NCERT Solutions For Class 12 Chemistry Chapter 7 The p Block Elements

7.1 Discuss the general characteristics of Group 15 elements with reference to their electronic configuration, oxidation state, atomic size, ionisation enthalpy and electronegativity.
Sol. In group 15 of the Periodic Table, the elements, nitrogen (₇N), phosphorus (₁₅P), arsenic (₃₃As), antimony (₅₁Sb) and bismuth (₈₃Bi) are present. The elements of this group can exhibit various oxidation states ranging between -3 to + 5. Negative oxidation state will be exhibited when they combine with less electronegative element andpositive oxidation state will be exhibited with more electronegative element. Positive oxidation state becomes more favourable as we more down the group due to increasing metallic character & electropositivity. Although due to inert pair effect the stability of +5 state will also decrease. The only stable compound of Bi (V) is BiF₅.
The atomic (covalent) and ionic radii (in a particular oxidation state) of the elements of nitrogen family (group 15) are smaller than the corresponding elements of carbon family (group 14). On moving down the group, the covalent and ionic radii (in a particular oxidation state) increase with increase in atomic number. There is a considerable increase in covalent radius from N to P. However, from As to Bi, only a small increase is observed.
As the size increases on moving down the group, the ionisation enthalpy increases. The ionisation enthalpy of nitrogen group elements is more than the corresponding elements of oxygen group. This is because of more stable half filled outermost p- subshell of nitrogen group elements. Electronegativity decreases down the group with increase in atomic size.

7.2 Why does the reactivity of nitrogen differ from phosphorus?
Sol. N₂ exist as a diatomic molecule containing triple bond.batoeen two N-atoms. Due to the presence —of triple bond betweep the two N-atoms, the bond dissociation energy is large (941 .4 kJ mol^-1 ). As aresult of this N₂ is inert and unreactive whereas, phosphorus exists as a tetratomic molecule, containg P – P single bond. Due to the presence of single bond, the bond dissociation energy is weaker (213 kJmol^-1 ) than N a N triple bond (941 .4 kJ mol^-1 ) and moreover due to presence of angular strain in P₄ tetrahedra. As a result of this, phosphorus is much more reactive than nitrogen.

7.3 Discuss the trends in chemical reactivity of group 15 elements.
Sol. Hydrides: All elements of group 15 form gaseous hydrides of the type MH₃.
In all the hydrides the central atom is sp³ hybridized and their shape is pyramidal due to presence of lone pair of electrons.
(a)The basic strength of the hydrides decreases as we move down the group.
Thus, NH₃ is the strongest base.
NH₃ > PH₃ > AsH₃ > SbH₃
(b)The thermal stability of the hydrides decreases as the atomic size increases, i.e., the M – H bond strength decreases which means reducing character increases.
(c)In the liquid state, the molecules of NH₃are associated due to hydrogen bonding. The molecules of other hydrides are not associated.
(d)NH₃ is soluble in water whereas other hydrides are insoluble.
(e)All the hydrides, except NH₃, are strong reducing agents and react with metal ions (Ag⁺, Cu²⁺, etc.) to form phosphides, arsenides or antimonides.
Halides: The elements of group 15 form two series of halides MX₃ and MX₅.
(a)All the elements of the group form trihalides. The ionic character of trihalides increases as we move down the group. Except NCl₃ all the trihalides are hydrolysed by water. This is due to the absence of d-orbitals in nitrogen.
(b)PF₃ is not hydrolysed because fluorine being more electronegative than oxygen forms more stable bonds with phosphorus than P – O bonds.
(c)N cannot form NX₅ because of non-availability of rforbitals. Bi cannot form BiX₃ because of reluctance of 6s electrons of Bi to participate in bond formation.
(d)The hybridisation of M in MX₃ is sp³ and shape is pyramidal. M in MX₅ is sp³ as hybridised and shape is trigonal pyramidal. The axial bonds in MX₅ are weaker and longer, So MX₅ are less stable and decompose on heating eg:

Oxides:
(a)Nitrogen forms a number of oxides. The rest of the members (P, As, Sb and Bi) of the group form two types of oxides : E₂0₃ and E₂O₅.
(b)The reluctance of P, As, Sb and Bi to enter into pπ -pπ multiple bonding leads to cage structures of their oxides and they exist as dimers, E₄O₆ and E₅O₁₀.
(c)The basic nature of die oxides increases with increase in atomic number of the element. Thus, the oxides of nitrogen (except N₂0 and NO), P (III) and As (III) are acidic, Sb (III) oxide is amphoteric and Bi (III) oxide is basic.

7.4 Why does NH₃ form hydrogen bond but PH₃ does not?
Sol. Nitrogen has an electronegativity value 3.0, which is much higher than that of H (2.1). As a result, N – H bond is quite polar and hence NH₃ undergoes intermolecular H – bonding.

Phosphorus have an electronegativity value 2-1. Thus, P – H bond is not polar and hence PH₃ does not undergo H – bonding.

7.5 How is nitrogen prepared in the laboratory? Write the chemical equations of the reactions . involved.
Sol. In laboratory, nitrogen is prepared by heating an equimolar aqueous solution of ammonium chloride and sodium nitrite. As a result of double decomposition reaction, ammonium nitrite is formed. Ammonium nitrite is unstable and decompose to form nitrogen gas.

7.6 How is ammonia manufactured industrially?
Sol. Commercially, by Haber’s process.

iron oxide, K₂O, Al₂0₃ The optimum conditions for the production of NH₃ are pressure of 200 atm and temperature of 100K.

7.7 Illustrate how copper metal can give different products on reaction with HN0₃.
Sol. On heating with dil HN0₃, copper gives copper nitrate and nitric oxide.

7.8 Give the resonating structures of N0₂ and N₂O₅.
Sol. Resonating structures of N0₂ are:

7.9 The HNH angle value is higher than HPH, H AsH and HSbH angles. Why?
(Hint: Can be explained on the basis of sp³ hybridisation in NH₃ and only s-p bonding , between hydrogen and other elements of the group).
Sol. In all these cases, the central atom is sp³ hybridized. Three of the four sp³ orbitals form three σ-bonds, while the fourth contains the lone pair of electrons. On moving down from N to Sb, the electronegativity of the central atom goes on decreasing. As a result of this, bond pairs of electrons lie away and away from the central atom. This is because of the force of repulsion between the adjacent bond pairs goes on decreasing and the bond angles keep on decreasing from NH₃ to SbH₃. Thus, bond angles are in the order:

7.10 Why does R₃P = O exist but R₃N = O does not (R = alkyl group)?
Sol. Nitrogen does not contains d-orbitals. As a result, it cannot expand its covalency beyond four and cannot form pπ – dπ multiple bonds. In constrast, P contains the d-orbitals, and can expand its covalency beyond 4 and can form pπ-dπ multiple bonds. Hence R₃P = O exist but R₃N = O does not.

7.11 Explain why NH₃ is basic while BiH₃ is only feebly basic.
Sol. In both NH₃ and BiH₃, N and Bi have a lone paif of electrons on the central atom and hence should behave as Lewis bases. But NH₃ is much more basic than BiH₃. Since the atomic size of N is much smaller than that of Bi, therefore, electron density on N-atom is much higher than that on Bi-atom. Thus, the tendency of N in NH₃ to donate its lone pair of electrons is much more in comparison to tendency of Bi in BiH₃. Hence, NH₃ is more basic than BiH₃.

7.12 Nitrogen exists as diatomic molecule and phosphorus as P₄. Why?
Sol. Nitrogen exists as a diatomic molecule having a triple bond between the two N-atoms, This is due its small size that it forms pπ-pπ multiple bonds with itself and with carbon /oxygen as well. On the other hand, phosphorus due to its larger size does not form multiple pπ-pπ bonds with itself. It prefers to form P – P single bonds and hence it exists as tetrahedral P₄ molecule.

7.13 Write main differences between the properties of white phosphorus and red phosphorus.
Sol.

Structure of white and red phosphorus are given below:

7.14 Why does nitrogen show catenation properties less than phosphorus?
Sol. The extent of catenation depends upon the strength of the element – element bond. The N – N bond strength (159 kJ mol^-1 ) is weaker than P – P bond strength (213 kJ mol^-1 ). Thus, nitrogen shows less catenation properties than phosphorus.

7.15 Give the disproportionation reaction of H₃ P0₃ .
Sol. On heating, H₃ P0₄ undergoes self – oxidation reduction, i.e:, disproportionation to form PH₃ .

7.16 Can PCl₅ act as an oxidising as well as a reducing agent Justify.
Sol. The oxidation state of P in PCl₅ is+5. Since P has five electrons in its valence shell, therefore, it cannot donate electron and cannot increase its oxidation state beyond + 5, Thus, PCl₅ cannot act as a reducing agent. It can act as oxidizing agent by itself undergoing reduction.

7.17 Justify the placement of O, S, Se, Te and Po in the same group’of the periodic table in terms of electronic configuration, oxidation state and hydride formation.
Sol. (1)Electronic configuration:
O (At. no. = 8) = [He] 2s² 2p⁴
S (At. no. = 16) = [Ne] 3s² 3p⁴
Se (At. no. = 34) = [Ar] 3d¹⁰ 4s² 4p⁴
Te (At. no. = 52) = [Kr] 4d¹⁰ 5s² 5p⁴ ,
Po (At. no. = 84) = [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p⁴ ,
Thus, all these elements have the same ns² np⁴ (n = 2 to 6) valence shell electronic configuration, hence are justified to be placed in group 16 of the Periodic Table.
(2)Oxidation state : Two more electrons are needed to acquire the nearest noble gas configuration. Thus, the minimum oxidation state of these elements should be – 2. O and to some extent S show – 2 oxidation state. Other element being more electropositive than O and S, do not show negative oxidation state. As these contain six electrons, thus, maximum oxidation state shown by them is+ 6. Other oxidation state shown by them are + 2 and + 4. O do not show+4 and + 6 oxidation state, due to the absence of d-orbitals. Thus, on the basis of maximum and minimum oxidation states, these elements are justified to be placed in the same group 16 of the periodic table.
(3)Hydride formation: All these elements share two of their valence electrons with 1 s- orbital of hydrogen to form hydrides of the general formula EH₂, i.e., H₂0, H₂S, H₂Se, H₂Te and H₂Po. Thus, on the basis of hydride formation, these elements are justified to be placed in the same group 16 of the Periodic Table.

7.18 Why is dioxygen a gas but sulphur a solid?
Sol. Due to the small size and high electronegativity, oxygen forms pπ- pπ multiple bonds. As a result, oxygen exists as diatomic (O₂) molecules. These molecules are held together by weak van der Waal’s forces of attraction which can be overcome by collisions of the molecules at room temperature. Therefore, O₂ is a gas at room temperature. Due to its bigger size and lower electronegativity, sulphur does not form pn-pn multiple bonds. It prefers to form S – S single bonds. S – S single bond is stronger then O-O single bond. Thus, sulphur has higher tendency for catenation than oxygen. Due to higher tendency for catenation and lower tendency for pπ – pπ multiple bonds sulphur exits as octa-atomic (S_g) molecule. Due to bigger size, the force of attraction holding the S_g molecules together are much stronger which cannot be overcome by collisions of molecules at room temperature. Therefore, sulphur is solid at room temperature.

7.19 Knowing the electron gain enthalpy values of O—>O^– and O—>O^2- as -141 and 702 kJ mol^-1 respectively, how can you account for [he formation of a large number of oxides having O^2- species and not O^–?
Sol. Let us consider the reaction of oxygen with monopositve metal, we can have two compounds. MO(O in -1 state) and M₂O (O in -2 state). The energy required for formation of O^-2 is compensated by increased coulombic attraction between M⁺ and O^-2. Coulombic force of attraction, F_A is proportional to product of charges on ions i.e.

where q₁ and q₂ are charges on ions and r is distance between ions. Same logic can be applied if metal is dispositive.

7.20 Which aerosols deplete ozone?
Sol. Aerosols like chlorofluorocarbons (CFC’s), i.e., freon (CCl₂F₂), depletes the ozone layer by supplying Cl* free radicals which convert O₃ to O₂

7.21 Describe the manufacture of H₂SO₄ by contact process?
Sol. Preparation of sulphuric acid:By Contact Process: Burning of sulphur or sulphide ores in presence of oxygen to produce SO₂. Catalytic oxidation of SO₂ with O₂to give SO₃ in the presence of V₂O₅.

Then SO₃ made to react with sulphuric acid of suitable normality to obtain a thick oily liquid called oleum.

Then oleum is diluted to obtain sulphuric acid of desired concentration.

The sulphuric acid obtained by contact process is 96-98% pure.

7.22 How is SO₂ an air pollutant?
Sol. (1) SO₂ dissolves in moisture present in air to form H₂SO₄ which damages building materials especially marble (acid – rain).- CaCO₃ + H₂SO₃ ——->CaSO₃ + H₂0 + CO₂
(2)It corrodes metals like Fe and steel. It also brings about fading and deterioration of fabrics, leather, paper, etc., and affecting the colour of paints.
(3)Even in low concentration (= 0.03 ppm), it has damaging effect on the plants. If exposed for a long time, i.e., a few days or weeks, it slows down the formation of chlorophyll i. e., loss of green colour. This is called chlorosis.
(4)It is strongly irritating to the respiratory track. It cause throat and eye irritation, resulting into cough, tears and redness in eyes. It also cause breathlessness and effects larynx i. e. „ voice box.

7.23 Why are halogens strong oxidising agents?
Sol. The halogens are strong oxidising agents due to
low bond dissociation enthalpy, high electronegativity and large negative electron gain enthalpy.

7.24 Explain why fluorine forms only one oxoacid, HOF.
Sol. Cl, Br and I form four series of oxo acids of general formula HOX, HOXO, HOXO₂, and H0XO₃. In these oxo-adds, the oxidation states of halogens are + 1, + 3, + 5, and + 7 respectively. However, due to high electronegativity, small size and absence of d-orbitals, F does not form oxo-acids with + 3, + 5 and + 7, oxidation states. It just forms one oxo-acid (HOF).

7.25 Explain why inspite of nearly the same electronegativity, nitrogen forms hydrogen bonding while chlorine does not.
Sol. Both .nitrogen (N) and chlorine (Cl) have electronegativity of 3.0. However, only nitrogen is involved in the hydrogen bonds (e.g., NH3) and not chlorine. This is due to smaller atomic size of nitrogen (atomic radius =70 pm) as compared to chlorine (atomic radius = 99) pm), therefore, N can cause greater polarisation of N-H bond than Cl in case of Cl—H bond.Consequently, N atom is involved in hydrogen bonding and not chlorine.

7.26 Write two uses of ClO₂
Sol. (1) ClO₂ is an excellent bleaching agent. It is 30 times stronger bleaching agent then the Cl₂. It is used as a bleaching agerit for paper pulp in paper industry and in textile industry. (2) ClO₂ is also a powerful oxidising agent and chlorinating agent. It acts as a germicide for disinfecting water. It is used for purifying drinking water.

7.27 Why are halogens coloured?
Sol. The halogens are coloured because their molecules absorb light in the visible region. As a result of which their electrons get excited to higher energy levels while the remaining light is transmitted. The color of halogens is the color of this transmitted light.

7.28 Write the reactions of F₂ and Cl₂ with water.
Sol.

7.29 How can-you prepare Cl₂ from HCl and HCl from CI₂? Write reactions only.
Sol.

7.30 What inspired N. Bartlett for carrying out reaction between Xe and PtF₆?
Sol. N. Bartlett observed that PtF₆ reacts with O₂to give an compound O₂+ [PtF₆]^–.
PtF₆ (g) + O₂ (g) ——–>O₂+[PtF₆]^–
Since the first ionization enthalpy of Xe (1170 kJ mol^-1 )is fairly close to that of 02 molecule (1175 kJ mol^-1 ), he thought that PtF₆ should also oxidise Xe to Xe⁺. This inspired Bartlett to carryout the reaction between Xe and PtF₆. When PtF₆ and Xe were made to react, a rapid reaction took place and a red solid, Xe+[PtF₆]^– was obtained.

7.31 What are the oxidation states of phosphorus in the following: –
(i) H₃PO₃ (ii)PCl₃
(iii) Ca₃P₂(iv)Na₃PO₄
(v) POF₃
Sol.

7.32 Write balanced equations for the following:
(i) NaCl is heated witlrsulphuric acid in the presence of MnO₂
(ii) Chlorine gas is passed into a solution of Nal in water.
Sol.

7.33 How are xenon fluorides XeF₂, XeF₄ and XeF₆ obtained?
Sol. XeF₂ , XeF₄ and XeF₆ are obtained by direct reaction between Xe and F₂ as follows:

7.34 With what neutral molecule is CIO^– isoelectronic? Is that molecule a Lewis base?
Sol. CIO^– has 17 + 8 +1 = 26 electrons.Also, OF₂ has (8 + 2 x 9) = 26 electrons, and ClF has (17 + 9) = 26 electrons.Out of these, ClF can act as Lewis base. The chlorine atom has three lone pair of electrons which it donates to form compounds like ClF₃, ClF₅ and ClF₇

7.35 How are XeO₃ and XeOF₄prepared?
Sol.

7.36 Arrange the following in the order of property indicated for each set: –
(i) F₂ , Cl₂ , Br₂ , I₂ – increasing bond dissociation enthalpy.
(ii) HF, HCI, HBr, HI – increasing acid . strength.
(iii) NH₃, PH₃, AsH₃, SbH₃, BiH₃ – increasing Sol. base strength.
Sol. (i) Bond dissociation enthalpy decreases as the bond distance increases from F₂ to I₂ due to increase in the size of the atom, on moving from F to I.
F – F bond dissociation enthalpy is smaller then the Cl – Cl and even smaller than Br – Br. This is because F atom is very small and have large electron-electron repulsion among the lone pairs of electrons in F₂ molecule where they are much closer to each other than in case of Cl₂. The increasing order of bond dissociation enthalphy is I, < F₂ < Br₂ < Cl₂
(ii) Acid strength of HF, HCI, HBr and HI depends upon their bond dissociation enthalpies. Since the bond dissociation enthalpy of H – X bond decreases from H – F to H-l as the size of atom increases from F to I.
Thus, the acid strength order is HF < HCI < HBr < HI
The weak acidic strength of HF is also due to H-bonding due to which release of H⁺ becomes difficult.
(iii) NH₃, PH₃, ASH₃, SbH₃ and BiH₃ behaves as Lewis bases due to the presence of lone pair of electrons on the central atom. As we move from N to Bi, size of atom increases. Electron density on central atom decreases and hence the basic strength decreases from NH₃ to BiH₃. Thus basic strength order is BiH₃<SbH₃<AsH₃<PH₃<NH₃

7.37 Which one of the following does not exist ?
(i)XeOF₄ (ii)NeF₂
(iii)XeF₄ (iv)XeF₆
Sol. NeF₂ does not exist. This is because the sum of first and second ionization enthalpies of Ne are much higher than those of Xe. Consequently, F₂ can oxidise Xe to Xe²⁺ but cannot oxidise Ne to Ne²⁺.

7.38 Give the formula and describe the structure of a noble gas species which is isostructural with: (i) ICI₄^– (ii) IBr₂^– (iii) Br0₃^–
Ans.(i) ICI₄^–: In ICI₄^–, central atom I has seven valence electrons and one due to negative charge. Four out of these 8 electrons are utilized in forming four single bonds with four Cl atoms. Four remaining electrons constitutes the two lone pairs. It is arranged in square planar structure. ICI₄^– has 36 valence electrons. A noble gas species having 36 valence electrons is XeF₄ (8 + 4 x 7 = 36). XeF₄ is also square planar.

(ii) IBr₂^–: In IBr₂^–, central atom I has eight electrons. Two of these are utilized in forming two single bonds with two Br atom. Six remaining electrons constitutes three lone pairs. It is arranged in linear structure.

IBr₂^– has 22 valence electrons. A noble gas species having 22 valence electrons is XeF₂ (8+2 x 7=22).
XeF₂ is also linear.
(iii) In Br0₃^– ion the central Br atom has 8 valence electrons (7 +1). Out of these, it shares 4 with two atoms of O forming Br = O bonds. Out of the remaining four .electrons, 2 are donated to the third O atom which accounts for its negative charge. The remaining 2 electrons constitute one lone pair. In order to minimise the force of repulsion, the structure of Br0₃^– ion must be pyramidal. Br0₃^– ion has (7 + 3 x 6 + 1) = 26 valence electrons and is isoelectronic as well as iso-structural with noble gas species Xe0₃ which has also 26(8 + 3 x 6) electrons.

7.39 Why do noble gases have comparatively large atomic sizes?
Sol. This is because noble gases have only van der Waal’s radii while others have covalent radii, van der Waal’s radii are larger than covalent radii.

7.40 List the uses of neoirand argon gases.
Sol. Uses of Neon
Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes. Glow’of different colours ‘neon signs’ can be produced by mixing neon with other gases. Neon bulbs and used in botanical gardens and in green’ houses.
Uses of Argon
Argon is used mainly to provide an inert atmosphere in high temperature metallurgical processes such as arc welding of metals and alloys. In the laboratory, it is used for handling substance which are air sensitive.
It is used in filling incandescent and fluorescent lamps where its presence retards the sublimation of the filament and thus increases the life of the lamp.It is also used in “neon signs” for obtaining lights of different colours.