Lewis Formulas (DP IB Chemistry): Revision Note

Alexandra Brennan

Written by: Alexandra Brennan

Reviewed by: Richard Boole

Updated on

Lewis Formulas

  • Lewis formulas are also known as Lewis structures or electron dot diagrams

  • They show all valence electrons in a covalently bonded species

    • This includes bonding pairs and lone pairs

  • A pair of electrons can be represented by:

    • Dots

    • Crosses

    • A combination of dots and crosses

    • A line

  • For example, chlorine can be shown as:

Lewis formulas for chlorine
Different Lewis formulas for chlorine molecules
  • Note: Cl–Cl is not a Lewis formula, since it does not show all the electron pairs

Steps for drawing Lewis formulas

  1. Count the total number of valence electrons:

    • This is for all atoms in the molecule

  2. Draw the skeletal positions:

    • This is to show how the atoms are connected

    • The central atom, usually the least electronegative, should be in the middle

  3. Add electron pairs:

    • First, add the bonding pairs between each pair of bonded atoms

    • Then, add lone pairs to complete octets (or duets for hydrogen), starting with outer atoms

  4. Check your structure:

    • Confirm that all valence electrons have been used

    • Ensure all atoms (except exceptions) have full octets, or duets for hydrogen

Worked Example

How many electrons are in the 2-aminoethanoic acid molecule?

2-aminoethanoic acid

Answer:

  • Add valence electrons for each atom:

    • C (×2) = 2 × 4 = 8

    • H (×5) = 5 × 1 = 5

    • O (×2) = 2 × 6 = 12

    • N (×1) = 5

    • Total = 30 electrons

  • Include both bonding and lone pairs

    • Remember to count lone pairs on O and N

    • Each single bond contains 2 electrons

    • Lone pairs are especially important on N and O, which follow the octet rule

  • Conclusion:

    • The total number of valence electrons is 30

Drawing Lewis formulas

  • Applying the steps to draw full Lewis formulas for some simple molecules:

Methane, CH4

  1. Count the total number of valence electrons:

    • C = 4 electrons

    • H = 1 electron

    • CH4 = (1 x C) + (4 x H) = (1 x 4) + (4 x 1) = 8 electrons total

  1. Draw the skeletal positions:

    • Carbon is the central atom surrounded by four hydrogen atoms

Step 2 of drawing the methane Lewis structure, placing the more electronegative carbon atom in the centre of 4 hydrogen atoms
  1. Add electron pairs:

    • First, add the bonding pairs

    • Each C-H bond contains one pair of electrons

    • Then, add lone pairs

    • The bonding pairs account for all electrons

    • So, there are no lone pairs

Step 3 of drawing the methane Lewis structure, placing the electrons
  1. Check your structure:

    • All valence electrons used

    • Carbon has a full octet

    • Hydrogen has duets

Ammonia, NH3

  1. Count the total number of valence electrons:

    • N = 5 electrons

    • H = 1 electron

    • NH3 = (1 x N) + (3 x H) = (1 x 5) + (3 x 1) = 8 electrons total

  1. Draw the skeletal positions:

    • Nitrogen is the central atom surrounded by three hydrogen atoms

Step 2 of drawing the ammonia Lewis structure, placing the more electronegative nitrogen atom in the centre of 3 hydrogen atoms
  1. Add electron pairs:

    • First, add the bonding pairs

      • Each N-H bond contains one pair of electrons

    • Then, add lone pairs

      • The bonding pairs account for 6 electrons

      • So, there is one lone pair on nitrogen

Step 3 of drawing the ammonia Lewis structure, placing the electrons
  1. Check your structure:

    • All valence electrons used

    • Nitrogen has a full octet

    • Hydrogen has duets

Tetrachloromethane, CCl4

  1. Count the total number of valence electrons:

    • C = 4 electrons

    • Cl = 7 electrons

    • CCl4 = (1 x C) + (4 x Cl) = (1 x 4) + (4 x 7) = 32 electrons total

  2. Draw the skeletal positions:

    • Carbon is the central atom surrounded by four chlorine atoms

Step 2 of drawing the tetrachloromethane Lewis structure, placing the carbon atom in the centre of 4 chlorine atoms
  1. Add electron pairs:

    • First, add the bonding pairs

      • Each C-Cl bond contains one pair of electrons

    • Then, add lone pairs

      • The bonding pairs account for 8 electrons and complete the octet for C

      • There are 24 remaining electrons

      • This means that there is a total of 12 lone pairs of electrons

      • Each chlorine atom has 3 lone pairs

Step 3 of drawing the tetrachloromethane Lewis structure, placing the electrons
  1. Check your structure:

    • All valence electrons used

    • Carbon has a full octet

    • Chlorine has octets

Worked Example

Draw the Lewis structures for:

  1. Water, H2O

  2. Carbon dioxide, CO2

  3. Hydrogen cyanide, HCN

Answers:

  1. Water, H2O

    • H2O = (2 x H) + (1 x O) = (2 x 1) + (1 x 6) = 8 electrons total

    • Oxygen is the central atom surrounded by two hydrogen atoms

    • Each O-H bond contains one pair of electrons

    • There are 2 lone pairs on the oxygen

Lewis structure of a water molecule, showing oxygen with two lone pairs and single bonds to two hydrogen atoms.
  1. Carbon dioxide, CO2

    • CO2 = (1 x C) + (2 x O) = (1 x 4) + (2 x 6) = 16 electrons total

    • Carbon is the central atom surrounded by two oxygen atoms

    • Each C=O bond contains two pairs of electrons

    • Each oxygen atom has 2 lone pairs

Lewis structure of carbon dioxide, with a central carbon atom doubly bonded to two oxygen atoms, each with four lone electrons.
  1. Hydrogen cyanide, HCN

    • HCN = (1 x H) + (1 x C) + (1 x N) = (1 x 1) + (1 x 4) + (1 x 5) = 10 electrons total

    • Carbon is the central atom surrounded by one hydrogen and one nitrogen atom

    • The C-H bond contains one pair of electrons

    • The Cidentical toN bond contains three pairs of electrons

    • The nitrogen atom has 1 lone pair

Lewis structure of hydrogen cyanide (HCN) showing single bond between H and C, triple bond between C and N, with electrons marked by dots.

When the octet rule does not apply

  • Most elements tend to follow the octet rule, aiming for 8 electrons in their valence shell to match a noble gas configuration

  • However, several elements can form stable compounds with fewer than 8 electrons:

    • Hydrogen (H):

      • Stable with 2 electrons (1s2), like helium

    • Lithium (Li):

      • Loses one electron to become Li+ (1s2)

    • Beryllium (Be):

      • Often forms compounds with just 4 valence electrons

    • Boron (B) and Aluminium (Al):

      • Can form stable compounds with only 6 valence electrons

  • Two examples of stable molecules with incomplete octets that do not follow the typical octet rule are:

Beryllium chloride, BeCl2

  • BeCl2 = (1 x Be) + (2 x Cl) = (1 x 2) + (2 x 7) = 16 electrons total

  • Beryllium is the central atom surrounded by two chlorine atoms

  • Each Be-Cl bond contains one pair of electrons

    • Be has only 4 electrons around it (2 bonding pairs)

    • Be does not achieve an octet, but the molecule is stable

  • Each chlorine atom has 3 lone pairs

    • Each chlorine has a complete octet

  • BeCl2 is stable despite beryllium not reaching 8 electrons

Lewis structure showing a beryllium atom bonded to two chlorine atoms, each with six additional dots representing valence electrons.

Boron trifluoride, BF3

  • BF3 = (1 x B) + (3 x F) = (1 x 3) + (3 x 7) = 24 electrons total

  • Boron is the central atom surrounded by three fluorine atoms

  • Each B-F bond contains one pair of electrons

    • B has only 6 electrons around it (3 bonding pairs)

    • B does not achieve an octet, but the molecule is stable

  • Each fluorine atom has 3 lone pairs

    • Each fluorine has a complete octet

  • BF3 is stable despite boron not reaching 8 electrons

Lewis dot structure of boron trifluoride (BF3) showing boron atom with three fluorine atoms and their valence electrons arranged in a trigonal planar shape.

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Alexandra Brennan

Author: Alexandra Brennan

Expertise: Chemistry Content Creator

Alex studied Biochemistry at Newcastle University before embarking upon a career in teaching. With nearly 10 years of teaching experience, Alex has had several roles including Chemistry/Science Teacher, Head of Science and Examiner for AQA and Edexcel. Alex’s passion for creating engaging content that enables students to succeed in exams drove her to pursue a career outside of the classroom at SME.

Richard Boole

Reviewer: Richard Boole

Expertise: Chemistry Content Creator

Richard has taught Chemistry for over 15 years as well as working as a science tutor, examiner, content creator and author. He wasn’t the greatest at exams and only discovered how to revise in his final year at university. That knowledge made him want to help students learn how to revise, challenge them to think about what they actually know and hopefully succeed; so here he is, happily, at SME.

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