Hydrides and alykls become less stable down the column.
The M-H mean bond enthalpies decrease down the column and consequently the hydrides and alkyls become thermodynamically less stable and kinetically more reactive. Carbon forms a wide range of hydrides, Silicon forms primarily SiH4 and Si2H6 which are spontaneously inflamable. Higher silanes decompose readily to Si2H6. Silanes are strong reducing agents. The germanes GeH4, Ge2H6, and Ge3H8 are less flammable than SiH4 and are resistant to hydrolysis. SnH4 decomposes at 0°C to Sn and PbH4 is extremely unstable.
The much higher stability of Carbon Hydrides is the difference in electronegativity. Hydrogen has an electronegativity of 2.2, and so Carbon is the one element in the group that forms a compound with Hydrogen in the +1 oxidation state, rather than the -1, which is much less stable.
Now we have separated out Carbon by its electronegativity, but the electronegativities of Silicon, Germanium and Tin are nearly identical, so what is the reason for the stability change. They are all less stable because of the negative oxidation state of the hydride, but the main difference is size. As we go down the group, the size increases while the electronegativity remains the same. The charge density is smaller, and there is less of a pull on the hydride by the central atom. This is shown in the bond lengths. The bond length in Methane is about 109 pm, 146 pm for Silane, 152 pm for Germane, and 171 pm in Stannane. This indicates that the bonds are progressively weaker down the group.
PbH4 is also very unstable because of the 6s inert pair effect, which makes the +2 oxidation state much more stable than the +4.