p-Block Trends

⚠In the NEET syllabus; removed from current NCERT.

11.1.6 Chemical Properties Oxidation state and trends in chemical reactivity Due to small size of boron, the sum of its first three ionization enthalpies is very high. This prevents it to form +3 ions and forces it to form only covalent compounds. But as we move from B to Al, the sum of the first three ionisation enthalpies of Al considerably decreases, and is therefore able to form Al3+ ions. In fact, aluminium is a highly electropositive metal. However, down the group, due to poor shielding effect of intervening d and f orbitals, the increased effective nuclear charge holds ns electrons tightly (responsible for inert pair effect) and thereby, restricting their participation in bonding. As a result of this, only p-orbital electron may be involved in bonding. In fact in Ga, In and Tl, both +1 and +3 oxidation states are observed. The relative stability of +1 oxidation state progressively increases for heavier elements: Al<Ga<In<Tl. In thallium +1 oxidation state is predominant whereas the +3 oxidation state is highly oxidising in character. The compounds in +1 oxidation state, as expected from energy considerations, are more ionic than those in +3 oxidation state. In trivalent state, the number of electrons around the central atom in a molecule of the compounds of these elements (e.g., boron in BF3) will be only six. Such electron deficient molecules have tendency to accept a pair of electrons to achieve stable electronic configuration and thus, behave as Lewis acids. The tendency to behave as Lewis acid decreases with the increase in the size down the group. BCl3 easily accepts a lone pair of electrons from ammonia to form BCl3⋅NH3.

⚠In the NEET syllabus; removed from current NCERT.

In tetravalent state the number of electrons around the central atom in a molecule (e.g., carbon in CCl4) is eight. Being electron precise molecules, they are normally not expected to act as electron acceptor or electron donor species. Although carbon cannot exceed its covalence more than 4, other elements of the group can do so. It is because of the presence of d orbital in them. Due to this, their halides undergo hydrolysis and have tendency to form complexes by accepting electron pairs from donor species. For example, the species like, SiF6 2–, [GeCl6]2–, [Sn(OH)6]2– exist where the hybridisation of the central atom is sp3d2.

⚠In the NEET syllabus; removed from current NCERT.

11.7 ALLOTROPES OF CARBON Carbon exhibits many allotropic forms; both crystalline as well as amorphous. Diamond and graphite are two well-known crystalline forms of carbon. In 1985, third form of carbon known as fullerenes was discovered by H.W.Kroto, E.Smalley and R.F.Curl. For this discovery they were awarded the Nobel Prize in 1996. 11.7.1 Diamond It has a crystalline lattice. In diamond each carbon atom undergoes sp3 hybridisation and linked to four other carbon atoms by using hybridised orbitals in tetrahedral fashion. The C–C bond length is 154 pm. The structure extends in space and produces a rigid three-dimensional network of carbon atoms. In this structure (Fig. 11.3) directional covalent bonds are present throughout the lattice. It is very difficult to break extended covalent bonding and, therefore, diamond is a hardest substance on the earth. It is used as an abrasive for sharpening hard tools, in making dyes and in the manufacture of tungsten filaments for electric light bulbs. 11.7.2 Graphite Graphite has layered structure (Fig.11.4). Layers are held by van der Waals forces and distance between two layers is 340 pm. Each layer is composed of planar hexagonal rings of carbon atoms. C—C bond length within the layer is 141.5 pm. Each carbon atom in hexagonal ring undergoes sp2 hybridisation and makes three sigma bonds with three neighbouring carbon atoms. Fourth electron forms a π bond. The electrons are delocalised over the whole sheet. Electrons are mobile and, therefore, graphite conducts electricity along the sheet. Graphite cleaves easily between the layers and, therefore, it is very soft and slippery. For this reason graphite is used as a dry lubricant in machines running at high temperature, where oil cannot be used as a lubricant. 11.7.3 Fullerenes Fullerenes are made by the heating of graphite in an electric arc in the presence of inert gases such as helium or argon. The sooty material formed by condensation of vapourised Cn small molecules consists of mainly C60 with smaller quantity of C70 and traces of fullerenes consisting of even number of carbon atoms up to 350 or above. Fullerenes are the only pure form of carbon because they have smooth structure without having ‘dangling’ bonds. Fullerenes are cage like molecules. C60 molecule has a shape like soccer ball and called Buckminsterfullerene (Fig. 11.5). It contains twenty six- membered rings and twelve five-membered rings. A six membered ring is fused with six or five membered rings but a five membered ring can only fuse with six membered rings. All the carbon atoms are equal and they undergo sp2 hybridisation. Each carbon atom forms three sigma bonds with other three carbon atoms. The remaining electron at each carbon is delocalised in molecular orbitals, which in turn give aromatic character to molecule. This ball shaped molecule has 60 vertices and each one is occupied by one carbon atom and it also contains both single and double bonds with C–C distances of 143.5 pm and 138.3 pm respectively. Spherical fullerenes are also called bucky balls in short.

⚠In the NEET syllabus; removed from current NCERT.

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. The ammonia molecule is trigonal pyramidal with the nitrogen atom at the apex. It has three bond pairs and one lone pair of electrons as shown in the structure. Ammonia gas is highly soluble in water. Its aqueous solution is weakly basic due to the formation of OH– ions. NH3(g) + H2O(l) ⇌ NH4+ (aq) + OH– (aq) … On a large scale, ammonia is manufactured by Haber’s process. N2(g) + 3H2(g) ⇌ 2NH3(g); ∆fH0 = – 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 × 10^5 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. The flow chart for the production of ammonia is shown in Fig. 7.1. Earlier, iron was used as a catalyst with molybdenum as a promoter. … 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)

⚠In the NEET syllabus; removed from current NCERT.

(i) Reactivity towards hydrogen: They all react with hydrogen to give hydrogen halides but affinity for hydrogen decreases from fluorine to iodine. Hydrogen halides dissolve in water to form hydrohalic acids. Some of the properties of hydrogen halides are given in Table 7.9. 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. Table 7.9: Properties of Hydrogen Halides (HF, HCl, HBr, HI) — Melting point/K: 190, 159, 185, 222; Boiling point/K: 293, 189, 206, 238; Bond length (H – X)/pm: 91.7, 127.4, 141.4, 160.9; ΔdissH/kJ mol–1: 574, 432, 363, 295; pKa: 3.2, –7.0, –9.5, –10.0.

⚠In the NEET syllabus; removed from current NCERT.

Ionisation enthalpy decreases down the group. It is due to increase in size. However, the elements of this group have lower ionisation enthalpy values compared to those of Group15 in the corresponding periods. This is due to the fact that Group 15 elements have extra stable half-filled p orbitals electronic configurations. Because of the compact nature of oxygen atom, it has less negative electron gain enthalpy than sulphur. However, from sulphur onwards the value again becomes less negative upto polonium. Next to fluorine, oxygen has the highest electronegativity value amongst the elements. Within the group, electronegativity decreases with an increase in atomic number. This implies that the metallic character increases from oxygen to polonium. Oxidation states and trends in chemical reactivity The elements of Group 16 exhibit a number of oxidation states (Table 7.6). The stability of -2 oxidation state decreases down the group. Polonium hardly shows –2 oxidation state. Since electronegativity of oxygen is very high, it shows only negative oxidation state as –2 except in the case of OF2 where its oxidation state is + 2. Other elements of the group exhibit + 2, + 4, + 6 oxidation states but + 4 and + 6 are more common. Anomalous behaviour of oxygen The anomalous behaviour of oxygen, like other members of p-block present in second period is due to its small size and high electronegativity. 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.

⚠In the NEET syllabus; removed from current NCERT.

All noble gases have general electronic configuration ns2np6 except helium which has 1s2 (Table 7.12). Many of the properties of noble gases including their inactive nature are ascribed to their closed shell structures. Due to stable electronic configuration these gases exhibit very high ionisation enthalpy. However, it decreases down the group with increase in atomic size. Atomic radii increase down the group with increase in atomic number. 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: (i) The noble gases except helium (1s2) have completely filled ns2np6 electronic configuration in their valence shell. (ii) They have high ionisation enthalpy and more positive electron gain enthalpy. … After this discovery, a number of xenon compounds mainly with most electronegative elements like fluorine and oxygen, have been synthesised. … The structures of the three xenon fluorides can be deduced from VSEPR and these are shown in Fig. 7.9. XeF2 and XeF 4 have linear and square planar structures respectively. XeF6 has seven electron pairs (6 bonding pairs and one lone pair) and would, thus, have a distorted octahedral structure as found experimentally in the gas phase. Fig. 7.9 The structures of (a) XeF2 (b) XeF4 (c) XeF6 (d) XeOF4 and (e) XeO3: (a) Linear (b) Square planar (c) Distorted octahedral (d) Square pyramidal (e) Pyramidal

⚠In the NEET syllabus; removed from current NCERT.

The absence of d orbitals in oxygen limits its covalency to four and in practice, rarely exceeds two. On the other hand, in case of other elements of the group, the valence shells can be expanded and covalence exceeds four. (i) Reactivity with hydrogen: All the elements of Group 16 form hydrides of the type H2E (E = O, S, Se, Te, Po). Some properties of hydrides are given in Table 7.7. Their acidic character increases from H2O to H2Te. The increase in acidic character can be explained in terms of decrease in bond enthalpy for the dissociation of H–E bond down the group. Owing to the decrease in enthalpy for the dissociation of H–E bond 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. Table 7.7: Properties of Hydrides of Group 16 Elements (H2O, H2S, H2Se, H2Te) — m.p/K: 273, 188, 208, 222; b.p/K: 373, 213, 232, 269; H–E distance/pm: 96, 134, 146, 169; HEH angle (°): 104, 92, 91, 90; ΔfH/kJ mol–1: –286, –20, 73, 100; ΔdissH (H–E)/kJ mol–1: 463, 347, 276, 238; Dissociation constant (aqueous solution, 298 K): 1.8×10–16, 1.3×10–7, 1.3×10–4, 2.3×10–3.

⚠In the NEET syllabus; removed from current NCERT.

7.22 Interhalogen Compounds When two different halogens react with each other, interhalogen compounds are formed. They can be assigned general compositions as 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 maximum number of atoms). Preparation The interhalogen compounds can be prepared by the direct combination or by the action of halogen on lower interhalogen compounds. The product formed depends upon some specific conditions, For example, Cl2 + F2 (equal volume) →(437 K) 2ClF; Cl2 + 3F2 (excess) →(573 K) 2ClF3; I2 + Cl2 (equimolar) → 2ICl; I2 + 3Cl2 (excess) → 2ICl3; Br2 + 3F2 (diluted with water) → 2BrF3; Br2 + 5F2 (excess) → 2BrF5 Properties Some properties of interhalogen compounds are given in Table 7.11. Table 7.11: Some Properties of Interhalogen Compounds (Type; Formula; Physical state and colour; Structure): XX′ — ClF colourless gas; BrF pale brown gas; IF detected spectroscopically (very unstable); BrCl gas; ICl ruby red solid (α-form), brown red solid (β-form); IBr black solid. XX′3 — ClF3 colourless gas, Bent T-shaped; BrF3 yellow green liquid, Bent T-shaped; IF3 yellow powder, Bent T-shaped (?); ICl3 orange solid, Bent T-shaped (?) (dimerises as Cl–bridged dimer (I2Cl6)). XX′5 — IF5 colourless gas but solid below 77 K, Square pyramidal; BrF5 colourless liquid, Square pyramidal; ClF5 colourless liquid, Square pyramidal. XX′7 — IF7 colourless gas, Pentagonal bipyramidal. These are all covalent molecules and are diamagnetic in nature. They are volatile solids or liquids at 298 K except ClF which is a gas. Their physical properties are intermediate between those of constituent halogens except that their m.p. and b.p. are a little higher than expected. Their chemical reactions can be compared with the individual halogens. 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), halate (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 (Example 7.19). The XX3 compounds have the bent ‘T’ shape, XX5 compounds square pyramidal and IF7 has pentagonal bipyramidal structures (Table 7.11). 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)