Question

test: covalent compounds
why is this lewis structure incorrect? choose all that apply.
:cl:
cl = si = cl
:cl:
there are too many bonds around si.
there should only be single bonds in this lewis structure.
the structure is missing a triple bond.
chlorine only has 6 electrons surrounding it.

Explanation:

• For silicon tetrachloride (\(SiCl_4\)), silicon has 4 valence electrons and chlorine has 7 valence electrons. The correct Lewis structure has single bonds (\(Si - Cl\)) as silicon forms 4 single bonds to complete its octet (by sharing electrons with 4 chlorine atoms). Chlorine needs one more electron to complete its octet (7 + 1 from single - bond = 8), and in the given structure, if there were double bonds (\(Si = Cl\)), chlorine would have 6 non - bonding electrons + 2 from the double bond = 8, but silicon would have more than 8 electrons (if double bonds are present, it's over - bonding for a typical octet - following structure in \(SiCl_4\)). Also, in \(SiCl_4\), single bonds are the correct bonding type as per the valence electron requirements (silicon 4, chlorine 1 per bond).
• Chlorine in the correct Lewis structure of \(SiCl_4\) has 6 non - bonding electrons and 2 bonding electrons (in a single bond) for a total of 8. In the given structure, if we assume the double - bond structure, chlorine would have 6 non - bonding electrons and 2 from the double bond (total 8), but the issue is with the bonding type (should be single bonds) and over - bonding of silicon (if double bonds are present, silicon would have more than 8 electrons in some interpretations, but mainly the incorrect bond type).

Answer:

• B. There should only be single bonds in this Lewis Structure.