Final answer:
Electrical conductivity differences among group 14 elements like Sn, C, Si, and Ge are a result of the variation in atomic size and the presence of accessible d-orbitals in Sn, which influence the mobility of electrons and the type of bonding, making Sn a conductor and C an insulator, with Si and Ge being semiconductors. So, the correct option is C. Presence of d-orbitals in Sn.
Step-by-step explanation:
The question revolves around why elements in group 14 of the periodic table exhibit different conductive properties; specifically, why Sn (tin) is a conductor, C (carbon) is an insulator, and Si (silicon) and Ge (germanium) are semiconductors. The answer lies in the atomic and electronic structures of these elements. While carbon is a nonmetal with electrons tightly bound to the nucleus, forming covalent networks that do not allow free motion of electrons, tin, on the other hand, displays metallic properties and can lose electrons to form positive ions or delve into a covalent bond with a formal 4+ oxidation state.
Silicon and germanium are metalloids, which have intermediate properties between those of metals and nonmetals, allowing them to conduct electricity under certain conditions, hence their categorization as semiconductors. The difference in electrical conductance between group 14 elements can be attributed primarily to variation in atomic size and the resulting difference in the ability of electrons to move freely. Larger atoms with more diffuse electrons tend to display more metallic character, which explains why Sn can act as a conductor. Moreover, the presence of accessible d-orbitals in Sn (option C) allows it to display both metallic and covalent bonding characteristics. Therefore, the presence of d-orbitals in Sn is a significant factor.