Why don't elements having the same valency as Carbon bond so readily?
It's not that they don't bond readily, it's that they can't form as stable of networks as carbon. Carbon has many different allotropes, while silicon has only a few. As you can see, carbon forms uniform hexagonal sheets with massive $\pi$ systems as in graphene and many of its other allotropes, while silicene is incapable of remaining flat and is bumpy. This is a result of the bond distances and p-orbital overlap.

$$
\begin{array}{|c|c|}\hline
\text{Bond} &amp;\ \Delta H ^\circ_f \text{ (kcal mol}^{-1}) &amp;\ \text{Bond Length (pm)} \\
\hline
\ce{C-C} &amp;\ 82.70 &amp;\ 154 \\
\hline
\ce{Si-Si} &amp;\ 53.06 &amp;\ 233 \\
\hline
\ce{Ge-Ge} &amp;\ 44.93 &amp;\ 241 ^{[1]} \\
\hline
\end{array}
$$

As you can see, as the atoms get larger, the bond length increases. $\pi$ bonding is completely reliant on the overlap of adjacent p-orbitals, so the farther away these orbitals are, the weaker the bonds.



Bond dissociation enthalpy also increases with atom size. This leads to an overall decrease in the stability of larger silicon or germanium compounds. Tin and lead are metals, which have much different chemistry altogether on account of their electronegativity and occupied d-orbitals, in addition to their size.




${[1]}$ Bond Enthalpy Table