Stabilization of the (+2) oxidation state relative to the (+4) oxidation state down the group.
The halides, oxides, and sulfides of the M2+ ions become more stable on descending the group. For example SiCl4, SiBr4 and SiI4 are all stable. PbCl4 decomposes at 105°C and PbI4 does not exist. Similarly the ease of oxidation of the M2+ halides decreases down the column.
PbCl2 may be converted to PbCl4 by heating in a stream of chlorine. Similarly PbO2 is an oxidizing agent, whereas SnO2, GeO2, and SiO2 are not. The compounds in the lower oxidation state are in general more ionic, less likely to form molecular structures, the halides are less readily hydrolized and the oxides are less acidic.Carbon can form both the +2 and +4 oxidation states due to the uniqueness principle and kinetic inertness, and its position on the Frost Diagram doesn't relate as much to the other elements of the group. The trend really starts with Silicon, which does not even form the +2 oxidation state, and has a very stable +4 oxidation state, to Lead which is very unstable in the +4 state. The major difference going down the group is the d orbital shielding. The shielding causes the valence electrons to be harder to remove. And in the case of Lead the 6s inert pair effect shows up and further shields 2 of the electrons so that they are considered part of the core.