Q.Select the member(s) of group 14 that

Group 14 (C, Si, Ge, Sn, Pb) share the outer configuration $ns^2np^2$, but — exactly as in Group 13 — inner-core complexity grows down the group (heavier members carry filled d, and for Pb, f subshells), disturbing otherwise-smooth trends. Covalent radius jumps sharply from C to Si but then increases only modestly through Ge, Sn, Pb, because the filled inner d/f orbitals of the heavier atoms shield poorly and limit further expansion — this is also why ionisation enthalpy falls steeply from C to Si (the "phenomenal decrease") but only gently after that, even rising slightly from Sn to Pb. Every Group 14 element shows a higher first ionisation enthalpy than the corresponding Group 13 element. The two common oxidation states are +4 and +2; because the combined first-four ionisation enthalpies are very large, +4-state compounds are generally covalent. Down the group, the inert-pair effect increasingly favours +2 over +4 in the order Ge < Sn < Pb, so Pb(II) is the stable state (Pb(IV) compounds are strong oxidisers) while C and Si are essentially +4-only — this explains why $PbCl_2$ reacts with $Cl_2$ to give the less-stable $PbCl_4$ (which itself decomposes readily on heating), and why $PbI_4$ cannot form at all (the Pb–I bond doesn't release enough energy to promote lead's inert $6s^2$ pair). This same inert-pair logic covers the classification of Group 14/13 oxides by acid–base character: $CO$, $B_2O_3$ (acidic in the boron sense) $SiO_2$, $CO_2$ are acidic; $Al_2O_3$, $PbO_2$, $Tl_2O_3$ are amphoteric or basic depending on the metal's position. The accessible d orbitals of Si, Ge, Sn and Pb (absent in carbon) let their halides expand coordination and hydrolyse — e.g. $SiF_6^{2-}$ forms readily because six small fluoride ions fit around silicon, whereas $SiCl_6^{2-}$ does not form because the larger chloride ions cannot be packed six-around-one without excessive steric crowding.