The pattern you observe for main group elements is due to the number of valence electrons present in the s and p subshells of the outermost shell. The most you can do is either lose all those s and p electrons (giving a ‘cation’[1] charged exactly according to the old main group number — new group number minus 10 for groups 13 through 18) or you can fill these up to a total of eight.
For transition metals, a d subshell is in the process of being filled. Since the d subshell of a lower shell is between the s and p subshells of a higher shell according to the aufbau principle, we can imagine that the highest number of ‘usable electrons’ would be 18. This is not typically the case, however.
For the first few transition metals, you can remove as many electrons as are present in the s and d valence subshells (e.g. 4s and 3d). In the 3d series, manganese is the last one that can do this to give manganese(VII), in the 4d series ruthenium can reach ruthenium(VIII) and in the 5d series iridium can reach iridium(IX).[2] For the remaining elements of a certain transition metal period, the highest accessable oxidation state again decreases. You can find an overview of the maximum oxidation states known for certain elements on Wikipedia.
The only elements that again follow a somewhat more simple scheme are the zinc group (zinc, cadmium and mercury). Since these have a completely filled d subshell in ground state, they are reluctant to give away any d electrons; hence their highest observed charge is $2+$, as if they were in group 2. (Indeed, they behave like group 2 metals in a lot of other respects, too; except mercury which is slightly different again.)
When it comes to the highest possible negative charge or oxidation state, things are more complicated yet again. Since all transition metals are metals, they are not commonly found as anions or anionic particles although some are more likely to than others. For some of these — typically the iron group and the two following groups — this corresponds to $12 - \text{group number}$. This small trend is more of an exception than of a rule, though. Note, for example, how zinc, cadmium and mercury are all known as in the $\mathrm{-II}$ state which does not make sense immediately just looking at the respective shells and subshells.
Many of the actually observable oxidation states can be rationalised by doing a complex molecular orbital analysis of the corresponding compounds. However, for most of these there is no simple and immediately obvious pattern that one can use — with the notable exception of the small bit of pattern outlined in the third paragraph.
Notes:
[1]: A better term here would be oxidation state. For many elements the highest possible free cation is not known but a compound with a corresponding oxidation state is.
[2]: Again, none of these are observable as free cations in standard chemistry; rather they are in compounds. Hence I have stopped using the term charge and started using oxidation state.
