Question

2. Describe polar covalent bonds using water as an example.

Answer 1

A polar covalent bond is a type of chemical bond in which two atoms share electrons unequally because one atom is more electronegative than the other. The more electronegative atom pulls the shared electron density closer to itself, acquiring a partial negative charge (δ–), while the less electronegative atom becomes partially positive (δ+). Each polar bond has a bond dipole (direction from δ+ to δ–); whether a whole molecule is polar depends on the vector sum of its bond dipoles.

Using water (H2O) as an example:
- Electronegativity: Oxygen is much more electronegative than hydrogen (O ≈ 3.5, H ≈ 2.1 on the Pauling scale). The O–H bonds are therefore polar covalent—oxygen pulls the shared electrons toward itself, so O is partially negative (δ–) and each H is partially positive (δ+).
- Lewis structure and shape: Oxygen has two lone pairs and forms two O–H bonds. The electron-pair geometry is tetrahedral, but with two lone pairs the molecular shape is bent (about 104.5°).
- Net molecular dipole: Because the molecule is bent, the two O–H bond dipoles do not cancel; their vector sum points toward the oxygen atom, giving water an overall dipole moment (water is a polar molecule).
- Consequences: The polarity of the O–H bonds and the molecular dipole allow strong intermolecular attractions (hydrogen bonds) between the δ+ hydrogens of one water molecule and the δ– oxygen of another. This explains many of water’s unusual properties: relatively high boiling point and melting point, high surface tension, high specific heat capacity, and excellent ability to dissolve other polar or ionic substances.

In short: in water, the unequal sharing of electrons in O–H bonds (polar covalent bonds) plus the bent shape produce a molecule with a permanent dipole and strong intermolecular hydrogen bonding, which underlies water’s characteristic physical properties.