Covalent Bonds: Dot & Cross Diagrams

A covalent bond is a shared pair of electrons held between two atoms — usually two non-metals. Each atom gives ONE electron to the shared pair, and that pair is attracted to both nuclei, holding the atoms together.

A dot-and-cross diagram is just a picture of where the outer electrons go. Follow four rules every single time:

1. Outer shell only. Draw the electrons in the outer shell of each atom — never the inner shells. (So oxygen shows 6 outer electrons, not 8.)
2. Dots for one atom, crosses for the other. This is only to show which atom each electron came from — the electrons are really identical. For a molecule made of two different elements, use • for one element and × for the other.
3. Shared pair sits in the overlap. Where two atoms bond, draw the shared pair between them — usually one dot and one cross together in the overlap region.
4. No brackets, no charges. Covalent molecules are neutral. Square brackets and + / − charges belong to ionic diagrams, not covalent ones.

Counting outer electrons (from the group number): H = 1, C = 4, N = 5, O = 6, F/Cl = 7.

The full-shell check. For each atom, count the electrons you can see around it (its own unshared electrons + every shared pair it touches). It must reach a full outer shell: 2 for hydrogen, 8 for everything else.

A single bond is just one shared pair of electrons. The six molecules you must be able to draw are below. For each, the key is: how many bonds does the central atom need, and how many lone pairs are left over?

Molecule	Outer electrons	Bonding pairs (bonds)	Lone pairs
H₂	H = 1, H = 1	1 (between the two H)	none
Cl₂	Cl = 7, Cl = 7	1 (between the two Cl)	3 on each Cl
HCl	H = 1, Cl = 7	1 (between H and Cl)	3 on Cl
H₂O	O = 6, H = 1 (×2)	2 (one O–H each)	2 on O
NH₃	N = 5, H = 1 (×3)	3 (one N–H each)	1 on N
CH₄	C = 4, H = 1 (×4)	4 (one C–H each)	none

Quick rule for the central atom: (number of bonds) + (number of lone pairs × 2) = its outer-electron count. For oxygen in water: 2 bonds + 2 lone pairs × 2 = 2 + 4 = 6 ✓. For nitrogen in ammonia: 3 bonds + 1 lone pair × 2 = 5 ✓.

For Cl₂ and HCl, remember the three lone pairs on every chlorine: chlorine has 7 outer electrons, uses 1 in the bond, leaving 6 = 3 lone pairs. Forgetting these loses a mark almost every time.

When one shared pair isn't enough to fill both atoms' shells, atoms share more than one pair with the same partner.

• Double bond = 2 shared pairs = 4 shared electrons (drawn as 2 dots + 2 crosses in the overlap). Written O=O or C=O in line form.

Oxygen, O₂. Each O has 6 outer electrons and needs 2 more to reach 8. Sharing one pair would only give each O seven electrons — not enough. So the two oxygen atoms share two pairs (a double bond). Each O then has 2 lone pairs left over (4 electrons) plus the 4 shared electrons = 8 ✓.

Carbon dioxide, CO₂. Carbon has 4 outer electrons; each oxygen has 6. Carbon forms a double bond to each oxygen (O=C=O), using all 4 of its electrons.

• Carbon: all 4 electrons shared in two double bonds → 8 electrons, no lone pairs.
• Each oxygen: 4 electrons shared (one double bond) + 2 lone pairs (4 electrons) = 8 ✓.

Mark-scheme tip. If the question says the answer should show a double bond, you MUST draw two shared pairs (four electrons) between the atoms — drawing only one pair is the most common error and scores zero for that bond.

A triple bond = 3 shared pairs = 6 shared electrons (3 dots + 3 crosses in the overlap). Written N≡N.

Nitrogen, N₂. Each N has 5 outer electrons and needs 3 more to reach 8. A single or double bond isn't enough, so the two N atoms share three pairs (a triple bond). Each N then keeps 1 lone pair (2 electrons) plus the 6 shared electrons = 8 ✓.

Why N₂ matters. Three shared pairs make the triple bond very strong — it is one of the hardest covalent bonds to break, which is why nitrogen gas is so unreactive.

Bond-strength order. triple > double > single — the more shared pairs, the stronger the bond.

Whether you are drawing a diagram or spotting the error in one the examiner gives you, run this two-step check:

1. Outer shell only? No inner shells should be drawn. If you see "2,8" rings the diagram is wrong.
2. Full shell on every atom? Count around each atom (its own electrons + every shared pair it touches). H must reach 2; everything else must reach 8.

Common faults to look for in a "spot the error" question:

• A double bond drawn as a single bond (only 2 electrons between O and O, or O and C) — the atom is left with 7, not 8.
• Missing lone pairs (e.g. forgetting the 3 lone pairs on chlorine, or the 2 on each oxygen).
• Charges or brackets added by habit — wrong for covalent.
• Too many electrons drawn around hydrogen (H can only ever hold 2).

A neat way to present any covalent dot-and-cross diagram is: central atom in the middle, outer atoms around it, lone pairs on the outside of the central atom, and the shared pairs clearly in each overlap.