Elements of Group-16 lie in P-block. These elements are oxygen, sulphur selenium, tellurium and polonium. They are sometimes known as chalcogens. This name is derived from the Greek word for brass and points to the association of sulphur and its congeners with Copper. Most Copper minerals contain either oxygen or sulphur and sometimes the other members of this group.
Oxygen is the most abundant element on earth. It forms about 46.6% by mass of earth crust. Dry air contains 20.95% oxygen by volume. Sulphur in the earth’s crust is only 0.03-0.1%. Sulphur exists also in combined state as sulphate as Gypsum (CaSO4.2H2O),Epsom salt (MgSO4.7HO), Baryte(BaSO4) and sulphides as Galena(PbS), Zinc blende (ZnS), Copper pyrite (CuFeS2). Selenium and Tellurium are also found as metal sellinides and tellurides in sulphide ores. Polonium exists in nature as decay products of Thorium and uranium minerals.
Important atomic and physical properties of group-16 elements:
(1) The elements of Group16 have six electrons in the outermost shell and have ns2np4 general electronic Configuration.
(2) The atomic and Ionic radii from top to bottom increases in the group due to increase in number Of shells .
(3) Ionisation enthalpy decreases down the group due to increase in the atomic size. But the elements of this group have lower Ionisation enthalpy compared to the elements of Group- 15 in the corresponding periods because group 15 elements have extra stable half-filled electronic Configuration.
(4) Due to smaller size and electronic repulsion, oxygen atom has less negative electron gain enthalpy than sulphur. However, from sulphur onwards the value again becomes less negative upto polonium.
(5) Within the group from top to bottom, Electronegativity decreases with an increase in atomic number. Next to Fluorine, oxygen has the highest electronegativity value amongst the elements.
(6) Oxygen and sulphur are non-metals, Selenium and polonium metalloids, whereas polonium is a metal. All these elements exhibit allotropy. The melting and boiling points increases down the group with an increase in atomic number. The large difference between m.p. and b.p. may be explained on the basis of their atomicity. Oxygen exists as diatomic (O2) whereas sulphur as polyatomic (S8).
Chemical properties:
1.Oxidation state: The elements of Group- 16 show a number of oxidation state. Oxygen shows -2, -1, +1 and +2. Sulphur ,selenium and Tellurium show -2, +2, +4 and +6. Polonium shows +2, and +4 oxidation state. Sulphur, Selenium and Tellurium usually show +4 oxidation state in their compounds with oxygen and +6 with Fluorine. The stability of +6 oxidation state decreases down the group and stability of +4 oxidation state increase due to inert pair effect.
2. Anomalous behaviour of oxygen: The anomalous behaviour of oxygen like other members of p-block elements in second period is due to its small size and high electronegativity. Due to small size and high electronegativity H2O is liquid and H2S is gas. Due to absence of d-orbitals in oxygen atom limits its covalency to four. On the other hand, in case of other elements elements of this group, the covalency exceeds four.
3. Reactivity with Hydrogen: All the elements of this group form Hydrides of the type H2E ( E = S, Se, Te, Po). The increase in acidic characters and decrease in thermal stability down the group from H2S to H2Po is due to decrease in bond dissociation enthalpy. All the Hydrides except water possess reducing property and this character increases from H2S to H2Te.
Properties of Hydrides of group 16 elements:
4. Reactivity with oxygen: All the elements of this group form oxides of the EO2 and EO3 types where E = S, Se, Te and Po. Both type of oxides are acidic in nature. Reducing property of dioxide decreases from SO2 to TeO2. SO2 is reducing while TeO2 is an oxidising agent.
5. Reactivity towards halogens:
Elements of this group form a large number of halides of the type EX6, EX4 and EX2. The stability of halides decreases in the order F– > Cl– > Br– >I–. Amongst hexahalides, hexafluorides are the only stable halides and they are gaseous in nature. They have octahedral structure and SF6 is exceptionally stable for steric reasons. Amongst tetrafluorides , SF4 is a gas, SeF4 is liquid and TeF4 is solid. All elements except Selenium form dichlorides and dibromides. The well known monohalides are dimeric in nature and undergo disproportionation as given below:–
2Se2Cl2 → SeCl4 + 3Se.
Preparation of Dioxygen :
(1) By heating oxygen containing salts such as chlorates, nitrates and permagnates.
2KClO3 → 2KCl + 3O2
(2) By thermal decomposition of the oxides metals low in the electrochemical series and higher oxides of some metals.
2Ag2O(s) → 4Ag(s) + O2(g)
2HgO(s) → 2Hg(s) + O2(g)
(3) Hydrogen peroxide is readily decomposed by catalyst such as finely divided metals and MnO2
2H2O2(aq) → 2H2O(l) + O2(g)
(4) On large scale it can be prepared from water or air. Electrolysis of water leads to release oxygen at anode. After liquification and fractional distillation, oxygen is obtained.
Properties of Dioxygen:
Dioxygen is a colourless and odourless gas. It is slightly soluble in water which is sufficient for the vital Support of marine and aquatic life. It liquifies at 90 K and freezes at 55 K. It has three stable isotopes: 16O, 17O and 18O. It is paramagnetic in nature.
Dioxygen directly reacts with nearly all metals and nonmetals except Au and Pt. Its reaction with other elements is often strongly exothermic which helps in sustaining the reaction. However some external heating is required as bond dissociation enthalpy of oxygen – oxygen double bond is high (493.4) kJ per mole.
Important reactions:
2Ca + O2 → 2CaO
4Al + 3O2 → 2Al2O3
P4 + 5O2 → P4O10
2ZnS + 3O2 → 2ZnO
Simple oxides:
A binary compounds of oxygen with other elements is called oxide. Oxides can be simple or mixed. Simple oxides can be classified as acidic, basic or amphoteric.
Acidic oxide: An oxide that combines with water to give an acid is called acidic oxide. As a general only non- metal oxides are acidic but oxides of some metals in high oxidation state are also acidic. Ex:- SO2, Cl2O7, CO2, Mn2O7, CrO3.
Basic oxides: The oxides which gives base with water are known as basic oxides. In general metallic oxides are basic. Ex. Na2O, CaO. Some metallic oxides exhibit a dual behaviour as both acidic and basic. Ex. Al2O3, ZnO. CO, NO and N2O are neutral oxides.
Ozone(O3)
Ozone is an allotrope of oxygen. It is formed from atmospheric oxygen in the presence of sunlight at a height of about 20 Km and the layer of ozone protects from UV radiation.
Preparation of ozone:
When a slow dry stream of oxygen is passed through a silent electrical discharge. 10% ozone is obtained and this product is known as ozonised oxygen.
3O2 → 2O3 ΔHo= +142 KJmol-1
Since formation of ozone is an endothermic process, it is essential to use a silent electrical discharge in its preparation to prevent its decomposition.
Properties of Ozone:
(1) pure ozone is a pale blue liquid and violet black solid.
(2) Due to the ease with which it liberates atoms of nascent oxygen, it acts as a powerful oxidising agent and oxides lead sulphide to lead sulphate and iodide ions to Iodine.
PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g)
2I– (aq) + H2O(l) + O3(g) → 2OH–(aq) + I2(s) + O2(g)
Test of Iodine
When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer (pH 9.2), Iodine is liberated which can be titrated against a standard solution of Sodium thiosulphate.
Depletion of ozone:
(1) Nitric oxide emitted from the exhaust systems of supersonic jet aeroplanes might be slowly depleting the concentration of ozone layer in the upper atmosphere.
NO(g) + O3(g) → NO2(g) + O2(g)
(2) Another threat to this ozone layer is probably posed by the use of freons which are used in aerosol sprays and as refrigerants.
Structure of ozone:
Allotropic forms of Sulphur
Sulphur forms numerous allotropes. Yellow rhombic (α- sulphur) and monoclinic (β- sulphur ) forms are the most important. Rhombic sulphur is stable at room temperature which can be transformed into monoclinic sulphur when heated above 369 K.
Rhombic sulphur:
This allotrope is yellow in Colour, m.p. 385.8 K and specific gravity 2.06. Rhombic sulphur crystals are formed on evaporating the roll sulphur in CS2. It is insoluble in water but it is readily soluble in CS2.
Monoclinic sulphur:
This form of sulphur is prepared by melting rhombic sulphur in a dish and cooling, till crust is formed. Two holes are made in the crust and the remaining liquid poured out. On removing the crust, colourless needle shaped crystals of B- sulphur are formed. Its m.p. is 393K and specific gravity is 1.98. It is stable above 369K and transform into α-sulphur below this temperature. At 369 K both the forms are stable.
Structure of sulphur:
Both rhombic and monoclinic sulphur have S8 molecules. These molecules are packed to give different crystal structure. The S8 ring in both the forms is puckered and has a crown shape. Several other modifications of sulphur contains 6-20 sulphur atoms per ring have been synthesised in the last two decades. In cyclo- S6, the ring adopts the chair form. At elevated temperature (~1000K), S2 is the dominant species and is paramagnetic like O2.
Sulphur dioxide (SO2)
Preparation of sulphur dioxide:
When sulphur is burnt in air or oxygen, sulphur dioxide is formed together with a little amount of sulphur trioxide.
S(s) + O2(g) → SO2(g)
In the laboratory, it is really generated by treating a sulphate with dilute Sulphuric acid.
SO32- (aq) + 2H+(aq) → H2O (l) + SO2(g)
Industrially, it is produced as a by product of the roasting of sulphide ores.
4FeS2(s) + 11O2(g) → 2Fe2O3(s) + 8SO2(g)
Properties of sulphur dioxide:
(1) Sulphur dioxide is a colourless gas with pungent smell and it is highly soluble in water.
(2) Sulphur dioxide when passed through water, forms a solution of sulphurous acid.
SO2(g) + H2O (l) → H2SO3(aq)
(3) It readily reacts with Sodium hydroxide solution forming sodium sulphite. which then reacts with more sulphur dioxide to form Sodium hydrogen sulphite.
2NaOH + SO2 → Na2SO3 + H2O
Na2SO3 + SO2 + H2O → 2NaHSO3
(4) Sulphur dioxide reacts with chlorine in the presence of charcoal to give sulphuryl chloride.
SO2(g) + Cl2(g) → SO2Cl2(l)
(5) When moist, sulphur dioxide behaves as a reducing agent and converts Fe(lll) ions to Fe(ll) ions and decolourises acidified KMnO4. This reaction is used to test this gas.
2Fe3+ + SO2 + 2H2O → 2Fe2+ SO42- + 4H+
5SO2 + 2MnO4- + 2H2O → 5SO42- + 4H+ + 2Mn2+
The molecules of SO2 is angular and has resonance structure like ozone.
Sulphuric acid: H2SO4
Manufacturing of Sulphuric acid:
Sulphuric acid is prepared by contact process in the following steps-
(1) Burning of sulphur or sulphide ores in air to generate SO2.
(2) Conversion of SO2 in SO3 by the reaction with oxygen in presence of a catalyst V2O5.
2SO2(g) + O2(g) → 2SO3(g) ΔfHo= -196.6 KJmol-1.
(3) Absorption of SO3 in H2SO4 to give oleum (H2S2O7) at 2 bar pressure and 720K temperature. Diffusion of oleum with water gives H2SO4.
SO3 + H2SO4 → H2S2O7
H2S2O7 + H2O → 2H2SO4
The Sulphuric acid obtained by contact process is 96-98% pure.
Properties of H2SO4
(1) Sulphuric acid is a colourless, dense oily liquid with a specific gravity of 1.84 at 298K.
(2) This acid freezes at 283K and boils at 611K. It dissolves in water with the evolution of a large quantity of heat hence during preparation, it must be added into water with constant stirring.
(3) Sulphuric acid is a dibasic acid hence it ionises in two steps in aqueous solution
H2SO4 (aq) + H2O (l) → H3O+(aq) + HSO4–(aq)
HSO4-(aq) + H2O (l) → H3O+(aq) + SO42- Where Ka>>>Kb.
(4) Due to low volatility of Sulphuric acid, it forms more volatile acid from their corresponding salts.
2MX + H2SO4 → 2HX + M2SO4
Where X = F, Cl, NO3.
(5) Concentrated Sulphuric acid is a strong dehydrating agent, hence it removes water from organic compounds.
C12H22O11 → 12C + 11H2O.
(6) Hot concentrated Sulphuric acid is a moderately strong oxidising agent between phosphoric acid and nitric acid. Hence metals and nonmetals are oxidised and itself reduced to SO2.
Cu + 2H2SO4 (conc.) → CuSO4 + SO2 + 2H2O
3S + 2H2SO4 (conc.) → 3SO2 + 2H2O
C + 2H2SO4 (conc.) → CO2 + 2SO2 + 2H2O
Oxaacids of sulphur:
