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P-BLOCK ELEMENTS

Group - 16 Elements

      Oxygen is very essential for living plants and animals. Ozone present in stratosphere protects us from harmful U.V. radiation coming from the sun. 46.6% Oxygen is present in earth's crust by mass and 20.946% in air by volume. S is found in wool, hair, onion, mustard, garlic, proteins, eggs, volcanoes, sulphates like gypsum (CaSO4 . 2 H2O), Epsom salt (MgSO4 . 7 H2O), baryte (BaSO4), sulphides like Zinc blende (ZnS), galena (PbS), copper pyrites (CuFeS2). 16th group consists O, S, Se, Te and Po. As these elements are forming ores, called "Chalcogens". O, S are non metals, Se, Te are metalloids, Po is radioactive metal. Oxygen is diatomic, other elements are octa atomic.

      General electronic configuration of these elements is ns2 np4. Atomic and ionic radii increases ionisation enthalpy, electron gain enthalpy, electronegativity decreases from top to bottom in the group. They show lower value of first ionisation enthalpy compared to group 15 elements (half - filled p - orbitals).
Due to compact nature of oxygen atom, it has less negative electron gain enthalpy than that of sulphur. Oxidation states exhibited by oxygen are -2 (in normal oxides: CaO, Na2O), −1 (in peroxides: H2O2, Na2O2), O (in O2, O3),  (in super oxides: KO2), +1 (in O2F2), +2 (in OF2), other elements exhibit +2, +4, +6 oxidation states. Due to small atomic size, high electronegativity, absence of d - orbitals, oxygen exhibit anomalous behaviour.

      These group elements can form MO2, MO3 types of oxides. Both these types of oxides are acidic in nature. Reducing property of oxides decreases from SO2 to TeO2.

     So SO2 is reducing agent, TeO2 is oxidising agent. Oxides can be mixed (Pb3O4, Fe3O4), acidic (CO2, N2O5, SO3, Cl2O7, Mn2O7, CrO3, V2O5) or basic (Na2O, CaO, BaO).

      The oxides neutral towards acids and bases (CO, NO, N2O) are called neutral oxides, reacts with both the acids and bases are called amphoteric oxides (H2O, Al2O3, ZnO, BeO, PbO, SnO2).

HALIDES: They form M2X2, MX4, MX6 type of halides.

These elements form mono, di, tetra and hexa halides.
Mono halides
    They are dimeric in nature and undergo disproportionation reaction. Structures are similar to H2O2 (Open Book).


DIHALIDES
They are formed by heating mono halides with halogens. S2Cl2 + Cl2  2 SCl2 (foul smelling, dark red liquid).
Their structures are similar to water i.e. angular or
V shape. Central atom undergo sp3
hybridization.   = 103°


TETRA HALIDES
As these halides (fluorides) preparation is difficult, are prepared indirectly.
    S + 4 CoF3 S F4 + 4 CoF2
Central atom in these halides undergo sp3d hybridization and gives see-saw shape.
Tetra halides can act both as lewis acids and lewis bases.
SF4 is highly reactive gas, powerful fluorinating agent. It is unstable, undergoes hydrolysis.

SF4 + 2 H2O  4 HF + SO2
SCl4 is unstable liquid, hydrolyse easily.
SCl4 + 4 H2O  S (OH)4 + 4 HCl
S(OH)4 H2SO4 + H2O
SeCl4 is solid that sublimes. TeCl4 is a solid and hygroscopic.     

 

HYDRIDES
These elements form H2M type hydrides.
Central atom M of these hydrides undergo sp3 hybridization.
They have angular shape.
Except H2O other hydrides are poisonous, foul smelling and highly volatile.
    Thermal stability, volatility, covalent nature, bond angles, bond strength decreases from H2O to H2Te.
Acidic and reducing character increases from H2O to H2Te.
    Due to hydrogen bonding in H2O, it exists in liquid state and possesses high B.P. than expected order of B.P. :   H2O  >  H2Te > H2Se  > H2S
Acidic/ reducing character:  H2Te >  H2Se > H2S > H2O

Bond angles:  H2O     >    H2S     >  H2Se     > H2Te     > H2Po
                 (104° 28')    (92° 30')     (91°)        (90°)         (90°)
* H2O is most stable less volatile, has high B.P., more bond angle and non poisonous. H2S has lowest B.P., so more volatile.
As H − S bond dissociation energy is lesser than that of H − Te bond dissociation energy. H2S is less acidic than H2Te.

 

O2 (Dioxygen)
      O2 is colurless, odourless gas. It is para magnetic. It can be prepared in the laboratory as follows.
      2 Ag2O  4 Ag + O2
      2 HgO  2 Hg + O2
      2 H2O2 2 H2O + O2
On large scale, O2 can be prepared either by electrolysis of water or by fractional distillation of liquid air. O2 reacts with metals and non metals, but not reacts with noble metals like Au, Ag.
4 Al + 3 O2 2 Al2O3
    P4 + 5 O2  P4O10
    CH4 + 2 O2  CO2 + 2 H2O
    C2H4 + 3 O2 2 CO2 + 2 H2O
   

Uses of O2:
* Used in oxyacetylene welding.
* Oxygen cylinders are used in hospitals, high altitude flying and mountaineering.
* Provides tremondous thrust in rockets by the combustion of Hydrazine in liquid oxygen.


O3 (OZONE)
     O3 is allotropic form of oxygen. O3 is a pale blue gas, O3 is diamagentic. If O3 is present in the troposphere, it causes green house effect. If the concentration of O3 exceeds 100 ppm, breathing becomes very difficult leads to nausea and headache. High concentration of O3 is explosive. (ΔH = −ve, ΔS = +ve, ΔG = −ve)

Preparation of O3:
O3 can be prepared from oxygen. By passing silent electric discharge through pure, dry, cold oxygen, 10% ozone is formed. Pure O3 can be condensed in a vessel surrounded by liquid oxygen.
        3 O2  2 O3            ΔH = +142 K.J./mole.

Properties of O3:

* Compared to oxygen O3 is thermodynamically, unstable as it easily decomposes to give nascent oxygen.
* O3 O2 + (O)
* O3 acts as powerful oxidising agent.
* O3 oxidises black PbS to white PbSO4.
    PbS + 4 O3 PbSO4 + 4 O2
* O3 oxidises moist KI to I2.
     2 KI + H2O + O3 I2 + O2 + 2 KOH
* O3 oxidises Ag to Ag2O.
     2 Ag + O3 Ag2O + O2

* Hg loses lustre and meniscus and sticks to glass surface when it reacts with O3. This phenomenon is called "Tailing of mercury". This reaction can be used to test O3. Also O3 turns starch iodide paper blue.
     2 Hg + O3 Hg2O + O2
Meniscus of Hg can be regained by shaking it with water (due to dissolution of Hg2O)
* Alkenes, Alkynes, benzene can form ozonoides with O3.
 
    
* O3 reacts with NO emitted by supersonic jet planes. So Ozone layer is depleted.
     NO + O3 NO2 + O2

Structure of O3:
O3 is angular. O undergoes sp2 hybridization.
Due to resonance, both the single and double bonds have bond length 128 pm. The bond angle is 117°.


Uses of O3:
* as a germicide, as a disinfectant.
* for sterilising water.
* for bleaching oils, starch, ivory.
* as oxidising agent in the manufacture of KMnO4.

 

SULPHUR
  Among the allotropic forms of S, yellow rhombic (α − sulphur) and monoclinic (β − sulphur) forms are very important. Rhombic sulphur is stable at room temperature and transforms to monoclinic sulphur at 369 K (transition temperature).
             
Rhombic sulphur has M.P. 385.8 K. It is insoluble in water but soluble in CS2, alcohol, benzene, ether. Its specific gravity is 2.06.

  Monoclinic sulphur has M.P. 392 K. It is soluble in CS2. Its specific gravity is 1.98. Above 1000 K, S2 is the dominant species and paramagnetic due to presence of 2 unpaired e− in Π* orbitals. S6 has chair form. Both α & β sulphur have S8 rings with crown (puckered) shape.

SO2
SO2 can be prepared either by burning S or sulphide ore.
 S + O2

 SO2
 4 FeS2 + 11 O2  2 Fe2O3 + 8 SO2
 It is colourless gas with pungent smell. Its B.P. is 263 K and highly soluble in water.
  SO2 + H2O  H2SO3

 Structure of SO2:
S undergoes sp2 hybridisation. SO2 has Angular (V) shape.
Due to resonance, S − O bonds have equal length & strength.

 

Uses of SO2:

* Used to bleach wool and silk.
* Used as antichlor, preservative, disinfectant.
* Used as solvent in its liquid state.
* Used in refining petroleum and sugar.


OXO ACIDS OF SULPHUR


 

H2SO4

It is a colourless, dense, oily liquid with a specific gravity 1.84, B.P.: 611 K,
M.P.: 283 K. As its heat of hydration is very high, while diluting the acid, it must be added to water slowly in small quantities.

Manufacture of H2SO4:

H2SO4 is manufactured by contact process in 3 stages as follows:
1) Burning of S or sulphide ore in air to give SO2.
     S + O2  SO2
     4 FeS2 + 11 O2  2 Fe2O3 + 8 SO2
2) Oxidation of SO2 to SO3 at optimum pressure 2 bar and optimum temperature 720 K (Lechatlier's principle) in presence of a catalyst V2O5.
   


3) Absorption of SO3 by H2SO4 to give Oleum.
    SO3 + H2SO4  H2S2O7 (Oleum)
    Oleum on dilution gives H2SO4
   

Properties of H2SO4:

It is dehydrating agent, moderately strong oxidising agent, less volatile.
 Ka1 (> 10) is larger than Ka2 (1.2 × 10−2).
 Cu + 2 H2SO4 
 CuSO4 + 2 H2O + SO2
 3 S + 2 H2SO4 
 2 H2O + 3 SO2
 C + 2 H2SO4 
 2 H2O + CO2 + 3 SO2
 CaF2 + H2SO4 
 2 HF + CaSO4

Charring takes place when reacts with carbohydrates.
 

Uses of H2SO2:

¤ Used in the manufacture of chemical fertilizers calcium super phopsate and ammonium sulphate.
¤ Used as a laboratory reagent.
¤ Used in lead batteries (storage batteries)
¤ Used in petroleum refining and detergent industry.
¤ Used in the manufacture of paints, dyes, pigments.
¤ Used in cleansing of metals. 

Group - 17 Elements

      Bleaching of cloths, chlorination of water, etching of glass, Teflon coated nonsticky utensils, tincture of iodine, photographic film .... are some of the applications of Group - 17 elements. F, Cl, Br, I, At are the members of the group. In greek 'halo' means salt, 'gen' means producer. As they form salts, they are called 'Halogens'. F2, Cl2 are gases. Br2 is liquid. I2 is solid. They exhibit different colours due to absorption of different quanta of radiation. F2 is yellow, Cl2 is greenish yellow, Br2 is red and I2 is violet.

     F & Cl are fairly abundant and F is found in Fluorspar (CaF2), Cryolite (Na3AlF6) and Fluorapatite [CaF2 . 3 Ca3(PO4)2], Cl is found in the form of NaCl in oceans (2.5% by mass), rock salt (NaCl), Carnallite (KCl . MgCl2 . 6 H2O).

      General electronic configuration of these elements is ns2 np5. Atomic & ionic radii increases from top to bottom in the group. Ionisation enthalpy, electronegativity, electron gain enthalpy decreases down the group.

F is having highest electronegativity, Cl is having highest electron gain enthalpy. Electron gain enthalpy of F2< Cl2 due to small atomic size of F, repulsions between electron pairs already present in it and the added electron, the tendency of addition of electron becomes less in F. So the electron gain enthalpy of F2 is less (−333 K.J./ mole) than that of Cl2 (−349 K.J./ mole).

The order of electron gain enthaply is: Cl2 > F2 > Br2 > I2.
The order of X − X bond dissociation enthalpy is: Cl2 > Br2 > F2 > I2.

The reason for the anomaly is due to small atomic size of F, less bond distance of F − F (1.48 A°), more lone pair - lone pair electron repulsions the bond dissociation of F2 is (158.8 K.J./ mole) less than those of Cl2 (242.6 K.J./ mole) and Br2 (192.8 K.J./ mole). F2 is strongest oxidising halogen. A halogen oxidises halide ions of higher atomic number.

F2 + 2 X−  2 F− + X2 (X = Cl, Br, I).
F has anamalous behaviour due to its small atomic size, highest electronegativity, low F − F bond dissociation enthalpy and non availability of d − orbitals in the valency shell. Though electron gain enthalpy of F2 is less electronegative than Cl2, F2 is stronger oxidising agent than Cl2. It is due to high hydration enthalpy of F− (515 K.J./ mole) and low enthalpy of dissociation of F − F bond (158.8 K.J./ mole).

All halogens react with H2. The order of acidic strength: HF < HCl < HBr < HI. They form OF2, O2F2 like oxides with oxygen. Both are strong fluorinating agents. OF2 is thermally stable at room temperature. Oxides of Cl are highly reactive oxidising agents and tend to explode. ClO2 is bleaching agent (for paper pulp and textiles) oxides of Br and I also very powerful oxidising agents. F exhibits −1 oxidation state where as other halogens exhibit +1, +3, +5 and +7 oxidation states. Due to non availability of d − orbitals and high electronegativity of F, it exhibits only −1 oxidation state.


CHLORINE

Properties of Cl2:



* With H2O: Cl2 reacts with water to liberate nascent oxygen, which is responsible for bleaching action (due to oxidation).
     
      HOCl  HCl + (O)
  Coloured substance + (O)  colourless substance. Bleaching effect of Cl2 is permanent.
* With hydro carbons: Cl2 gives substituted compounds with alkanes & addition compounds with alkenes & alkynes.


    
Uses of Cl2:
* for bleaching cotton & wood pulp.
* in the extraction of gold & Pt.
* in the sterilisation of water.
* in the preparation of poisonous gases like Phosgene (COCl2), tear gas (CCl3NO2), Mustard gas (Cl CH2 CH2SCH2 CH2Cl).
* in the manufacture of DDT, dyes, drugs, CHCl3, CCl4 etc.


HCl
It was prepared by Glauber in 1648 as follows.
     
HCl is colourless, pungent smelling gas. Its B.P. 189 K, M.P. 159 K. It is strong acid. Conc. HCl & conc. HNO3 when mixed in 3 : 1 ratio, 'AQUA REGIA' is formed, which dissolves gold & Pt due to the formation of their complex compounds. [Tetra chloro aurate (III) & Hexa chloro platinate (IV)].
3 HCl + HNO3  NOCl (Nitrosyl chloride) + 2 H2O + 2 (Cl)
Au + 3 Cl  AuCl3                  AuCl3 + HCl  H [AuCl4]
Pt + 4 Cl  PtCl4                   PtCl4 + 2 HCl  H2 [PtCl6]

HCl decomposes carbonates, sulphites, sulphides etc.

Uses of HCl:
* as a laboratory reagent.
* in the manufacture of Cl2, NH4Cl, glucose.
* in the purification of gold & Pt.


Oxoacids of Halogens:
F forms only one oxoacid HOF, due to its small size and high electronegativity. Br & I can form HBrO3, HIO3, HBrO4 and HIO4. Cl can form HClO, HClO2, HClO3 and HClO4.


INTERHALOGEN COMPOUNDS
     The compounds which are formed between 2 different halogens are called interhalogen compounds. The general formula of these compounds is AXn. Where n = 1, 3, 5 or 7. A is larger halogen, X is smaller halogen. Oxidation state of X is always −1, but A could be +1, +3, +5 or +7. They can be prepared as follows.

F2 + ClF3 ClF5
Br2 + 3 F2 2 BrF3
Br2 + 5 F2  2 BrF5
         (excess)
I2 + Cl2  2 ICl
(equimolar)
I2 + 3 Cl2

 2 ICl3
      (excess)
AX type interhalogens: ClF, BrF BrCl, ICl
AX3 type interhalogens: ClF3, BrF3, ICl3
AX5 type interhalogens: BrF5, IF5
AX7 type interhalogens: IF7.
     AX type interhalogen compunds are linear, AX3 compounds are T − shape (sp3d hybridisation), AX5 compounds are square pyramidal (sp3d2 hybridisation), AX7 compounds are pentagonal bipyramidal (sp3d3 hybridisation).

Interhalogen compounds are more reactive than Cl, Br, I. This is due to weak bond in interhalogen compound (A − X). As X − X bond is stronger compared to A − X, ICl is more reactive than I2. They are used as non − aqueous solvents & fluorinating agents (ClF3, BrF3).
 

Group - 18 Elements

          He (Helium), Ne (Neon), Ar (Argon), Kr (Krypton), Xe (Xenon), Rn (Radon)     belongs to 18th group in periodic table. As these elements found in air (except Rn), they are called "Aerogens". Their abundance in air is 1% by weight. Due to stable ns2 np6 octet (He = 1s2) configuration, they have high ionization enthalpy and positive electron gain enthalpy. They neither lose nor gain nor share the electrons and called "Noble (inert) gases". Due to inert nature of these elements, discovery of these elements took more than 100 years.
           P.J.C. Janssen & J.N. Lockyer found a new element Helium (Helios = Sun) in 1868 during total solar eclipse. Henry Cavendish found that N2 separated from air contain (1/125th part by volume of N2) a gas "Argon" (Argoes = Lazy). Abundance of Ar is maximum (0.934%) in air among other noble gases. Ramsay & Travers identified a gas "Neon" (Neos = New) during fractional distillation of liquid Argon. By repeated fractional distillation of liquid air, Ramsay found another gas "Krypton" (Krypton = Hidden). He also found "Xenon" (Xenon = Stranger) from Krypton. Radon is obtained during the decay of Ra − 226.
                          88Ra226 86Rn222 + 2He4
Due to inertness, they are very useful. Helium is used in gas cooled atomic reactors (as a heat transfer agent). 'He' is used for creating inert atmosphere during welding of Mg & Al. O2 − He mixture (20% O2 + 80% He) is used in the treatment
of Asthma and also for artificial respiration in deep sea diving (if air is used instead of 'He', N2 present in air dissolved in blood under high pressure and causes pain called "Caisson Sickness" or "bends"). 'He' is used as cryogenic agent (B.P.: 4.2 K) in NMR (Non Magnetic Resonance) spectrometers and MRI (Magnetic Resonance Imaging) systems used for clinical diagnosis. 'He' is non inflammable and light gas (but heavier than H2), used in filling baloons for meteorological observations.
             Neon (Ne) bulbs are used in botanical gardens. Ne is used in fluorescent bulbs for advertisement display purposes. 'Ne' is used in warning signals (e.g.: air ports) as 'Ne' lights are visible from longer distances even in fog and mist. When 'Ne' is mixed with other gases, they produce different colours, used in Neon signs (in advertisements).
             Argon (Ar) is used in filling fluorescent lamps and electric bulbs to increase the life of the filament (lamp). Ar is used to provide inert atmosphere during arc welding of metals and alloys.
Krypton is used in electric bulbs, in measuring thickness of metal sheets, Kr - 85 is used in electronic tubes.
         Xenon is used in electric bulbs, in high speed photographs, in bubble chamber to identify neutral measons and gamma photons.
         Radon is radioactive and used in the treatment of cancer, in the detections of defects (cracks, void space) in metals and solids, in industrial radiography.
         Noble gases are mono atomic (Cp/Cv = 1.67), odourless, colourless, inert.
They have very low M.P. & B.P. due to weak dispersion forces (Vanderwaals forces).
These gases are slightly soluble in water. Except Helium, all the noble gases are absorbed by activated coconut charcoal. Due to stable electronic configuration, they exhibit very high ionization enthalpies and positive electron gain enthalpies (earlier electron affinity of noble gases was taken as zero). Upto 1962 people thought that noble gases do not combine with other elements, so they does not form compounds.
Xe has low ionization enthalpy (large atomic size) among the noble gases. In 1962 Neil Bartlett observed that PtF6 reacts with O2.
            O2 + PtF6 O2+[PtF6]-
'He' also observed that first ionization enthalpy of Oxygen (1175 K.J./ mole) is very closer to first ionization enthalpy of Xe (1170 K.J./ mole) and thought that PtF6 also reacts with Xe and prepared first noble gas compound XePtF6.
          
         Due to high electronegativity values of F & O, they can form compounds with noble gases (like Xe). Xe can form XeF2, XeF4, XeF6, XeO3, XeO4, XeO2F2 and XeOF4 compounds.

 

Xenon Compounds
Xe directly reacts with F2 under different conditions to form Xenon Fluorides.

XeF2, XeF4 & XeF6 are crystalline & sublime at 298 K. They are readily hydrolysed.  



XeF6 + RbF Rb+ [XeF7]−
XeO3 is a colourless explosive solid.
XeOF4 is a colourless volatile liquid.



Posted Date : 06-08-2021

గమనిక : ప్రతిభ.ఈనాడు.నెట్‌లో కనిపించే వ్యాపార ప్రకటనలు వివిధ దేశాల్లోని వ్యాపారులు, సంస్థల నుంచి వస్తాయి. మరి కొన్ని ప్రకటనలు పాఠకుల అభిరుచి మేరకు కృత్రిమ మేధస్సు సాంకేతికత సాయంతో ప్రదర్శితమవుతుంటాయి. ఆ ప్రకటనల్లోని ఉత్పత్తులను లేదా సేవలను పాఠకులు స్వయంగా విచారించుకొని, జాగ్రత్తగా పరిశీలించి కొనుక్కోవాలి లేదా వినియోగించుకోవాలి. వాటి నాణ్యత లేదా లోపాలతో ఈనాడు యాజమాన్యానికి ఎలాంటి సంబంధం లేదు. ఈ విషయంలో ఉత్తర ప్రత్యుత్తరాలకు, ఈ-మెయిల్స్ కి, ఇంకా ఇతర రూపాల్లో సమాచార మార్పిడికి తావు లేదు. ఫిర్యాదులు స్వీకరించడం కుదరదు. పాఠకులు గమనించి, సహకరించాలని మనవి.

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