I was wondering about this in class, drawing the structures of carbon dioxide and silicon dioxide.
Carbon and silicon are both in Group 4/14, but coming up with oxygen, one can only form a simple molecular structure while the other can form a giant covalent structure. I found no satisfactory in my chemistry textbook and online. What difference between carbon and silicon is this caused by?
Silicon is larger than Carbon, so Silicon's electron cloud is more diffused. Thus, if a hypothetical bond was formed between Silicon and Oxygen, the p - orbitals would not be able to come close enough to overlap enough to form a sufficiently strong bond. The bond would be too weak, and not enough energy would be released in this hypothetical bond formation, so this is energetically unfavourable.
Compare it to the analogical case $\ce{Nā”N}$ versus tetrahedral $\ce{P4}$ with 3 single bonds per atom.
Carbon forms in $\ce{CO2}$ or $\ce{CS2}$ 2 double bonds, while silicon in $\ce{SiO2}$ or $\ce{SiS2}$ 4 single bonds.
Generally, elements from the 2nd period (C, N, O) have a stronger ( C,N much stronger) tendency to form double/triple bonds than their counterparts in the 3rd period ( Si, P, S), mostly because of better overlapping of p orbitals at forming $\pi$ bonds.
If there is bad orbital overlapping, eventual $\pi$ bonds are weak and multiple single bonds are energetically preferred to 1 multiple bond. By other words, forming hypothetical $\ce{O=Si=O}$ would release much less energy than forming polymeric $\ce{(SiO2)x}$
