3b. How to draw electronic bonding diagrams of small covalent molecules (Lewis dot and cross) and how to deduce a formula from valencies of elements

How do we draw electron diagrams of molecules?

   Cl-Cl

The diagrams above illustrate two of the many 'dot and cross' styles of representing the electronic structure of molecules.

The Venn diagram style (above right) is the best style to use because it most clearly shows the sharing of a pair of electrons to form a covalent bond.

The dots and crosses represent electrons, typically x from one atom and o from the other. (but not always!)

Two chlorine atoms combine to form the diatomic chlorine molecule (Cl₂).

The atomic number of chlorine is 17, therefore the atom has 17 protons holding 17 electrons, arranged in three shells, with a 2.8.7 electron configuration.

Chlorine needs to share one electron with another atoms to give itself a stable outer shell of 8 electrons.

Therefore two chlorine atoms share a pair of electrons, one from each atom, to form the single covalent bond.

The left-hand diagram shows all the inner and outer shells of electrons and the electron pair sharing indicated in a 'Venn diagram' overlap sharing style (as used in maths to indicate a shared grouping - and I think this is the best style, but others are accepted in exams - but take care).

The  right-hand diagram only shows the electrons and the pair of bonding electrons on the intersection of the two outer shells.

There are lots of dot and cross electron diagrams including simplified Lewis diagrams of covalent bonding in molecules

Lewis diagrams are quite minimalist, all they show is a duplet of electrons associated with hydrogen and for simple molecules, the complete octet of outer shell electrons for the other atoms.

It is important to appreciate that dot and cross Lewis diagrams for covalent molecules do not show the structure of the compound in terms of ...

(i) The 3D arrangement of the atoms in covalent molecules (e.g. H₂O compound)

(ii) the relative size of molecules in a covalent molecule

The simplest molecules are formed from two atoms and examples of their formation are shown below.

The electrons are shown as dots and crosses to indicate which atom the electrons come from, though all electrons are the same!

The diagrams may only show the outer electron arrangements for atoms that use two or more electron shells.

The electron structures are given in parentheses ().

I'll now describe lots of examples of simple covalent molecules, starting with the most 'simplest' of molecules of two atoms (diatomic) and finishing with molecules with up to 8 atoms!

I've illustrated each molecule describe lots of different styles of diagram you should be familiar with.

STYLES OF DIAGRAMS illustrated by ammonia  (full details on ammonia)

Consider molecules of >2 atoms, the minimum to consider the molecules as having a 'shape'.

(1) dot and cross electronic diagram, (2) 3D ball and stick model, (3) 3D space-filling model and (4) 2D displayed formula and all represent ammonia!

Its worth making the following comments on the different representations of simple covalent molecules of >2 atoms

So, reminders:

(1) Dot and cross diagrams are good for showing the electronic detail of the structure, and whether the bonds are single (ox) or double (ox) etc. However, it gives no idea on the shape of the molecule i.e. the 3D spatial arrangement of the bonds and atoms (its effectively a 2D diagram, but the molecule might not be flat!) and no information about the relative size of the atoms.

(2) It needs a 3D ball and stick model diagram to give an idea of the spatial arrangement of the atoms, but not the relative size of the atoms.

(3) A space filling model would give the shape of the molecule and the relative size of the atoms. However, neither of these two models show any electronic details of how the covalent bond is formed.

(4) Displayed formulae clearly shows how the atoms are bonded together, e.g. the arrangement of the three N-H single bonds, but only gives a 2D view of the molecule.

Some key points and extra notes about electronic diagram of covalent molecules

How to work out a covalent compound formula

Selected valencies of elements

The valency of an element is the combing power of its atoms.

Note that some elements can have more than one (valency):

Make the connection between the valency of an element and the number of outer electrons in its atoms ('valency electrons')

Hydrogen  H (1)

Chlorine Cl and other halogens (often 1 in simple molecular compounds)

Oxygen O (2 in most compounds)

Sulphur S (2, 4 and 6)

Boron B and aluminium Al (3 in all compounds)

Nitrogen N (3, 4 and 5)

Carbon C and silicon Si (4 in nearly all compounds)

Phosphorus P (usually 3 or 5)

To work out a covalent compound formula by combining 'A' with 'B' the rule is

number of atom 'A' x valency of atom 'A' = number of atom 'B' x valency of atom 'B'

Three examples are shown below

 

The diagram on the left illustrates the three covalent examples above for methane CH4 nitrogen trichloride NCl3 carbon dioxide CO2

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Key Concepts about Using Valency to Work Out Formulae with examples

• Valency definition: The number of electrons an atom needs to gain, lose, or share to achieve a full outer shell.

• Covalent bonding rule: Non-metals share electrons so that each atom achieves a stable outer shell.

• Steps to determine formula:

  1. Write down the valency of each element.

     • Hydrogen (1), Oxygen (2), Nitrogen (3), Carbon (4), Chlorine (1).

  2. Balance the valencies so that all atoms achieve full shells.

     • Example: Carbon (4) + Hydrogen (1) → CH₄ (carbon shares 4 electrons, each hydrogen shares 1).

  3. Check the total number of bonds matches the valency requirements.

     • Oxygen (2) + Hydrogen (1) → H₂O.

     • Nitrogen (3) + Hydrogen (1) → NH₂.

     • Carbon (4) + Oxygen (2) → CO₂ (two double bonds).

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Student Exam Tips

• Memorise common valencies:

  • H = 1, O = 2, N = 3, C = 4, Cl = 1.

• Check balance:

  • Multiply atoms until valencies balance (e.g. H₂O: 2 hydrogens × valency 1 = 2, balances oxygen’s valency 2).

• Confirm with diagrams:

  • Draw dot-and-cross diagrams to check electron sharing matches valency.

• Practice double/triple bonds:

  • O₂ (double bond), N₂ (triple bond), CO₂ (two double bonds).

• Exam technique:

  • Always show working (valency balancing).

  • Use correct subscripts in formulae (e.g. H₂O not H₂O).

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Typical Misconceptions

• “Valency is the same as group number” → Not always. e.g. Oxygen is in Group 6 but has valency 2.

• “Hydrogen needs 8 electrons” → False. Hydrogen only needs 2 (a full first shell).

• “Covalent bonds are weak” → Incorrect. The bonds are strong; it’s the intermolecular forces that are weak.

• “Forget multiple bonds” → Students often miss double/triple bonds when balancing valencies.

• “Formulae can be guessed” → Wrong. Formulae must be deduced systematically using valency.

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Quick Revision Overlay

• Valency method: Balance electron requirements.

• Examples:

  • H₂ (1+1), H₂O (2+1+1), CH₄ (4+1+1+1+1), NH₃ (3+1+1+1), CO₂ (4+2+2).

• Diagrams: Dot-and-cross, displayed formulae.

• Properties: Simple molecules (low mp/bp, poor conductivity), giant covalent (high mp/bp, hardness, conductivity in graphite).

• Exam focus: Deduce formulae, draw diagrams, explain properties, avoid misconceptions.