Lewis structures of diatomic molecules – single bonds
A diatomic molecule contains two atoms only. In this topic we focus on
species where the two atoms are joined by a single covalent bond, written
as X–Y. Examples include homonuclear molecules like
H–H, F–F, Cl–Cl, and heteronuclear
molecules such as H–F, H–Cl, Cl–F,
I–Cl, and I–Br.
General electron-counting pattern
To draw a Lewis structure for a neutral diatomic molecule with a single bond:
- Count the total number of valence electrons for both atoms.
- Place one bond between the atoms (this uses \(2\) electrons).
-
Distribute the remaining electrons as lone pairs so that each atom
(except hydrogen) has an octet of electrons.
For molecules like \(\mathrm{H_2}\) or \(\mathrm{HF}\) the single bond completes
the duet on hydrogen, while the remaining electrons appear as lone pairs on the
non-hydrogen atom.
Lone pairs and bond pairs
In all these single-bond diatomics:
- There is one bonding pair (the shared pair in the X–Y bond).
-
The remaining valence electrons form lone pairs around each atom so
that non-hydrogen atoms reach eight electrons in their valence shell.
Examples:
Polarity: nonpolar vs polar bonds
If the two atoms are the same (\(\mathrm{H_2}\), \(\mathrm{F_2}\),
\(\mathrm{Cl_2}\), \(\mathrm{Br_2}\), \(\mathrm{I_2}\)), the bonding pair is
shared evenly and the bond is nonpolar.
If the atoms are different (\(\mathrm{HF}\), \(\mathrm{HCl}\),
\(\mathrm{ClF}\), \(\mathrm{ICl}\), \(\mathrm{IBr}\)), the more electronegative
atom attracts the bonding pair more strongly, giving a polar covalent
bond with partial charges \(\delta^+\) and \(\delta^-\).
Typical examples in this tool
The calculator follows exactly this counting logic and then plots the single
bond and lone pairs in a schematic circle-and-dot diagram for each molecule.
