Sulfur tetrachloride, \(SF_4\), is an interesting molecule because it shows some of the weird concepts in chemical bonding, so we will explore it further here.
First, recall that fluorine generally forms single bonds in covalently bonded compounds, as fluorine atoms are one electron away from a full octet. It uses a single bonding pair to make up that one-electron deficiency. \(SF_4\) is no exception here. If the fluorine atoms were replaced with any other halogen, the general trend would remain the same.
Here is where formal charge is important: if we draw a single bond from sulfur to each of the four fluorine atoms, the formal charges of all four fluroine atoms are zero. As for the sulfur, however, we will show that calculation here:
$$F.C. = valence e^{-} - \left(\frac{1}{2} \times bonding e^{-} + lone pair e^{-}\right) = 6 - \left(\frac{1}{2} \times 8) = 6 - 4 = +2$$
The sulfur atom has a formal charge of \(+2\). Since the charge of the molecule is the sum of all formal charges, we have not made \(SF_4\), but \(SF_4^{2+}\)! How do we fix this?
To fix this, we must violate the Octet Rule, giving the central sulfur atom an expanded valence shell, room for more electrons to lower the formal charge of the sulfur atom. Forming double bonds with the fluorine is not the way to go, as now fluorine is also making irregular numbers of bonds. Thus we must add lone pairs to the central atom. But how many?
We must think of the formal charge equation as one we must solve for the number of lone pair electrons. The formal charge equation states
$$F.C. = valence e^{-} - \left(\frac{1}{2} \times bonding e^{-} + lone pair e^{-}\right)$$
We want the formal charge to be \(0\), and the number of bonding electrons is \(8\) (coming from the four single bonds). Sulfur atoms have six valence electrons by themselves. Plug all this into the equation:
$$F.C = 6 - (4 + lone pair e^{-})$$
It is obvious from solving this equation that we need two extra electrons, or one lone pair.
Thus the sulfur atom has five electron domains, four in the form of single bonds and one as a lone pair. The charge of the whole molecule is zero. In a similar manner, if we wanted a polyatomic anion with these same atoms, we could squeeze up to one more lone pair onto the sulfur to get \(SF_4^{2-}\), but \(SF_4\) is more important.