p Block Elements

General Configuration: ns2 np1-6

Chemistry

Chapter-7

The P-Block elements

Across a period: Covalent radii and metallic character decreases, but electro negativity, electron affinity, oxidizing power and ionization energy increases. Down the group: Covalent radii and metallic character increases, but electro negativity, electron affinity, oxidizing power and ionization energy decreases.

Inert pair effect: While going down the group, the ns2 electrons become more and more reluctant to participate in bond formation. This is because down the group bond energy decreases and so the energy required to un-pair ns2 electrons is not compensated by the energy released in forming two additional bonds.

Group 15 Elements:

N 2s2 2p3 Non-Metal P 3s2 3p3 Non-Metal As 4s2 4p3 Metalloids Sb 5s2 5p3 Metalloids Bi 6s2 6p3 Metal

Trends of some of the atomic and physical Properties:

Electronic Configuration:

The valence shell electronic configuration of these elements is ns2 np3.The s orbital in these elements is

completely filled and p orbitals are half-filled, making their electronic configuration extra stable.

Occurrence:

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 (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.

Atomic and Ionic Radii:

Covalent and ionic radii increase in size down the group.

Ionisation Enthalpy:

Ionisation enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half-filled p orbitals electronic configuration and smaller size, the ionisation enthalpy of the group 15 elements is much greater than that of group 14 elements in the corresponding periods.

Electronegativity:

The electronegativity value, in general, decreases down the group with increasing atomic size.

Physical Properties:

All the elements of this group are polyatomic.

Dinitrogen is a diatomic gas while all others are solids.

Metallic character increases down the group. Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids and bismuth is a metal. This is due to decrease in ionisation enthalpy and increase in atomic size.

The boiling points, in general, increase from top to bottom in the group but the melting point increases up to arsenic and then decreases up to bismuth.

Except nitrogen, all the elements show allotropy.

Chemical Properties:

Oxidation states:

N -3, +3 P -3, +3, +5 As -3, +3, +5 Sb +3, +5 Bi +3, +5

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

The tendency to exhibit –3 oxidation state decreases down the group due to increase in size and metallic character.

The stability of +5 oxidation state decreases and that of +3 state increases (due to inert pair effect) down the group.

Nitrogen exhibits + 1, + 2, + 4 oxidation states.

Nitrogen is restricted to a maximum covalency of 4 since only four (one s and three p) orbitals are available for bonding. 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 PF– 6.

Nitrogen exhibits +1, +2 and +4 O.S also when it reacts with oxygen. All these O.S tend to disproportionate in acid solution. e.g. 3HNO2 → HNO3 + H2O + 2NO

Anomalous properties of nitrogen

Nitrogen differs from the rest of the members of this group due to its smaller size, high electro negativity, high ionization enthalpy and non-availability of d-orbitals.

Nitrogen can form pπ-pπ multiple bondswith itself and with other elements having small size and high electronegativity (e.g., C, O).

Nitrogen exists as diatomic molecule with a triple bond.

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

P, As and Sb form P-P, As-As and Sb-Sb single bonds whereas Bi forms metallic bonds. However, N-N

single bond is weaker than P-P single bond, because of high inter electronic repulsion of non-bonding electrons owing to small bond length.

Catenation tendency is weaker in N as N-N bond is much weaker than P-P, As-As and Sb-Sb due to inter electronic repulsions because of small bond length.

Except nitrogen, the heavier elements can form dπ-pπ bonds, e.g. R3P=O or R3P=CH2 and also when transition elements like P(C2H5)3 and As(C6H5)3 act as ligands; they form dπ-dπ bonds.

Reactivity towards hydrogen:

Form hydrides of formula EH3

Structure pyramidal

Bond angle decreases down the group due to decrease in electro negativity.

Stability decreases due to increase in size.

Reducing character increases due to decrease in stability and decrease in bond dissociation enthalpy. NH3 is a mild reducing agent while BiH3 is strongest.

Basic character decreases in the order: NH3> PH3> AsH3> SbH3 > BiH3.

Reactivity towards oxygen:

Form two types of oxides of the formula: E2O3 and E2O5.

The oxide in the higher oxidation is more acidic than that of lower oxidation state.

Acidic character decreases down the group.

(i) N2O3, P2O3 Acidic

(ii) As2O3, Sb2O3 Amphoteric

(iii) Bi2O3 Basic oxide

Reactivity towards halogens:

They form halides of the formula: EX3 and EX5.

Nitrogen does not form pentahalide due to non-availability of the d orbitals in its valence shell.

Pentahalides are more covalent than trihalides.

All trihalides except those of nitrogen are stable. Only NF3 is stable. Trihalides except BiF3 are predominantly covalent.

Reactivity 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).

Dinitrogen

Preparation

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).

In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite. NH4CI(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl (aq)

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

(NH4)2Cr2O7⎯⎯⎯→ N2 + 4H2O + Cr2O3

Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide. Ba(N3)2→ Ba + 3N2 2NaN3→ 2Na + 3N2

Properties

Dinitrogen is a colourless, odourless, tasteless and non-toxic gas.

Dinitrogen is rather inert at room temperature because of the high bond enthalpy of N≡N bond. Reactivity, however, increases rapidly with rise in temperature.

At higher temperatures, it directly combines with some metals to form predominantly ionic nitrides and with non-metals, covalent nitrides. A few typical reactions are:

6Li + N2 ⎯⎯⎯→ Heat 2Li3N

3Mg + N2 ⎯⎯⎯→ Heat Mg3N2

It combines with hydrogen at about 773 K in the presence of a catalyst (Haber’s Process) to form ammonia: 773K N2 (g) + 3H2 (g) ----------2NH3 (g); ∆f H y = –46.1 kJmol–1

Dinitrogen combines with dioxygen only at very high temperature (at about 2000 K) to form nitric oxide, NO. N2 + O2(g) -----2NO(g)

Uses:

The main use of dinitrogen is in the manufacture of ammonia and other industrial chemicals containing nitrogen, (e.g., calcium cyanamide).

It also finds use where an inert atmosphere is required (e.g., in iron and steel industry, inert diluent for reactive chemicals).

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

Ammonia

On a large scale, ammonia is manufactured by Haber’s process.

N2 (g) + 3H2 (g) -----------2NH3 (g); ∆f H0 = – 46.1 kJ mol–1

In accordance with Le Chatelier’s principle, high pressure would favour the formation of ammonia. The optimum conditions for the production of ammonia are a pressure of 200 × 105 Pa (about 200 atm), a temperature of

~ 700 K and the use of a catalyst such as iron oxide with small amounts of K2O and Al2O3 to increase the rate of attainment of equilibrium.

Properties

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.

Structure:

The presence of a lone pair of electrons on the nitrogen atom of the ammonia molecule makes it a Lewis base. It donates the electron pair and forms linkage with metal ions and the formation of such complex compounds finds applications in detection of metal ions such as Cu2+, Ag+:

Cu2+ (aq) + 4 NH3(aq) ---[Cu(NH3 )4]2+(aq)

(blue) (deep blue)

Ag+(aq) + Cl-(aq) -----AgCl (s)

(colourless) (white ppt)

AgCl (s) + 2NH3(aq) ----- [Ag (NH3)2] Cl (aq)

(white ppt)(colourless)

Uses:

Ammonia is used to produce various nitrogenous fertilisers (ammonium nitrate, urea, ammonium phosphate and ammonium sulphate) and in the manufacture of somemost important one being nitric acid. Liquid ammonia is also used as a refrigerant.

Nitric Acid

Preparation

In the laboratory, nitric acid is prepared by heating KNOretort.

NaNO3+H2SO4----- NaHSO4 + HNO

Ostwald’s process.

On a large scale it is prepared mainly by Ostwald’s process.

This method is based upon catalytic oxidation of NH

Nitric oxide thus formed combines with oxygen

Nitrogen dioxide so formed, dissolves in water to give HNO

Catalytic conditions: Pt /Rh gauge

NO thus formed is recycled and the aqueous HNOmass. Further concentration to 98% can be achieved by dehydration with concentrated H

Properties

It is a colourless liquid (f.p. 231.4 K and b.p. 355.6 K).

Laboratory grade nitric acid contains ~ 68% of the HNO3 by mass and has a specific gravity of 1.504.

In the gaseous state, HNO3 exists as a planar molecule:

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

HNO3(aq) + H2O(l) → H3O+ (aq) + NO

Concentrated nitric acid is a strong oxidising agent 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.

Ammonia is used to produce various nitrogenous fertilisers (ammonium nitrate, urea, ammonium phosphate and ammonium sulphate) and in the manufacture of some inorganic nitrogen compounds, the most important one being nitric acid. Liquid ammonia is also used as a refrigerant.

In the laboratory, nitric acid is prepared by heating KNO3 or NaNO3 and concentrated H

+ HNO3

On a large scale it is prepared mainly by Ostwald’s process.

This method is based upon catalytic oxidation of NH3 by atmospheric oxygen.

Nitric oxide thus formed combines with oxygen giving NO2.

Nitrogen dioxide so formed, dissolves in water to give HNO3.

Pt /Rh gauge catalyst, Temperature: 500K, Pressure: 9 bar

NO thus formed is recycled and the aqueous HNO3 can be concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated H

It is a colourless liquid (f.p. 231.4 K and b.p. 355.6 K).

Laboratory grade nitric acid contains ~ 68% of the HNO3 by mass and has a specific gravity of

In the gaseous state, HNO3 exists as a planar molecule:

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

(aq) + NO3 – (aq)

Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals

The products of oxidation depend upon the concentration of the acid, temperature and the nature of the material undergoing oxidation.

Ammonia is used to produce various nitrogenous fertilisers (ammonium nitrate, urea, ammonium inorganic nitrogen compounds, the

most important one being nitric acid. Liquid ammonia is also used as a refrigerant.

and concentrated H2SO4 in a glass

9 bar

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

Laboratory grade nitric acid contains ~ 68% of the HNO3 by mass and has a specific gravity of

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

and attacks most metals except noble metals

The products of oxidation depend upon the concentration of the acid, temperature and the nature

3Cu + 8 HNO3(dilute) → 3Cu(NO3)2 + 2NO + 4H2O Cu + 4HNO3(conc.) → Cu(NO3)2 + 2NO2 + 2H2O

Zinc reacts with dilute nitric acid to give N2O and with concentrated acid to give NO2.

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.

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 → 2HIO3 + 10 NO2 + 4H2O

C + 4HNO3 → CO2 + 2H2O + 4NO2

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

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

Brown Ring Test:

The test is usually carried out by adding dilute ferrous sulphate solution to an aqueous solution containing nitrate ion, and then carefully adding concentrated sulphuric acid along the sides of the test tube. A brown ring at the interface between the solution and sulphuric acid layers indicate the presence of nitrate ion in solution.

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

[Fe (H2O)6]2++ NO → [Fe(H2O)5 (NO)]2+ + H2O

(brown)

Uses:

The major use of nitric acid is in the manufacture of ammonium nitrate for fertilisers and other nitrates for use in explosives and pyrotechnics.

It is also used for the preparation of nitro-glycerine, trinitrotoluene and other organic nitro compounds. Other major uses are in the pickling of stainless steel, etching of metals and as an oxidiser in rocket fuels.

Group 16 Elements:

O 2s2 2p4 Non-Metal S 3s2 3p4 Non-Metal Se 4s2 4p4 Metalloids Te 5s2 5p4 Metalloids Po 6s2 6p4 Metal

Occurrence

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.

The abundance of sulphur in the earth’s crust is only 0.03-0.1%.

Combined sulphur exists primarily 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.

Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores.

Polonium occurs in nature as a decay product of thorium and uranium minerals

Atomic and Ionic Radii

Due to increase in the number of shells, atomic and ionic radii increase from top to bottom in the group. The size of oxygen atom is, however, exceptionally small.

Ionisation Enthalpy

Ionisation enthalpy decreases down the group due to increase in size.

The elements of this group have lower ionisation enthalpy values compared to those of Group15 in the corresponding periods. Because Group 15 elements have extra stable half-filled p orbitals electronic configurations.

Electron Gain Enthalpy

Because of the smaller size of oxygen atom, it has less negative electron gain enthalpy than sulphur. And from sulphur onwards the value again becomes less negative up to polonium.

Electronegativity

Next to fluorine, oxygen has the highest electronegativity value amongst the elements. Within the group, electronegativity decreases with an increase in atomic number.

Physical Properties

All the elements exhibit allotropy. The melting and boiling points increase with an increase in atomic number down the group.

The large difference between the melting and boiling points of oxygen and sulphur may be explained on the basis of their atomicity; oxygen exists as diatomic molecule (O2) whereas sulphur exists as polyatomic molecule (S8).

Chemical Properties

Oxidation states

Anomalous behaviour of oxygen

Anomalous behaviour of oxygen: The anomalous behaviour of oxygen, due to its small size and high electronegativity and the absence of d orbitals.

One typical example of effects of small size and high electronegativity is the presence of strong hydrogen bonding in H2O which is not found in H2S, due to which H2O is a liquid but H2S is a gas.

Trends in chemical reactivity

(i) Reactivity with hydrogen:

All the elements of Group 16 form hydrides of the type H2E.

Acidic character 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.

The thermal stability of hydrides also decreases from H2O to H2Po.

All the hydrides except water possess reducing property and this character increases from H2S to H2Te.

(ii) Reactivity with oxygen:

All these elements form oxides of the EO2 and EO3 types.

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.

Besides EO2 type, sulphur, selenium and tellurium also form EO3 type oxides (SO3, SeO3, TeO3). Both types of oxides are acidic in nature.

(iii) Reactivity towards the 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 decreases in the order F–> Cl–> Br–> I–. Amongst hexahalides, hexafluorides are the only stable halides.

All hexafluorides are gaseous in nature.

They have octahedral structure. Sulphur hexafluoride, SF6 is exceptionally stable for steric reasons.

Dioxygen

Preparation

Dioxygen can be obtained in the laboratory by the following ways: (i) By heating oxygen containing salts such as chlorates, nitrates and permanganates.

2KClO3 --- 2KCl + 3O2

(ii) 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) 2Pb3O4(s) → 6PbO(s) + O2(g) 2HgO(s) → 2Hg(l) + O2(g) 2PbO2(s) → 2PbO(s) + O2(g)

(iii) Hydrogen peroxide is readily decomposed into water and dioxygen by catalysts such as finely divided metals and manganese dioxide. 2H2O2(aq) → 2H2O(l) + O2(g)

(iv) 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.

Properties

Dioxygen is a colourless and odourless gas.

It liquefies at 90 K and freezes at 55 K.

It has three stable isotopes: 16O, 17O and 18O. Molecular oxygen, O2 is unique in being paramagnetic in spite of having even number of electrons.

Dioxygen directly reacts with nearly all metals and non-metals except some metals (e.g., Au, Pt) and some noble gases.

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

2Ca+ O2 -- 2CaO

4Al + 3O2-- 2Al2O3

Uses:

In addition to its importance in normal respiration and combustion processes, 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 mountaineering.

The combustion of fuels, e.g., hydrazines in liquid oxygen, provides tremendous thrust in rockets.

Simple Oxides

A binary compound of oxygen with another element is called oxide.

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, a base.

CaO + H2O ---Ca(OH)2

In general, metallic oxides are basic. Some metallic oxides exhibit a dual behaviour. They show characteristics of both acidic as well as basic oxides. Such oxides are known as amphoteric oxides. They react with acids as well as alkalies.For example, Al2O3 reacts 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)

There are some oxides which are neither acidic nor basic. Such oxides are known as neutral oxides. Examples of neutral oxides are CO, NO and N2O.

Ozone

Ozone is an allotropic form of oxygen. It is too reactive to remain for long in the atmosphere at sea level. At a height of about 20 kilometres, it is formed from atmospheric oxygen in the presence of sunlight.

This ozone layer protects the earth’s surface from an excessive concentration of ultraviolet (UV) radiations.

Preparation

When a slow dry stream of oxygen is passed through a silent electrical discharge, conversion of oxygen to ozone (10%) occurs.

The product is known as ozonised oxygen.

3O2 → 2O3 ΔHV (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

Properties

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 about 100 parts per million, breathing becomes uncomfortable resulting in headache and nausea.

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).

Due to the ease with which it liberates atoms of nascent oxygen

O3 → O2 + O

it acts as a powerful oxidising agent. For e.g.,

it oxidises lead sulphide to lead sulphate and iodide ions to iodine.

PbS(s) + 4O3(g) → PbSO4(s) + 4O2(g)

2I– (aq) + H2O(l) + O3(g) → 2OH– (aq) + I2(s) + O2(g)

When ozone reacts with an excess of potassium iodide solution buffered with a borate buffer (pH 9.2), iodine is liberated which can be titrated against a standard solution of sodium thiosulphate. This is a quantitative method for estimating O3 gas.

The two oxygen-oxygen bond lengths in the ozone molecule are identical (128 pm) and the molecule is angular as expected with a bond angle of about 117o. It is a resonance hybrid of two main forms:

Uses:

It is used as a germicide, disinfectant and for sterilising water.

It is also used for bleaching oils, ivory, flour, starch, etc. It acts as an oxidising agent in the manufacture of potassium permanganate.

Sulphur — Allotropic Forms

Sulphur forms numerous allotropes, but the two most important allotropes of sulphur are:

Rhombic sulphur (α-sulphur)

Monoclinic sulphur (β-sulphur)

The most interesting feature is their thermal stability, the allotropes of sulphur are inter-convertible i.e. rhombic sulphur when heated above 369K gives monoclinic sulphur.

Rhombic sulphur (α-sulphur)

Rhombic sulphur is crystalline in nature and has octahedral shape. On heating the solution of roll sulphur in CS2 we get rhombic sulphur. It is yellow with a melting point of 385.8K and specific gravity 2.06. Rhombic sulphur cannot be dissolved in water but can be dissolved in benzene, ether, alcohol etc.

Monoclinic sulphur (β-sulphur)

When we take a dish and melt rhombic sulphur in that dish, we obtain monoclinic sulphur after cooling it. In this process we make two holes in the crust and pour out the remaining liquid. After this we get colourless needle-shaped crystals of β-sulphur when the crust is removed.

369K is called transition temperature because both the allotropes of sulphur are stable at this temperature. In other words, we can conclude that α sulphur is stable below 369K and it becomes β-sulphur above that temperature.

Rhombic and monoclinic sulphur, both have S8 molecules. The alternative packing of S8 molecules gives different crystal structures.

Sulphur Dioxide

Preparation

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)

The gas after drying is liquefied under pressure and stored in steel cylinders.

Properties

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

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

Uses:

Sulphur dioxide is used

(i) in refining petroleum and sugar (ii) in bleaching wool and silk and

(iii) as an anti-chlor, disinfectant and preservative. (iv) Sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite (industrial

chemicals) are manufactured from sulphur dioxide. (v) 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, H2S2O

Some of these acids are unstable and cannot be isolated

Sulphuric Acid

Manufacture

Sulphuric acid is one of the most important

Contact Process

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.

a number of oxoacids such as

O5, H2SxO6 (x = 2 to 5), H2SO4, H2S2O7, H2SO

Some of these acids are unstable and cannot be isolated.

Sulphuric acid is one of the most important industrial chemicals worldwide.

Sulphuric acid, sodium hydrogen sulphite and calcium hydrogen sulphite (industrial

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

SO5, H2S2O8.

Physical Properties

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

The chemical reactions of sulphuric acid are as a result of the following characteristics: (a) low volatility (b) strong acidic character (c) strong affinity for water and (d) ability to act as an oxidising agent.

In aqueous solution, sulphuric acid ionises in two steps.

The larger value of Ka1 (Ka1>10) means that H2SO4 is largely dissociated into H+ and HSO4–. Greater

the value of dissociation constant (Ka), the stronger is the acid. The acid forms two series of salts: normal sulphates (such as sodium sulphate and copper sulphate) and acid sulphates (e.g., sodium hydrogen sulphate).

Sulphuric acid, because of its low volatility can be used to manufacture more volatile acids from their corresponding salts.

Concentrated sulphuric acid is a strong dehydrating agent. Many wet gases can be dried by passing them through sulphuric acid, provided the gases do not react with the acid. Sulphuric acid removes water from organic compounds; it is evident by its charring action on carbohydrates.

Hot concentrated sulphuric acid is a moderately strong oxidising agent. In this respect, it is intermediate between phosphoric and nitric acids. Both metals and non-metals are oxidised by concentrated sulphuric acid, which is reduced to SO2.

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

3S + 2H2SO4(conc.) → 3SO2 + 2H2O

C + 2H2SO4(conc.) → CO2 + 2 SO2 + 2 H2O

Uses:

Sulphuric acid is a very important industrial chemical. A nation’s industrial strength can be judged by the quantity of sulphuric acid it produces and consumes.

It is needed for the manufacture of hundreds of other compounds and also in many industrial processes.

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

Other uses are in:

(a) Petroleum refining

(b) Manufacture of pigments, paints and dyestuff intermediates

(c) Detergent industry

(d) Metallurgical applications (e.g., cleansing metals before enamelling, electroplating and galvanising

(e) Storage batteries

(f) In the manufacture of nitrocellulose products and

(g) As a laboratory reagent.

Group 17 Elements

F 2s2 2p5 Non-Metal Cl 3s2 3p5 Non-Metal Br 4s2 4p5 Metalloids I 5s2 5p5 Metalloids At 6s2 6p5 Metal

These are collectively known as the halogens (Greek halo means salt and genes means born i.e., salt producers). The halogens are highly reactive non-metallic elements.

Occurrence

Fluorine and chlorine are fairly abundant while bromine and iodine less so. Fluorine is present mainly as insoluble fluorides (fluorspar CaF2, cryolite Na3AlF6 and fluorapatite 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 lifecontain iodine in their systems; various seaweeds, for example, contain up to 0.5% of iodine and Chile saltpetre contains up to 0.2% of sodium iodate.

Atomic and Ionic Radii

The halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge. The atomic radius of fluorine like the other elements of second period is extremely small. Atomic and ionic radii increase from fluorine to iodine due to increasing number of quantum shells.

Ionisation Enthalpy

They have little tendency to lose electron. Thus, they have very high ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy decreases down the group.

Electron Gain Enthalpy

Halogens have maximum negative electron gain enthalpy in the corresponding periods. This is due to the fact that the atoms of these elements have only one electron less than stable noble gas configurations.

Electron gain enthalpy of the elements of the group becomes less negative down the group. However, the negative electron gain enthalpy of fluorine is less than that of chlorine. It is due to small size of fluorine atom. As a result, there are strong interelectronic repulsions in the relatively small 2p orbitals of fluorine and thus, the incoming electron does not experience much attraction.

Electronegativity

They have very high electronegativity. The electronegativity decreases down the group. Fluorine is the most electronegative element in the periodic table.

Physical Properties

Physical state:Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.

Melting and boiling point: Their melting and boiling points steadily increase with atomic number.

Colours:All halogens are coloured. This is due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level. By absorbing different quanta of radiation, they display different colours. For example, F2, has yellow, Cl2, greenish yellow, Br2, red and I2, violet colour.

Solubility:Fluorine and chlorine react with water. Bromine and iodine are only sparingly soluble in water but are soluble in various organic solvents such as chloroform, carbon tetrachloride, carbon disulphide and hydrocarbons to give coloured solutions.

Bond dissociation enthalpies

The smaller enthalpy of dissociation of F2 compared to that of Cl2 whereas X-X bond dissociation enthalpies from chlorine onwards show the expectedtrend: Cl – Cl > Br – Br > I – I. A reason for this anomaly is the relatively large electron-electron repulsion among the lone pairs in F2 molecule where they are much closer to each other than in case of Cl2.

Chemical Properties

Oxidation states

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

The higher oxidation states of chlorine, bromine and iodine are realised 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. Being the most electronegative, it exhibits only –1 oxidation state.

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

Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids.

X2(g) + H2O (l) ---- HOX(aq) + HX (aq)

Anomalous behaviour of fluorine

Fluorine is anomalous in many properties. For example, ionisation enthalpy, electronegativity, enthalpy of bond dissociation and electrode potentials are all higher for fluorine than expected from the trends set by other halogens.

Ionic and covalent radii, m.p. and b.p. and electron gain enthalpy are quite lower than expected.

The anomalous behaviour of fluorine is due to its small size, highest electronegativity, low F-F bond dissociation enthalpy, and 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.

Hydrogen fluoride is a liquid (b.p. 293 K) due to strong hydrogen bonding. Other hydrogen halides are gases.

(i) Reactivity towards hydrogen: They all react with hydrogen to give hydrogen halides but affinity for hydrogen decreases from fluorine to iodine. They dissolve in water to form hydrohalic acids. The acidic strength of these acids varies in the order: HF < HCl < HBr < HI. The stability of these halides decreases down the group due to decrease in bond (H–X) dissociation enthalpy in the order: H–F > H–Cl > H–Br > H–I.

(iv) Reactivity 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. These oxidesare essentially oxygen fluorides because of the higher electronegativity of fluorine than oxygen. Both are strong fluorinating agents. Chlorine, bromine and iodine form oxides in which the oxidation states of these halogens range from +1 to +7. A combination of kinetic and thermodynamic factors leads to the generally decreasing order of stability of oxides formed by halogens, I > Cl > Br. The higher oxides of halogens tend to be more stable than the lower ones. Chlorine oxides, Cl2O, ClO2, Cl2O6 and Cl2O7 are highly reactive oxidising agents and tend to explode. ClO2 is used as a bleaching agent for paper pulp and textiles and in water treatment. The bromine oxides, Br2O, BrO2, BrO3 are the least stable halogen oxides (middle row anomaly) and exist only at low temperatures. They are very powerful oxidising agents. The iodine oxides, 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.

(v) Reactivity towards metals: Halogens react with metals to form metal halides. For e.g., bromine reacts with magnesium to give magnesium bromide. Mg (s)+ Br2 ---MgBr (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 e.g., SnCl4, PbCl4, SbCl5 and UF6 are more covalent than SnCl2, PbCl2, SbCl3 and UF4 respectively.

(iv) Reactivity of halogens: towards other halogens: Halogens combine amongst themselves to form a number of compounds known as interhalogens of the types XX′, XX’3, XX5 ′ and XX7 ′ where X is a larger size halogen and X′ is smaller size halogen.

Chlorine

Preparation

It can be prepared by any one of the following methods:

(i) 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.

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

(ii) By the action of HCl on potassium permanganate. 2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2

Manufacture of chlorine

(i) 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

(ii) 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.

Properties

It is a greenish yellow gas with pungent and suffocating odour. It is about 2-5 times heavier than air. It can be liquefied easily into greenish yellow liquid which boils at 239 K. It is soluble in water. 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. It reacts with compounds containing hydrogen to form HCl.

H2 + Cl2 --2HCl

H2S + Cl2 ----2HCl + S

C10 H16 + 8Cl2--16HCl+ 10C

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

2NaOH + Cl2 → NaCl + NaOCl + H2O (cold and dilute)

6 NaOH + 3Cl2 → 5NaCl + NaClO3 + 3H2O (hot and conc.)

With dry slaked lime it gives bleaching powder.

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

Chlorine reacts with hydrocarbons and gives substitution products with saturated hydrocarbons and addition products with unsaturated hydrocarbons.

For example,

CH4 + Cl2 → CH3Cl + HCl

Methane Methyl chloride

C2H4 + Cl2 ⎯⎯→ C2H4Cl2

Ethene 1,2-Dichloroethane

Chlorine water on standing loses its yellow colour due to the formation of HCl and HOCl.

Hypochlorous acid (HOCl) so formed, gives nascent oxygen which is responsible for oxidising and bleaching properties of chlorine.

(i) It oxidises ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid and iodine to iodic acid.

2FeSO4 + H2SO4 + Cl2 → Fe2(SO4)3 + 2HCl Na2SO3 + Cl2 + H2O → Na2SO4 + 2HCl SO2 + 2H2O + Cl2 → H2SO4 + 2HCl I2 + 6H2O + 5Cl2 → 2HIO3 + 10HCl

(ii) It is a powerful bleaching agent; bleaching action is due to oxidation. Cl2 + H2O → 2HCl + O (iii) Coloured substance + O → Colourless substance

It bleaches vegetable or organic matter in the presence of moisture. Bleaching effect of chlorine is permanent.

Uses:

It is used (i) for bleaching wood pulp (required for the manufacture of paper and rayon), bleaching cotton and textiles,

(ii)in the extraction of gold and platinum

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

(iv) in sterilising drinking water and

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

Hydrogen Chloride

Preparation

In laboratory, it is prepared by heating sodium chloride with concentrated sulphuric acid.

NaCl + H2SO4⎯⎯⎯⎯ 420K → NaHSO4 + HCl

NaHSO4 + NaCl ⎯⎯⎯⎯ 823K → Na2SO4 + HCl

HCl gas can be dried by passing through concentrated sulphuric acid.

Properties

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.

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 NH

NH3 + HCl → NH4Cl

When three parts of concentrated HCl and one part of concentrated HNO3 are mixed, aqua regia is formed which is used for dissolving noble metals, e.g., gold, platinum.

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

Na2CO3 + 2HCl → 2NaCl + H2O + CO2

NaHCO3 + HCl → NaCl + H2O + CO2

Na2SO3 + 2HCl → 2NaCl + H2O + SO2

Uses:

It is used (i) in the manufacture of chlorine, NH4Cl and glucose (from corn starch),

(ii)for extracting glue from bones and purifying bone black,

(iii)in medicine and as a laboratory reagent.

Oxoacids of Halogens

If the acid contains oxygen (called an oxoacid), then the suffixes the lower and higher number of oxygens in the acid formula.

Fluorine has a very small size and highwhich is known as fluoric(I) acid or hypofluorous acid. The other elements of the halogen family form several oxoacids. They cannot be isolated in the pure state. They are stable in aqueous solution or in the form of salts.

Halogens generally form four series of oxoacids namely hypohalous acids (+1 oxidation state), halous acids (+3 oxidation state), halic acids (

Structures of the Oxoacids of Halogens

Interhalogen Compounds

When two different halogens react with each other, interhalogen compounds are formed. They can be assigned general compositions as XXof smaller size and X is more electropositive than Xnumber of atoms per molecule also increases. Thus, iodine (VII) fluoride shoulof atoms as the ratio of radii between I and F should be maximum. That is why its formula is IF7 (having maximum number of atoms).

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 NH

concentrated HCl and one part of concentrated HNO3 are mixed, aqua regia is formed which is used for dissolving noble metals, e.g., gold, platinum.

Hydrochloric acid decomposes salts of weaker acids, e.g., carbonates, hydrogen carbonates

→ 2NaCl + H2O + CO2

→ NaCl + H2O + CO2

→ 2NaCl + H2O + SO2

It is used (i) in the manufacture of chlorine, NH4Cl and glucose (from corn starch),

for extracting glue from bones and purifying bone black,

in medicine and as a laboratory reagent.

If the acid contains oxygen (called an oxoacid), then the suffixes –ous and –ic is used again, representing the lower and higher number of oxygens in the acid formula.

mall size and high electronegativity. Therefore, it forms only one oxoacid, HOF which is known as fluoric(I) acid or hypofluorous acid. The other elements of the halogen family form

oxoacids. They cannot be isolated in the pure state. They are stable in aqueous solution or in the

Halogens generally form four series of oxoacids namely hypohalous acids (+1 oxidation state), halous acids (+3 oxidation state), halic acids (+5 oxidation state) and perhalic acids (+7 oxidation state).

Structures of the Oxoacids of Halogens

When two different halogens react with each other, interhalogen compounds are formed. They can be XX′, XX3 ′, XX5 ′ and XX7 ′ where X is halogen of larger size and X′

of smaller size and X is more electropositive than X′. As the ratio between radii of X and X′ increases, the number of atoms per molecule also increases. Thus, iodine (VII) fluoride should have maximum number of atoms as the ratio of radii between I and F should be maximum. That is why its formula is IF7 (having

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.

concentrated HCl and one part of concentrated HNO3 are mixed, aqua regia is

hydrogen carbonates, sulphites,

It is used (i) in the manufacture of chlorine, NH4Cl and glucose (from corn starch),

ic is used again, representing

Therefore, it forms only one oxoacid, HOF which is known as fluoric(I) acid or hypofluorous acid. The other elements of the halogen family form

oxoacids. They cannot be isolated in the pure state. They are stable in aqueous solution or in the

Halogens generally form four series of oxoacids namely hypohalous acids (+1 oxidation state), halous +5 oxidation state) and perhalic acids (+7 oxidation state).

When two different halogens react with each other, interhalogen compounds are formed. They can be ′, XX3 ′, XX5 ′ and XX7 ′ where X is halogen of larger size and X′

′. As the ratio between radii of X and X′ increases, the d have maximum number

of atoms as the ratio of radii between I and F should be maximum. That is why its formula is IF7 (having

In general, interhalogen compounds are more reactive than halogens (except fluorine). This is because X–X′ bond in interhalogens is weaker than X–X bond in halogens except F–F bond.

All these undergo hydrolysis giving halide ion derived from the smaller halogen and a hypohalite (when XX′), halite (when XX′3), halite (when XX′5) and perhalate (when XX′7) anion derived from the larger halogen.

XX’ + H2O -- HX’ + HOX

Their molecular structures are very interesting which can be explained on the basis of VSEPR theory. The XX3 compounds have the bent ‘T’ shape, XX5 compounds square pyramidal and IF7 has pentagonal bipyramidal structures.

Uses:

These compounds can be used as non-aqueous solvents. Interhalogen compounds are very useful fluorinating agents. ClF3 and BrF3 are used for the production of UF6 in the enrichment of 235U.

U(s) + 3ClF3(l) → UF6(g) + 3ClF(g)

Group 18 Elements

He 2s2 2p6 Ne 3s2 3p6 Ar 4s2 4p6

Kr 5s2 5p6 Xe 6s2 6p6

All these are gases and chemically unreactive. They form very few compounds. Because of this they are termed noble gases.

Occurrence

All the noble gases except radon occur in the atmosphere. Their atmospheric abundance in dry air is ~ 1% by volume of which argon is the major constituent. Helium and sometimes neon are found in minerals of radioactive origin e.g., pitchblende, monazite, cleveite. The main commercial source of helium is natural gas. Xenon and radon are the rarest elements of the group. Radon is obtained as a decay product of 226Ra.

Ionisation 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

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.

Physical Properties

All the noble gases are monoatomic.

They are colourless, odourless and tasteless.

They are sparingly soluble in water.

They have very low melting and boiling points because the only type of interatomic interaction in these elements is weak dispersion forces.

Helium has the lowest boiling point (4.2 K) of any known substance.

It has an unusual property of diffusing through most commonly used laboratory materials such as rubber, glass or plastics.

Chemical Properties

In general, noble gases are least reactive. Their inertness to chemical reactivity is attributed to the following reasons:

The noble gases except helium (1s2 ) have completely filled ns2 np6 electronic configuration in their valence shell.

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

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) [673K, 1bar]

(xenon in excess)

Xe (g) + 2F2 (g) ⎯⎯⎯⎯⎯⎯ → XeF4 (s) [873 K, 7 bar]

(1:5 ratio)

Xe (g) + 3F2 (g) ⎯⎯⎯ − → XeF6(s) [573 K, 60 70bar]

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

XeF4 + O2F2 - XeF6 +O2

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 fluor 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 Xe03.

6XeF4 + 12 H2O → 4Xe + 2Xe03 + 24 HF + 3 O2

XeF6 + 3 H2O → XeO3 + 6 HF

Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2.

XeF6 + H2O → XeOF4 + 2 HF

XeF6 + 2 H2O → XeO2F2 + 4HF

XeO3 is a colourless explosive solid and has a pyramidal molecular structure. XeOF4 is a colourless volatile liquid and has a square pyramidal molecular structure.

The structures of the three xenon fluorides can be deduced from VSEPR

Uses:

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. There are no significant uses of Xenon and Krypton. They are used in light bulbs designed for special purposes.

NCERT IN-TEXT QUESTIONS

Page No. 174

Question 1. Why are pentahalides more covalent than trihalides in the members of the nitrogen family? Answer: The electronic configuration of the elements of nitrogen family (group 15) is ns2p3. Because of the inert pair effect, the valence s-electrons cannot be released easily for the bond formation. This means that the elements can form trivalent cation (E3+) by releasing valence p-electrons while it is difficult to form pentavalent cation (E5+). Under the circumstances, if all the five valence electrons are to be involved in the bond formation, the compounds showing Penta valency or +5 oxidation state must be of covalent nature. This is particularly the case in the higher members (Sb, Bi) of the family where the inert pair effect is quite prominent.

Question 2. Why is BiH3 the strongest reducing agent amongst all the hydrides of Group 15 elements? Answer: This is because as we move down the group, the size increases, as a result, length of E-H bond increases and its strength decreases, so that the bond can be broken easily to release H2 gas. Hence, BiH3 is the strongest reducing agent.

Page No. 175

7.3. Why is N2 less reactive at room temperature? Answer:

Due to presence of triple bond between two N-atoms (N ≡ N), the bond dissociation energy of N2 is very high. As a result, N2 becomes less reactive at room temperature.

Page No. 177

7.4. Mention the conditions required to maximise the yield of ammonia. Answer:

Ammonia is prepared by Haber’s process as given below:

7.5. How does ammonia react with a solution of Cu2+? Answer:

Page No. 179

7.6. What is the covalence of nitrogen in N2O5? Answer:

In N2O5, each N-atom has four shared pairs of e-1 s as shown:

Page No.182

7.7. Why is bond angle in PH+4 ion higher than in PH3?

Answer:

In both PH3 and PH+4 ion, the phosphorus atom is sp3 hybridised. However, in PH3 the central atom has a

pyramidal structure due to the presence of lone electron pair on the phosphorus atom.

Because of lone pair: shared pair repulsion which is more than that of shared pair: shared pair repulsion, the bond angle in PH3 is nearly 93-6°. In PH+4 ion, there is no lone electron pair on the phosphorus atom. It has a tetrahedral structure with bond angle of 109°-28′. Thus, the bond angle in PH+4 ion is higher than in PH3.

7.8. What happens when white phosphorus is heated with concentrated NaOH solution in an inert atmosphere of CO2? Answer:

Page No. 184

7.9. What happens when PCl5 is heated? Answer:

7.10. Write a balanced equation for the hydrolytic reaction of PC is in heavy water. Answer:

Page No. 185

7.11. What is the basicity of H3POAnswer:

7.12. What happens when H3PO4 is heated?Answer:

On heating, H3PO3 disproportionate

Page No. 189

7.13. List the important sources of sulphur.Answer

Sulphur mainly occurs in the combined states in earth’s crust in the form of sulphates and sulphides.Sulphates: gypsum (CaSO4.2H2O); Epsom (MgSOSulphides: Galena (PbS); zinc blende (ZnS);sulphur occur ‘as H2S and in organic materials such as eggs, proteins, garlic, onion, mustard, hair and wool.

7.14. Write the order of thermal stability of the Answer:

The thermal stability of hydrides of group 16 elements decreases down the group. This is because down the group, size of the element (M) increases, Mdecreases so that it can be broken down easH2Se > H2Te > H2Po

7.15. Why is H2O a liquid and H2S a gas?Answer:

Due to high electronegativity of O than S, Hresult, H2O exists as an associated molecule in which each O is tetrahedrally surrounded by four Hmolecules. Therefore, H2O is a liquid at room temperature.On the other hand, H2S does not undergo Htogether by weak van der Waals forces of attraction. A small amount of energy is required to break these forces of attraction. Therefore, H2S is a gas at room temperature.

PO4?

is heated?

to form PH3 and H3PO4 with O.S. of-3and + 5.

important sources of sulphur.

Sulphur mainly occurs in the combined states in earth’s crust in the form of sulphates and sulphides.O); Epsom (MgSO4.7H2O); baryte (BaSO4), etc.

Galena (PbS); zinc blende (ZnS); copper pyrites (CuFeS2); iron pyrites (FeSS and in organic materials such as eggs, proteins, garlic, onion, mustard, hair and

7.14. Write the order of thermal stability of the – hydrides of Group 16 elements.

The thermal stability of hydrides of group 16 elements decreases down the group. This is because down the group, size of the element (M) increases, M-H bond length increases and thus, stability of Mdecreases so that it can be broken down easily. Hence, we have order of thermal stability as H

S a gas?

Due to high electronegativity of O than S, H2O undergoes extensive intermolecular Hassociated molecule in which each O is tetrahedrally surrounded by four H

molecules. Therefore, H2O is a liquid at room temperature. S does not undergo H- bonding. It exists as discrete molecules which are held

er Waals forces of attraction. A small amount of energy is required to break these S is a gas at room temperature.

3and + 5.

Sulphur mainly occurs in the combined states in earth’s crust in the form of sulphates and sulphides. ), etc.

); iron pyrites (FeS2), etc. Traces of S and in organic materials such as eggs, proteins, garlic, onion, mustard, hair and

hydrides of Group 16 elements.

The thermal stability of hydrides of group 16 elements decreases down the group. This is because down H bond length increases and thus, stability of M-H bond

ily. Hence, we have order of thermal stability as H2O > H2S >

O undergoes extensive intermolecular H-bonding. As a associated molecule in which each O is tetrahedrally surrounded by four H2O

bonding. It exists as discrete molecules which are held er Waals forces of attraction. A small amount of energy is required to break these

Page No.190

7.16. Which of the following does not react with oxygen directly? Zn, Ti, Pt, Fe Answer:

Platinum (Pt) is a noble metal and does not react with oxygen directly.

7.17. Complete the following reactions: (i)C2H2 + O2 -> (ii) 4Al + 3 O2 -> Answer

Page No. 192

7.18. Why does O3 act as a powerful oxidising agent? Answer:

On heating, O3 readily decomposes to give O2 and nascent oxygen.

Since nascent oxygen is very reactive, therefore, O3 acts as a powerful oxidising agent.

7.19. How is O3 estimated quantitatively? Answer:

When O3 is treated with excess of KI solution buffered with borate buffer (pH = 9.2), I2 is liberated quantitatively.

The I2 thus liberated is titrated against a standard solution of sodium thiosulphate using starch as an indicator.

Page No.194

7.20. What happens when sulphur dioxide is passed through an aqueous solution of Fe (III) salt? Answer:

SO2 acts as a reducing agent and reduces aqueous solution of Fe (III)salt to Fe (II) salt.

7.21. Comment on the nature of two S-O bonds formed in S02 molecule. Are the two S-O bonds in this molecule equal? Answer:

SO2 exists as an angular molecule with OSO bond angle of 119.5°. It a resonance hybrid of two canonical-forms:

7.22. How is the presence of SO2 detected? Answer:

So2 is a colourless and pungent smell gas. It can be detected by using potassium permanganate solution. When So2 is passed through an acidified potassium permanganate solution, then it decolonises the solution as it reduces MnO4- ions to Mn2+ ions.

Page No. 197

7.23. Mention three areas in which H2SO4 plays an important role. Answer:

(i) Sulphuric acid is used for the manufacture of a number of chemicals like hydrochloric acid, phosphoric acid, nitric acid along with a large number of organic compounds. (ii) A mixture of concentrated nitric acid and concentrated sulphuric acid is used in the manufacture of explosives like picric acid, T.N.T, dynamite etc. (iii) Dilute solution of acid is employed in petroleum refining in order to remove the unwanted impurities of sulphur.

Question 24. Write the conditions to maximise the yield of H2SO4 by Contact process. Answer: The key step in the manufacture of sulphuric acid is oxidation of SO2 to SO3 in presence of V2O5 catalyst.

The reaction is exothermic and reversible. Hence, low temperature and high pressure are the favourable conditions for maximum yield of SO3. In practice a pressure of 2 bar and temperature of 720 K is maintained.

Question 25. Why is Ka2 « Ka1 for H2SO4 in water? Answer: H2SO4 is a very strong acid in water largely because of its first ionisation to H3O

+ and HSO4– The

ionisation of HSO4– to H3O

+ and SO42- is very small. That is why, Ka2« Ka1.

Page No. 202

Question 26. Considering the parameters such as bond dissociation enthalpy, electron gain enthalpy and

hydration enthalpy, compare the oxidising powers of F2 and Cl2. Answer: The oxidising powers of both the members of halogen family are expressed in terms of their electron accepting tendency and can be compared as their standard reduction potential values. F2 + 2e– → 2F–; E° = 2-87 V,

Cl2 + 2e– → 2Cl–; E° = 1-36 V Since the E° of fluorine is more than that of chlorine, it is a stronger oxidising agent. Three factors contribute towards the oxidation potentials of both the halogens. These are : (i) Bond dissociation enthalpy: Bond dissociation enthalpy of F2 (158 kJ mol-1) is less compared to that of Cl2 (242·6 kJ mol-1). (ii) Electron gain enthalpy: The negative electron gain enthalpy of F (- 332·6 kJ mol-1) is slightly less than of Cl (-348·5 kJ mol-1). (iii) Hydration enthalpy: The hydration enthalpy of F- ion (515 kJ mol-1) is much higher than that of Cl- ion (381 kJ mol-1) due to its smaller size. From the available data, we may conclude that lesser bond dissociation enthalpy and higher hydration enthalpy compensate lower negative electron gain enthalpy of fluorine as compared to chlorine. Consequently, F2 is a more powerful oxidising agent than Cl2.

Question 27. Give two examples to show the anomalous behaviour of fluorine. Answer:

Ionisation enthalpy, electro-negativity and electrode potential are higher for fluorine than the expected trends of another halogen.

Fluorine does not show any positive oxidation state except in HOF.

Question 28. Sea is the greatest source of some halogens. Comment. Answer: Sea water contains chlorides, bromides and iodides of sodium, potassium, magnesium and calcium but sodium chloride being the maximum makes sea water saline. Various sea weeds contain up to 0.5% iodine.

Page No.204

Question 29. Give the reason for bleaching action of Cl2. Answer: Chlorine bleaches by oxidation Cl2 + H2O → HCl + HOCl → HCl + [O] The nascent oxygen reacts with dye to make it colourless.

Question 30. Name two poisonous gases which can be prepared from chlorine gas. Answer: COCl2 (phosgene), CCl3NO2 (tear gas)

Page No.207

Question 31. Why is ICI more reactive than l2? Answer: In general, interhalogen compounds are more reactive than halogens due to weaker X-X’ bonding than X-X bond. Thus, ICI is more reactive than I2.

Question 32. Why is helium used in diving apparatus? Answer: Helium along with oxygen is used in the diving apparatus by the sea divers. Since it is very little soluble in blood, it reduces decompression and causes less discomfort to the diver in breathing. A mixture of helium and oxygen does not cause pain due to very low solubility of helium in blood as compared to nitrogen.

Question 33. Balance the following equation: XeF6 + H2O → XeO2F2 + 4HF Answer:

XeF6 + 2H2O → XeO2F2 + 4HF Question 34. Why has it been difficult to study the chemistry of radon? Answer: Radon is radioactive with very short half-life which makes the study of chemistry of radon difficult.

NCERT EXERCISES

Q 1:Discuss the general characteristics of Group 15 elements with reference to their electronic configuration, oxidation state, atomic size, ionisation enthalpy and electronegativity Answer General trends in group 15 elements (i) Electronic configuration: There are 5 valence electrons for all the elements in group 15. ns2np3 is their general electronic configuration. (ii) Oxidation states: All these elements require three or more electrons to complete their octets and have 5 valence electrons. It is difficult in gaining electrons as the nucleus will have to attract three more electrons. This happens only with nitrogen as it is the smallest in size and the distance between the nucleus and the valence shell is relatively small. The remaining elements of this group show a formal oxidation state of −3 in their covalent compounds. In addition to the −3 state, N and P also show −1 and −2 oxidation states. All the elements present in this group show +3 and +5 oxidation states. However, the stability of +5 oxidation state decreases down a group, whereas the stability of +3 oxidation state increases. This happens because of the inert pair effect.

(iii) Ionization energy and electronegativity Ionization decreases as we move down the group. This happens because of increase in atomicsizes. Moving down the group, electronegativity decreases due to increase in size. (iv) Atomic size: As we move down the group atomic size increases. This increase inthe atomic size is attributed to an increase in the number of shells.

Q 2: Why does the reactivity of nitrogen differ from phosphorus? Answer

Nitrogen is chemically less reactive. This is because of the high stability of its molecule, N2. InN2, the two nitrogen atoms form a triple bond. This triple bond has very high bond strength, which is very difficult to break. It is because of nitrogen’s small size that it isable to form pπ−pπ bonds with itself. This property is not exhibited by atoms such as phosphorus. Thus, phosphorus is more reactive than nitrogen.

Q 3: Discuss the trends in chemical reactivity of group 15 elements. Answer

General trends in chemical properties of group − 15

(i) Reactivity towards hydrogen: The elements of group 15 react with hydrogen to form hydrides of typeEH3, where E = N,P, As, Sb, or Bi. The stability of hydrides decreases on moving down fromNH3toBiH3.

(ii)Reactivity towards oxygen: The elements of group 15 form two types of oxides:E2O3 and E2O5, where

E = N, P, As, Sb, or Bi. The oxide with the element in the higher oxidation state is more acidic than the other. However, the acidic character decreases on moving down a group.

(iii) Reactivity towards halogens: The group 15 elements react with halogens to form two series of saltsEX3andEX5. However, nitrogen does not formNX5 as it lacks the d-orbital. All trihalides (exceptNX3) are stable.

(iv) Reactivity towards metals: The group 15 elements react with metals to form binary compounds in which metals exhibit−3 oxidation states.

Q 4: Why does NH3 form hydrogen bond but PH3 does not?

Answer

When compared to phosphorus nitrogen is highly electronegative. This results in a greaterattraction of electrons towards nitrogen inNH3 than towards phosphorus inPH3. Hence, the extent of hydrogen bonding inPH3 is very less as compared toNH3.

Q 5: How is nitrogen prepared in the laboratory? Write the chemical equations of the reactions involved Answer

An aqueous solution of ammonium chloride is treated with sodium nitrite.

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

NO andHNO3 are produced in small amounts. These are impurities that can be removed on passing nitrogen gas through aqueous sulphuric acid, containing potassium dichromate.

Q 6: How is ammonia manufactured industrially? Answer: Ammonia is prepared on a large-scale by the Haber’s process.

N2(g) + 3H2(g) ⇌ 2NH3(g)

The optimum conditions for manufacturing ammonia are: (i) Pressure (around 200×105 Pa) (ii) Temperature (4700 K) (iii) Catalyst such as iron oxide with small amounts ofAl2O3 andK2O

Q 7: Illustrate how copper metal can give different products on reaction with

Answer

Concentrated nitric acid is a strong oxidizing agent. It is used for oxidizing most metals.oxidation depend on the temperature, concentration of the acid, and alsoon the material undergoing oxidation.

3Cu + 8HNO3(dil.) → 3Cu(NO3)2 + 2NO + 4HCU + 4HNO3(conc.) →Cu(No3)2 + 2NO

Q 8: Give the resonating structures ofNOAnswer (1)

(2)

Q 9: The HNH angle value is higher than HPH, HAsH and HSbH angles. Why? [Hint: Can

beexplained on the basis of sp3 hybridisation inother elements of the group]. Answer

HydrideNH3PH3AsH3SbH3 H−M−H angle 107°92°91° 90° The above trend in the H−M−H bond angle can be explained on the basis of the electronegativity of the central atom. Since nitrogen is highly electronegative, there is highelectron density around nitrogen. Thiscauses greater repulsion between the electron pairsaround nitrogen, resulting in maximum bond angle.

Illustrate how copper metal can give different products on reaction with

Concentrated nitric acid is a strong oxidizing agent. It is used for oxidizing most metals.oxidation depend on the temperature, concentration of the acid, and alsoon the material undergoing

+ 2NO + 4H2O + 2NO2 +2H2O

Give the resonating structures ofNO2 and N2O5.

The HNH angle value is higher than HPH, HAsH and HSbH angles. Why? [Hint: Can

hybridisation in NH3 and only s−p bonding between hydrogen and

−M−H bond angle can be explained on the basis of the electronegativity of the central atom. Since nitrogen is highly electronegative, there is highelectron density around nitrogen. Thiscauses greater repulsion between the electron pairsaround nitrogen, resulting in maximum bond angle.

Illustrate how copper metal can give different products on reaction with HNO3.

Concentrated nitric acid is a strong oxidizing agent. It is used for oxidizing most metals.The products of oxidation depend on the temperature, concentration of the acid, and alsoon the material undergoing

The HNH angle value is higher than HPH, HAsH and HSbH angles. Why? [Hint: Can

−p bonding between hydrogen and

−M−H bond angle can be explained on the basis of the electronegativity of the central atom. Since nitrogen is highly electronegative, there is highelectron density around nitrogen. This causes greater repulsion between the electron pairsaround nitrogen, resulting in maximum bond angle.

We know that electronegativity decreasesinteractions between theelectron pairs decrease, thereby decr

Q 10: Why doesR3P=O exist but RAnswer

N (unlike P) lacks the d-orbital. This restricts nitrogen to expand its coordination numberbeyond four. Hence,R3N=O does not exist.

Q 11: Explain whyNH3 is basic whileAnswer

NH3 is distinctly basic while BiH3 is feebly basic.Nitrogen has a small size due to which the lone pair of electrons is concentrated in a smallregion. This means that the charge density per unit volumeatom increases and the charge gets distributed over a largearea decreasing the electron density. Hence, the electronelement hydrides decrease on moving down the group.

Q 12: Nitrogen exists as diatomic molecule and phosphorus asPAnswer Nitrogen owing to its small size has a tendency to formNitrogen thus forms a very stable diatomic molecule,form pπ−pπ bonds decreases (because of the large size of heavier elements). Therefore, phosphorus (like other heavier metals) exists in theP4

Q 13: Write the main differences between the properties of whiteAnswer:

electronegativity decreases on moving down a group. Consequently, the repulsive interactions between theelectron pairs decrease, thereby decreasing the H−M−H bond angle.

R3N=O does not (R = alkyl group)?

orbital. This restricts nitrogen to expand its coordination numberbeyond four.

is basic while BiH3 is only feebly basic.

is feebly basic. Nitrogen has a small size due to which the lone pair of electrons is concentrated in a smallregion. This means that the charge density per unit volume is high. On moving down agroup, the size of the central atom increases and the charge gets distributed over a large area decreasing the electron density. Hence, the electron-donating capacity of group 15element hydrides decrease on moving down the group.

Nitrogen exists as diatomic molecule and phosphorus asP4. Why?

Nitrogen owing to its small size has a tendency to form pπ−pπ multiple bonds with itself.Nitrogen thus forms a very stable diatomic molecule, N2. On moving down a group, the te

bonds decreases (because of the large size of heavier elements). Therefore, phosphorus (like other heavier metals) exists in theP4 state.

Write the main differences between the properties of white phosphorus

on moving down a group. Consequently, the repulsive −M−H bond angle.

orbital. This restricts nitrogen to expand its coordination numberbeyond four.

Nitrogen has a small size due to which the lone pair of electrons is concentrated in a smallregion. This is high. On moving down agroup, the size of the central

donating capacity of group 15

multiple bonds with itself. . On moving down a group, the tendency to

bonds decreases (because of the large size of heavier elements). Therefore, phosphorus (like

phosphorus and red phosphorus.

Q 14: Why does nitrogen show catenation properties less than phosphorus?

Answer

Catenation is much more common in phosphorous compounds than in nitrogen compounds. This is because of the relative weakness of the N−N single bond as compared to the P−P single bond. Since nitrogen atom is smaller, there is greater repulsion of electron density of two nitrogen atoms, thereby weakening the N−N single bond.

Q 15:Give the disproportionation reaction ofH3PO3. Answer

On heating, orthophosphorus acid H3PO3disproportionate to give orthophosphoric acid H3PO4 and phosphine PH3. The oxidation states of P in various species involved in the reaction are mentioned below. 4H3P

-3O3→ 3H3P+5O4 + P-3H3

Q 16: CanPCl5 act as an oxidising as well as a reducing agent? Justify. Answer

PCl5 can only act as an oxidizing agent. The highest oxidation state that P can show is +5. InPCl5, phosphorus is in its highest oxidation state (+5). However, it can decrease its oxidation state and act as an oxidizing agent.

Q 17: Justify the placement of O, S, Se, Te and Po in the same group of the periodic table interms of electronic configuration, oxidation state and hydride formation. Answer

The elements of group 16 are collectively called chalcogens. (i) Elements of group 16 have six valence electrons each. The general electronic configuration of these elements isns2np4, where nvaries from 2 to 6

(ii) Oxidation state: As these elements have six valence electrons ns2np4, they should display an oxidation state of −2. However, only oxygen predominantly shows the oxidation state of −2 owing to its high electronegativity. It also exhibits the oxidation state of −1 (H2O2), zero O2, and +2 (OF2). However, the stability of the −2 oxidation state decreases on moving down a group due to a decrease in the electronegativity of the elements. The heavier elements of the group show an oxidation state of +2, +4, and +6 due to the availability of d-orbitals.

(iii) Formation of hydrides: These elements form hydrides of formulaH2E, where E = O, S, Se, Te, PO. Oxygen and sulphur also form hydrides of typeH2E2. These hydrides are quite volatile in nature.

Q 18: Why is dioxygen a gas but sulphur a solid? Answer

Oxygen is smaller in size when compared to sulphur. Since its size is small, it can form pπ−pπ bonds and form O2(O==O) molecule. Also, the intermolecular forces in oxygen are weak van der Wall’s, which cause it to exist as gas. On the other hand, sulphurdoes not formM2 molecule but exists as a puckered structure held together by strong covalent bonds. Hence, it is a solid.

Q 19: Knowing the electron gain enthalpy values for O → O -1 and O → O 2- as −141 and 702kJmol−1 respectively, how can you account for the formation of a large number of oxides having O2−speciesandnotO−? (Hint: Consider lattice energy factor in the formation of compounds). Answer

More the lattice energy of a compound, more stable it will be. Stability of an ionic compound depends on its lattice energy.

Lattice energy is directly proportional to the charge carried by an ion. When a metalcombines with oxygen, the lattice energy of the oxide involvingOHence, the oxide havingO2− ions are more stable than oxides havingofO2− is energetically more favourable thanformation of

Q 20. Which aerosols deplete ozone?

Answer:

The aerosol which is responsible for the depletion of ozone is: Freons or chlorofluorocarbons (CFCs)

The molecules of CFS breaks down when there is presence of ultraviolet radiations and forms chlorine free radicals which then combines with ozone to form oxygen.

Q 21. Describe the manufacture of

Answer:

Manufacture

Sulphuric acid is one of the most important industrial chemicals worldwide

the manufacture of H2 SO4by contact process

In practice, the plant is operated at 2 bar (pressure) and 720 K (temperature). The sulphuric acid thus obtained is 96-98% pure.

More the lattice energy of a compound, more stable it will be. Stability of an ionic compound depends on

rgy is directly proportional to the charge carried by an ion. When a metalcombines with oxygen, the lattice energy of the oxide involvingO2− ion is much more thanthe oxide involving

ions are more stable than oxides havingO−. Hence, we can say that formation is energetically more favourable thanformation of O−.

Q 20. Which aerosols deplete ozone?

The aerosol which is responsible for the depletion of ozone is: Freons or chlorofluorocarbons (CFCs)

of CFS breaks down when there is presence of ultraviolet radiations and forms chlorine free radicals which then combines with ozone to form oxygen.

Q 21. Describe the manufacture of H2 SO4by contact process?

e most important industrial chemicals worldwide.

by contact process

In practice, the plant is operated at 2 bar (pressure) and 720 K (temperature). The sulphuric acid thus

More the lattice energy of a compound, more stable it will be. Stability of an ionic compound depends on

rgy is directly proportional to the charge carried by an ion. When a metalcombines with ion is much more thanthe oxide involving O− ion.

. Hence, we can say that formation

The aerosol which is responsible for the depletion of ozone is: Freons or chlorofluorocarbons (CFCs)

of CFS breaks down when there is presence of ultraviolet radiations and forms chlorine

In practice, the plant is operated at 2 bar (pressure) and 720 K (temperature). The sulphuric acid thus

Q 22: How is SO2 an air pollutant?

Answer:

The environment is harmed by sulphur dioxide in many ways:

Sulphuric acid is formed, when it is combined with water vapour present in the atmosphere. This causes acid that damages plants, soil, buildings (those made of marble are more prone) etc.

SO2 causes irritation in respiratory tract, throat, eyes and can also affect the larynx to cause breathlessness.

The colour of the leaves of the plant gets faded when it is exposed to sulphur dioxide for a long time. This defect is known as chlorosis. The formation of chlorophyll is affected by the presence of sulphur dioxide.

Q 23: Why are halogens strong oxidising agents?

Answer:

Halogens have an electronic configuration of np5, where n =2-6. Thus, halogens require only one more electron to complete their octet and to attain the stable noble gas configuration. Moreover, halogens have high negative electron gain enthalpies and are highly electronegative with low dissociation energies. As a result, they have a high tendency to gain an electron. Hence, they act as strong oxidising agents.

Q 24: Explain why fluorine forms only one oxoacid, HOF.

Answer:

Fluorine has high electronegativity and small size, hence it forms only one oxoacid i.e., HOF.

Q 25: Explain why in spite of nearly the same electronegativity, nitrogen forms hydrogen bonding while chlorine does not.

Answer:

Oxygen has a smaller size and due to which a higher electron density per unit volume. Hence, oxygen forms hydrogen bonds while chlorine does not despite having similar electronegative values.

Q 26. Write two uses of ClO2. Answer:

Applications of ClO2 ( a )Used for purification of water. ( b ) Used for bleaching.

Q 27. Why are halogens coloured? Answer:

Halogens are coloured because they take in radiations from the visible spectrum. This excites the valence electrons to a higher energy level. The amount of energy required for excitation differs from halogen to halogen, thus they exhibit different colours.

Q28. Write the reactions of F2 and Cl2 with water Answer:

( i ) Cl2 + H2O → HCl + HOCl ( ii ) 2 F2 + 2H2O → 4H+ + 4F- + O2 + 4HF

Q29. How can you prepare Cl2 from HCl and HCl from Cl2? Write reactions only Answer:

( i ) HCl is prepared from Cl2by reacting it with water. Cl2 + H2O → HCl + HOCl

( ii ) Cl2 is prepared by Deacon’s process from HCl 4HCl + O2 → 2Cl2 + 2H2O

Q30. What inspired N. Bartlett for carrying out reaction between Xe and PtF6? Answer:

N. Barlett observed that PtF6 and O2 react to produce a compound O2+[ PtF6]–.

As the first ionization enthalpy of Xe( 1170 kJ/mol ) is very close to that of O2 , he figured that PtF6 could also oxidize Xe to Xe+. Thus, he reacted PtF6 and Xe to form a red coloured compound Xe+ [ PtF6]

–.

Q31. What are the oxidation states of phosphorus in the following: ( a ) H3PO3 ( b ) PCl3 ( c ) Ca3P2 ( d ) Na3PO4 ( e ) POF3?

Answer:

Let the oxidation state of phosphorous be x (a) H3PO3 3 + x + 3( -2) = 0 x -3 = 0 x =3

(b) PCl3 x + 3( -1) = 0 x = 3

(c) Ca3P2 3(2) + 2 (x) = 0 2x = -6 x = -3

(d) Na3PO4 3(1) + x + 4(-2) = 0 x -5 =0 x =5

(e)POF3 x + ( -2) + 3( -1) = 0 x -5 = 0 x = 5

Q 32. Write balanced equations for the following: (i) NaCl is heated with sulphuric acid in the presence of MnO2. (ii) Chlorine gas is passed into a solution of NaI in water. Answer:

(a) 4NaCl + MnO2 + 4H2SO4→ MnCl2 + 4NaHSO4 + 2H2O +Cl2 (b) Cl2 + NaI → 2NaCl + I2

Q33. How are xenon fluorides XeF2, XeF4 and XeF6 obtained? Answer:

XeF2, XeF4 and XeF6are obtained through direction reactions between Xe and F2. The product depends upon the conditions of the reaction: Xe + F2→ XeF2 (excess)

Xe + 2F2→ XeF4 ( 1:5 ratio)

Xe (g) + 3F2 (g) ⎯⎯⎯ − → XeF6(s) [573 K, 60 70bar]

Q34. With what neutral molecule is ClO– isoelectronic? Is that molecule a Lewis base? Answer:

ClO– is isoelectronic with ClF. Total electrons in ClO– = 17 + 8 + 1 =26 Total electrons in ClF = 17 + 9 = 26 As ClF accepts electrons from F to form ClF3, ClF behaves like a Lewis base.

Q35. How are XeO3 and XeOF4 prepared? Answer:

XeO3 can be obtained using two methods: ( 1 ) 6XeF4 + 12H2O → 4Xe + 2XeO3 + 24HF + 3O2 ( 2 ) XeF6 + 3H2O → XeO3 + 6HF XeOF4 is obtained using XeF6 XeF4 + H2O → XeOF4 + 2HF

Q36. Arrange the following in the order of property indicated for each set: (i) F2, Cl2, Br2, I2 – increasing bond dissociation enthalpy. (ii) HF, HCl, HBr, HI – increasing acid strength. (iii) NH3, PH3, AsH3, SbH3, BiH3 – increasing base strength. Answer:

(1) Bond dissociation energy normally lowers on moving down a group because of increase in the atomic size. However, F2has a lower bond dissociation energy than Cl2 and Br2. This is because the atomic size of fluorine is very small. Therefore, the increasing order for bond dissociation enthalpy is: I2< F2< Br2< Cl2 (2) Bond dissociation energy of a H-X molecule (where X = F, Cl, Br, I) lowers with an increase in the size of an atom. As, H-I bond is the weakest it will be the strongest acid. Therefore, the increasing order acidic strength is : HF <HCl<HBr< HI

(3) BiH3≤ SbH3<AsH3< PH3< NH3 On moving from nitrogen to bismuth, the atomic size increases but the electron density of the atom decreases. Hence, the basic strength lowers.

Q37. Which one of the following does not exist? (i) XeOF4 (ii) NeF2 (iii) XeF2 (iv) XeF6 Answer:

The one that does not exist is NeF2.

Q38. Give the formula and describe the structure of a noble gas species which is isostructural with: ( a ) ICl4–

( b ) IBr2– ( c ) BrO3– Answer:

(a) XeF4 is isoelectronic toICl4–. And it square planar in geometry:

(b) XeF2 is isoelectronic with IBr2–. It has a linear structure.

(c)XeO3 is isoelectric and isostructural to BrO3–. It has a pyramidal structure.

Q39. Why do noble gases have comparatively large atomic sizes?

Answer:

Noble gases have atomic radii that corresponds to van der Waal’s radii. Whereas, other elements have a covalent radius. Now, by definition, van der Waal’s radii are bigger than covalent radii. This is the reason why noble gases have relatively bigger atomic sizes.

Q40. List the uses of neon and argon gases. Answer:

Uses of Argon gas: (a)Argon is used to keep an inert atmosphere in high temperature metallurgical operations like arc welding. (b)It is used in fluorescent and incandescent lamps where it is required to check the sublimation of the filament. Thereby, increasing the life of the lamp. (c) Argon is used in laboratories to handle substances that are air-sensitive.

Uses of neon gas: (a) Neon is filled in discharge tubes for advertising or decoration. (b) Neon is used for making beacon lights. (c) It is used alongside helium to protect electrical equipment against high voltage.