How to Draw Dot-and-Cross Diagrams

Dot-and-cross diagrams are one of the first tools you meet for showing chemical bonding. They show the outer-shell electrons of each atom, using dots for one atom’s electrons and crosses for the other’s, so you can see exactly which electrons are shared or transferred. They appear in almost every chemistry exam at this level. This guide gives you a reliable step-by-step method, then works through examples of increasing difficulty.

What the symbols mean

• A dot (•) represents an electron from one atom.
• A cross (×) represents an electron from another atom.
• Electrons are identical in reality — the dots and crosses only show where each electron came from.
• Usually, only the outer shell electrons are drawn. Some courses ask for all shells; check your syllabus.
• Circles around atoms represent outer shells. Where circles overlap in covalent diagrams, the electrons in the overlap are shared.

Before you start: find the outer electrons

For main-group elements, the number of outer electrons equals the last digit of the group number (or the group number in the 1–8 system):

Element	Group	Outer electrons
H	1	1
C	14	4
N	15	5
O	16	6
F, Cl	17	7
Na	1	1
Mg	2	2

You can also read this from the electron configuration: chlorine is 2,8,7, so it has seven outer electrons (see the electron configuration tool).

Method for covalent molecules

1. Identify the atoms and count their outer electrons.
2. Decide the central atom (usually the one that forms the most bonds; never hydrogen).
3. Work out how many bonds each atom needs to complete its outer shell (8, or 2 for hydrogen).
4. Draw overlapping circles — one per atom — with each bonded pair of atoms overlapping.
5. Place the shared pairs in the overlaps: one electron from each atom per bond (one dot + one cross).
6. Add the remaining outer electrons as lone pairs in the non-overlapping part of each circle.
7. Check: count the electrons in each atom’s circle (8 for most atoms, 2 for H) and make sure the total number of dots and crosses equals the total number of outer electrons you started with.

Worked example 1: hydrogen, H₂

• Each H has 1 electron.
• Draw two overlapping circles. Place one dot and one cross in the overlap.
• Check: each circle contains 2 electrons. ✓ Total electrons: 2. ✓

That shared pair is a single covalent bond, H–H.

Worked example 2: hydrogen chloride, HCl

• H has 1 outer electron (cross); Cl has 7 (dots).
• One bond needed. In the overlap: one cross from H, one dot from Cl.
• Chlorine’s remaining 6 dots go around the outside of its circle as three lone pairs.
• Check: H circle has 2; Cl circle has 8. Total: 1 + 7 = 8 electrons drawn. ✓

Worked example 3: water, H₂O

• O has 6 outer electrons (dots); each H has 1 (crosses).
• Oxygen needs 2 more electrons, so it forms two bonds, one with each H.
• Draw O in the middle, overlapping with an H circle on each side. In each overlap, put one dot and one cross.
• Oxygen’s remaining 4 dots form two lone pairs.
• Check: O circle has 8; each H circle has 2. Total: 6 + 1 + 1 = 8. ✓

(A dot-and-cross diagram doesn’t show the true shape. Water is actually bent at about 104.5° — see VSEPR and molecular geometry.)

Worked example 4: methane, CH₄

• C has 4 outer electrons (dots); each H has 1 (cross).
• Carbon forms four bonds. Draw C in the centre with four overlapping H circles.
• Each overlap: one dot + one cross.
• Carbon has no lone pairs.
• Check: C has 8; each H has 2. Total 4 + 4 = 8. ✓

Worked example 5: ammonia, NH₃

• N has 5 (dots); three H with 1 each (crosses).
• Three bonds, each with one dot and one cross.
• The remaining 2 dots on N form one lone pair.
• Check: N has 8; each H has 2. Total 5 + 3 = 8. ✓

Worked example 6: oxygen, O₂ (a double bond)

• Each O has 6 outer electrons: one atom as dots, one as crosses.
• Each O needs 2 more, so they share two pairs — a double bond.
• In the overlap, place four electrons: two dots and two crosses (conventionally as two dot-cross pairs).
• Each O has 4 remaining electrons as two lone pairs.
• Check: each circle contains 4 shared + 4 lone = 8. Total: 12. ✓

Worked example 7: carbon dioxide, CO₂

• C has 4 (crosses); each O has 6 (dots).
• Carbon needs 4 more electrons; each oxygen needs 2. So carbon forms a double bond with each oxygen: O=C=O.
• Each overlap contains 2 crosses (from C) and 2 dots (from O).
• Each O has two lone pairs; C has none.
• Check: C has 8 (4 + 4 shared); each O has 8 (4 shared + 4 lone). Total: 4 + 6 + 6 = 16. ✓

Worked example 8: nitrogen, N₂ (a triple bond)

• Each N has 5. Each needs 3 more, so they share three pairs — a triple bond.
• The overlap contains six electrons: three dots and three crosses.
• Each N has one lone pair.
• Check: each circle has 6 shared + 2 lone = 8. Total 10. ✓

Method for ionic compounds

Ionic diagrams look different because electrons are transferred, not shared. The circles don’t overlap.

1. Identify the ions and their charges (metal loses electrons, non-metal gains) — see ionic bonding explained.
2. Draw each ion separately, with its outer shell after the transfer.
3. Show the transferred electron(s) using the symbol of the atom they came from.
4. Put each ion in square brackets with its charge at the top right.
5. Check the charges balance.

Worked example 9: sodium chloride, NaCl

• Na (2,8,1) loses its 1 outer electron; Cl (2,8,7) gains it.
• Sodium ion: draw [Na]⁺ with an empty outer shell, or with its full second shell of 8 dots if your course asks for it.
• Chloride ion: draw [Cl]⁻ with 8 outer electrons: 7 dots (chlorine’s own) and 1 cross (from sodium).
• Charges: +1 and −1 balance. ✓

Worked example 10: magnesium oxide, MgO

• Mg (2,8,2) loses 2 electrons → Mg²⁺. O (2,6) gains 2 → O²⁻.
• Draw [Mg]²⁺ and [O]²⁻, the oxide ion with 6 dots + 2 crosses.
• Charges balance. ✓

Worked example 11: magnesium chloride, MgCl₂

• Mg loses 2 electrons, but each Cl accepts only 1. So you need two chloride ions.
• Draw [Cl]⁻ [Mg]²⁺ [Cl]⁻, each chloride with 7 dots and 1 cross.
• Charges: +2 and 2 × (−1). ✓ (Alternatively, write 2[Cl]⁻.)

Common mistakes

• Drawing ionic compounds with overlapping circles. Overlaps mean sharing — use separate ions in brackets.
• Forgetting the charges or brackets on ions.
• Missing lone pairs, especially on O, N and halogens. Always do the final electron count.
• Wrong number of shared electrons in double and triple bonds: a double bond has 4 electrons in the overlap; a triple has 6.
• Using dots for everything. The point of dots and crosses is to show origin.
• Giving hydrogen 8 electrons. Hydrogen’s shell holds only 2.
• Thinking the diagram shows shape. It doesn’t; use VSEPR for shapes.

Quick practice

Draw dot-and-cross diagrams for: (a) F₂, (b) CCl₄, (c) HCN, (d) CaO, (e) Na₂O.

Answers (in words): (a) one shared pair, three lone pairs on each F. (b) C in the centre, four single bonds, three lone pairs on each Cl. (c) H–C single bond, C≡N triple bond (6 shared electrons), one lone pair on N. (d) [Ca]²⁺ [O]²⁻ with 6 dots + 2 crosses on O. (e) two [Na]⁺ and one [O]²⁻ with 6 dots + 2 crosses.

Key takeaways

• Dots and crosses show where electrons came from; draw only outer shells unless told otherwise.
• Covalent: overlapping circles; each bond = one pair in the overlap; add lone pairs; check 8 (or 2 for H).
• Ionic: separate ions in square brackets with charges; transferred electrons shown with the donor’s symbol.
• Always do a final electron count and charge check.

For the next step, try Lewis dot structures and covalent bonding explained.