How to draw dot and cross diagrams for GCSE chemistry

Dot and cross diagrams show how electrons are arranged in the bonds between atoms. They are one of the most commonly tested drawing skills in GCSE Chemistry, and once you learn the method you can apply it to any compound the exam throws at you.

This guide walks through what dot and cross diagrams actually represent, then gives you a clear step-by-step process for both ionic and covalent compounds. There are worked examples for every molecule you are likely to meet, plus the mistakes that cost students marks every year.

A dot and cross diagram is a model that shows the electrons in the outer shells of bonding atoms. You draw one atom's electrons as dots and the other atom's electrons as crosses. This makes it easy to see which electrons came from which atom.

For ionic compounds, the diagram shows electrons being transferred from one atom to another. The resulting ions are drawn in square brackets with their charges written outside.

For covalent compounds, the diagram shows electrons being shared between atoms. The shared pairs sit in the overlap between the two outer shells.

Dot and cross diagrams only show outer-shell electrons. You do not need to draw the inner shells or the nucleus in detail. Some mark schemes accept a simple circle for each atom with the element symbol in the centre.

Ionic bonding happens between metals and non-metals. The metal atom transfers its outer-shell electrons to the non-metal atom so that both end up with full outer shells. Follow these steps every time.

1. Write down the electronic configuration of each atom. Work out how many outer-shell electrons each one has.

2. Decide how many electrons need to be transferred. The metal needs to lose its outer electrons; the non-metal needs to gain enough to fill its outer shell.

3. Draw the metal atom with its outer electrons shown as dots (or crosses). Draw the non-metal atom with its outer electrons shown as the opposite symbol.

4. Show the electron(s) moving from the metal to the non-metal. You can use an arrow to indicate the transfer.

5. Draw the resulting ions. Put square brackets around each ion and write the charge outside the bracket – for example [Na]⁺ and [Cl]⁻.

6. Check that both ions now have full outer shells (usually eight electrons, except for atoms that fill a shell of two like lithium or beryllium losing all outer electrons).

Sodium has the electronic configuration 2,8,1, so it has one outer-shell electron. Chlorine has the configuration 2,8,7, so it has seven outer-shell electrons and needs one more.

Draw sodium's single outer electron as a dot. Draw chlorine's seven outer electrons as crosses. Transfer sodium's dot to chlorine, giving chlorine eight outer electrons.

The sodium ion is drawn in square brackets with a + charge: [Na]⁺. It now shows the full second shell (eight electrons) beneath the empty outer shell. The chloride ion is drawn in square brackets with a – charge: [Cl]⁻. It now has eight electrons in its outer shell – seven original crosses plus the one transferred dot.

Magnesium is 2,8,2, with two outer electrons. Oxygen is 2,6, with six outer electrons. Magnesium transfers both of its outer electrons to oxygen.

Draw magnesium with two dots in its outer shell. Draw oxygen with six crosses. Transfer both dots to oxygen, giving it a full shell of eight.

The magnesium ion is [Mg]²⁺ and the oxide ion is [O]²⁻. Notice the 2+ and 2– charges because two electrons have been transferred.

Magnesium has two outer electrons, and each chlorine atom needs one electron to complete its outer shell. This means one magnesium atom bonds with two chlorine atoms.

Draw the magnesium atom in the centre with two dots. Draw one chlorine atom on each side, each with seven crosses. Transfer one dot to each chlorine atom.

The result is one [Mg]²⁺ ion and two [Cl]⁻ ions. Each chlorine now has eight outer electrons, and magnesium has lost its entire outer shell to reveal the full shell below.

Covalent bonding happens between non-metal atoms. Instead of transferring electrons, the atoms share pairs of electrons. Each shared pair counts as one covalent bond.

1. Write down the electronic configuration of each atom and count the outer-shell electrons.

2. Work out how many more electrons each atom needs to achieve a full outer shell. This tells you how many bonds each atom will form.

3. Draw the outer shells of the atoms overlapping where bonds form. One atom's electrons should be dots, the other's should be crosses.

4. Place one electron from each atom into the overlap for each bond. A single bond has one shared pair (two electrons in the overlap). A double bond has two shared pairs (four electrons). A triple bond has three shared pairs (six electrons).

5. Fill in any remaining lone pairs (non-bonding pairs) on each atom.

6. Check that every atom has a full outer shell – eight electrons for most atoms, or two for hydrogen.

Oxygen has six outer electrons and needs two more, so it forms two bonds. Each hydrogen has one outer electron and needs one more, so each forms one bond.

Draw oxygen in the centre with six crosses. Overlap it with two hydrogen atoms, each contributing one dot. Place one dot and one cross into each overlap to form two single bonds.

Oxygen ends up with eight electrons in its outer shell – two bonding pairs shared with hydrogen and two lone pairs. Each hydrogen has two electrons (one dot, one cross) and a full outer shell.

Carbon has four outer electrons and needs four more, so it forms four bonds. Each oxygen has six outer electrons and needs two, so each forms two bonds. Carbon forms a double bond with each oxygen atom.

Draw carbon in the centre with four dots. Draw one oxygen on each side with six crosses. In each overlap, place two dots from carbon and two crosses from oxygen to form a double bond (four electrons in total per overlap).

Carbon now has eight electrons around it (four shared pairs). Each oxygen also has eight (two shared pairs plus two lone pairs).

Carbon has four outer electrons and needs four more, so it forms four single bonds. Each hydrogen has one outer electron and forms one bond.

Draw carbon in the centre with four dots, arranged evenly around the atom. Overlap four hydrogen atoms, each contributing one cross. Each overlap contains one dot and one cross – a single shared pair.

Carbon ends up surrounded by eight electrons (four bonding pairs). Each hydrogen has two electrons and a full shell.

Each nitrogen atom has five outer electrons and needs three more to fill its outer shell, so each forms three bonds. The two nitrogen atoms form a triple bond – three shared pairs of electrons.

Draw one nitrogen with five dots and the other with five crosses. In the overlap, place three dots and three crosses (six electrons in total). Each nitrogen has one lone pair sitting outside the overlap.

Both nitrogen atoms end up with eight electrons: Three bonding pairs in the triple bond and one lone pair each.

Forgetting the square brackets and charges on ionic diagrams is the single most frequent error. Examiners will not award full marks for an ionic dot and cross diagram unless the ions are clearly shown in brackets with the correct charges.

Drawing inner-shell electrons when the question only asks for the outer shell wastes time and can introduce errors. Unless the question specifically says otherwise, only draw the outer-shell electrons.

Using all dots or all crosses defeats the purpose of the diagram. The whole point is to show where each electron originated. Make sure one atom's electrons are dots and the other's are crosses.

Forgetting lone pairs on covalent diagrams is another common slip. Oxygen in water has two lone pairs, and each nitrogen in N₂ has one. These must appear in your diagram.

Drawing the wrong number of bonds happens when students do not count the outer electrons correctly. Always start by writing the electronic configuration and working out how many electrons each atom needs.

Start every diagram by writing the electronic configuration. This takes five seconds and prevents the most common errors.

Draw your diagrams large enough that the dots and crosses are clearly distinguishable. Cramped diagrams are hard for the examiner to read and hard for you to check.

If the question asks you to draw a dot and cross diagram for a compound you have not practised, do not panic. The method is always the same: Count the outer electrons, decide how many need to be transferred or shared, and draw accordingly.

Look at the number of marks available. A one-mark question might only need the shared pair shown. A three-mark question usually expects the full diagram with all outer electrons, correct symbols, and (for ionic compounds) brackets and charges.

Practise drawing the common examples until they are automatic. NaCl, MgO, H₂O, CO₂, CH₄ and N₂ cover every pattern you are likely to be asked about at GCSE.