[AdiDex]

I recently read that hybridization is just like a operation which is done to get knowledge about the shape of molecule . It has nothing to do with reality

("all bonding theories, orbital hybridization is a model, and should not be taken to be a real phenomenon." Alan G.Sharpe Inorganic chemistry page no. 100 )

I've seen a illustration of CH4 molecule , actually how carbon is bonded to 4 Hydrogen atoms  .[shown in attached picture]

But I am confused about how it will look for H₂O molecule ?? as there are 2 lone pairs , say one in s orbital and other one p orbital .
Then how all overlappings will take place ?? As s orbital will not participate in overlaping as it is already filled .

Is it ok to say that two lone pairs are in SP³ hybridised obritals ??
Is it all about where is the highest electron density (lone pairs) ??

Is it true I am confused because of my previous knowledge of bonding ("overlapping")??

[AWK]

Each orbital, hybridized or not has room for two electrons.
I, doing X-ray crystallography, in which X-ray are difracted on electrons, very often can see these electrons (not directly - after calculations), as electron density about 0.6 Angstroem (60 pm) from oxygen in two positions.

[Enthalpy]

[In H₂O] s orbital will not participate in overlaping as it is already filled.

But in CH₄ it does.

As soon as the nuclei are close enough (say, as compared with the radii of atomic orbitals) overlapping takes place. It creates bonding and antibonding molecular orbitals; already full atomic orbitals can imply that the electrons will fill the antibonding molecular orbital too, and the net effect may be energetically defavourable - but it can happen if, for instance, other electrons keep the nuclei together.

One detail more: your book shows slim 2p orbitals. Don't be fooled. They are very thick.