Group 16 Elements and Properties of Oxygen

General Properties:

(i) Ionisation Enthalpy: Ionisation enthalpy of these elements decreases down the group. It is due to an increase in size. However, the elements of this group have lower ionisation enthalpy values compared to those of Group 15 elements. This is due to the fact that Group 15 elements have extra stable half-filled p orbitals electronic configuration.
The oxygen atom has less negative electron gain enthalpy than sulphur because of the compact nature of its shells due to which the electronic repulsion is greater.

(ii) Oxidation State: They show -2, +2, +4, +6 oxidation states. Oxygen does not show +6 oxidation state due to absence of d – orbitals. Po does not show +6 oxidation state due to inert pair effect.
The stability of -2 oxidation state decreases down the group due to increase in atomic size and decrease in electronegativity. Since electronegativity of oxygen is very high, it shows only –2 oxidation state (except in the case of OF2 where its oxidation state is + 2).
+ 4 and + 6 are more common. Sulphur, selenium and tellurium usually show + 4 oxidation state in their compounds with oxygen and + 6 with fluorine. Down the group, the stability of + 6 oxidation state decreases and that of + 4 oxidation state increases (due to inert pair effect).

(iii) Electronegativity: It is the tendency of an atom to attract shared paired electrons towards itself. The order of electronegativity is O > S > Se > Te > Po (regular trend)

(iv) Electron gain enthalpy: Oxygen has less negative electron gain enthalpy than S because of the small size of O.
From S to Po electron gain enthalpy becomes less negative to Po because of an increase in atomic size.  The order of Electron Gain Enthalpy is S > Se > Te > Po > O.

(v) Melting and boiling point: It increases with an increase in atomic number. Oxygen has much lower melting and boiling points than Sulphur because oxygen is diatomic (O2) and Sulphur is Octatomic (S8).

Group 16 Elements and Properties of Oxygen

Chemical Properties:

(i) Hydrides: All the elements of group 16 from hydrides of the general formula H2E where E is the element belonging to group 16. Following are some of the characteristics of these hydrides: H2O, H2S, H2Se, H2Te.

  • Their acidic character increases from H2O to H2Te. This is due to the decrease in bond (H–E) dissociation enthalpy down the group. H2Te > H2Se > H2S > H2O
  • The thermal stability also decreases down the group due to an increase in bond length.
    H2O > H2S > H2Se > H2Te
  • All the hydrides except water possess the reducing property and this character varies in order. H2Te > H2Se > H2S.
  • Bond angles H2O > H2S > H2Se > H2Te. This is due to repulsion between lone pairs.

(ii) Reactivity with Oxygen: All these elements form oxides of the EO2 and EO3 types where E = S, Se, Te or Po.

  • There are two types of oxides – simple oxides(e.g., MgO, Al2O3) and mixed oxides (Pb3O4, Fe3O4).
  • Simple oxides can be further classified on the basis of their acidic, basic or amphoteric character. An oxide that combines with water to give an acid Five more states sue OxyContin maker Purdue Pharma for opioid epidemic british dragon nutraceuticals where to buy anabolic steroids in australia, buy anavar poland – ตลาดปลากัด ปลากัดออนไลน์ ซื้อ ขายปลากัด is called acidic oxide. For example, Acidic oxides. CO2, SO2, SO3, N2O5, N2O3, P4O6, P4O10, Cl2O7, CrO3, Mn2O7, V2O5.
    Cl2O7 + H2O → 2 HClO4
  • Generally, non-metal oxides are acidic but oxides of some metals in higher oxidation states also have acidic character (e.g., Mn2O7, CrO3, V2O5).
    Mn2O7 + H2O → 2 HMnO4
  • The oxide which gives an alkali on dissolved in water is known as basic oxide(e.g., Na2O, CaO, BaO). Generally, metallic oxides are basic in nature. OR They either dissolve in water to form alkalies or combine with acids to form salts and water or combine with acidic oxides to form salts; e.g., Na2O, CaO. CuO, FeO, BaO etc.
    Na2O + H2O → 2 NaOH
    CaO + H2O → Ca(OH)2
    CuO + H2SO4 → CuSO4 + H2O
  • Besides MO2type oxides sulphur, selenium and tellurium also form MO3 type oxides (SO3, SeO3, TeO3). Both types of oxides are acidic in nature. Increasing order of acidic nature of oxides is TeO3 < SeO3 < SO3.
  • Some metallic oxides exhibit a dual behaviour. They show the characteristics of both acidic and basic oxides. Such oxides are known as amphoteric oxides. They react with acids as well as to alkalis. E.g.: Al2O3, Ga2O3 OR These can combine with acids as well as bases for eg., ZnO, Al2O3, BeO, Sb2O3, Cr2O3, PbO etc.
    PbO + 2 NaOH → Na2PbO2 + H2O

    PbO + H2SO4 → PbSO4 + H2O
    Cr2O3 + 2 NaOH → Na2Cr2O4 + H2O
    Cr2O3 + 3 H2SO4 → Cr2(SO4)3 + 3 H2O
  • There are some oxides that are neither acidic nor basic. Such oxides are known as neutral oxides. Examples of neutral oxides are CO, NO and N2
  • Mixed Oxides: They behave as a mixture of two simple oxides, e.g., Pb3O4 (2PbO + PbO2), Fe3O4 (FeO + Fe2O3), Mn3O4 (2MnO + MnO2).
  • SO2is a gas whereas SeO2 is solid. This is because SeO2 has a chain polymeric structure whereas SO2 forms discrete units.
  • Reducing character of dioxides decreases down the group because oxygen has a strong positive field that attracts the hydroxyl group and removal of H+ becomes easy.

(iii) Reactivity Towards Halogens: Elements of group 16 form a larger number of halides of the type EX6, EX4 and EX2 where E is an element of the group and X is a halogen.

  • The stability of halides decreases in the order F> Cl > Br > I. This is because E-X bond length increases with an increase in size.
  • Among Hexa halides, fluorides are the most stable because of steric reasons.
  • Dihalides are sp3hybridized and so, are tetrahedral in shape.
  • Hexafluorides are only stable halides that are gaseous and have sp3d2hybridization and an octahedral structure.

Group 16 Elements and Properties of Oxygen

Dioxygen (O2)  

(i) Preparation:

  • By heating chlorates, nitrates and permanganates.
    2 KClO3 → 2 KCl + 3O2 (laboratory method)
    2 KMnO4 → K2MnO4 + MnO2 + O2
  • By the thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some.
    2Ag2O(s) → 4Ag(s) + O2(g)
    2Pb3O4(s) → 6PbO(s) + O2(g)
    2HgO(s) → 2Hg(l) + O2(g)
    2PbO2(s) → 2PbO(s) + O2(g)
  • By the decomposition of Hydrogen peroxide (H2O2) in presence of manganese dioxide. 2H2O2 → 2H2O + O2
  • On large scale, it can be prepared from water or air. Electrolysis of water leads to the release of hydrogen at the cathode and oxygen at the anode. It is also obtained by the fractional distillation of air.

Group 16 Elements and Properties of Oxygen

Properties of Oxygen:

  • Colourless, odourless and tasteless gas.
  • It is paramagnetic and exhibits allotropy.
  • Three isotopes of oxygen are 168O, 178O and 188
  • Oxygen does not burn but is a strong supporter of combustion.
  • Dioxygen directly reacts with metals and non-metals(except with some metals like Au, Pt etc and with some noble gases). e.g.
    2Ca + O2 → 2CaO
    P4 + O2 → P4O10

    4Al + 3O2 → 2Al2O3
    C + O2 → CO2

Group 16 Elements and Properties of Oxygen

Uses:

  • Oxygen is used in oxyacetylene welding, in the manufacture of many metals, particularly steel.
  • Oxygen cylinders are widely used in hospitals, high altitude flying and in
  • Liquid O2is used in rocket fuels.

 Ozone

Ozone is an allotropic form of oxygen.
(i) Preparation: It is prepared by passing a silent electric discharge through pure and dry oxygen 10 – 15 % oxygen is converted to ozone.
3 O2(g)→ 2 O3(g); ∆H= +142 kJ/mol

Since the formation of ozone from oxygen is an endothermic process, a silent electric discharge should be used, unless the ozone formed undergoes decomposition.

(ii) Properties:

  • Pure ozone is a pale blue gas.
  • Dark blue liquid and violet-black
  • Ozone has a characteristic smell.
  • It is slightly soluble in water but more in turpentine oil, glacial acetic acid or CCl4.
  • O3molecule is diamagnetic but O3 is paramagnetic.
  • Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen results in the liberation of heat (∆H is negative) and an increase in entropy (∆S is positive). So the Gibbs energy change (∆G) for this process is always negative (∆G = ∆H – T∆S).
  • Due to the ease with which it liberates nascent oxygen (O3→ O2 + O), it acts as a powerful oxidising agent. For e.g.,
  • It oxidises lead sulphide to lead sulphate
  • PbS(s) + 4O3(g) → PbSO4(s) +4O2(g)
  • It oxidises H2S to S
  • H2S + O3→ H2O + S ¯ (yellow)

Oxides of nitrogen (particularly nitric oxide) combine very rapidly with ozone and deplete it. Thus, nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes, slowly depleting the concentration of the ozone layer in the upper atmosphere.
NO(g)+ O3(g) → NO2(g) + O2(g)

Group 16 Elements and Properties of Oxygen

Bleaching Action:
O3 also bleaches coloured substances through oxidation.
O3 → O2 + [O]
                 Nascent Oxygen
Colouring Matter +[O] → Decolourised Matter

Tests for Ozone:
i) A filter paper soaked in alcoholic benzidine becomes brown when brought in contact with O3(this is not shown by H2O2)
ii)Tailing of mercury Pure mercury is a mobile liquid but when brought in contact with O3its mobility decreases and it starts sticking to glass surface forming a type of tail due to the dissolution of Hg2O (mercury sub-oxide) in Hg.
2Hg + O3 → Hg2O + O2

Estimation of Ozone: When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer, iodine is liberated. The liberated iodine can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating O3 gas.

Uses:
(i) It is used as a germicide, disinfectant and for sterilising water.
(ii) It is also used for bleaching oils, ivory, flour, starch, etc.
(iii) It acts as an oxidising agent in the manufacture of potassium permanganate.
(iv) For detecting the position of a double bond in the unsaturated organic compounds.

 

Group 16 Elements and Properties of Oxygen

 

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