CBSE NOTES CLASS 12 CHEMISTRY CHAPTER 7

p BLOCK ELEMENTS

GROUP 15 ELEMENTS

The 15th group of the Periodic Table consists of nitrogen. phosphorus. arsenic, antimony and bismuth. These elements are known as pnicogens and their compounds as pniconides.

Occurrence of group 15 elements

Molecular nitrogen comprises 78% by volume of the atmosphere. In the earth’s crust, it occurs as sodium nitrate, NaNO3 (called Chile saltpetre) and potassium nitrate, KNO3 (Indian saltpetre). It is found in the form of proteins in plants and animals.

Phosphorus occurs in minerals of the apatite family, Ca9(PO4)6.CaX2 (X = F, Cl or OH) (e.g., fluorapatite Ca9(PO4)6.CaF2) which are the main components of phosphate rocks. Phosphorus is an essential constituent of animal and plant matter. It is present in bones as well as in living cells. Phosphoproteins are present in milk and eggs.

Arsenic, antimony and bismuth are found mainly as sulphide minerals.

PHYSICAL PROPERTIES OF GROUP 15 ELEMENTS

 Property Nitrogen (N) Phosphorus (P) Arsenic (As) Antimony (Sb) Bismuth (Bi) Atomic number 7 15 33 51 83 Atomic Mass 14.01 30.97 74.92 121.76 208.98 Electronic configuration [He] 2s2 2p3 [Ne]3s2 3p3 [Ar] 3d10 4s2 4p3 [Kr] 4d10 5s2 5p3 [Xe] 4f145d106s26p3 Ionisation enthalpy (ΔH/(kJ mol–1) I 1402 1012 947 834 703 II 2856 1903 1798 1595 1610 III 4577 2910 2736 2443 2466 Electronegativity 3.0 2.1 2.0 1.9 1.9 Covalent radius/pm 70 110 121 141 148 Ionic radius/pm 171 212 222 76 103 Melting point/K 63 317 1089 904 544 Boiling point/K 77.2 554 888 1860 1837 Density/[g cm-3 (at 298 K)] 0.879 1.823 5.778 6.697 9.808 Common Physical Form(s) Colorless Gas White Solid / Red Solid/black Yellow Solid / Grey Solid Yellow Solid / Silver-White Metallic Solid Pink-White Metallic Solid

(i) Electronic configuration

The valence shell electronic configuration of all the elements of this group is ns2np3

They are smaller than the corresponding group 14 elements. They have higher effective nuclear charge compared to group 14 elements and and the electron enters the same shell. Also electrons in the same cell do not screen each other. As a result the electrons are more strongly bound to the nucleus.

The atomic radii increase with increase in atomic number as we go down the group. This is due to increase in number of shells.

The covalent radius increases considerably from N to P but not so much from As to Bi.

From N to P, the increase is due to (a) addition of new shell, (b) strong shielding effect of s- and p-electrons in case of P.

From As to Bi, the d- and f- electrons have poor shielding effect, which reduces the effect of increase due to additional shell.

(iii) Ionization enthalpies

The ionization enthalpies of group 15 elements are considerably higher than those of group 14 elements. This is due to the increase nuclear charge, reduced atomic radii and half filled p electrons.

The ionization energy decreases down the group, because of increasing atomic size.

(iv) Electronegativity

Due to smaller atomic sizes and smaller number of electrons needed to attain noble gas configuration, the elements of group 15 are more electronegative than corresponding group 14 elements.

As we move down a group, electronegativity decreases, owing to increase in size of the atom.

(v) Metallic character

N and P are non-metals, As and Sb are metalloids and Bi is metal.

Due to increase in nuclear charge, higher electronegativity, the elements of group 15 are less metallic than corresponding group 14 elements.

On moving down the group, the metallic character increases due to increase in size, decrease and electronegativity.

(v) Physical state

Nitrogen is a diatomic gas in native form. All other elements in the group are tetra-atomic solids (E4).

(vi) Melting and boiling points

The melting point increases from nitrogen to arsenic. The boiling points increase regularly on moving down the group.

(vii) Density

Increases down the group.

(viii) Allotropy

All the elements (except nitrogen and bismuth) exhibit allotropy.

• Phosphorus – White, red, black

• Arsenic – Grey, yellow

• Antimony – Metallic yellow (explosive), grey

CHEMICAL PROPERTIES OF GROUP 15 ELEMENTS

Catenation

They exhibit the property of catenation but to lesser extent due to weak E – E bond than 14 group elements.

Reactivity

Elemental nitrogen is highly unreactive because of its strong triple bond (almost as inert as noble gases). Because of small size of nitrogen, it is able to form pπ−pπ bonds with itself. This property is not exhibited by atoms such as phosphorus. Thus, phosphorus is more reactive than nitrogen.

White phosphorus is extremely reactive and kept in water. It is inflammable and can be ignited at 45°C.

Oxidation states

 Element N P As Sb Possible Oxidation State -3 to +5 -3, +3, +4, +5, +6 +3, +5, +6 +3, +5, +6

The common oxidation states of these elements are –3, +3 and +5.

As we move down the group, the tendency to exhibit -3 oxidation state decreases. This is due to the increase in atomic size and metallic character.

The stability of +5 oxidation state decreases down the group. The only well characterized Bi (V) compound is BiF5.

Due to the inert pair effect, the stability of +5 oxidation state decreases down the group while that of +3 oxidation state increases.

Nitrogen also exhibits +1, +2, +4 oxidation states when it reacts with oxygen.

In the case of nitrogen, all oxidation states from +1 to +5 tend to disproportionate in acid solution.

5HNO2 → HNO3 + 3H2O + N2O + 2NO + NO2

Nitrogen has only s- and p-orbitals, but no d-orbitals in its valance shell. Therefore, nitrogen can show a maximum covalency of 4. A covalency of four is obtained by sharing its lone pair of electron with another atom or ion.

The heavier elements have vacant d orbitals in the outermost shell which can be used for bonding (covalency) and hence, expand their covalence as in PF6.

Anomalous properties of nitrogen

Nitrogen, the first member of group 15, differs from rest of the group members because of,

• Small size

• High value of electronegativity

• High ionization enthalpy

• Absence of d-orbitals

Differences

• Tendency to form multiple bonds (pπ –pπ)

Nitrogen has ability to form pπ -pπ multiple bonds with itself and with other elements having small size and high electronegativity (e.g., C, O).

Heavier elements of this group do not form pπ -pπ bonds as their atomic orbitals are so large and diffused that they cannot have effective overlapping.

• dπ –pπ bond

Due to absence of d orbital, nitrogen cannot form dπ –pπ bond as the heavier elements can e.g., R3P = O or R3P = CH2 (R = alkyl group). Phosphorus and arsenic can form dπ –dπ bond also with transition metals when their compounds like P(C2H5)3 and As(C6H5)3 act as ligands.

• Catenation and atomicity

Nitrogen exists as a diatomic molecule with a triple bond (one s and two p) between the two atoms. That is why its bond enthalpy (941.4 kJ mol–1) is very high. Other elements form tetratomic molecules such as P4, As4 and Sb4. Phosphorus, arsenic and antimony form single bonds as P–P, As–As and Sb–Sb while bismuth forms metallic bonds in elemental state.

The single N–N bond is weaker than the single P–P bond because of high interelectronic repulsion of the non-bonding electrons, owing to the small bond length. As a result the catenation tendency is weaker in Nitrogen.

• Nature of oxides

Nitrogen forms five oxides of monomeric nature. N2O4 exists in dimeric from and is diamagnetic. Others can form at the most three types of oxides, X4O6, X4O8, X4O10 of dimeric nature.

• Nature of hydrides

Hydride of nitrogen is stable while the hydrides of other elements are not stable and act as reducing agent. Hydrogen bonding is present in ammonia but not present in other hydrides.

• Nature of trihalides

Except NF3, the halides of nitrogen are unstable and explosive. The halides of other elements are stable. Unlike P, As and Sb, nitrogen does not form penta-halides.

• Nitrogen can form tri-negative ion N3-. This tendency is less in P and absent in other elements.

Trends in chemical properties of group 15 elements

Hydrides of group 15 elements

All the elements of this group form hydrides of the type EH3, which are covalent and pyramidal in shape.

 Property NH3 PH3 AsH3 SbH3 BiH3 Melting point/K 195.2 139.5 156.7 185 – Boiling point/K 238.5 185.5 210.6 254.6 290 (E–H) Distance/pm 101.7 141.9 151.9 170.7 – HEH angle (°) 107.8 93.6 91.8 91.3 – Δf H°/kJ mol–1 –46.1 13.4 66.4 145.1 278 ΔdissH° (E–H) / kJ mol–1 389 322 297 255 –

Some properties follow the order in the following order

These properties are

• Thermal stability,

• Basic strength,

• Solubility in water,

• Bond angle NH3(107.4°), PH3 (92°), AsH3 (91°), SbH3(90°)

• Strength of E–H bond

These properties are,

• Reducing character,

• Covalent character,

• Rate of combustion

Halides of group 15 elements

All the elements of this group form tri-halides, EX3 and except nitrogen all form penta-halides, EX5. Trihalides of these elements except those of nitrogen are stable. In case of nitrogen, only NF3 is known to be stable. Trihalides except BiF3 are predominantly covalent in nature.

Trihalides (except those of N) behave as Lewis acid and the order of their strength is

PCl3 > AsCl3 > SbCl3.

Tri-halides of N behave as Lewis base and have the following order of strength,

NF3 < NCl3 < NBr3 < NI3

Oxides of group 15 elements

All the elements of this group form oxides of the type E2O3 and E2O5. Nitrogen forms five oxides. The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state. Their acidic character decreases down the group.

 Oxides of Nitrogen N2O NO N2O3 NO2 N2O5 Neutral Acidic Strongly Acidic Oxides of Phosphorus. P4O10 P2O5 Acidic

The acidic strength of pentaoxide and trioxides decrease on moving down the group, i.e.,

N2O5 > P2O5 > As2O5 > Sb2O5

Reactivity of group 15 elements towards metals

All these elements react with metals to form their binary compounds exhibiting –3 oxidation state, such as, Ca3N2 (calcium nitride) Ca3P2 (calcium phosphide), Na3As2 (sodium arsenide), Zn3Sb2 (zinc antimonide) and Mg3Bi2 (magnesium bismuthide).

NITROGEN AND ITS COMPOUNDS

DINITROGEN (N2)

Preparation of dinitrogen (N2)

Dinitrogen is produced commercially by the liquefaction and fractional distillation of air. Liquid dinitrogen (b.p. 77.2 K) distils out first leaving behind liquid oxygen (b.p. 90 K).

Laboratory method

1. In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.

NH4Cl(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq)

Small amounts of NO and HNO3 are also formed in this reaction; these impurities can be removed by passing the gas through aqueous sulphuric acid containing potassium dichromate.

2. It can also be obtained by the thermal decomposition of ammonium dichromate.

(NH4)2Cr2O7 N2 + 4H2O + Cr2O3

3. Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide.

Ba(N3)2 Ba + 3N2

Physical properties of N2

Dinitrogen is a colourless, odourless, tasteless and non-toxic gas. It has two stable isotopes: 14N and 15N. It has a very low solubility in water (23.2 cm3 per litre of water at 273 K and 1 bar pressure) and low freezing and boiling points.

Chemical properties of N2

Reaction of dinitrogen with metals

Dinitrogen is quite inert at room temperature because of the high bond enthalpy of N≡N bond. Reactivity increases rapidly with rise in temperature. At higher temperatures, it directly combines with some metals to form ionic nitrides.

6Li + N2 2Li3N

3Mg + N2 Mg3N2

Reaction of dinitrogen with dihydrogen

It combines with hydrogen at about 700 K in the presence of a catalyst to form ammonia (Haber’s Process)

N2(g) + 3H2(g) 2NH3(g); ΔfH = – 46.1 kJ mol–1

Reaction of dinitrogen with dioxygen

Dinitrogen combines with dioxygen only at very high temperature (at about 2000 K) to form nitric oxide, NO.

N2 + O2(g) 2NO(g)

Reaction of dinitrogen with non-metals

At high temperature it reacts with non-metals to form covalent nitrides.

2B + N2 2BN

Uses of dinitrogen

• N2 is used in the manufacture of HNO3, NH3, CaCN2 (calcium cyanamide) and other nitrogenous compounds.

• It is used for filling electric bulbs.

• Liquid dinitrogen is used as a refrigerant to preserve biological materials, food items and in cryosurgery.

AMMONIA (NH3)

Ammonia is produced in the atmosphere by decay of urea,

NH2CONH2 + 2H2O → (NH4)2CO3 ↔ 2NH3 + H2O + CO2

Preparation of ammonia

1. Lab method

2NH4Cl + Ca (OH)2 → 2NH3 + 2H2O + CaCl2

(NH4)2SO4 + 2NaOH → 2NH3 + 2H2O + Na2SO4

2. Haber’s process

Ammonia is prepared at high pressure (200 atm) and 700 K, in the presence of catalysts.

N2(g) + 3H2(g) 2NH3(g); ΔfH = – 46.1 kJ mol–1

The optimum conditions for manufacturing ammonia are:

1. Pressure (around 200 × 105 Pa)

2. Temperature (700 K)

3. Catalyst such as iron oxide with small amounts of Al2O3 and K2O

Properties of ammonia

(i) Physical properties of ammonia

Ammonia is a colourless gas with a pungent odour. Its freezing and boiling points are 198.4 and 239.7 K respectively. In the solid and liquid states, it is associated through hydrogen bonds as in the case of water and that accounts for its higher melting and boiling points than expected on the basis of its molecular mass.

(ii) Geometry of ammonia molecule

The molecule is trigonal pyramidal with the nitrogen atom at the apex. It has three bond pairs and one lone pair of electrons.

(iii) Solubility of ammonia

It is extremely soluble in water due to H–bonding. Its aqueous solution is weakly basic due to the formation of OH ions.

NH3(g) + H2O(l) → NH4+ (aq)+ OH- (aq)

(iv) Reaction of ammonia with acids

It forms ammonium salts with acids, e.g., NH4Cl, (NH4)2SO4, etc. As a weak base, it precipitates the hydroxides of many metals from their salt solutions. For example,

Due to presence of a lone pair of electrons on the nitrogen atom of the ammonia molecule, it behaves as a Lewis base. It donates the electron pair and forms linkage with metal ions. Formation of complex compounds is used for detection of metal ions such as Cu2+, Ag+:

(v) Reaction of ammonia with chlorine

When NH3 is in excess, N2 is the main product

8NH3 + 3Cl2 → 6NH4CI + N2

When Cl2 is in excess, NCl3 is the main product

NH3 + 3Cl2 → NCl3 + 3HCl

Uses of ammonia

1. Ammonia is used to produce various nitrogenous fertilizers (ammonium nitrate, urea, ammonium phosphate and ammonium sulphate) and in the manufacture of some inorganic nitrogen compounds, like nitric acid.

2. Liquid ammonia is also used as a refrigerant.

OXIDES OF NITROGEN

 Name Formula OS of N Common methods of preparation State & nature Dinitrogen oxide [Nitrogen(I) oxide] N2O +1 NH4NO3 N2O+2H2O colourless gas, neutral Nitrogen monoxide [Nitrogen(II) oxide] NO +2 2NaNO2 + 2FeSO4 + 3H2SO4 → FeSO4 + 2NaHSO4 + 2H2O + 2NO colourless gas, neutral Dinitrogen trioxide [Nitrogen(III) oxide] N2O3 +3 2NO + N2O4 2N2O3 blue solid, acidic Nitrogen dioxide [Nitrogen(IV) oxide] NO2 +4 2Pb(NO3)2 4NO2+2PbO brown gas, acidic Dinitrogen tetroxide [Nitrogen(IV) oxide] N2O4 +4 2NO2 N2O4 colourless solid/ liquid, acidic Dinitrogen pentoxide [Nitrogen(V) oxide] N2O5 +5 4HNO3 + P4O10 → HPO3+2N2O5 colourless solid, acidic

STRUCTURES OF OXIDES OF NITROGEN

• NO2 contains odd number of valence electrons. That is why it is converted to stable N2O4 molecule with even number of electrons, on dimerisation.

NITRIC ACID (HNO3)

Nitrogen forms oxoacids such as,

• H2N2O2 (hyponitrous acid),

• HNO2 (nitrous acid)

• HNO3 (nitric acid)

Acids containing oxygen and/or –OH group are called oxoacids.

Preparations of HNO3

1. Lab method (heating of KNO3 or NaNO3 with conc H2SO4)

NaNO3 + H2SO4 (conc) → NaHSO4 + HNO3

2. Ostwald’s process

2NO (g) + O2 (g) 2NO2 (g)

3NO2 (g) + H2O (l) → 2HNO3 (g) + NO (g)

NO formed here is recycled.

The aqueous HNO3 can be concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated H2SO4.

Physical properties of HNO3

It is a syrupy, colourless, pungent liquid usually available as 68%. 15.7 M aqueous solution is often yellow due to small concentrations of NO2. At concentration the specific gravity is 1.504.

Structure of HNO3

In the gaseous state, HNO3 exists as a planar molecule with the structure as shown.

CHEMICAL REACTIONS OF HNO3

Acidic character of HNO3

In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions.

HNO3(aq) + H2O(l) → H3O+(aq) + NO3- (aq)

Reactions of HNO3 with metals

Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. The products of oxidation depend upon the concentration of the acid, temperature and the nature of the material undergoing oxidation.

3Cu + 8HNO3 (dilute) → 3Cu(NO3)2 + 2NO + 4H2O

Cu + 4HNO3 (conc.) → Cu(NO3)2 + 2NO2 + 2H2O

4Zn + 10HNO3 (dilute) → 4Zn(NO3)2 + 5H2O + N2O

Zn + 4HNO3 (conc.) → Zn(NO3)2 + 2H2O + 2NO2

Some metals (e.g., Cr, Al) do not dissolve in concentrated nitric acid because of the formation of a passive film of oxide on the surface.

Reaction of HNO3 with non-metals

Concentrated nitric acid also oxidises non–metals and their compounds. Iodine is oxidised to iodic acid, carbon to carbon dioxide, sulphur to H2SO4, and phosphorus to phosphoric acid.

I2 + 10HNO3 (conc.) → 2HIO3 + 10NO2 + 4H2O

C + 4HNO3 (conc.) → CO2 + 2H2O + 4NO2

S8 + 48HNO3 (conc.) → 8H2SO4 + 48NO2 + 16H2O

P4 + 20HNO3 (conc.) → 4H3PO4 + 20 NO2 + 4H2O

Brown ring test of nitrate

Fe2+ reduces nitrates to nitric oxide, which reacts with Fe2+ to form a brown coloured complex. Dilute ferrous sulphate solution is added to an aqueous solution containing nitrate ion, and then adding concentrated sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and sulphuric acid layers indicates the presence of nitrate ion in solution.

FeSO4 + 6H2O → [Fe(H2O)6]SO4

2HNO3 + 3H2SO4 + 6FeSO4 → 3Fe2(SO4)3 + 2NO + 4H2O

[Fe(H2O)6]SO4+ NO → [Fe(H2O)5(NO)]SO4 + H2O

or

NO3- + 3Fe2+ + 4H+ → NO + 3Fe3+ + 2H2O

Uses of nitric acid

1. Used in the manufacture of ammonium nitrate for fertilisers and other nitrates for use in explosives and pyrotechnics.

2. Used for the preparation of nitroglycerin, trinitrotoluene and other organic nitro compounds.

3. Used in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels.

Pickling is a metal surface treatment used to remove impurities, such as stains, inorganic contaminants, rust or scale from ferrous metals, copper, precious metals and aluminum alloys.

4. Used as nitrating reagent

PHOSPHORUS AND ITS COMPOUNDS

Allotropic forms of phosphorus

• White phosphorus

• Red phosphorus

• Black phosphorus

White Phosphorus is less stable and therefore, more reactive than the other solid forms under normal conditions because of angular strain in the P4 molecule where the angles are only 60°. It readily catches fire in air to give dense white fumes of P4O10.

P4 + 5O2 → P4O10

It dissolves in boiling NaOH solution in an inert atmosphere giving PH3.

P4 + 3NaOH + 3H2O → PH3 + NaH2PO2 (sodium hypophosphite)

With HNO3

P4 + 20HNO3 + 3H2O → 4H3PO4 + 4H2O + 20NO2

Red phosphorus is obtained by heating white phosphorus at 573K in an inert atmosphere for several days. Red phosphorus possesses iron grey lustre. Chemically, red phosphorus is much less reactive than white phosphorus. It does not glow in the dark.

Black phosphorous

When red phosphorus is heated under high pressure, a series of phases of black phosphorus is formed. Black phosphorus has two forms α-black phosphorus and β-black phosphorus. α-black phosphorus is formed when red phosphorus is heated in a sealed tube at 803K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals. It does not oxidise in air. β-Black phosphorus is prepared by heating white phosphorus at 473K under high pressure. It does not burn in air upto 673K. It is a good conductor of electricity.

Match box side contains red P or P2S3 + glue and tip of match stick contains red P, KIO3, chalk and glue.

Uses: It is used in match boxes, explosives, as rat poison, in fertilizers and alloys

Distinction between white and red phosphorus

 Property White Phosphorous Red Phosphorous Structure P4 molecules are held by weak vander Waal’s forces. P4 molecules are held by covalent bonds in polymeric structure. Colour and Nature White phosphorus is a translucent white waxy solid. Possesses iron grey lustre. Odour It has garlic like odour. It is odourless. Physiological Action White phosphorus is poisonous. It is non poisonous. Solubility It is insoluble in water but soluble in carbon disulphide Insoluble in water as well as in carbon disulphide. Reactivity/Reaction with KOH/NaOH Highly reactive due to angular strain. Less reactive. No reaction Hardness It is soft waxy solid and can be easily cut with a knife. It is hard and crystalline solid. Phosphorescence White phosphorus glows in dark (chemiluminescence). It does not glow in the dark. Conductivity Bad conductor Semi conductor Ignition Ignition temperature is low (303 K) so burns easily in air. Ignition temperature is high (543 K), so does not catch fire easily. Reaction with Cl2 Burns easily in Cl2 forming PCl3 and PCl5. Combines with Cl2 only on heating.

PHOSPHINE (PH3)

Preparation of phosphine

Phosphine is prepared by the reaction of calcium phosphide with water or dilute HCl.

Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3

Ca3P2 + 6HCl → 3CaCl2 + 2PH3

In the laboratory, it is prepared by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO2.

P4 + 3NaOH + 3H2O → PH3 + 3NaH2PO2 (sodium hypophosphite)

Pure PH3 is non inflammable but becomes inflammable owing to the presence of P2H4 or P4 vapours. To purify it from the impurities, it is absorbed in HI to form phosphonium iodide (PH4I), which on treating with KOH gives phosphine.

PH4I + KOH → KI + H2O + PH3

Properties of phosphine

1. It is a colourless gas with rotten fish like smell and is highly poisonous. It explodes in contact with traces of oxidising agents like HNO3, Cl2 and Br2 vapours.

2. It is slightly soluble in water. The solution of PH3 in water decomposes in presence of light giving red phosphorus and H2. When absorbed in copper sulphate or mercuric chloride solution, the corresponding phosphides are obtained.

3CuSO4 + 2PH3 → Cu3P2 + 3H2SO4

3HgCl2 + 2PH3 → Hg3P2 + 6HCl

3. Phospine is weakly basic

PH3 + HBr → PH4Br

Uses of phosphine

• It is used to prepare smoke screens in warfare.

• A mixture of CaC2 and Ca3P2 is used in Holme’s signals. Containers containing calcium carbide and calcium phosphide are pierced and thrown in the sea when the gases evolved burn and serve as a signal.

PHOSPHORUS TRICHLORIDE (PCl3)

Preparation of PCl3

It is obtained by passing dry chlorine over heated white phosphorus.

P4 + 6Cl2→ 4PCl3

It is also obtained by the action of thionyl chloride with white phosphorus.

P4 + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2

Properties of PCl3

It is a colourless oily liquid, having pyramidal shape [sp3 – hybridised].

It hydrolyses in the presence of moisture.

PCl3 + 3H2O → H3PO3 + 3HCl

It reacts with organic compounds containing –OH group such as CH3COOH, C2H5OH.

3CH3COOH + PCl3 → 3CH3COCl + H3PO3

3C2H5OH + PCl3 → 3C2H5 Cl + H3PO3

PHOSPHORUS PENTACHLORIDE (PCl5)

Preparation of PCl5

Phosphorus pentachloride is prepared by the reaction of white phosphorus with excess of dry chlorine.

P4 + 10Cl2 → 4PCl5

It can also be prepared by the action of SO2Cl2 on phosphorus.

P4 + 10 SO2Cl2 → 4PCl5 + 10 SO2

Structure of PCl5

PCl5 in gaseous and liquid phases has sp3d – hybridization and its shape is trigonal bipyramidal. The three equatorial P – Cl bonds are equivalent while the two axial bonds are longer equatorial bonds.

In solid state, PCl5 exists as an ionic solid, [PCl4]+ [PCl6]- in which, the cation, [PCl4]+is tetrahedral and the anion [PCl6]- is octahedral.

Properties of PCl5

It is a yellowish white powder and in moist air, it hydrolyses to POCl3 and finally gets converted to phosphoric acid.

PCl5 + H2O → POCl3 + 2HCl

POCl3 + 3H2O → H3PO4 + 3HCl

When heated, it sublimes but decomposes on stronger heating.

PCl5 PCl3 + Cl2

It reacts with organic compounds containing –OH group converting them to chloro derivatives.

C2H5OH + PCl5 → C2H5Cl + POCl3 + HCl

CH3COOH + PCl5 →CH3COCl + POCl3 +HCl

Finely divided metals on heating with PCl5 give corresponding chlorides.

2Ag + PCl5 +2AgCl + PCl3

Sn + PCl5 +SnCl2 + PCl3

Uses of PCl5

It is used in the synthesis of some organic compounds, e.g., C2H5Cl, CH3COCl.

OXOACIDS OF PHOSPHORUS

Acids containing oxygen and/or –OH group are called oxaocids.

In oxoacids, phosphorus is tetrahedrally surrounded by other atoms. All these acids contain at least one P=O bond and one P–OH bond.

The oxoacids in which phosphorus has lower oxidation state (less than +5) contain, in addition to P=O and P–OH bonds, either P–P (e.g., in H4P2O6) or P–H (e.g., in H3PO2) bonds but not both.

 Name Formula Oxidation state of P Characteristic bonds Preparation Hypophosphorous (Phosphinic) H3PO2 +1 One P – OH Two P – H One P = O white P4 + alkali Orthophosphorous (Phosphonic) H3PO3 +3 Two P – OH One P – H One P = O P2O3 + H2O Pyrophosphorous H4P2O5 +3 Two P – OH Two P – H Two P = O PCl3 + H3PO3 Hypophosphoric H4P2O6 +4 Four P – OH Two P = O One P – P red P4 + alkali Orthophosphoric H3PO4 +5 Three P – OH One P = O P4O10 + H2O Pyrophosphoric H4P2O7 +5 Four P – OH Two P = O One P – O – P heat phosphoric acid Metaphosphoric* (HPO3)n +5 Three P – OH Three P = O Three P – O – P Heat phosphorus acid + Br2 in a sealed tube

Structures of oxoacids of phosphorus

 Hypophosphorous acid Orthophosphorous acid Pyrophosphorous acid Hypophosphoric acid Orthophosphoric acid Pyrophosphoric acid Cyclotrimetaphosphoric acid, (HPO3)3 Polymetaphosphoric acid (HPO3)n
• Polymetaphosphoric acid, (HPO3)n exists in polymeric forms only. Characteristic bonds of (HPO3)3 have been given here.

• The compositions of the oxoacids are interrelated in terms of loss or gain of H2O molecule or O-atom.

• These acids in +3 oxidation state of phosphorus tend to disproportionate to higher and lower oxidation states. For example, orthophophorous acid (or phosphorous acid) on heating disproportionates to give orthophosphoric acid (or phosphoric acid) and phosphine.

4H3PO3 → 3H3PO4 + PH3

• The acids which contain P–H bond have strong reducing properties. Thus, hypophosphorous acid is a good reducing agent as it contains two P–H bonds and reduces, for example, AgNO3 to metallic silver.

4AgNO3+2H2O+H3PO2 → 4Ag + 4HNO3 + H3PO4

• These P–H bonds are not ionisable to give H+ and do not play any role in basicity. Only those H atoms which are attached with oxygen in P–OH form are ionisable and cause the basicity. Thus, H3PO3 and H3PO4 are dibasic and tribasic respectively as the structure of H3PO3 has two P–OH bonds and H3PO4 three.

GROUP 16 ELEMENTS

The elements oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po) belong to group 16 of the Periodic Table. These elements are known as chalcogens, i.e., ore forming elements.

GENERAL PHYSICAL PROPERTIES OF GROUP 16 ELEMENTS

 Property O S Se Te Po Atomic number 8 16 34 52 84 Atomic Mass 16.00 32.06 78.96 127.60 210.00 Electronic configuration [He] 2s2 2p4 [Ne]3s2 3p4 [Ar] 3d10 4s2 4p4 [Kr] 4d10 5s2 5p4 [Xe] 4f145d106s26p4 Ionisation enthalpy ΔfH1 (kJ mol–1) 1314 1000 941 869 813 Electron gain enthalpy -141 -200 -195 -190 -174 Electronegativity 3.50 2.44 2.48 2.01 1.76 Covalent radius/pma 66 104 117 137 146 Ionic radius/pm 140 184 198 221 230b Density/[g cm-3 (at 298 K)] 1.32c 2.06d 4.19e 6.25 – Melting point/K 55 393f 490 725 520 Boiling point/K 90 718 958 1260 1235 Oxidation states* –2,–1,1,2 –2,2,4,6 –2,2,4, 6 –2,2,4,6 2,4

a - Single bond;

b - Approximate value;

c - At the melting point;

d - Rhombic sulphur;

e - Hexagonal grey;

f - Monoclinic form, 673 K.

* Oxygen shows oxidation states of +2 and +1 in oxygen fluorides OF2 and O2F2 respectively.

Oxidation States

The stability of –2 oxidation state decreases down the group. Polonium hardly shows –2 oxidation state.

Since electronegativity of oxygen is very high, it shows only negative oxidation state as –2 except that it shows oxidation states of +2 and +1 in oxygen fluorides OF2 and O2F2 respectively.

Other elements of the group exhibit + 2, + 4, + 6 oxidation states but + 4 and + 6 are more common. 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 while stability of +4 oxidation state increases (inert pair effect). Bonding in +4 and +6 oxidation states are primarily covalent.

Abundance

Oxygen is the most abundant of all the elements on earth. Oxygen forms about 46.6% by mass of earth’s crust. Dry air contains 20.946% oxygen by volume.

Sulphur in the earth’s crust is only 0.03-0.1%. Combined sulphur exists as sulphates such as gypsum CaSO4.2H2O, epsom salt MgSO4.7H2O, baryte BaSO4 and sulphides such as galena PbS, zinc blende ZnS, copper pyrites CuFeS2.

Traces of sulphur occur as hydrogen sulphide in volcanoes. Organic materials such as eggs, proteins, garlic, onion, mustard, hair and wool contain sulphur.

Density

It increases down the group regularly

Melting point and boiling point

Both show a regular increase down the group due to increase in molecular weight and van der Waals’ forces.

Ionization enthalpy

Ionization energy of group 16 elements is large which decreases gradually from O to Po due to increase in size of atoms and increase in screening effect.

Electron gain enthalpy

They have high electron gain enthalpy, which decrease from O to Po. As the size of the atom increases, the extra added electron feels lesser attraction by nucleus and hence, electron gain enthalpy decreases.

Electronegativity

It decreases down the group due to decrease in effective nuclear charge down the group.

Metallic and non-metallic character

The metallic character increases from oxygen to polonium.

Catenation

Group 16 elements follow the order in catenation as shown below

S-S > Se-Se > O-O > Te-Te

Atomicity

Oxygen is diatomic, sulphur and selenium are octa-atomic with puckered ring structure.

They increase regularly from O to Po.

CHEMICAL PROPERTIES OF GROUP 16 ELEMENTS

Anomalous behaviour of oxygen

The anomalous behaviour of oxygen is due to its small size and high electronegativity. Example of effects of small size and high electronegativity is the presence of strong hydrogen bonding in H2O which is not found in H2S.

Absence of d orbitals in oxygen limits its covalency to four and in practice, rarely exceeds two. On the other hand, in case of other elements of the group, the valence shells can be expanded and covalence can exceed four.

Reactivity of group 16 elements with hydrogen

All the elements of Group 16 form hydrides of the type H2E (E = O, S, Se, Te, Po).

H2O is a liquid due to hydrogen bonding, while others are colourless gases with unpleasant smell.

The acidic character of hydrides of group 16 elements increases from H2O to H2Te. The increase in acidic character can be explained in terms of decrease in bond (H–E) dissociation enthalpy down the group.

Due to the same reason, the thermal stability of hydrides also decreases down the group.

All the hydrides except water possess reducing property and this character increases down the group.

 Property H2O H2S H2Se H2Te m.p/K 273 188 208 222 b.p/K 373 213 232 269 H–E distance/pm 96 134 146 169 HEH angle (°) 104 92 91 90 Δf H/kJ mol–1 –286 –20 73 100 ΔdissH (H–E) /kJ mol–1 463 347 276 238 Dissociation constant at 298K 1.8×10–16 1.3×10–7 1.3×10–4 2.3×10–3

Reactivity of group 16 elements with oxygen

All these elements form oxides of the types EO2 and EO3 where E = S, Se, Te or Po.

Ozone (O3) and sulphur dioxide (SO2) are gases while selenium dioxide (SeO2) is solid.

Reducing property of dioxide decreases from SO2 to TeO2; SO2 is reducing while TeO2 is an oxidising agent.

Sulphur, selenium and tellurium also form EO3 type oxides (SO3, SeO3, TeO3).

Both types of oxides are acidic in nature.

Reactivity of group 16 elements towards halogens

Elements of Group 16 form a large 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 the halides is in the order F > Cl> Br > I–.

Amongst hexahalides, hexafluorides are the only stable halides. All hexafluorides are gaseous in nature. Sulphur hexafluoride, SF6 is exceptionally stable for steric reasons. Both of these have octahedral structure.

Amongst tetrafluorides, SF4 is a gas, SeF4 a liquid and TeF4 a solid.

These fluorides have sp3d hybridisation and have trigonal bipyramidal structures in which one of the equatorial positions is occupied by a lone pair of electrons. This geometry is also regarded as see-saw geometry.

All elements except oxygen` form dichlorides and dibromides. These dihalides are formed by sp3 hybridisation and thus, have tetrahedral structure.

The monohalides (S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2) are dimeric in nature.

These dimeric halides undergo disproportionation as follows,

2Se2Cl2 → SeCl4 + 3Se

OXYGEN AND ITS COMPOUNDS

DIOXYGEN

Preparation of dioxygen

• By heating oxygen containing salts such as chlorates, nitrates and permanganates.

2KClO3 $\stackrel{\mathrm{H}\mathrm{e}\mathrm{a}\mathrm{t}/\mathrm{M}\mathrm{n}{\mathrm{O}}_{2}}{\to }$ 2KCl + 3O2

• By the thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some metals.

2Ag2O(s) → 4Ag(s) + O2(g);

2HgO(s) → 2Hg(l) + O2(g)

2Pb3O4(s) → 6PbO(s) + O2(g)

2PbO2(s) → 2PbO(s) + O2(g)

• Hydrogen peroxide readily decomposes into water and dioxygen by catalysts such as finely divided metals and manganese dioxide.

2H2O2(aq) → 2H2O(l) + O2(g)

• 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.

Industrially, dioxygen is obtained from air by first removing carbon dioxide and water vapour and then, the remaining gases are liquefied and fractionally distilled to give dinitrogen and dioxygen.

Physical properties of dioxygen

It is colourless, odourless, tasteless, slightly heavier than air and sparingly soluble in water. It liquefies at 90 K and freezes at 55 K. Oxygen atom has three stable isotopes: 16O, 17O and 18O. Molecular oxygen, O2 is unique in being paramagnetic inspite of having even number of electrons.

Chemical properties of dioxygen

Dioxygen directly reacts with nearly all metals and non-metals except some metals (e.g., Au, Pt) and some noble gases. In many cases one element forms two or more oxides.

Its combination with other elements is strongly exothermic which helps in sustaining the reaction. However, to initiate the reaction, some external heating is required as bond dissociation enthalpy of oxgyen-oxygen double bond is high (493.4 kJ mol–1).

2Ca + O2 → 2CaO

4Al + 3O2 → 2Al2O3

P4 + 5O2 → P4O10

C + O2 → 2CO2

2ZnS + 3O2 → 2ZnO + 2SO2

CH4 + 2O2 → CO2 + 2H2O

Some compounds are catalytically oxidized,

2SO2 + O2 $\stackrel{{\mathrm{V}}_{2}{\mathrm{O}}_{5}}{\to }$2SO3

4HCl + O2 $\stackrel{\mathrm{C}\mathrm{u}{\mathrm{C}\mathrm{l}}_{2}}{\to }$ 2Cl2 + 2H2O

Uses of dioxygen

• In respiration and combustion

• In welding and cutting using oxy-hydrogen or oxy-acetylene torch

• In iron and steel industry to increase the content of blast

• In life support systems e.g., in hospitals, for divers, miners and mountaineers

• In combustion of rocket fuels, e.g., hydrazines in liquid oxygen, which provides tremendous thrust in rockets.

Tests of O2

• With NO it gives reddish brown fumes of NO2.

• It is adsorbed by alkaline pyrogallol and the solution turns brown.

(pyrogalol)

The oxides vary widely in their nature and properties. Oxides can be simple (e.g., MgO, Al2O3) or mixed (Pb3O4, Fe3O4).

Simple oxides can be classified on the basis of their acidic, basic or amphoteric character. An oxide that combines with water to give an acid is termed acidic oxide (e.g., SO2, Cl2O7, CO2, N2O5). For example, SO2 combines with water to give H2SO3, an acid.

SO2 + H2O → H2SO3

As a general rule, only non-metal oxides are acidic but oxides of some metals in high oxidation state also have acidic character (e.g., Mn2O7, CrO3, V2O5).

The oxides which give a base with water are known as basic oxides (e.g., Na2O, CaO, BaO). For example, CaO combines with water to give Ca(OH)2.

CaO + H2O → Ca(OH)2

In general, metallic oxides are basic. Metallic oxides which exhibit characteristics of both acidic as well as basic oxide, are known as amphoteric oxides. They react with acids as well as alkalies.

Al2O3(s) + 6HCl (aq) + 9H2O(l) → 2 [Al(H2O)6]3+ (aq) + 6Cl- (aq)

Al2O3(s) + 6NaOH (aq) + 3H2O(l) + 2Na3[(Al(OH)6] (aq)

The oxides which are neither acidic nor basic are known as neutral oxides. Examples are CO, NO and N2O.

OZONE (O3)

Ozone is an allotropic form of oxygen. It is too reactive to remain for long in the atmosphere at sea level. It is formed from atmospheric oxygen in the presence of sunlight at a height of about 20 kilometres,. This ozone layer protects the earth’s surface from an excessive concentration of ultraviolet (UV) radiations.

Preparation of ozone

By passing a slow dry stream of oxygen through a silent electrical discharge, oxygen gets converted to ozone (10%). The product is known as ozonised oxygen.

3O2 → 2O3, ΔHo (298 K) = +142 kJ mol–1

Since the formation of ozone from oxygen is an endothermic process, it is necessary to use a silent electrical discharge in its preparation to prevent its decomposition.

If concentration of ozone greater than 10 per cent is required, a battery of ozonisers can be used, and pure ozone (b.p. 385 K) can be condensed in a vessel surrounded by liquid oxygen.

Physical properties of ozone

Pure ozone is a pale blue gas, dark blue liquid and violet-black solid. Ozone has a characteristic smell and in small concentrations it is harmless. However, if the concentration rises above 100 parts per million, breathing becomes uncomfortable resulting in headache and nausea.

Chemical reactions of ozone

(i) Decomposition of ozone

Ozone is thermodynamically unstable with respect to oxygen since its decomposition into oxygen results in the liberation of heat (ΔH is negative) and an increase in entropy (ΔS is positive). These two effects reinforce each other, resulting in large negative Gibbs energy change (ΔG) for its conversion into oxygen. Therefore, high concentrations of ozone can be dangerously explosive.

2O3 → 3O2; ΔH = -284 kJ/mol

(ii) Oxidising action

Due to the ease with which it liberates atoms of nascent oxygen (O3 → O2 + [O]), it acts as a powerful oxidising agent.

PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g)

2FeSO4 + H2SO4 + O3 → Fe2(SO4)2 + H2O + O2

It liberates iodine from neutral KI solution and the liberated I2, turns starch paper blue.

2KI (aq) + H2O(l) + O3(g) → 2KOH(aq) + I2(s) + O2(g)

I2 + Starch → Blue Colour

Nitrogen oxides (particularly nitric oxide) combine very rapidly with ozone and there is, thus, the possibility that nitrogen oxides emitted from the exhaust systems of supersonic jet aeroplanes might be slowly depleting the concentration of the ozone layer in the upper atmosphere.

NO(g) + O3(g) → NO2 (g) + O2 (g)

(iii) Resonance of O3

The two oxygen-oxygen bond lengths in the ozone molecule are identical (128 pm) and the molecule is angular with a bond angle of about 117o.

Uses of ozone

• As a germicide and disinfectant for sterilizing water.

• As a bleaching agent for oils, ivory wax and delicate fibres.

• For detecting, the position of double bond in unsaturated compounds.

• In destroying odours coming from cold storage room, slaughter houses and kitchen of hotels.

• In the manufacture of potassium permanganate.

ALLOTROPES OF SULPHUR

Sulphur forms many allotropes of which the yellow rhombic (α-sulphur) and monoclinic (β-sulphur) forms are the most important.

The stable form at room temperature is rhombic sulphur, which transforms to monoclinic sulphur when heated above 369 K.

Rhombic sulphur (α-sulphur)

It is yellow in colour with m.p. 385.8 K and specific gravity 2.06. Rhombic sulphur crystals are formed on evaporating the solution of roll sulphur in CS2. It is insoluble in water but dissolves to some extent in benzene, alcohol and ether. It is readily soluble in CS2.

Roll sulphur is sulphur in the form of rods or sticks made by casting molten sulphur

Monoclinic sulphur (β-sulphur)

Colourless, m.p. is 393 K and specific gravity 1.98. It is soluble in CS2.

It 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 β-sulphur are formed.

It is stable above 369 K and transforms into α-sulphur below it. Conversely, α-sulphur is stable below 369 K and transforms into β-sulphur above this. At 369 K both the forms are stable. This temperature is called transition temperature.

α-sulphur β-sulphur

Both rhombic and monoclinic sulphur have S8 molecules. These S8 molecules are packed to give different crystal structures. The S8 ring in both the forms is puckered and has a crown shape.

Several other modifications of sulphur containing 6-20 sulphur atoms per ring have been synthesized. In cyclo-S6, the ring adopts the chair form.

At elevated temperature (~1000 K), S2 is the dominant species and is paramagnetic like O2.

SULPHUR DIOXIDE (SO2)

Method of preparation of sulphur dioxide

Sulphur dioxide is formed together with a little (6-8%) sulphur trioxide when sulphur is burnt in air or oxygen:

S(s) + O2(g) → SO2(g)

In the laboratory it is readily generated by treating a sulphite 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) + 8 SO2 (g)

The gas after drying is liquefied under pressure and stored in steel cylinders, because it is a strong oxidising and reducing agent.

Physical properties of sulphur dioxide

It is a colourless gas with pungent smell, highly soluble in water. It liquefies at room temperature under a pressure of two atmospheres and boils at 263 K.

Chemical reactions of sulphur dioxide

Structure of dulphur dioxide

The molecule of SO2 is angular. It is a resonance hybrid of the two canonical forms:

Reaction of sulphur dioxide with water

Sulphur dioxide, when passed through water, forms a solution of sulphurous acid.

SO2(g) + H2O(l) H2SO3 (aq)

Reaction of sulphur dioxide with lime water

It turns lime water milky due to the formation of calcium sulphite.

Ca(OH)2 + SO2 → CaSO3 (milky) + H2O

However, in excess of SO2 milkiness disappears due to the formation of calcium bisulphite.

CaSO3 + SO2 + H2O → Ca(HSO3)2

Reaction of sulphur dioxide with sodium hydroxide

It reacts readily with sodium hydroxide solution, forming sodium sulphite, which then reacts with more sulphur dioxide to form sodium hydrogen sulphite.

NaOH + SO2 → Na2SO3 + H2O;

2Na2SO3 + SO2 + H2O→ NaHSO3

Reaction of sulphur dioxide with chlorine

Sulphur dioxide reacts with chlorine in the presence of charcoal (which acts as a catalyst) to give sulphuryl chloride, SO2Cl2.

SO2 (g) + Cl2 (g) → SO2Cl2

Reaction of sulphur dioxide with oxygen

It is oxidised to sulphur trioxide by oxygen in the presence of vanadium(V) oxide catalyst.

SO2 (g) + O2 (g) 2SO3 (g)

Reducing properties of sulphur dioxide

Moist sulphur dioxide behaves as a reducing agent. For example, it converts iron(III) ions to iron(II) ions and decolourises acidified potassium permanganate(VII) solution; the latter reaction is a convenient test for the gas.

2Fe3+ + SO2 + 2H2O → 2Fe2+ + SO42- + 4H+

5SO2 + 2MnO4- + 2H2O → 5SO42- + 4H+ + 2Mn2+

Uses of sulphur dioxide

• in refining petroleum and sugar

• in bleaching wool and silk and

• as an anti-chlor, disinfectant and preservative

• Sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite (industrial chemicals) are manufactured from sulphur dioxide.

• Liquid SO2 is used as a solvent to dissolve a number of organic and inorganic chemicals.

Oxoacids of sulphur

Sulphur forms a number of oxoacids such as H2SO3, H2S2O3, H2S2O4, H2S2O5, H2SxO6 (x = 2 to 5), H2SO4, H2S2O7, H2SO5, H2S2O8.

Some of these acids are unstable and cannot be isolated. They are known in aqueous solution or in the form of their salts.

Sulphuric acid (H2SO4)

Sulphuric acid is one of the most important industrial chemicals.

It is manufactured by lead chamber process or contact process.

Contact process involves three steps:

(i) Burning of sulphur or sulphur ores in air to generate SO2.

(ii) Conversion of SO2 to SO3 by the reaction with oxygen in the presence of a catalyst V2O5).

SO2 (g) + O2 (g) 2SO3 (g), ΔrHo = -196.6 kJmol-1

(iii) Absorption of SO3 in H2SO4 to give oleum (H2S3O7) which upon hydrolysis gives H2SO4.

Physical properties of sulphuric acid

Sulphuric acid is a colourless, dense, oily liquid with a specific gravity of 1.84 at 298 K. The acid freezes at 283 K and boils at 611 K.

It dissolves in water with the evolution of a large quantity of heat. Hence, care must be taken while preparing sulphuric acid solution from concentrated sulphuric acid. The concentrated acid must be added slowly into water with constant stirring.

Chemical properties of sulphuric acid

The chemical reactions of sulphuric acid are as a result of the following characteristics:

1. low volatility

2. strong acidic character

3. strong affinity for water and

4. ability to act as an oxidising agent.
• In aqueous solution, sulphuric acid ionises in two steps.

H2SO4 (aq) + H2O (l) → H3O+ (aq) + HSO4 (aq); Ka1 = very large (>10)

HSO4 (aq) + H2O (l) → H3O+ (aq) + SO42- (aq); Ka2 = 1.2 × 10–2

The larger value of Ka1 means that H2SO4 is largely dissociated into H+ and HSO4. Greater the value of dissociation constant (Ka), the stronger is the acid.

• Because of its low volatility, sulphuric acid can be used to manufacture more volatile acids from their corresponding salts.

• The acid forms two series of salts: normal sulphates (such as sodium sulphate and copper sulphate) and acid sulphates (e.g., sodium hydrogen sulphate).

MX + H2SO4 → HX + MHSO4

2MX + H2SO4 → 2HX + M2SO4

• Concentrated H2SO4 is a strong dehydrating agent.

C12H22O11 12C + 11H2O

The burning sensation of concentrated H2SO4 on skin is due to the dehydrating reaction.

• Hot concentrated sulphuric acid is a moderately strong oxidising agent. It is intermediate between phosphoric acid and nitric acid.

Cu + 2H2SO4(conc.) → CuSO4 + SO2 + 2H2O

S + 2H2SO4(conc.) → 3SO2 + 2H2O

C + 2H2SO4(conc.) → CO2 + 2SO2 + 2H2O

Uses of sulphuric acid

The bulk of sulphuric acid produced is used in the manufacture of fertilizers (e.g., ammonium sulphate, superphosphate).

Other uses,

1. petroleum refining

2. manufacture of pigments, paints and dyestuff intermediates

3. detergent industry

4. metallurgical applications (e.g.,cleansing metals before enameling, electroplating and galvanizing)

5. storage batteries

6. in the manufacture of nitrocellulose products

7. as a laboratory reagent.

GROUP 17 ELEMENTS

The 17 group of Periodic Table contains five elements fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) collectively known as halogens (salt forming elements). Astatine is artificially prepared radioactive element.

Abundance

Being very reactive in nature, they are not found free in nature. Their presence in earth’s crust follows the order.

F2 > Cl2 > Br2 > I2

Fluorine is present mainly as insoluble fluorides (fluorspar CaF2, cryolite Na3AlF6 and fluoroapatite 3Ca3(PO4)2.CaF2) and small quantities are present in soil, river water plants and bones and teeth of animals. Sea water contains chlorides, bromides and iodides of sodium, potassium, magnesium and calcium, but is mainly sodium chloride solution (2.5% by mass). The deposits of dried up seas contain these compounds, e.g., sodium chloride and carnallite, KCl.MgCl2.6H2O. Certain forms of marine life contain iodine in their systems; various seaweeds, for example, contain upto 0.5% of iodine and Chile saltpetre contains upto 0.2% of sodium iodate.

GENERAL PHYSICAL PROPERTIES OF GROUP 17 ELEMENTS

 Property F Cl Br I At Atomic number 9 17 35 53 85 Atomic mass/g mol–1 19.00 35.45 79.90 126.90 210 Electronic configuration [He] 2s22p5 [Ne] 3s23p5 [Ar] 3d104s24p5 [Kr] 4d105s25p5 [Xe] 4f145d106s26p5 Covalent radius/pm 64 99 114 133 – Ionic radius X–/pm 133 184 196 220 – Ionisation enthalpy /kJ mol–1 1680 1256 1142 1008 – Electron gain enthalpy/kJ mol–1 –333 –349 –325 –296 – Electronegativityb 4 3.2 3.0 2.7 2.2 ΔHydH(X–)/kJ mol–1 515 381 347 305 – F2 Cl2 Br2 I2 – Melting point/K 54.4 172.0 265.8 386.6 – Boiling point/K 84.9 239.0 332.5 458.2 – Density/gcm–3 1.5(85K) 1.66(203K) 3.19(273)K 4.94(293K) – Distance X – X/pm 143 199 228 266 – Bond dissociation enthalpy /(kJ mol–1) 158.8 242.6 192.8 151.1 – Eo/V 2.87 1.36 1.09 0.54 – Physical State Gas Gas Liquid Solid Atomicity Diatomic Diatomic Diatomic Diatomic Flame Colour Pale yellow Yellowish green Reddish brown Deep Violet

Bond energy and bond length

The bond length increases from fluorine to iodine and in the same order bond energy decreases However, the bond dissociation energy of F2 is lesser due to its smaller size and larger electron-electron repulsion. The order of bond dissociation energy is

Cl2 > Br2 > F2 > I2

The bond length follows the order,

F-F < Cl-Cl < Br-Br < I-I

Electron gain enthalpy of group 17 elements

Due to small size of fluorine (hence, high electron density), the extra electron to be added feels more electron-electron repulsion. Therefore fluorine has less electron gain enthalpy than chlorine.

Reduction potentials and oxidising nature of group 17 elements

E°/V of halogens are positive and decrease from F to I. Therefore, halogens act as strong oxidising agents and their oxidising power decreases from fluorine to iodine. Fluorine is the strongest oxidising agent and is most reactive. That’s why it is prepared by the electrolysis of a mixture of KHF2 and anhydrous HF using Monel(nickel-copper) metal as a catalyst.

Solubility of group 17 elements

Halogens are soluble in water which follows the order F2 > Cl2 > Br2 > I2.

The process involves 3 steps.

Therefore,

The solubility of iodine in water is enhanced in the presence of KI.

KI + I2 ↔ KI3 ↔ K+ + I3-

• I2 forms blue colour complex with starch.

Anomalous behaviour of fluorine

Like other elements of p-block present in second period of the periodic table, fluorine shows anomalous behaviour. For example, ionization enthalpy, electronegativity, and electrode potentials are all higher for fluorine than expected from the trends set by other halogens. Also, ionic and covalent radii, m.p. and b.p., enthalpy of bond dissociation and electron gain enthalpy are much lower than expected.

The anomalous behaviour of fluorine is due to,

• small size,

• highest electronegativity,

• low F-F bond dissociation enthalpy,

• non availability of d orbitals in valence shell.

Most of the reactions of fluorine are exothermic (due to the small and strong bond formed by it with other elements). It forms only one oxoacid while other halogens form a number of oxoacids.

Chemical properties of group 17 elements

Oxidation states of group 17 elements

All the halogens exhibit –1 oxidation state. However, chlorine, bromine and iodine also exhibit +1, +3, +5 and +7 oxidation states as explained below.

• The higher oxidation states of chlorine, bromine and iodine are realized mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms. e.g., in interhalogens, oxides and oxoacids.

• The oxidation states of +4 and +6 occur in the oxides and oxoacids of chlorine and bromine

• The fluorine atom has no d orbitals in its valence shell and therefore cannot expand its octet and being the most electronegative, exhibits only -1 oxidation state.

Reactivity of group 17 elements

All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

The ready acceptance of an electron is the reason for the strong oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the solid phase. In general, a halogen oxidises halide ions of higher atomic number.

F2 + 2X → 2F + X2 (X = Cl, Br or I)

Cl2 + 2X → 2Cl + X2 (X = Br or I)

Br2 + 2I→ 2Br + I2

Reaction of halogens with water

Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids. The reaction of iodine with water is non-spontaneous and, I can be oxidised by oxygen in acidic medium; just the reverse of the reaction observed with fluorine.

2F2(g) + 2H2O(l) → 4H+(aq) + 4F- + (aq) + O2 (g)

X2(g) + H2O(l) → HX(aq) + HOX(aq) where X = Cl or Br

4I- (aq) + 4H+(aq) + O2 (g) → 2I2 s 2H2O

Reaction of halogens with hydrogen

HF is a low boiling liquid (293K) due to intermolecular hydrogen bonding, while HCI, HBr, HI are gases. The boiling point follows the trend

HF > HI > HBr > HCl

Other properties show the following trend,

 HI > HBr > HCl > HF acid strength, reducing character, bond length HI < HBr < HCl < HF thermal stability, dipole moment, bond strength, stability

Reactivity of halogens towards oxygen

Halogens form many oxides with oxygen but most of them are unstable. Fluorine forms two oxides OF2 and O2F2. However, only OF2 is thermally stable at 298 K. At 298K O2F2 oxidises plutonium to PuF6 and the reaction is used for removing plutonium as PuF6 from spent nuclear fuel.

Chlorine forms a number of oxides such as, Cl2O, Cl2O3, Cl2O5, Cl2O7, ClO2.

ClO2 is used as a bleaching agent for paper pulp, textiles and in water treatment.

Br2O, BrO2, BrO3 are the least stable bromine oxides and exist only at low temperatures. This anomalous behavior of bromine is due to middle row anomaly. They are very powerful oxidising agents.

The middle row anomaly is due to the transition metal contraction.

Bromine has the electron configuration [Ar]4s23d104p5. The 3d orbital has no radial nodes and hence quite close to the nucleus. Due to this there is relatively little repulsion between the 3d electrons and the 4p electrons. Therefore the ionization energies are higher than expected from the periodic trends.

The iodine oxides, i.e., I2O4, I2O5, I2O7 are insoluble solids and decompose on heating. I2O5 is a very good oxidising agent and is used in the estimation of carbon monoxide.

Reaction of halogens with metals

Halogens react with metals to form metal halides. For example, bromine reacts with magnesium to give magnesium bromide.

Mg (s) + Br2 (l) → MgBr2 (s)

The ionic character of the halides decreases in the order

MF > MCl > MBr > MI

where M is a monovalent metal.

If a metal exhibits more than one oxidation state, the halides in higher oxidation state will be more covalent than the one in lower oxidation state. For example, SnCl4, PbCl4, SbCl5 and UF6 are more covalent than SnCl2, PbCl2, SbCl3 and UF4 respectively.

Reactivity of halogens towards other halogens

Halogens combine amongst themselves to form a number of compounds known as inter-halogens of the types XX′ , XX’3, XX’5 ′ and XX’7 where X is a larger size halogen and X′ is smaller size halogen.

Some inter halogen compounds

 Type Formula Physical state and colour Structure XX′1 ClF colourless gas – BrF pale brown gas – IFa detected spectroscopically – BrClb gas ICl ruby red solid (α-form) – Brown red solid (β-form) – IBr black solid – XX′3 ClF3 colourless gas Bent T-shaped BrF3 yellow green liquid Bent T-shaped IF3 yellow powder Bent T-shaped ICl3c orange solid Bent T-shaped XX′5 IF5 colourless gas but solid below 77 K Square pyramidal BrF5 colourless liquid Square pyramidal ClF5 colourless liquid Square pyramidal XX′7 IF7 colourless gas Pentagonal bipyramidal

aVery unstable;

bThe pure solid is known at room temperature;

cDimerises as Cl–bridged dimer (I2Cl6)

Reaction of halogens with alkali

Fluorine reacts as follows,

2F2 + → 2NaF + OF2 + H2O

2F2 + $4\underset{\begin{array}{c}\mathrm{h}\mathrm{o}\mathrm{t}\\ \mathrm{c}\mathrm{o}\mathrm{n}\mathrm{c}\end{array}}{\underset{⏟}{\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}}}$ → 4NaF + O2 + 2H2O

Other halogens form hypo-halite with dilute NaOH and halate with conc. NaOH.

X2 + → NaX + NaOX + H2O

2X2 + $6\underset{\begin{array}{c}\mathrm{h}\mathrm{o}\mathrm{t}\\ \mathrm{c}\mathrm{o}\mathrm{n}\mathrm{c}\end{array}}{\underset{⏟}{\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}}}$ → 5NaX + NaXO3 + 3H2O

CHLORINE AND ITS COMPOUNDS

Occurrence

Common salt, NaCl is the major source of chlorine. It is also present in sea water and as rock salt.

Preparation of Chlorine

1. By heating manganese dioxide with concentrated hydrochloric acid.

MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O

However, a mixture of common salt and concentrated H2SO4 is used in place of HCl (Weldon’s Process).

4NaCl + MnO2 + 4H2SO4 → MnCl2 + 4NaHSO4 + 2H2O + Cl2

1. By the action of HCl on potassium permanganate.

2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2

Manufacture of chlorine

1. Deacon’s process: By oxidation of hydrogen chloride gas by atmospheric oxygen in the presence of CuCl2 (catalyst) at 723 K.

4HCl + O2 2Cl2 + 2H2O

2. Electrolytic process: Chlorine is obtained by the electrolysis of brine (concentrated NaCl solution). Chlorine is liberated at anode. It is also obtained as a by–product in many chemical industries.

2NaCl + 2H2O ↔ 2NaOH + Cl2 + H2

Physical properties of chlorine

It is yellowish green gas, collected by upward displacement of (it is heavier than) air, poisonous in nature, soluble in water. It’s aqueous solution is known as chlorine water. Boiling point is 239K.

Chemical properties of chlorine

• Chlorine reacts with a number of metals and non-metals to form chlorides.

2Al + 3Cl2 → 2AlCl3

P4 + 6Cl2 → 4PCl3

2Na + Cl2 → 2NaCl

S8 + 4Cl2 → 4S2Cl2

2Fe + 3Cl2 → 2FeCl3

• It has great affinity for hydrogen,

H2 + Cl2 → 2HCl

H2S + Cl2 → 2HCl + S

C10H16 + 8Cl2 → 16HCl + 10C

• Action of water

Cl2 + H2O → HOCl + HCl

HOCL → HCl + [O] (Nascent Oxygen)

• The bleaching action of chlorine is due to oxidation and is permanent.

Cl2 + H2O → 2HCl + [O]

Coloured matter + [O] → colourless matter.

With dry slaked lime it gives bleaching powder.

2Ca(OH)2 + 2Cl2 → Ca(OCl)2 + CaCl2 + 2H2O

The composition of bleaching powder is Ca(OCl)2.CaCl2.Ca(OH)2.2H2O.

• It oxidises ferrous to ferric and sulphite to sulphate. Chlorine oxidises sulphur dioxide to sulphur trioxide and iodine to iodate. In the presence of water they form sulphuric acid and iodic acid respectively.

2FeSO4 + H2SO4 + Cl2 → Fe2(SO4)3 + 2HCl

Na2SO3 + Cl2 + H2O → Na2SO4 + 2HCl

SO2 + 2H2O + Cl2 → H2SO4 + 2HCl

I2 + 6H2O + 5Cl2 → 2HIO3 + 10HCl

• With cold and dilute alkalies chlorine produces a mixture of chloride and hypochlorite but with hot and concentrated alkalies it gives chloride and chlorate.

+ Cl2 → NaCl + NaOCl + H2O

$6\underset{\begin{array}{c}\mathrm{h}\mathrm{o}\mathrm{t}\\ \mathrm{c}\mathrm{o}\mathrm{n}\mathrm{c}\end{array}}{\underset{⏟}{\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}}}$ + 3Cl2 → 5NaCl + NaClO3 + 3H2O

• Chlorine reacts with hydrocarbons and gives substitution products with saturated hydrocarbons and addition products with unsaturated hydrocarbons. For example,

CH4 + Cl2 CH3Cl + HCl

C2H4 + Cl2 C2H4Cl2

Uses of chlorine

• for bleaching wood pulp (required for the manufacture of paper and rayon), bleaching cotton and textiles,
• in the extraction of gold and platinum

• in the manufacture of dyes, drugs and organic compounds such as CCl4, CHCl3, DDT, refrigerants, etc.

• In sterilising drinking water and

• preparation of poisonous gases such as phosgene (COCl2), tear gas (CCl3NO2), mustard gas (ClCH2CH2SCH2CH2Cl).

Hydrochloric acid (HCl)

Preparation of HCl

NaCl + H2SO4 NaHSO4 + HCl

NaHSO4 + NaCl Na2SO4 + HCl

HCl gas can be dried by passing through concentrated sulphuric acid

Properties of HCl

• It is a colourless and pungent smelling gas. It is easily liquefied to a colourless liquid (b.p.189 K) and freezes to a white crystalline solid (f.p. 159 K). It is extremely soluble in water and ionises as follows,

HCl(g) + H2O ↔ H3O+(aq) + Cl(aq); Ka = 107

Its aqueous solution is called hydrochloric acid. High value of dissociation constant (Ka) indicates that it is a strong acid in water.

• It reacts with NH3 and gives white fumes of NH4Cl.

NH3 + HCl → NH4Cl

• Noble metals like gold, platinum can dissolve in aqua-regia [three part conc HCl and one part of conc HNO3].

Au + 4H+ + NO3- + 4Cl- → AuCl4- + NO + 2H2O

3Pt +16H+ + 4NO3- +18Cl- → 3PtCl62- + 4NO + 8H2O

• Hydrochloric acid decomposes salts of weaker acids, e.g., carbonates, hydrogencarbonates, sulphites, etc.

Na2CO3 + 2HCl → 2NaCl + H2O + CO2

NaHCO3 + HCl → NaCl + H2O + CO2

Na2SO3 + 2HCl → 2NaCl + H2O + SO2

Uses of HCl

It is used

1. in the manufacture of chlorine, NH4Cl and glucose (from corn starch),

2. for extracting glue from bones and purifying bone black,

3. in medicine and as a laboratory reagent.

IODINE (I2)

Major source of iodine is deep sea weeds of laminaria variety. Their ashes which are called kelp contain 0.5% iodine as iodides.

Another source of I2 is caliche or crude chile saltpetre (NaNO3) which contains 0.2%, NaIO3

Iodine is purified by sublimation.

It shows no reaction with water. Tincture of iodine is a mixture of I2 and KI dissolved in rectified spirit.

Oxoacids of halogens

Due to high electronegativity and small size and absence of d-orbitals; fluorine forms only one oxoacid, HOF known as fluoric (I) acid or hypofluorous acid.

The other halogens form several oxoacids. Most of them cannot be isolated in pure state. They are stable only in aqueous solutions or in the form of their salts.

+3 oxidation state of bromine and iodine are unstable due to inert pair effect, therefore, HBrO2 and HIO2 do not exist.

The central atom in the oxoacids is sp3 hybridized. Every oxoacid has essentially one X-OH bond. Whereas most oxoacids have X = O bonds present in them. This double bond between oxygen and halogen is dπ-pπ in nature.

In the series of oxoacids, the first member possesses high acidic strength. This is due to high electronegativity and small size of the halogen atom. The acidic strength increases with increase in the oxidation number of halogens.

 Halic (I) acid (Hypohalous) HOF (Hypofluorous) HOCl (Hypochlorous) HOBr (Hypobromous) HOI (Hypoiodous) Halic (III) acid (Halous acid) – HOCIO (chlorous acid) – – Halic (V) acid (Halic acid) – HOCIO2 (chloric acid) HOBrO2 (bromic acid) HOIO2 (iodic acid) Halic (VII) acid (Perhalic acid) – HOCIO3 (perchloric acid) HOBrO3 (perbromic acid) HOIO3 (periodic acid)

GROUP 18 ELEMENTS

The 18th group of the Periodic Table consists of colourless, odourless gases at room temperature.

General and physical properties of group 18 elements

 Propery He Ne Ar Kr Xe Rn* Atomic number 2 10 18 36 54 86 Atomic mass/g mol–1 4.00 20.18 39.95 83.80 131.30 222.00 Electronic configuration 1s2 [He] 2s22p6 [Ne] 3s23p6 [Ar] 3d104s24p6 [Kr] 4d105s25p6 [Xe] 4f 145d106s26p6 Atomic radius/pm 120 160 190 200 220 – Ionisation enthalpy/kJmol-1 2372 2080 1520 1351 1170 1037 Electron gain enthalpy/kJmol-1 48 116 96 96 77 68 Density (at STP)/gcm–3 1.8×10–4 9.0×10–4 1.8×10–3 3.7×10–3 5.9×10–3 9.7×10–3 Melting point/K – 24.6 83.8 115.9 161.3 202 Boiling point/K 4.2 27.1 87.2 119.7 165.0 211 Atmospheric content (% by volume) 5.24×10–4 – 1.82×10–3 0.934 1.14×10–4 8.7×10–6

Electronic configuration

Their valence shell electronic configuration is ns2np2 except helium which has 1s2. The properties of noble gases including their inactive nature are ascribed to their closed shell structures.

Ionization Enthalpy

Due to stable electronic configuration these gases exhibit very high ionisation enthalpy. However, it decreases down the group with increase in atomic size.

Atomic radii increase down the group with increase in atomic number.

Electron gain enthalpy

Since noble gases have stable electronic configurations, they have no tendency to accept the electron and therefore, have large positive values of electron gain enthalpy.

Density

The density of noble gases increases down the group.

Heat of vaporization

They have very low values of heat of vaporisation due to weak van der Waals’ forces of attraction and monoatomic nature. The value increases down the group.

Solubility in water

They are slightly soluble in water and solubility increases from He to Rn.

Liquefaction

It is extremely difficult to liquify inert gases due to weak van der Waals’ forces of attraction among their molecules. Hence, they posses low value of critical temperature also.

Chemical properties of group 18 elements

• The noble gases are inert in nature due to following reasons
1. The noble gases except helium (1s2) have completely filled ns2np6 electronic configuration in their valence shell.

2. They have high ionisation enthalpy and more positive electron gain enthalpy.

• In 1962, the first compound of noble gases was prepared. Neil Bartlett initially carried out a reaction between oxygen and PtF6. This resulted in the formation of a red compound, [O2]+[PtF6]-.

Later, he realized that the first ionization energies of oxygen (1175 kJ/mol) and Xe (1170 kJ/mol) are almost the same. Thus, he tried to prepare a compound with Xe and PtF6. He was successful and a red-coloured compound, was formed. It was hexafluoroplatinate.

Xe + PtF6 → Xe[PtF6]

Now, many compounds of Xe and Kr are known with fluorine and oxygen.

Preparation of compounds of xenon

Xenon-fluorine compounds

Xenon forms three binary fluorides, XeF2, XeF4 and XeF6 by the direct reaction of elements under appropriate experimental conditions.

Xe (g) + F2 (g) XeF2(s) (xenon in excess)

Xe (g) + 2F2 (g) XeF4(s) (1:5 ratio)

Xe (g) + 3F2 (g) XeF6(s) (1:20 ratio)

XeF6 can also be prepared by the interaction of XeF4 and O2F2 at 143K.

XeF4 + O2F2 XeF6 + O2

Properties of xenon-flourine compounds

XeF2, XeF4 and XeF6 are colourless crystalline solids and sublime readily at 298 K. They are powerful fluorinating agents.

They are readily hydrolysed even by traces of water. For example, XeF2 is hydrolysed to give Xe, HF and O2.

2XeF2 (s) + 2H2O(l) → 2Xe (g) + 4 HF(aq) + O2(g)

Xenon fluorides react with fluoride ion acceptors to form cationic species and fluoride ion donors to form fluoro anions.

XeF2 + PF5 → [XeF]+ [PF6] ;

XeF4 + SbF5 → [XeF3]+[SbF6]

XeF6 + MF → M+[XeF7] (M = Na, K, Rb or Cs)

Xenon-oxygen compounds

Hydrolysis of XeF4 and XeF6 with water gives XeO3.

6XeF4 + 12H2O → 4Xe + 2XeO3 + 24HF + 3O2

XeF6 + 3H2O → XeO3 + 6HF

Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2.

XeF6 + H2O → XeOF4 + 2HF

XeF6 + 2H2O → XeO2F2 + 4HF

XeO3 is a colourless explosive solid and XeOF4 is a colourless volatile liquid.

Structures of xenon compounds

The structures of the three xenon fluorides can be deduced from VSEPR theory. XeF2 and XeF4 have linear and square planar structures respectively.

XeF6 has seven electron pairs (6 bonding pairs and one lone pair) and would, thus, have a distorted octahedral structure as found experimentally in the gas phase.

XeO3 has a pyramidal molecular structure and XeOF4 has a square pyramidal molecular structure.

Uses of xenon compounds

• Helium is a non-inflammable and light gas. Hence, it is used in filling balloons for meteorological observations. It is also used in gas-cooled nuclear reactors.

• Liquid helium (b.p. 4.2 K) finds use as cryogenic agent for carrying out various experiments at low temperatures. It is used to produce and sustain powerful superconducting magnets which form an essential part of modern NMR spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical diagnosis.

• It is used as a diluent for oxygen in modern diving apparatus because of its very low solubility in blood.

• Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes.

• Neon bulbs are used in botanical gardens and in green houses.

• 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.

• Xenon and Krypton are used in light bulbs designed for special purposes.