Chemistry · Year 12 · Bonding and Molecular Geometry · Autumn Term

Covalent Bonding and Lewis Structures

Drawing Lewis structures to represent shared electron pairs and formal charges.

• National Curriculum Attainment Targets
• A-Level: Chemistry - Covalent Bonding
• A-Level: Chemistry - Lewis Structures

About This Topic

This topic explores the subtle forces that exist between molecules, which determine whether a substance is a gas, liquid, or solid at room temperature. Students learn about electronegativity and how it leads to polar bonds, and then how those polar bonds (or lack thereof) result in London forces, permanent dipole-dipole interactions, or hydrogen bonding.

In the UK curriculum, this is a critical area for explaining the 'anomalous' properties of water and the trends in boiling points of organic compounds. Students must be able to identify which forces are present in a given molecule and predict their relative strengths. This understanding is vital for everything from distillation in organic chemistry to the structure of DNA in biology.

Students grasp these concepts faster through structured comparison tasks and 'predict and explain' activities, where they must justify why two molecules of similar mass have vastly different boiling points.

Key Questions

1. Construct Lewis structures for simple molecules and polyatomic ions.
2. Explain the concept of formal charge and its use in determining the most stable Lewis structure.
3. Differentiate between single, double, and triple covalent bonds.

Learning Objectives

• Construct Lewis structures for molecules and polyatomic ions containing up to four electron domains.
• Calculate formal charges for atoms within a Lewis structure to identify the most plausible arrangement of electrons.
• Differentiate between single, double, and triple covalent bonds based on electron pair sharing and bond length.
• Evaluate the stability of different resonance structures for a given molecule or ion.

Before You Start

Atomic Structure and Electron Configuration

Why: Students must understand electron shells and how to determine the number of valence electrons for each element.

Periodic Trends

Why: Knowledge of electronegativity is helpful for understanding bond polarity, which relates to the distribution of electrons in covalent bonds.

Key Vocabulary

Covalent Bond	A chemical bond formed by the sharing of one or more pairs of electrons between atoms.
Lewis Structure	A diagram showing the bonding between atoms of a molecule or polyatomic ion, with dots representing valence electrons.
Valence Electrons	Electrons in the outermost shell of an atom that are available for forming chemical bonds.
Formal Charge	A hypothetical charge assigned to an atom in a molecule, assuming all bonds are purely covalent and electron pairs are shared equally.
Octet Rule	The tendency of atoms to prefer having eight electrons in their valence shell, achieved by gaining, losing, or sharing electrons.

Watch Out for These Misconceptions

Common Misconception

Hydrogen bonds are bonds *inside* a molecule (intramolecular).

What to Teach Instead

Hydrogen bonds are *inter*molecular forces between molecules. A 'draw the dotted line' activity helps students distinguish between the strong covalent bond within a water molecule and the weaker hydrogen bond between two different water molecules.

Common Misconception

London forces only exist in non-polar molecules.

What to Teach Instead

London forces exist between *all* molecules, but they are the *only* force in non-polar ones. Using a Venn diagram activity helps students categorise which molecules have which combinations of forces.

Active Learning Ideas

Inquiry Circle: Boiling Point Mystery

Groups are given a list of compounds with their molar masses and boiling points. They must identify the intermolecular forces in each and explain the trends, specifically focusing on why water and ethanol are outliers.

40 min·Small Groups

Simulation Game: The Static Charge Test

Students use a charged rod to deflect a stream of different liquids (e.g., water, hexane, ethanol). They observe which streams bend and use their knowledge of molecular dipoles to explain why polar liquids are attracted to the rod.

20 min·Small Groups

Think-Pair-Share: The Importance of Hydrogen Bonding

Pairs discuss what would happen to life on Earth if water didn't have hydrogen bonding (e.g., ice sinking, oceans evaporating). They share their most significant 'consequence' with the class.

15 min·Pairs

Real-World Connections

• Pharmaceutical chemists use Lewis structures to predict the reactivity of drug molecules, helping to design new medications with specific therapeutic effects.
• Materials scientists at companies like DuPont analyze the bonding in polymers using Lewis structures to develop new plastics and composites with enhanced strength and durability for applications ranging from aerospace to consumer goods.
• Environmental chemists draw Lewis structures to understand the bonding in atmospheric pollutants, such as ozone (O3) and nitrogen dioxide (NO2), to model their formation and reactions.

Assessment Ideas

Quick Check

Provide students with the chemical formulas for CO2, NH3, and SO4^2-. Ask them to draw the Lewis structure for each, including all valence electrons and formal charges. Collect and review for accuracy in electron placement and formal charge calculation.

Peer Assessment

In pairs, students exchange Lewis structures they have drawn for a given molecule (e.g., H2O, NO3^-). One student explains their structure, focusing on electron counting and formal charge. The other student acts as a verifier, asking clarifying questions and identifying any discrepancies. They then switch roles.

Exit Ticket

On an index card, students write the Lewis structure for HCN. They must also identify the formal charge on each atom and indicate whether the bonds are single, double, or triple. This checks their ability to apply the rules to a new molecule.

Frequently Asked Questions

What is electronegativity and how does it create dipoles?

Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond. When two atoms with different electronegativities bond, the electrons are pulled closer to the more electronegative atom, creating a partial negative charge (delta minus) and a partial positive charge (delta plus) on the other, forming a dipole.

Why is hydrogen bonding the strongest intermolecular force?

Hydrogen bonding occurs when hydrogen is bonded to a very electronegative atom (N, O, or F). This creates a very large dipole. Because hydrogen is so small, the delta-plus charge is highly concentrated, allowing it to get very close to the lone pair on a neighbouring molecule, resulting in a particularly strong attraction.

How does active learning help students understand polarity?

Active learning strategies like 'molecular tug-of-war' or using 3D simulations allow students to visualise the 'pull' of electrons. By physically representing the atoms and the electron density, students can better understand why some molecules with polar bonds (like CO2) are actually non-polar overall due to their symmetry.

How do London forces change with molecular size?

As molecules get larger, they have more electrons. This increases the chance of temporary dipoles forming and makes the electron cloud more 'polarisable'. Larger molecules therefore have stronger London forces, which explains why the boiling points of noble gases and alkanes increase with atomic or molecular mass.