p-Block Groups 13 to 18

⚠In the NEET syllabus; removed from current NCERT.

In p-block elements the last electron enters the outermost p orbital. As we know that the number of p orbitals is three and, therefore, the maximum number of electrons that can be accommodated in a set of p orbitals is six. Consequently there are six groups of p–block elements in the periodic table numbering from 13 to 18. Boron, carbon, nitrogen, oxygen, fluorine and helium head the groups. Their valence shell electronic configuration is ns2np1-6(except for He). The inner core of the electronic configuration may, however, differ. The difference in inner core of elements greatly influences their physical properties (such as atomic and ionic radii, ionisation enthalpy, etc.) as well as chemical properties. Consequently, a lot of variation in properties of elements in a group of p-block is observed. The maximum oxidation state shown by a p-block element is equal to the total number of valence electrons (i.e., the sum of the s- and p-electrons). Clearly, the number of possible oxidation states increases towards the right of the periodic table. In addition to this so called group oxidation state, p-block elements may show other oxidation states which normally, but not necessarily, differ from the total number of valence electrons by unit of two. The important oxidation states exhibited by p-block elements are shown in Table 11.1. In boron, carbon and nitrogen families the group oxidation state is the most stable state for the lighter elements in the group. However, the oxidation state two unit less than the group oxidation state becomes progressively more stable for the heavier elements in each group. The occurrence of oxidation states two unit less than the group oxidation states are sometime attributed to the ‘inert pair effect’.

⚠In the NEET syllabus; removed from current NCERT.

11.1 GROUP 13 ELEMENTS: THE BORON FAMILY This group elements show a wide variation in properties. Boron is a typical non-metal, aluminium is a metal but shows many chemical similarities to boron, and gallium, indium, thallium and nihonium are almost exclusively metallic in character. … 11.1.1 Electronic Configuration The outer electronic configuration of these elements is ns2np1. A close look at the electronic configuration suggests that while boron and aluminium have noble gas core, gallium and indium have noble gas plus 10 d-electrons, and thallium has noble gas plus 14 f- electrons plus 10 d-electron cores. … 11.1.2 Atomic Radii On moving down the group, for each successive member one extra shell of electrons is added and, therefore, atomic radius is expected to increase. However, a deviation can be seen. Atomic radius of Ga is less than that of Al. This can be understood from the variation in the inner core of the electronic configuration. The presence of additional 10 d-electrons offer only poor screening effect (Unit 2) for the outer electrons from the increased nuclear charge in gallium. Consequently, the atomic radius of gallium (135 pm) is less than that of aluminium (143 pm). 11.1.3 Ionization Enthalpy The ionisation enthalpy values as expected from the general trends do not decrease smoothly down the group. The decrease from B to Al is associated with increase in size. The observed discontinuity in the ionisation enthalpy values between Al and Ga, and between In and Tl are due to inability of d- and f-electrons ,which have low screening effect, to compensate the increase in nuclear charge. The order of ionisation enthalpies, as expected, is ∆iH1 <∆iH2 <∆iH3. The sum of the first three ionisation enthalpies for each of the elements is very high. Effect of this will be apparent when you study their chemical properties. 11.1.4 Electronegativity Down the group, electronegativity first decreases from B to Al and then increases marginally (Table 11.2). This is because of the discrepancies in atomic size of the elements. Table 11.2 (B, Al, Ga, In, Tl): Ionization enthalpy ∆iH1 (kJ mol–1) 801, 577, 579, 558, 589; Electronegativity (Pauling scale) 2.0, 1.5, 1.6, 1.7, 1.8; Melting point / K 2453, 933, 303, 430, 576.

⚠In the NEET syllabus; removed from current NCERT.

11.5.1 Electronic Configuration The valence shell electronic configuration of these elements is ns2np2. The inner core of the electronic configuration of elements in this group also differs. 11.5.2 Covalent Radius There is a considerable increase in covalent radius from C to Si, thereafter from Si to Pb a small increase in radius is observed. This is due to the presence of completely filled d and f orbitals in heavier members. 11.5.3 Ionization Enthalpy The first ionization enthalpy of group 14 members is higher than the corresponding members of group 13. The influence of inner core electrons is visible here also. In general the ionisation enthalpy decreases down the group. Small decrease in ∆iH from Si to Ge to Sn and slight increase in ∆iH from Sn to Pb is the consequence of poor shielding effect of intervening d and f orbitals and increase in size of the atom. 11.5.4 Electronegativity Due to small size, the elements of this group are slightly more electronegative than group 13 elements. The electronegativity values for elements from Si to Pb are almost the same. 11.5.5 Physical Properties All members of group14 are solids. Carbon and silicon are non-metals, germanium is a metalloid, whereas tin and lead are soft metals with low melting points. Melting points and boiling points of group 14 elements are much higher than those of corresponding elements of group 13. 11.5.6 Chemical Properties Oxidation states and trends in chemical reactivity The group 14 elements have four electrons in outermost shell. The common oxidation states exhibited by these elements are +4 and +2. Carbon also exhibits negative oxidation states. Since the sum of the first four ionization enthalpies is very high, compounds in +4 oxidation state are generally covalent in nature. In heavier members the tendency to show +2 oxidation state increases in the sequence Ge<Sn<Pb. It is due to the inability of ns2 electrons of valence shell to participate in bonding. The relative stabilities of these two oxidation states vary down the group. Carbon and silicon mostly show +4 oxidation state. Germanium forms stable compounds in +4 state and only few compounds in +2 state. Tin forms compounds in both oxidation states (Sn in +2 state is a reducing agent). Lead compounds in +2 state are stable and in +4 state are strong oxidising agents.

⚠In the NEET syllabus; removed from current NCERT.

Group 15 includes nitrogen, phosphorus, arsenic, antimony, bismuth and moscovium. As we go down the group, there is a shift from non-metallic to metallic through metalloidic character. Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids, bismuth and moscovium are typical metals. … 7.1.2 Electronic Configuration The valence shell electronic configuration of these elements is ns2np3. The s orbital in these elements is completely filled and p orbitals are half-filled, making their electronic configuration extra stable. 7.1.3 Atomic and Ionic Radii Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and/or f orbitals in heavier members. 7.1.4 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. The order of successive ionisation enthalpies, as expected is ∆iH1 < ∆iH2 < ∆iH3 (Table 7.1). 7.1.5 Electronegativity The electronegativity value, in general, decreases down the group with increasing atomic size. However, amongst the heavier elements, the difference is not that much pronounced. 7.1.6 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 upto arsenic and then decreases upto bismuth. Except nitrogen, all the elements show allotropy. Table 7.1 (N, P, As, Sb, Bi): Ionisation enthalpy I (kJ mol–1) 1402, 1012, 947, 834, 703; Electronegativity 3.0, 2.1, 2.0, 1.9, 1.9.

⚠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 these elements have seven electrons in their outermost shell (ns2np5) which is one electron short of the next noble gas. 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. 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. 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. Table 7.8 (F, Cl, Br, I): Electronegativity 4, 3.2, 3.0, 2.7; Bond dissociation enthalpy /(kJ mol–1) 158.8, 242.6, 192.8, 151.1. … They have very high electronegativity. The electronegativity decreases down the group. Fluorine is the most electronegative element in the periodic table. … One curious anomaly we notice from Table 7.8 is the smaller enthalpy of dissociation of F2 compared to that of Cl2 whereas X-X bond dissociation enthalpies from chlorine onwards show the expected trend: 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. Oxidation states and trends in chemical reactivity All the halogens exhibit –1 oxidation state. However, chlorine, bromine and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also as explained below: … 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.

⚠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.

Chlorine can be prepared by heating manganese dioxide with concentrated hydrochloric acid, MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O. A mixture of common salt and concentrated H2SO4 is used in place of HCl: 4NaCl + MnO2 + 4H2SO4 → MnCl2 + 4NaHSO4 + 2H2O + Cl2.