Why is #H_2# a nonpolar covalent bond?

In a polar covalent bond, one atom is substantially more electronegative than the other, and strongly polarizes electron density towards itself, i.e.

#""^(+delta)H-X^(delta-)#

Now while the bond is still covalent, the more electronegative atom polarizes electron density....and with the hydrogen halides, this often leads to acidic behaviour...

#HX(aq) + H_2O(l) rarr H_3O^+ +X^(-)#

And in the acid, charge polarization is so great that the #H-X# bond breaks.

But with the dihydrogen molecule, #H-H#, there can be no question that the participating ATOMS have EQUAL electronegativity...and as a consequence there is NO charge separation, and thus NO POLARITY... The same situation applies for other homonuclear diatomic molecules and bonds: #X-X#; #R_3C-CR_3; N-=N; R_3Si-SiR_3; R_3C-SiR_3#.

A hydrogen atom has a certain electronegativity (how much it pulls electrons to itself in a compound). However, #"H"_2# involves two identical atoms, each having an identical pull on the electrons being shared. Since each pull is equal and opposite, the electrons are pretty much distributed equally, meaning it is nonpolar.