What are dot-and-cross diagrams in the context of Cambridge International A Levels Chemistry?

Dot-and-cross diagrams are a visual representation used in Cambridge International A Levels Chemistry to illustrate the bonding between atoms in a molecule. They show how valence electrons are shared or transferred in covalent and ionic bonds, using dots and crosses to represent electrons from different atoms.

How do dot-and-cross diagrams help in understanding chemical bonding?

Dot-and-cross diagrams help students visualize the arrangement of electrons in molecules, making it easier to understand the nature of chemical bonds. By showing which electrons are shared or transferred, these diagrams clarify the difference between covalent and ionic bonding, which is a key concept in AS-Level Physical Chemistry.

What is the difference between covalent and ionic bonds in dot-and-cross diagrams?

In dot-and-cross diagrams, covalent bonds are depicted by shared pairs of electrons between atoms, often shown as overlapping dots and crosses. Ionic bonds, on the other hand, are represented by the complete transfer of electrons from one atom to another, resulting in ions that are shown separately with their respective charges.

Can dot-and-cross diagrams be used for complex molecules?

Yes, dot-and-cross diagrams can be used for complex molecules, although they are most effective for simple molecules. For larger molecules, these diagrams can become cumbersome, but they still provide valuable insights into the electron distribution and bonding nature, which is crucial for understanding molecular geometry and reactivity in AS-Level Chemistry.

Why are dot-and-cross diagrams important for Cambridge International A Levels exams?

Dot-and-cross diagrams are important for Cambridge International A Levels exams because they help students demonstrate their understanding of chemical bonding concepts. These diagrams are often used in exam questions to test a student's ability to represent and explain the electron arrangements in molecules, which is a fundamental skill in AS-Level Physical Chemistry.

Why is "Omitting lone pairs from covalent diagrams." a common mistake in Dot-and-cross diagrams?

Why it happens: Only drawing the bonds. How to avoid it: Draw every lone pair — they are marked and they determine shape. Source: 9701 Examiner Reports 2022-2024

Why is "Forgetting brackets/charges on ions." a common mistake in Dot-and-cross diagrams?

Why it happens: Treating ions like neutral atoms. How to avoid it: Enclose each ion in [ ] with its charge top-right.

Why is "Drawing one shared pair for a double bond." a common mistake in Dot-and-cross diagrams?

Why it happens: Not counting bond order. How to avoid it: Double = two shared pairs; triple = three.

Why is "Showing shared (overlapping) electrons in an ionic diagram." a common mistake in Dot-and-cross diagrams?

Why it happens: Mixing up ionic and covalent representations. How to avoid it: Ionic = full transfer (the cation's outer shell is empty); only covalent diagrams overlap electrons.

Why is "Forcing exactly 8 electrons onto a Period 3 central atom." a common mistake in Dot-and-cross diagrams?

Why it happens: Over-applying the octet rule. How to avoid it: P and S can expand their octet (PCl₅ has 10, SF₆ has 12 around the central atom).

Chemical Bonding(AS-Level Physical Chemistry)Cambridge International A Levels Chemistry (9701)

Dot-and-cross diagrams

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Dot-and-Cross Diagrams — Cambridge International AS & A Level Chemistry 9701 (2025-2027 syllabus)

How to draw dot-and-cross diagrams for ionic, covalent and coordinate (dative) bonding, including molecules with expanded octets and species with an odd number of electrons.

What you’ll learn

Mapped to the Cambridge IGCSE 9701 syllabus (2025-2027).

3.7.1 — Use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding, including the compounds in 3.4 and 3.5 (may include species with an expanded octet or an odd number of electrons).

How to draw a dot-and-cross diagram

Outer electrons only; dots and crosses to show which atom each electron came from.

A dot-and-cross diagram shows how the outer-shell (valence) electrons are arranged in a bond.

The rules:

Show only the outer-shell electrons (inner shells are ignored).
Use dots for one atom's electrons and crosses for the other's — this shows where each electron came from. (Electrons are identical in reality; the symbols are just for bookkeeping.)
Covalent bond = a shared pair in the overlap region (one dot + one cross).
Lone pairs must be drawn too (they affect shape — link to 3.5).
Ions are drawn in square brackets with the charge written outside, top-right.

Outer-shell electrons only.
Dots vs crosses = which atom the electron came from.
Covalent = shared pair in overlap; draw lone pairs.
Ions in [ ] with charge.

Ionic dot-and-cross

Transfer electrons, then show each ion separately in brackets with its charge.

For ionic compounds, show the transfer of electrons from metal to non-metal, then draw the separate ions.

Sodium chloride (NaCl). Na ( $2,8,1$ ) gives its single outer electron (a cross) to Cl ( $2,8,7$ ):

$[\text{Na}]^+$ — outer shell now empty (the next shell down is full).
$[\text{Cl with 8 outer electrons}]^-$ — the donated electron is shown as a cross among Cl's dots.

Magnesium oxide (MgO). Mg loses 2 electrons → $[\text{Mg}]^{2+}$ ; O gains 2 → $[\text{O with 8}]^{2-}$ .

Calcium fluoride (CaF₂). $[\text{Ca}]^{2+}$ with two separate $[\text{F}]^-$ ions, each having gained one electron.

Always include the square brackets and charges — they are required for the mark.

Show electron transfer.
Draw each ion in [ ] with its charge.
Cation outer shell usually empty; anion gains the electrons.

Covalent dot-and-cross

Shared pairs in the overlap; show lone pairs; double/triple bonds = 2/3 shared pairs.

For covalent molecules, place each shared pair in the overlap between the atoms and draw all lone pairs.

HCl — one shared pair (1 dot + 1 cross); Cl also has 3 lone pairs.
H₂O — two O–H shared pairs and two lone pairs on oxygen.
CO₂ — two C=O double bonds (each is two shared pairs); O atoms have lone pairs.
N₂ — a triple bond (three shared pairs) plus one lone pair on each N.

Water: each O–H bond is a shared pair (O's dot + H's cross); oxygen keeps two lone pairs.

Shared pair in the overlap (dot + cross).
Always draw lone pairs.
Double = 2 shared pairs; triple = 3.

Coordinate bonds, expanded octets and odd electrons

Dative bonds show both electrons from one atom; some atoms exceed (or fall short of) an octet.

Coordinate (dative covalent) bonds. Both shared electrons come from one atom, so they are drawn as two of the same symbol (e.g. two dots).

NH₄⁺ — three normal N–H bonds plus one dative bond (the N lone pair shared with H⁺); the whole ion is in brackets with a $+$ charge.
Al₂Cl₆ — bridging Cl atoms donate a lone pair into the Al of the other unit.

Expanded octets. Period 3 atoms can have more than 8 outer electrons:

PCl₅ — phosphorus shares 5 pairs (10 electrons around P).
SF₆ — sulfur shares 6 pairs (12 electrons around S).

Odd-electron species. Some molecules have an unpaired electron (a free radical) and cannot give every atom a full octet:

NO and NO₂ have an odd total number of electrons → one atom is left with an unpaired electron.

The syllabus explicitly allows these expanded-octet and odd-electron diagrams — don't force every atom to exactly 8.

Dative bond: both shared electrons the same symbol.
NH₄⁺ in brackets with + charge.
Expanded octets: PCl₅ (10 e⁻), SF₆ (12 e⁻).
Odd-electron species (NO, NO₂) have an unpaired electron.

How it’s examined

Dot-and-cross drawing appears on Paper 2 (2-3 marks) for a named molecule or ion. Examiner reports flag: missing lone pairs, missing brackets/charges on ions, drawing the wrong number of shared pairs in double/triple bonds, and not showing the dative bond clearly in NH₄⁺.

Sources: Cambridge International AS & A Level Chemistry 9701 syllabus (2025-2027), section 3.7; 9701 Examiner Reports 2022-2024; 9701/21 May/Jun 2024 question paper and mark scheme. Last reviewed 2026-05-20.

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Step-by-step worked examples — Dot-and-cross diagrams

Step-by-step solutions to past-paper-style questions on dot-and-cross diagrams, written exactly the way a tutor would explain them at the board.

1Dot-and-cross for H₂O (2 marks)
Getting started• covalent
Question
Describe the correct dot-and-cross structure of a water molecule.
Step-by-step solution
1. Step 1
  Two O–H bonds: each shared pair = 1 dot (O) + 1 cross (H).
2. Step 2
  Oxygen keeps two lone pairs (6 outer electrons − 2 bonding = 4 = 2 lone pairs).
Answer
Two shared O–H pairs (dot + cross each) and two lone pairs on oxygen.
Examiner tip
Lone pairs are marked — leave them off and you lose a mark even if the bonds are right.
2Ionic dot-and-cross for NaCl (3 marks)
Getting started• ionic
Question
Describe the dot-and-cross diagram for sodium chloride.
Step-by-step solution
1. Step 1
  Na loses its 1 outer electron → [Na]⁺ with an empty outer shell shown.
2. Step 2
  Cl gains it → [Cl with 8 outer electrons]⁻, the extra electron drawn as a cross.
3. Step 3
  Both ions in square brackets with charges (+ and −) outside.
Answer
[Na]⁺ (empty outer shell) and [Cl with 8 e⁻, one a cross]⁻, each in brackets with its charge.
Examiner tip
For ionic diagrams there are NO shared pairs — electrons are transferred, not overlapped.
3Dot-and-cross for CO₂ (3 marks)
Building confidence• covalent, double bond
Question
Describe the correct dot-and-cross structure of carbon dioxide, O=C=O.
Step-by-step solution
1. Step 1
  Two C=O double bonds, each with TWO shared pairs (each pair = 1 dot + 1 cross).
2. Step 2
  Each oxygen has two lone pairs remaining.
3. Step 3
  Carbon has no lone pairs (all 4 outer electrons are shared).
Answer
C shares two pairs with each O (two double bonds); each O carries two lone pairs; C has none.
Examiner tip
Each double bond must show two shared pairs — a single shared pair loses the mark.
4Ionic dot-and-cross for MgCl₂ (3 marks)
Building confidence• ionic, ratio
Question
Describe the dot-and-cross diagram for magnesium chloride.
Step-by-step solution
1. Step 1
  Mg (2,8,2) loses both outer electrons → [Mg]²⁺ with an empty outer shell.
2. Step 2
  Each electron goes to a different Cl → two [Cl with 8 e⁻]⁻ ions (one cross each).
3. Step 3
  All three ions bracketed with charges: one 2+ and two 1−.
Answer
[Mg]²⁺ with two [Cl]⁻ ions; Mg's two electrons go one to each chlorine.
Examiner tip
Show the 1 : 2 ratio explicitly — two separate chloride ions, not one with both electrons.
5Dative bond in NH₄⁺ (3 marks)
Stretch• dative
Question
How is the dative covalent bond shown in the dot-and-cross diagram of NH₄⁺?
Step-by-step solution
1. Step 1
  Three N–H bonds are normal shared pairs (1 dot + 1 cross).
2. Step 2
  The fourth N–H bond uses N's lone pair — both electrons are dots (from N), shared with H⁺.
3. Step 3
  Whole ion in square brackets with a + charge.
Answer
The fourth bond shows two dots (both from N's lone pair); the ion is bracketed with a + charge.
Examiner tip
An arrow (N→H) may also be used for the dative bond — but the two dots make the donor clear.
6Expanded octet: SF₆ (3 marks)
Stretch• expanded octet
Question
Describe the dot-and-cross diagram of sulfur hexafluoride, SF₆.
Step-by-step solution
1. Step 1
  Six S–F bonds: sulfur shares all 6 of its outer electrons, one with each F.
2. Step 2
  Sulfur now has 12 electrons around it — an expanded octet (allowed for Period 3).
3. Step 3
  Each fluorine has three lone pairs; sulfur has none.
Answer
Six shared S–F pairs (12 e⁻ on S, expanded octet); three lone pairs on each F.
Examiner tip
Period 3 central atoms can exceed 8 outer electrons — don't force an octet on S.

Key Formulae — Dot-and-cross diagrams

The formulae you need to memorise for dot-and-cross diagrams on the Cambridge IGCSE paper, with every variable defined in plain English and a note on when to use it.

Lone pairs on a central atom
$\text{lone pairs} = \frac{(\text{outer electrons}) - (\text{electrons used in bonding})}{2}$
$outer electrons$
the atom's own valence electrons (= group number)
$electrons used in bonding$
one per single bond, two per double, etc.
When to use
Checking how many lone pairs to draw, e.g. O in H₂O: (6 − 2)/2 = 2 lone pairs.

Key Definitions and Keywords — Dot-and-cross diagrams

Definitions to memorise and the exact keywords mark schemes credit for dot-and-cross diagrams answers — sharpened from recent examiner reports for the 2026 Cambridge IGCSE sitting.

Dot-and-cross diagram
Examiner keyword
A diagram showing the outer-shell electrons of bonded atoms, using dots and crosses to show which atom each electron came from.
Shared (bonding) pair
Examiner keyword
A covalent bond shown as one dot + one cross between two atoms; double = two such pairs, triple = three.
Lone pair (in diagrams)
Examiner keyword
A non-bonding pair of outer electrons, which must be drawn on dot-and-cross diagrams.
Dative bond (in diagrams)
Examiner keyword
A shared pair where both electrons come from the same atom — shown as two dots (or an arrow from the donor).
Expanded octet
Examiner keyword
More than eight outer electrons around a central atom, possible for Period 3 elements (e.g. PCl₅, SF₆).

Common Mistakes and Misconceptions — Dot-and-cross diagrams

The traps other students keep falling into on dot-and-cross diagrams questions — taken from recent Cambridge IGCSE examiner reports and mark schemes — and how to avoid them.

✕Omitting lone pairs from covalent diagrams.
9701 Examiner Reports 2022-2024
Why it happens
Only drawing the bonds.
How to avoid it
Draw every lone pair — they are marked and they determine shape.
✕Forgetting brackets/charges on ions.
Why it happens
Treating ions like neutral atoms.
How to avoid it
Enclose each ion in [ ] with its charge top-right.
✕Drawing one shared pair for a double bond.
Why it happens
Not counting bond order.
How to avoid it
Double = two shared pairs; triple = three.
✕Showing shared (overlapping) electrons in an ionic diagram.
Why it happens
Mixing up ionic and covalent representations.
How to avoid it
Ionic = full transfer (the cation's outer shell is empty); only covalent diagrams overlap electrons.
✕Forcing exactly 8 electrons onto a Period 3 central atom.
Why it happens
Over-applying the octet rule.
How to avoid it
P and S can expand their octet (PCl₅ has 10, SF₆ has 12 around the central atom).

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

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Dot-and-cross diagrams — frequently asked questions

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Dot-and-cross diagrams

Short Study Notes in the form of Slides

Detailed Notes

What you’ll learn

How it’s examined

1Dot-and-cross for H₂O (2 marks)

2Ionic dot-and-cross for NaCl (3 marks)

3Dot-and-cross for CO₂ (3 marks)

4Ionic dot-and-cross for MgCl₂ (3 marks)

5Dative bond in NH₄⁺ (3 marks)

6Expanded octet: SF₆ (3 marks)

Lone pairs on a central atom

Dot-and-cross diagram

Shared (bonding) pair

Lone pair (in diagrams)

Dative bond (in diagrams)

Expanded octet

✕Omitting lone pairs from covalent diagrams.

✕Forgetting brackets/charges on ions.

✕Drawing one shared pair for a double bond.

✕Showing shared (overlapping) electrons in an ionic diagram.

✕Forcing exactly 8 electrons onto a Period 3 central atom.

Practice questions

Past paper style quiz

Assess your understanding

Dot-and-cross diagrams — frequently asked questions