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GROUP 6thA

in , , , on Saturday, June 20, 2009

Introduction
  • Oxygen , sulphur , selenium , tellurium and polonium belong to VIA group of periodic table . O, S, Se and Te are also known as chalcogens (ore forming elements). Polonium is the only radioactive element in the group. It was discovered by Curie. Most abundant element in earth crust is oxygen , Oxygen occurs freely in the atmosphere of which it comprises 20.8% by volume . With carbon , nitrogen and hydrogen , it is essential to living matter. Water contains 89% by weight of combined oxygen . Oxygen is perhaps , the most abundant element on the surface of moon also . The approximate abundance of oxygen in the crystal rocks of the earth is about 45% by weight .

  • Sulphur occurs uncombined in many parts of the world. Selenium and tellurium are comparatively rare elements. Polonium, the radioactive element, has 27 isotopes, which are all radioactive. Of these 210Po occurs naturally.

  • Oxygen is the second most electronegative element after fluorine, but in the VIA group oxygen is the most electronegative element. Oxygen was discovered by Priestley. In laboratory it is prepared by heating KClO3 in presence of MnO2 as catalyst. Oxygen is a principal component of water and of the rocks and minerals in the earth's crust. About 21% of the components of dry air consist of O2 molecules. The second form of the element oxygen, called ozone is present in very small amounts in the upper atmosphere and protects us from harmful effects of ultraviolet rays from the sun. Oxygen is obtained as a result of thermal decomposition of salts containing anions (e.g. KNO3, KClO3, KMnO4 etc. Certain higher oxides, such as PbO2, Pb3O4 etc also liberate oxygen on thermal on thermal decomposition. On large scale, oxygen is prepared by fractional distillation of liquid air. Oxygen is collected by downward displacement of water. Any drying agent can be used for absorption of O2. But alkaline pyrogallol is most suitable. Oxygen molecule is paramagnetic and best explained by MO theory. The normal Oxidation State of oxygen is -2. Oxygen can not form more than two bonds because of non availability of d-orbitals in its outer shell . The density, melting point (K) and boiling point (K) of oxygen are 1.14 g/ cc, 54.20 and 90.0 respectively. The ionisation energy of oxygen is 314-k. cals/ mole.

  • The molecules O2 and O3 called allotropes, because these are different forms of the same element oxygen. On absorbing UV radiation, ozone decomposes to diatomic O2 molecules. Among the elements in the periodic table, only the noble gases, halogens of elements with oxygen are exothermic and hence release heat. Many of such reactions are slow at room temperatures but the rate of reaction increases as the temperature is increased. With metallic elements, oxygen forms ionic compounds containing either O2-(oxide ion), the O22- (peroxide ion) or the O2- (supper oxide ion). Oxygen reacts with metals to form oxides, many of which are not soluble in water. For example, CuO and Ag2O are insoluble in water. Oxygen reacts with non-metals to form a large number of molecular compounds, such as CO, CO2, NO, NO2, SO2 , SO3 , H2O, P2O5 (P4O10) etc. Oxygen always exists as diatomic (O2) molecule. Oxygen is an essential component of many organic compounds, such as fats, carbohydrates, proteins etc. The outer electronic configuration of the elements of VIA group is ns2np4.

  • The metallic character increases from oxygen to polonium. The group VI a elements show almost purely non-metallic chemistry with feeble metallic properties becoming just precipitable as the atomic number increases. Oxygen is highly electronegative, sulphur is less electronegative than oxygen, and Both O and S are non-metals. Se and Te are metalloids and Po is weakly electropositive metal which is rare as well as radioactive . The general increase in metallic character on descending the group is demonstrated by (a) Increasing basic and ionic character of the oxides (b) Increasing tendency to form complexes such as SeBr62- TeBr62- and PoI62- (c) Decreasing stability of hydrides (d) Increase in density (e) The decrease in ionisation energy , electronegativity and electron affinity from O to Po, an (f) A more negative reduction potential for the conversion of the elements to the -2 oxidation state. The metallic properties therefore begin to appear in Te in a less stable hydride, Te H2 and in amphoteric character in the oxide, TeO2. I.P. and electronegativity decreases form O to Po.

  • Oxygen shows -2 oxidation state in most of its compounds, but +2 in OF2 , + 1 in O2F2 and -1 in peroxides . All the other elements show oxidation states of -2, +2, +4 and +6. It is due to the presence of d-orbitals in them.

  • Since atom of these elements are two electrons short of the noble gas configuration, they are expected to have tendencies to achieve noble gas configuration by gaining or sharing of two electrons. The electronegativity of oxygen is very high. So its compounds with most metallic elements can be more but for other elements, only those compounds with the most electropositive elements can be more than 50% ionic. Essentially covalent bonding resulting results with a variety of less electronegative metals and non-metals and in these the VIA group elements may be considered to have a formal charge of -2

  • Tendency to form chain is less evident with Oxygen. It forms O-O links in peroxides and in O2F2; in Ozone and O4F2 the chain contains three O atoms and in O4F2 there are 4 oxygen atoms. All species with O3 and O4 chain are stable.Sulphur exhibits strong tendency to catenation. The enhanced tendency of catenation in S is due to fairly large bond energy of the S-S bond .

  • Due to its high electronegativity, oxygen shows an oxidation state of 2- in all of its compounds except (a) In OF2, where it has +2 oxidation state (b) In peroxides, where it has -1 oxidation state (c) In super oxides (e.g.) KO2), where it exhibits -1/2 oxidation state.

  • The positive oxidation states of +2, +4 and +6 become increasingly more common as we move from S to Po. The +4 and +6 states are more common and stable. The +4 State is most stable for Te and the VI or II state is most stable for Po. In +4 oxidation state, these elements show oxidising and reducing properties , but in +6 state, they are predominantly oxidising ( such as H2SO4 ) .

  • Compounds of S to Po in +2 state are few and comparatively unstable. For example, SO , SeO , TeO , are unstable . Odd positive oxidation states, such as +3 or +5 are rare. Bonding in positive oxidation states is covalent.

  • Due to fairly large bond energy of the S - S bond, sulphur exhibits a strong tendency to catenation. Although bond energies of Se-Se and Te -Te bonds are smaller, the tendency to form chains persists in the crystallised form of these compounds.

  • The catenation property of sulphur is maximum in the group. Tendency to form chains is less evident with oxygen. It forms O-O links in peroxides and in O2F2 ; in ozone and O3F2 the chain contains three O atoms and in O4F2 there are 4 oxygen atoms . All species with O3 and O4 chains are stable. Sulphur exhibits tendency to catenation. The enhanced tendency of catenation in S is due to fairly large bond energy of the S-S bond.

  • Density, melting point and boiling point increase from O to Po. Oxygen is diatomic molecule, sulphur, selenium and tellurium are octa atomic. For example, sulphur molecule is octa atomic (S8) and it has a puckered ring structure. All the elements in oxygen family exist in allotropic forms. Two allotropic forms of oxygen are O2 and O3. Sulphur exist in forms like rhombic monoclinic, plastic etc. Rhombic sulphur is stable upto 95.6°C. Monoclinic sulphur is stable above 95.6C°. At 95.6°C both forms exist in equilibrium. This temperature is called transition temperature. Most abundant and common form of sulphur is rhombic sulphur. Oxygen and sulphur are non-conductors, selenium and tellurium are semi conductors and polonium is conductor .

  • All the elements of VIA group form hydrides of the type H2M, (H2O, H2S, H2Se, H2T2 and H2Po). The stability of hydrides decreases from H2O to H2Po.The acidic nature of hydrides decreases from H2O to H2Po , ( H2O<>2S <>2Te <>2Po). The reducing power of the hydrides increases from H2O to H2Po , (H2O <>2S <>2Se <>2Te 0°C ).The volatility of these hydrides decreases from H2S to H2Te .Decreasing order of bond angles is H2O ( 104°) > H2S ( 92.5 °) > H2 Se (90°) > H2 Te ( 89°)H2O ( 0.95A°) <>2 S ( 1.45A° ) <>2 Te ( 1.72 A°) .All the hydrides , except water are gases at room temperature .All hydrides , except water are volatile , poisonous and bad smelling .In acid solution , the acidity of hydrides increase from H2S to H2Te. All the hydrides have bent structure. The bond angle decrease from H2O to H2 Te .Due to intermolecular hydrogen bonding, water is liquid and has high melting and boiling point compared to other hydrides of the group elements. .Increasing order of H-A bond length is :

  • In the formation of water molecule, oxygen atom undergoes sp3 Hybridisation. Because of the repulsion of non-bonding electron pairs, the tetrahedral bond angle is decreased from 109° 28 ' to 104° 5'. In the case of other hydrides the bond angle is nearly 90°. This indicates that in these hydrides only pure valence p-orbitals re involved in covalent bonding.

  • Thermal stability, bond angle , bond strength and basic character decreases from oxygen to polonium .

  • All the elements of this form halides of the type M2 X2 , MX2 , MX2 , MX4 and MX6 ( M = S, Se, Te and X = halogens ) in which oxidation states of chalcogens are +1, +2 , +4 an +6 respectively .

  • Oxygen combine with chlorine to form a number of oxides such as Cl2O (chlorine monoxide), ClO2 (chlorine dioxide), Cl2O62O7(chlorine heptoxide) in which the oxidation states of chlorine are +1, +4, +6, and +7 respectively. (chlorine hexoxide) and Cl

  • Elements other than oxygen generally form tetrachlorides of the type SCl4, SeCl4, TeCl4 and PoCl4, SCl4 is unstable liquid while other are stable solids.

  • The tetrahalides have distorted trigonal bipyramid structure. S, Se and Te form fluorides of the type SF6, SeF6 and TeF6. These can be prepared by direct combination of these elements (S, Se and Te) with fluorine. All these fluorides undergo hydrolysis and degree of hydrolysis increases with increase in atomic number. (TeF6, SeF6 and SF6). These fluorides have octahedron structure due to sp3d2 Hybridisation of the central atom.

  • All these elements form oxides of the type of MO2 and MO3 (M=S, Se, Te and Po). Sulphur Forms SO2 and So3.

  • Due to its high electronegativity oxygen shows an oxidation state of -2 in all of its compounds expect (a) In OF2, where it has +2 oxidation state (b) In peroxides, where it has -1 oxidation state (c) In super oxides (e.g KO2) where it exhibits -1/ 2 oxidation state.

  • The Positive oxidation states of +2, +4 and +6 become increasingly more common as we move from S to Po. The +4 and +6 states are more common and stable. The +4 state is most stable for Te and the VI or II stage is mo0st stable for Po. In +4 oxidation state, these elements show oxidation and reducing properties, but in +6 state they are predominantly oxidising (such as H2SO4).

  • Compounds of S to Po in +2, states are few and comparatively unstable. For example So, SeO, TeO are unstable. Odd positive oxidation states such as +3 or +5 are rare. Bonding in positive oxidation states is covalent.

  • Due to fairly large bond energy of the S S bond sulphur exhibits a strong tendency to catenation. Although bond energies of Se Se and Te Te bond are smaller the tendency to form chains persists in the crystallised form of these compounds.

  • As the metallic character increases, the acidic nature of oxides gradually increases. Reducing power gradually decreases and oxidising power increases.

  • S. Se and Te produce trioxides of the type SO3, SeO3 and TeO3. The acidic nature of these oxides is in the order SO3 > SeO3 > TeO3. Te and Po also form monoxides of the type TeO, PoO.

  • The ionic character of dioxide increases and acidic character decreases will an increase in atomic number. The solubility of dioxides in water decreases from SO2 to TeO2.

  • The strength of resulting oxyacids decreases in the order H2SO3 > H2SeO3 > H2TeO3 .

  • Thermal stabilities of trioxides is in the order TeO3 > SeO3 > SO3. The trioxides are acidic in nature. They dissolve in water to form ic acids. The strength of ic acids decreases in the order H2SO4 > H2SeO4 > H2TeO4.

  • Sulphur forms a number of oxyacids. These are sulphoxylic acid (H2SO2), sulphurous acid (H2SO3) sulphuric acid (H2SO4), pyrosulphurous acid (H2S2O5), thiosulphuric acid (H2S2O3), pyrosulphuric acid (H2S2O7) dithionic acid (H2S2O6) , peroxymonosulphuric acid or Caro's acid (H2SO5) and peroxydisulphuric acid or Marshall's acid (H2S2O8).

  • Peroxymonosulphuric acid (H2SO5) and peroxydisulphuric acid (H2S2O8) are the peroxy acids with O O linkage.

  • The dioxides of VIA group elements dissolve lin water to give ous acids of the type (H2MO3). PoO2 in insoluble in water but soluble in acids.In the 'ous acids the oxidation state of elements is +4.

  • The trioxides of VIA group elements form ic acids of the type H2MO4. The strength of ic acids follows the order H2SO4 > H2SrO42TeO4. In these acids the oxidation number of elements is +6. > H

  • Sulphurous acid is unstable and acts as oxidising and reducing agent. Sulphuric acid acts as oxidising dehydrating and pickling agent. King of chemicals is H2SO4. It is hygroscopic.

  • Ozone an allotrope of oxygen is largely present in the upper layer of atmospher (stratosphere) and prevents the powerful UV radiation from the sun. Ozone is bluish (Pale blue) gas with fishy odour. In liquid state it is blue in colour. It is more soluble in water than oxygen. It is soluble in turpentine oil. It acts as oxidising reducing and bleaching agent. It tails mercury by oxidation. Its bleaching action is due to oxidation. It turns benzidine paper brown, starch iodine paper blue and tetraethyl base paper violet. A mixture of ozone and cyanogen, (CN2) is used ar rocket fuel. Ozone is diamagnetic and has angular structure. The OO bond length is 1.27 A°. O3 reduces H2O2 to water. J Ozone was first noticed by Van Marum. J Schenbien named it ozone because of its peculiar smell. Soret (1866) consider it as an allotrope of oxygen and he also studied its triatomic nature . Ozone is largely present in the upper layer of the atmosphere, where it is formed by the action of U.V rays on O2. Ozone can be prepared byn subjecting pure dry O2 to electric discharge in an apparatus known as ozonizer. There are two important types of ozonizers - Siemen ozonizer and Brodie's ozonizer. When F2 is passed into water a mixture of O2 and O3 is formed . Ozone is pale blue gas at room temperature with a characteristic fishy odour. In liquid state ozone is blue in colour. Ozone is slightly more soluble in water than oxygen. Ozone is soluble in terpentine oil and aborbed by it . Solubility of O3 in water is more than that of O22 is an endothermic reaction and is accompanied by decrease in volume. O3 readily decomposes to O2 at room temperature. Ozone acts as a powerful oxidising agent, reducing agent and bleaching agent. Ozone bleaches vegetable colours by oxidation. Ozone tails mercury by oxidation In some cases all the three atoms of ozone molecule are used up for oxidation. For example,
     because of its polarity. J Formation of ozone from O

  • In general in the oxidation of SO2, SeO2, TeO2 etc to their higher oxides, all the three O atoms of O3 are used up. Ozone is known as dry bleach as it bleaches in the absence of moisture. Ozone forms addition compounds with organic compounds containing carbon- carbon multiple bonds and these addition compounds are known as ozonides. Thus ozone is used to located the position of carbon-carbon multiple bond in the organic compounds.. It is used as a germicide and disinfectant for purification of air and sterilisation of water. It is used to purify air in underground railways, zoos, mines etc. Ozone is diamagnetic and has angular structure. OO bond length in O3 is 1.27 A° which lies between single (1.48A°) and double bond lengths (1.10 A°) due to resonance. The bond angle is 116°49'.O3 bleaches coloured matter by oxidation.

  • H2SO4 removes water from oxalic and formic acids. H2SO4 is manufactured by lead chamber process and contact process. In lead chamber process oxides of nitrogen act as catalyst and act as carriers of oxygen from air SO2. In contact process V2O5 is used as a catalyst. H2SO4 forms hydrates with explosive violence. Its affinity for water is very high. Metals like Zn, Mg, Na, K, Al Sn etc. Which lie above hydrogen in electrochemical series react with dilute H2SO4 to liberate H2. Metals like Pb Au, Pt do not react with H2SO4. Charring of sugar in presence of H2SO4 is due to dehydration. The dehydrating action of H2SO4 can be explained by the formation of monohydrate (H2SO4.H2O) and dihydrate (H2SO4. 2H2O).

  • Compounds consisting of oxygen and a second element are called oxides. Example are : K2O,CaO,Fe2O3,CO2,P2O5,SO3,Cl2O7,OsO4 etc. All chemical elements except He, Ne and Ar form oxides. The chemical bond between oxygen and the second element in oxides may be ionic or covalent. According to chemical properties, oxides are classified as salt forming and non salt forming. N2O , NO and SiO2 are non- salt forming oxides. The salt forming oxides are subdivided into basic acid and amphoteric oxides. Basic oxides are those whosw hydrates are bases. For example, Na2O and CuO are basic oxides because their corresponding hydrates, NaOH and Cu(OH)2 are bases. Mono and divalent metal oxides belong generally to basic oxides. The chemical bonding in this case is ionic. Oxides of alkali metals (Li, Na, K, Rb, Cs, Fr) and those of alkaline earth metals (Ca, Sr, Ba, Ra) react with water to form bases. The remaining basic oxides do not practically react with water. All basic oxides react with acids to from salts and water. Basic oxides also react with oxides to give salts.

  • Acid oxides are those whose hydrates are acidic. Non metal and mental oxides belong to his class oxides. Examples are N2O3, P2O5, CrO3, Mn2O7, CO2, V2O5 , SO3, Cl2O7. These are acid oxides because their corresponding hydrates HNO2, H3PO4, H2CrO4, HMnO4, H2CO3, H3VO4, HClO4 are acid. The chemical bonding in the acid oxides may be ionic and covalent. Most acid oxides react with water to yield acids. SiO2 does not react with water. Acid oxides with bases or alkalies to form salts and water

  • The metal oxides while depending upon the medium, exhibit either basic or acidic properties, that is , react with both acids and bases are called aqmphoteric oxides. Hydrates of amphoteric oxides can exist both in acidic and basic forms.

  • Amphoteric oxides react with both acids and bases

  • The Amphoteric oxides do not directly add to water. Generally oxides of metals are basic and those of non- metals are acidic. The basic nature of oxides in a group increases from top to bottom and acidic nature of oxides decreases. The basic nature of oxides decreases in a period from left to right and acidic nature of oxides increases.

  • The most acidic oxide in second period is N2O5, formed by nitrogen. The increasing order of acidic nature of oxides in third period is Na2O <>2O3 <>2 <>2O5 <>3 <>2O7.

  • Gun powder is a mixture of sulphur, charcoal and KNO3 .

  • H2S is weak dibasic acid. It is dried over P2O5 or CaCl2. So2 bleached due to reducation. SO2 act as an oxidising and reducing agent. In SO2, S is sp2 hybridised. Thus bgond angle in SO2 is 120°. SO2 molecule has an angular shape. Shape of SO3 molecule is planar triangular. The Hybridisation of S is sp2 and bond angle is 120°. SO42- ion is tetrahedral because of sp3 hybridization. O3, NO, NO2, KO2 etc. have three electron bonds. O2 molecule is paramagnetic due to unpaired electrons.

  • Sulphur and iodine do not combine to form any compound. However some solid solution of sulphur and iodine are known.



This entry was posted on Saturday, June 20, 2009 at 4:02 PM and is filed under , , , . You can follow any responses to this entry through the RSS 2.0. You can leave a response.

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    # by Rajiv - June 21, 2009 at 2:11 PM

    Can i get notes for inorganic chemistry in detail for free download....for iit-jee

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