The world-class physical organic chemist Paul von Ragué Schleyer has pointed out that experimental H-C-H angles tend to be about 111-113° regardless of "hybridization" unless there is some other reason, such as molecular symmetry (e.g. methane, CH₄, or silane, SiH₄) or steric crowding, to have a different angle. (This information is taken from an invited talk I heard him give in about 1993; I'm sure he's published this, but I can't give a reference.)

Actually, VSEPR is likely to be more correct than hybridization theory. For example, EF₅ compounds exist for all Group 15 elements except nitrogen. But if you run a molecular-orbital calculation on NF₅, the molecular orbitals and bonding patterns are almost identical to those in PF₅. So why doesn't NF₅ exist? Five fluorine atoms don't fit around nitrogen (the axial N-F bonds are calculated to be 20 picometers longer than the equatorial bonds), but they fit quite well around phosphorus (axial and equatorial P-F bonds are about the same length). Again, see Kutzelnigg's paper.

Examine bond angles in NH₃ (ca. 108°) vs. PH₃ (ca. 95°) or NF₃ (ca. 103°) vs. PF₃ (ca. 97°). The three hydrogens or fluorines can only get so close to each other before they touch--and no closer. Around a big central atom, this will correspond to a smaller angle than around a small central atom.

Of course this begs the question of why NF₃ has a smaller bond angle than NH₃, to which the hand-waving answer is: electronegativity effects! The bonding atomic orbitals in NF₃ will have more p-character and therefore a smaller bond angle than those in NH₃.

Based on Kutzelnigg's paper and Schleyer's talk, hybridization is not likely to be important for bond angles at sulfur--or, for that matter, at oxygen! Of course, this is quite controversial and you're likely to find lots of chemists who don't agree with me... so read the literature for yourself.

For more references, see this previous MadSci answer and thi s one.