\/sub>\u00a0gas is used in the manufacture of ammonia.<\/p>\n

\u2022 Nitrogen shows variable oxidation states.<\/p>\n

-3 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0-2 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 -1 \u00a0 \u00a0 \u00a0 \u00a00 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 +1 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0+ 2 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0+3 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0+4 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0+5<\/p>\n

NH3<\/sub>\u00a0 \u00a0 \u00a0 \u00a0 \u00a0\u00a0N2<\/sub>H4<\/sub>\u00a0 \u00a0 \u00a0NH2<\/sub>OH \u00a0 \u00a0N2<\/sub>\u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0N2<\/sub>O \u00a0 \u00a0 \u00a0 \u00a0 \u00a0NO \u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0N2<\/sub>O3<\/sub>\u00a0 \u00a0 \u00a0 \u00a0 \u00a0 \u00a0N2<\/sub>O4<\/sub>\u00a0 \u00a0 \u00a0 \u00a0 \u00a0N2<\/sub>O5<\/sub><\/p>\n

Oxo acids of nitrogen<\/span><\/strong><\/p>\n

Nitrous acid (HNO2<\/sub> ) is unstable except in dilute solutions. HNO2<\/sub> is formed when acid reacts with metal nitrites.<\/p>\n

NaNO2<\/sub>(s) + HCl(aq)\u00a0\u2192 NaCl(aq) + HNO2<\/sub>(aq)<\/p>\n

Nitric acid (\u0397\u039d\u039f3<\/sub> ) is a colourless liquid which boils at 86 \u00b0C. \u0397\u039d\u039f3<\/sub> a stable, strong acid as well as a strong oxidizing agent. Concentrated \u0397\u039dO3<\/sub> acid is usually yellow. In the light, it decomposes to form nitrogen dioxide and oxygen.<\/p>\n

4HNO3<\/sub>(l) \u2192 4NO2<\/sub>(g) + O2<\/sub>(g) + 2H2<\/sub>O(l)<\/em><\/p>\n

Compounds<\/strong><\/p>\n

The two most common compounds of nitrogen are Potassium Nitrate (KNO3<\/sub>) and Sodium Nitrate (NaNO3<\/sub>). These two compounds are formed by decomposing organic matter that has potassium or sodium present and are often found in fertilizers and byproducts of industrial waste.<\/p>\n

Nitrides<\/h4>\n

Nitrides are compounds of nitrogen with a less electronegative atom; in other words it’s a compound with atoms that have a less full valence shell. These compounds form with lithium and Group 2 metals. Nitrides usually has an oxidation state of -3<\/p>\n

3<\/span>M<\/span>g<\/span>+\u00a0N2<\/sub><\/span>\u2192<\/span>M<\/span>g<\/span>3<\/sub>N2<\/sub><\/span><\/span><\/p>\n

Ammonium Ions<\/h4>\n

Nitrogen goes through fixation by reaction with hydrogen gas over a catalyst. This process is used to produce ammonia. As mentioned earlier, this process allows us to use nitrogen as a fertilizer because it breaks down the strong triple bond held by N2<\/sub>.<\/sub>\u00a0The famous Haber-Bosch process for synthesis of ammonia looks like this:<\/p>\n

\u00a0N2<\/sub>+<\/span>3<\/span>H<\/span>2<\/sub><\/span><\/span>\u2192<\/span>2<\/span>NH3<\/sub><\/span><\/p>\n

Ammonia is a base and is also used in typical acid-base reactions.<\/p>\n

<\/div>\n

\u039dH3<\/sub> acts as an oxidizing agent and as well as an acid<\/p>\n

2Na(s) + 2NH3<\/sub>(l) \u2192 2NaNH2(s) + H2(g)<\/p>\n

\u039dH3<\/sub> acts as a weak reducing agent.<\/p>\n

3NH3<\/sub>(l) + 2Cl2<\/sub>(g)\u00a0\u2192 2\u00a0N2<\/sub>(g) + 6HCl(g)<\/p>\n

3CuO(s) + 2NH3<\/sub>(l)\u00a0\u2192 \u00a0N2<\/sub>(g) + 3Cu(s) + 3H2<\/sub>O(l)<\/p>\n

\u039dH3<\/sub> acts as a weak reducing agent.<\/p>\n

2<\/span>NH3<\/sub><\/span>(<\/span>a<\/span>q<\/span>)<\/span><\/span><\/span><\/span>+<\/span>H<\/span>2<\/sub><\/span><\/span>S<\/span>O<\/span>4<\/sub><\/span><\/span>\u2192<\/span>(<\/span>N<\/span>H<\/span>4<\/sub><\/span><\/span>)2<\/sub><\/span><\/span>S<\/span>O4<\/sub><\/span>(<\/span>a<\/span>q<\/span>)<\/span><\/span><\/span><\/span><\/span><\/span><\/p>\n

Ammonium salts<\/strong><\/span><\/p>\n

Ammonium salts decompose quite readily on heating.<\/p>\n

(NH4<\/sub>)2<\/sub>CO3<\/sub>(s) \u2192\u0394 \u00a02NH3<\/sub>(g) + CO2<\/sub>(g) + H2<\/sub>O(l)<\/p>\n

NH4<\/sub>Cl(s) \u2192\u0394 \u00a0NH3<\/sub>(g) + HCl(g)<\/p>\n

NH4<\/sub>NO2<\/sub>(s) \u2192\u0394 \u00a0N2<\/sub>(g) + 2H2<\/sub>O(l)<\/p>\n

NH4<\/sub>NO3<\/sub>(s)\u00a0\u2192\u0394 \u00a0N2<\/sub>O(g) + 2H2<\/sub>O(l)<\/p>\n

(NH4<\/sub>)2<\/sub>Cr2<\/sub>O7<\/sub>(s)\u00a0\u2192\u0394 \u00a0N2<\/sub>(g) + Cr2<\/sub>O3<\/sub>(s) + 4H2<\/sub>O(s)<\/p>\n

Oxides of Nitrogen<\/h4>\n

Nitrides use a variety of different oxidation numbers from +1 to +5 to for oxide compounds.\u00a0 Almost all the oxides that form are gasses, and exist at 25 degrees Celsius. Oxides of nitrogen are acidic and easily attach protons.<\/p>\n

N2<\/sub>O5<\/sub><\/span><\/span>+<\/span>H2<\/sub><\/span><\/span>O<\/span>\u2192<\/span>2<\/span>H<\/span>N<\/span>O3<\/sub><\/span>(<\/span>a<\/span>q<\/span>)<\/span><\/span><\/span><\/span><\/p>\n

The oxides play a large role in living organisms. They can be useful, yet dangerous.<\/p>\n

    \n
  • Dinitrogen monoxide (N2<\/sub>O) is a anesthetic used at the dentist as a laughing gas.<\/li>\n
  • Nitrogen dioxide (NO2<\/sub>) is harmful. It\u00a0 binds to hemoglobin molecules not allowing the molecule to release oxygen throughout the body. It is released from cars and is very harmful.<\/li>\n
  • Nitrate (NO3<\/sub>–<\/sup>) is a polyatomic ion.<\/li>\n
  • The more unstable nitrogen oxides allow for space travel.<\/li>\n<\/ul>\n
    \n
    \n

    Hydrides<\/h4>\n

    Hydrides of nitrogen include ammonia (NH3<\/sub>) and hyrdrazine (N2<\/sub>H4<\/sub>).<\/p>\n

      \n
    • In aqueous solution, ammonia forms the ammonium ion which we described above and it has special amphiprotic properties.<\/li>\n
    • Hyrdrazine is commonly used as rocket fuel<\/li>\n<\/ul>\n<\/div>\n<\/div>\n
      \n

       <\/p>\n

      Applications of Nitrogen<\/h3>\n
        \n
      • Nitrogen provides a blanketing for our atmosphere for the production of chemicals and electronic compartments.<\/li>\n
      • Nitrogen is used as fertilizer in agriculture to promote growth.<\/li>\n
      • Pressurized gas for oil.<\/li>\n
      • Refrigerant (such as freezing food fast)<\/li>\n
      • Explosives.<\/li>\n
      • Metals treatment\/protectant via exposure to nitrogen instead of oxygen<\/li>\n<\/ul>\n

        \u00a0Oxygen<\/span><\/h3>\n

        Allotropes of Oxygen<\/strong><\/p>\n

        There are two allotropes of oxygen; dioxygen (O2<\/sub>) and trioxygen (O3<\/sub>) which is called ozone. The reaction of converting dioxygen into ozone is very endothermic causing it to occur rarely and only\u00a0in the\u00a0low atmosphere. The reaction is shown below:<\/p>\n

        3O2<\/sub><\/span>(<\/span>g<\/span>)<\/span><\/span><\/span><\/span>\u2192<\/span>2O3<\/sub><\/span>(<\/span>g<\/span>)<\/span><\/span><\/span><\/span><\/p>\n

        Ozone is unstable and quickly decomposes back to oxygen but is a great oxidizing agent<\/p>\n

        Group 17 elements (halogens) fluorine, chlorine, bromine, and iodine react with oxygen to form oxides. Fluorine forms two oxides with oxygen which are F2<\/sub>O and F2<\/sub>O2<\/sub>. Both\u00a0fluorine\u00a0oxides are called oxygen fluorides because\u00a0fluorine\u00a0is the more electronegative element. One of the fluorine reactions is shown below.<\/p>\n

        O2<\/sub>(<\/span>g<\/span>)<\/span><\/span><\/span><\/span>+<\/span>F<\/span>2<\/sub><\/span>(<\/span>g<\/span>)<\/span><\/span><\/span><\/span>\u2192<\/span>F2<\/sub>O2<\/sub><\/span><\/span>(<\/span>g<\/span>)<\/span><\/span><\/span><\/span><\/p>\n<\/div>\n

         <\/p>\n

        Sulphur<\/h3>\n
        \n

        Allotropes of Sulphur<\/h4>\n
        \n

        Rhombic Sulphur (Octahedral or Alpha Sulphur)<\/span><\/strong><\/p>\n

        Rhombic sulphur is prepared by dissolving roll sulphur in carbon disulphide, and then evaporating the solution slowly, at room temperature.Rhombic or octahedral sulphur consists of rings of 8 atoms of sulphur. It is the most stable of all the allotropes of sulphur. It is soluble in carbon disulphide, benzene, chloroform, etc., but is insoluble in water. It is non-conductor of heat and electricity. It is transparent and pale yellow in color.<\/span><\/i><\/p>\n<\/div>\n<\/div>\n

        <\/div>\n
        \n

        Monoclinic Sulphur (Prismatic or Beta Sulphur)<\/span><\/strong><\/p>\n

        This allotrope of sulphur is prepared by melting roll sulphur in a dish. The molten sulphur is allowed to cool slowly. The top layer solidifies first and forms a crust. Two holes are made in the crust with the help of a heated nail. The molten sulphur is poured out through one of the holes. Then with the help of a knife the crust is carefully peeled off<\/span><\/p>\n

        Pale-yellow, transparent needle shaped crystals are seen projecting out form the inner surface of the dish. These are the crystals of monoclinic sulphur. Monoclinic sulphur also consists of 8 atom rings. It is stable only above 96<\/span>o<\/span><\/sup>C. When it cools down below 96<\/span>o<\/span><\/sup>C, it changes to rhombic sulphur i.e., 96<\/span>o<\/span><\/sup>C is the transition temperature of this sulphur.<\/span><\/p>\n

        Plastic Sulphur<\/span><\/strong><\/p>\n

        On heating sulphur, till almost the boiling point and suddenly cooling it by pouring into cold water a viscous mass is formed. This sudden cooling does not allow sufficient time to the molecules to rearrange themselves to form monoclinic or rhombic forms of sulphur. Hence the molecules form an interwined mass, consisting of both rhombic and monoclinic varieties of sulphur. This is called plastic sulphur\u00a0<\/span><\/i><\/p>\n

        This type of sulphur is a dark brown or even black, sticky substance. It is elastic. It has no sharp melting point. It does not dissolve in carbon disulphide. On standing, it slowly changes to the rhombic forms, as it gains the eight atom ring structure.<\/span><\/p>\n

        Colloidal Sulphur<\/strong><\/p>\n

        This type of sulphur is prepared by passing hydrogen sulphide through a cooled saturated solution of sulphur dioxide in water, or by adding a solution of sulphur and alcohol in water. Colloidal sulphur is soluble in carbon disulphide. It is used in medicine.<\/span><\/i><\/p>\n<\/div>\n

        <\/h4>\n

        Sulphur dioxide<\/h4>\n
        Properties of sulphurdioxide<\/i><\/span><\/strong><\/div>\n
        \n
          \n
        • Sulphur dioxide ‘SO<\/span>2<\/span><\/sub>‘ a colorless gas with a pungent and suffocating odor. It is readily soluble in cold water, sparingly soluble in hot water, and soluble in alcohol, acetic acid, and sulfuric acid.<\/span><\/li>\n<\/ul>\n
            \n
          • It can be produced by the reaction of sulphur with oxygen, by burning sulphur in air, and it is often produced during the roasting of sulfide ores, e.g, zinc smelting.<\/span><\/li>\n<\/ul>\n
              \n
            • It is a non-combustible and non-supporter of combustion.<\/span><\/li>\n<\/ul>\n

              Magnesium burns in sulphur di oxide giving magnesium oxide and sulphur.<\/span><\/p>\n

              <\/span><\/i><\/span><\/p>\n

              Potassium reacts with sulphur dioxide forming potassium sulphite and potassium thiosulphate.<\/span><\/i><\/span><\/p>\n

               <\/p>\n

                \n
              • An aqueous solution of sulphur dioxide is acidic in nature<\/span><\/li>\n<\/ul>\n
                  \n
                • Sulphur dioxide is called an \u201c acid anhydride of sulphurous acid\u201d<\/span><\/li>\n<\/ul>\n
                    \n
                  • Sulphur dioxide dissolves in water to give unstable sulphurous acid<\/span><\/li>\n<\/ul>\n

                    <\/span><\/i><\/span><\/p>\n

                     <\/p>\n<\/div>\n

                    \n

                    Sulphur dioxide reacts with alkali solution to give salt and water. [sulphites and bisulphites] When sulphur di oxide is\u00a0<\/span><\/i>bubbled through sodium hydroxide solution no precipitate is formed since sodium sulphite is a soluble salt.When sulphur dioxide is passed through Ca (OH)2<\/span><\/sub> solution, a white precipitate of insoluble calcium sulphite is\u00a0<\/span>formed. If excess sulphur dioxide is passed, the precipitate disappears forming soluble calcium bi sulphite.<\/span><\/p>\n

                    <\/span><\/i><\/span><\/p>\n

                    <\/span><\/i><\/span><\/p>\n

                     <\/p>\n

                      \n
                    • Reacts with carbonates and liberates carbon dioxide.<\/span><\/li>\n<\/ul>\n

                      <\/span><\/p>\n

                      <\/span><\/i><\/p>\n<\/div>\n

                      <\/div>\n
                      \n
                        \n
                      • Reacts with basic oxide to give sulphites.<\/span><\/li>\n<\/ul>\n

                        <\/span><\/i><\/span><\/p>\n

                        <\/i><\/span><\/p>\n

                        <\/span><\/i><\/span><\/p>\n

                          \n
                        • Sulphur dioxide is a strong oxidizing agent.<\/span><\/li>\n<\/ul>\n

                          Oxidises magnisium to magnisium ion(Mg<\/span>+<\/span><\/sup>2<\/span><\/sup>)<\/span><\/p>\n

                          <\/span><\/i>Oxidises hydrogen silphide to sulphur<\/span><\/p>\n

                          <\/span><\/i><\/p>\n

                            \n
                          • Sulphur dioxide is a strong reducing agent.<\/span><\/li>\n<\/ul>\n

                            In the presence of moisture SO2<\/sub> liberated nascent hydrogen and reduction takes place by addition of hydrogen.<\/span><\/i><\/span><\/p>\n

                            Example:<\/span><\/i><\/span><\/p>\n

                            <\/span><\/i><\/span><\/p>\n

                            Potassium dichromate, potassium permanganate and nitric acid are reduced by the action of SO2<\/sub> by removal of\u00a0<\/span><\/i>oxygen.Example:<\/p>\n

                            <\/span><\/i><\/span><\/p>\n

                            <\/span><\/i>Sulphur dioxide gas exhibits bleaching properties in presence of moisture. It dissolves in water liberating nascent\u00a0<\/span>hydrogen. Coloring matter is bleached by reaction with nascent hydrogen. Nascent hydrogen removes oxygen\u00a0<\/span>atoms from the coloring matter (reduces coloring matter) and it loses its color. This bleaching is temporary\u00a0<\/span>because the bleached product on exposure to atmospheric oxygen adds on oxygen atoms from air and regains its\u00a0<\/span>original color.<\/span><\/p>\n

                            Uses of sulphur dioxide<\/i><\/span><\/strong><\/p>\n

                            The various uses of sulphur dioxide are:<\/span><\/p>\n

                            1) In the manufacturing of sulphuric acid, sulphites, and hydrogen sulphite.<\/span><\/p>\n

                            2) In the sugar industry for refining and decolorizing sugar.<\/span><\/i>3) For refining kerosene, and other petroleum products.<\/span><\/p>\n

                            4) As a disinfectant.<\/span><\/i>5) As a fumigant.<\/span><\/p>\n

                            6) For bleaching delicate articles.<\/span><\/i>7) As antichlor, to remove the excess chlorine from substances that have been bleached by chlorine.<\/span><\/p>\n

                            8) As a solvent for glue.<\/span><\/i>9) As a refrigerant in household refrigerators.<\/span><\/p>\n

                            10) As a preservative for wines, meat, dry fruits etc.<\/span><\/i><\/p>\n<\/div>\n

                            <\/div>\n

                            <\/h4>\n

                            Sulphuric acid<\/i><\/span><\/h4>\n
                            \n
                            <\/div>\n
                            \n

                            Sulphuric acid, H2<\/sub>SO4<\/sub> is a colorless, odorless, extremely corrosive, oily liquid. It was initially called oil of vitriol.<\/span><\/p>\n

                              \n
                            • It occurs in the free state and combined state.<\/span><\/li>\n<\/ul>\n
                                \n
                              • It is also formed near sulphide beds.<\/span><\/li>\n<\/ul>\n<\/div>\n

                                 <\/p>\n

                                Uses of sulphuric acid<\/i><\/span><\/strong><\/p>\n

                                  \n
                                • In the manufacture of fertilizers, ammonium phosphate and calcium super phosphate.<\/span><\/li>\n<\/ul>\n
                                    \n
                                  • In the manufacture of rayon and nylon and also in the preparation of dyes and drugs from coal tar derivatives.<\/span><\/li>\n<\/ul>\n
                                      \n
                                    • In the manufacture of the explosives such as Tri-nitro toluene , Tri-nitro glycerine and picric acid.<\/span><\/li>\n<\/ul>\n
                                        \n
                                      • In the manufacture of nitric acid, hydrochloric acid and phosphoric acid.<\/span><\/li>\n<\/ul>\n
                                          \n
                                        • In the manufacture of sodium sulphate for glass industry and ferrous sulphate for ink industry.<\/span><\/li>\n<\/ul>\n
                                            \n
                                          • In the purification of petrol, kerosene, and lubricants.<\/span><\/li>\n<\/ul>\n
                                              \n
                                            • It is used in metallurgy for extraction of metals. Leaching of metallic compounds gives sulphates which on electrolysis gives the metal in pure form .It is used for pickling of metals.<\/span><\/li>\n<\/ul>\n
                                                \n
                                              • It is used in storage of batteries.<\/span><\/li>\n<\/ul>\n<\/div>\n

                                                 <\/p>\n

                                                Halogens<\/h3>\n
                                                  \n
                                                • Fluorine is such a powerful oxidising agent that you can’t reasonably do solution reactions with it.<\/li>\n
                                                • Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine can’t get those electrons back from the chloride ions formed.That means that chlorine is a more powerful oxidising agent than either bromine or iodine.<\/li>\n
                                                • Similarly bromine is a more powerful oxidising agent than iodine. Bromine can remove electrons from iodide ions to give iodine – and the iodine can’t get them back from the bromide ions formed.<\/li>\n<\/ul>\n

                                                  This all means that oxidising ability falls as you go down the Group.<\/p>\n

                                                   <\/p>\n

                                                    \n
                                                  • Fluoride and chloride ions won’t reduce concentrated sulphuric acid.<\/li>\n
                                                  • Bromide ions reduce the sulphuric acid to sulphur dioxide. In the process, the bromide ions are oxidised to bromine.<\/li>\n
                                                  • Iodide ions reduce the sulphuric acid to a mixture of products including hydrogen sulphide. The iodide ions are oxidised to iodine.<\/li>\n
                                                  • Reducing ability of the halide ions increases as you go down the Group.<\/li>\n<\/ul>\n

                                                     <\/p>\n

                                                     <\/p>\n

                                                    Noble gases<\/h3>\n

                                                    Position of noble gas in periodic table<\/span><\/i><\/strong><\/p>\n

                                                    Noble gases are also known as inert gases and do not take part in chemical reactions. They have their outermost shell complete and thus remain stable. They do not generally combine with other substances, nor are they affected by oxidising agents or by reducing agents. They are placed in the 18 or VIIIA group. Since, the outermost shell is complete, the valency is zero, hence VIIIA group is also referred to as zero group.<\/span><\/p>\n

                                                    occurrence of noble gas<\/span><\/i><\/strong><\/p>\n

                                                    Noble gases always occur in free state because of their inert nature. All the noble gases, except radon are present in air in small amounts. The relative abundance of the noble gases in air is = 1%. Helium is present in natural gas up to the extent of 10 per cent. It is also present in small quantities in the minerals of radioactive elements. Water from certain springs also gives off gases, which are rich in helium and argon.<\/span><\/i><\/p>\n

                                                    Uses of noble gas<\/span><\/i><\/strong><\/p>\n

                                                    Argon is used mainly to provide an inert atmosphere in high temperature metallurgical processes (arc welding of metals or alloys) and for filling electric bulbs. It is also used in the laboratory for handling substances that are air-sensitive.<\/span><\/i><\/p>\n

                                                    Helium is a non-flammable and light gas. Hence, it is used in filling balloons for meteorological observations. It is also used in gas-cooled nuclear reactors. Liquid helium (boiling point 4.2K) finds use as cryogenic agent for carrying out various experiments at low temperatures. It is used to produce and sustain powerful super conducting magnets, which form essential part of modern NMR spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical diagnosis.<\/span><\/p>\n

                                                    Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes. There are no significant uses of xenon and krypton. They are used in light bulbs designed for special purposes.<\/span><\/i><\/p>\n","protected":false},"excerpt":{"rendered":"

                                                    Aluminium Aluminum (also called Aluminium) is the third most abundant element in the earth’s crust. \u00a0It is commonly used in the household as aluminum foil, in crafts such as dyeing and pottery, and also in construction to make alloys. In its purest form the metal is bluish-white and very ductile. 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