#"And polarity, i.e. charge separation, decisively influences......"# the properties of the water molecule.
You are probably aware that the water molecule, which is very low molecular weight, i.e. #18*g*mol^-1#, has a disproportionately high melting and boiling point. Compare (as a block) the normal boiling points of water, #OH_2, 100# #""^@C#, ammonia, #NH_3, -33.3# #""^@C#, and hydrogen fluoride, #HF, +19.5# #""^@C#, with the normal boiling points of the hydrides of the lower Group members: cf. #SH_2, -60# #""^@C#, phosphine, #PH_3, -87.7# #""^@C#, and hydrogen chloride, #HCl, -85.1# #""^@C#.
Ordinarily, we would expect that as the molecule becomes larger (as the central atoms become larger), the boiling points would increase due to increased dispersion forces (which depend primarily on the number of electrons, and thus on #Z,"the atomic number"#).
For the first row hydrides, the #""^(-delta)"X"-"H"^(delta+)# dipole is very pronounced, and this manifests in the disproportionately high normal boiling points of these first row hydrides. The difference can be attributed to hydrogen-bonding, which acts as a potent intermolecular force BETWEEN molecules, i.e. the dipoles tend to line up, and requires more energy to break. With the lower group hydrides, while dispersion forces tend to increase, the hydrogen-bonding interaction between molecules is diminished, because of the reduction in bond polarity.
And we can compare this to a scenario where hydrogen-bonding is not conceived to operate due to the less polar #"element"-"hydrogen"# bond in the hydrides of the #"Group 14"# #"elements."#
The #"normal boiling point"# of #"methane"#, #CH_4#, is #-164# #""^@C#; for #"silane"#, #SiH_4#, it is #-112# #""^@C#; and for #"germane"#, #GeH_4#, it is #-88# #""^@C#. The larger molecular gives rise to a more polarizable electron cloud, and thus higher boiling points.
