The typical $\ce{Si-Si}$ single bond length in a silane is around 2.33 angstroms.  This is much longer than a typical $\ce{C-C}$ single bond (~1.53 Å) and helps explain why silicon-silicon single bonds are so much weaker than carbon-carbon single bonds (bond dissociation energy: ~53 kcal/mol for silicon silicon *vs.* ~83 kcal/mol for carbon-carbon single bonds).  When we try to form a silicon-silicon double or triple bond, the large separation between the two silicon atoms results in even **less effective p-orbital overlap and therefore even weaker pi bonds**.  The result is that while [disilenes][1] and [disylynes][2] are known, they are **extremely** reactive.  Only when bulky substituents are placed around the double or triple bonds can the molecules be isolated and characterized.          

  [1]: http://chemistry.stackexchange.com/questions/27839/why-does-si2h4-readily-polymerize-at-rtp-but-ethene-does-not-requires-catalyst/27854#27854
  [2]: http://chemistry.stackexchange.com/questions/16538/why-is-disilyne-bent/16543#16543