The central atom X uses $sp$ hybrid orbitals for bonding.

The molecule is linear.

The central atom X has no lone pairs of electrons.

The molecule must be polar if X and Y have different electronegativities.

1

2

3

4

Non-polar bond

Ionic bond

Polar covalent bond

Pure covalent bond

Both atoms involved in the bond contribute one electron each to the shared pair.

One atom donates both electrons of the shared pair.

The bond is formed between a Lewis acid and a Lewis base.

The bond is essentially identical to a regular covalent bond once formed.

The formation of a covalent bond always results in a net decrease in potential energy, leading to greater stability compared to isolated atoms.

Overlapping of atomic orbitals leads to increased electron density between the nuclei, contributing to bond formation.

The strength of a covalent bond is inversely proportional to the extent of orbital overlap.

The stability of a covalent bond arises from the balance between attractive forces between nuclei and electrons, and repulsive forces between nuclei and between electrons.

$4$

$5$

$6$

$7$

13, 5

12, 6

14, 4

11, 7

$s - s$ orbitals

$p - p$ orbitals

$s - p$ orbitals

$s - p - s$ orbitals

$3$

$4$

$6$

$5$

Ionic bond

Polar covalent bond

Non-polar covalent bond

None of the above

Transfer of electrons from one atom to another

Sharing of electrons between atoms

Electrostatic attraction between oppositely charged ions

Interaction between a metal and a non-metal