Hey mrlucky0 and refid
I must say, refid, that i do not agree with you - methylgroups in amines, and ammonium-ions for that matter, are electrondonating and not electronwithdrawing.
The basicity trends of amines are most easily explained by looking at the acidity of ammonium-ions, which (in solution) follows the trend given below:
NH4+ > (CH3)3NH+ > CH3NH3+ > (CH3)2NH2+
That is NH4+-ions are the most acidic (least basic) ones, whereas (CH3)2NH2+-ions are the least acidic (most basic) ones.
This trend can be explained by two conflicting effects - polarizability and solvation.
If you look at gas phase acidities the trend is as follows:
NH4+ > CH3NH3+ > (CH3)2NH2+ > (CH3)3NH+
Which is what we would normally expect since more methyl-groups donates more electrons spreading out the positive charge and therefore stabilizing the ammonium-ion (making it less acidic, and the conjugate base more basic).
On the other hand solvation effects leads to the following trend:
(CH3)3NH+ > (CH3)2NH2+ > CH3NH3+ > NH4+
This trend can be explained by the size of the ammonium-ion - the larger, the least well solvated, therefore least stable and hence more acidic (less basic).
A combination of these two trends gives the acidity (and therefore basicity) trend of ammonium-ions (amines) in solution.
I hope this answers your question!