Why is negative charge more stable on sulphur than on oxygen?
There is more to it than meets the electronegative eye.
Get into the habit of transferring attributes into energy.  In this case the relevant energies are the electron affinity of the base money ($\ce{RO}$ vs $\ce{RS}$) versus the bond energy of the moiety to hydrogen.  A strong anionic Bronsted-Lowry base will have a low value for the first versus a high value for the latter.
Surprise!  Sulfur has a higher electron affinity than oxygen!  Add the fact that the oxygen-hydrogen bond is stronger than the hydrogen-sulfur bond and by both measures, $\ce{RO^-}$ appears as the stronger base.
The higher electron affinity of sulfur and even selenium, compared with oxygen, has another curious effect.  We know about the aromaticity of certain ring ketones in which the electronegative oxygen atom imparts a positive charge to a three- or seven-membered carbon ring.  It turns out that despite what an elctronegativity table might suggest, the higher electron affinity of sulfur and selenium versus oxygen causes this characteristic to carry over to the corresponding $\ce{C=S}$ and $\ce{C=Se}$ compounds.