Lesson Notes By Weeks and Term v5 - Grade 11

Lesson Video

This page supports the lesson note with a companion video and a short classroom-ready summary.

For class groups and homework, share this lesson page so learners also get the summary, objectives, and full lesson context.

Performance objectives

• Define and differentiate types of intermolecular forces (IMFs).
• Relate IMFs to physical properties like boiling point.
• Draw Lewis diagrams to determine molecular polarity.
• Explain how IMFs influence the state of matter.
• Predict the relative strength of IMFs between molecules.

Lesson summary

This week, we delve deeper into the world of matter and materials, focusing specifically on molecular structure and intermolecular forces. Understanding these concepts is crucial as they determine the physical properties of substances we encounter daily, from the water we drink to the plastics used in packaging and the materials used in constructing our homes. In South Africa, understanding material properties is crucial for industries like mining (extracting metals with specific properties), agriculture (designing effective fertilizers and pesticides), and manufacturing (creating durable and safe products).

Lesson notes

2.1 Molecular Structure: The Building Blocks Before we can understand intermolecular forces, we need to appreciate the structure of molecules. Molecules are formed when atoms bond together through intramolecular forces (covalent bonds). The arrangement of these atoms, and the type of bonds between them, determines the molecule's shape and polarity.

Lewis Diagrams (Structural Formulas): These diagrams show how atoms are connected in a molecule and the distribution of valence electrons. Remember to satisfy the octet rule (or duet for hydrogen) as much as possible. Shared electrons represent covalent bonds. Lone pairs are pairs of electrons not involved in bonding.

Example: Water (H₂O). Oxygen has 6 valence electrons, and each hydrogen has

1. Oxygen needs two more electrons to complete its octet. Each hydrogen shares one electron with the oxygen, forming two single covalent bonds. The oxygen atom also has two lone pairs of electrons.

Steps:

1. Count total valence electrons.

2. Draw the skeletal structure (least electronegative atom in the center).

3. Distribute electrons as lone pairs to outer atoms first, satisfying octets.

4. Place remaining electrons on the central atom.

5. If the central atom does not have an octet, form multiple bonds.

Polarity: Polarity arises due to differences in electronegativity between atoms in a bond. Electronegativity is the ability of an atom to attract electrons in a chemical bond. If one atom is significantly more electronegative than another, the bonding electrons will be pulled closer to the more electronegative atom, creating a partial negative charge (δ-) on that atom and a partial positive charge (δ+) on the less electronegative atom. This creates a polar bond.

Example: In HCl, chlorine is more electronegative than hydrogen.

Therefore, the bonding electrons are pulled closer to the chlorine atom, making it partially negative (δ-) and the hydrogen atom partially positive (δ+). The molecule has a dipole moment.

Molecular Polarity: The overall polarity of a molecule depends on both the polarity of its individual bonds and its molecular geometry. If the bond dipoles cancel each other out due to symmetry, the molecule is nonpolar, even if it contains polar bonds. If the bond dipoles do not cancel, the molecule is polar.

Example 1: Carbon Dioxide (CO₂): Each C=O bond is polar.

However, the molecule is linear, so the bond dipoles point in opposite directions and cancel each other out.

Therefore, CO₂ is nonpolar.

Example 2: Water (H₂O): The O-H bonds are polar. The molecule is bent (not linear), so the bond dipoles do not cancel out.

Therefore, water is polar. 2.2 Intermolecular Forces: Attractive Interactions Intermolecular forces (IMFs) are attractive forces between molecules. These forces are much weaker than the intramolecular forces (covalent bonds) that hold atoms together within a molecule. IMFs are responsible for many of the physical properties of matter, such as boiling point, melting point, viscosity, and surface tension.

Van der Waals Forces: A general term encompassing all intermolecular forces except for hydrogen bonding.

London Dispersion Forces (LDF): Present in all molecules (polar and nonpolar). These forces arise from temporary, instantaneous dipoles that occur due to the random movement of electrons. The larger the molecule (i.e., the more electrons it has), the stronger the LD

F. Example: Noble gases (He, Ne, Ar, Kr, Xe). They are nonpolar, but they can be liquefied at low temperatures due to LDFs. Xenon (Xe) has a higher boiling point than helium (He) because it has more electrons and therefore stronger LDFs.

Dipole-Dipole Forces: Occur between polar molecules. The positive end of one polar molecule is attracted to the negative end of another polar molecule. These forces are stronger than LDFs for molecules of similar size and shape.

Example: Propanone (acetone, CH₃COCH₃). It has a permanent dipole moment because the carbonyl group (C=O) is polar.

Therefore, propanone experiences dipole-dipole forces.

Hydrogen Bonding: A special type of dipole-dipole interaction that occurs when a hydrogen atom bonded to a highly electronegative atom (N, O, or F) is attracted to a lone pair of electrons on another N, O, or F atom in a neighboring molecule. Hydrogen bonds are particularly strong intermolecular forces.

Example: Water (H₂O). The hydrogen atoms are bonded to oxygen, and they are attracted to the lone pairs on oxygen atoms in other water molecules. This explains why water has a relatively high boiling point compared to other molecules of similar size. 2.3 Intermolecular Forces and Physical Properties Boiling Point and Melting Point: Substances with stronger intermolecular forces have higher boiling points and melting points. This is because more energy is required to overcome the attractive forces holding the molecules together.

Viscosity: Viscosity is a measure of a fluid's resistance to flow. Liquids with stronger intermolecular forces tend to be more viscous.

Worked example

Question: Which of the following substances has the highest boiling point: methane (CH₄), ammonia (NH₃), or water (H₂O)? Explain your answer in terms of intermolecular forces.

Solution:

Methane (CH₄) is nonpolar and experiences only London dispersion forces.

Ammonia (NH₃) is polar and experiences London dispersion forces and dipole-dipole forces. It also exhibits hydrogen bonding (N-H bond).

Water (H₂O) is polar and experiences London dispersion forces, dipole-dipole forces, and strong hydrogen bonding (O-H bond).

Since hydrogen bonding is the strongest intermolecular force, and water forms stronger hydrogen bonds than ammonia due to the higher electronegativity of oxygen, water will have the highest boiling point.

Question: Explain why ethanol (CH₃CH₂OH) is soluble in water, while ethane (CH₃CH₃) is not.

Solution:

Ethanol contains an -OH group, which allows it to form hydrogen bonds with water molecules. Water is also polar and can form hydrogen bonds.

Therefore, ethanol and water are miscible (soluble in all proportions) due to these strong intermolecular attractions. "Like dissolves like."

Ethane is a nonpolar molecule and only experiences London dispersion forces. It cannot form strong interactions with water molecules.

Therefore, ethane is not soluble in water.

Question: Arrange the following substances in order of increasing boiling point: pentane (C₅H₁₂), hexane (C₆H₁₄), butane (C₄H₁₀). Explain your reasoning.

Solution:

All three substances are alkanes and therefore nonpolar. The only intermolecular forces present are London dispersion forces.

The strength of London dispersion forces increases with molecular size (number of electrons).

Therefore, the order of increasing boiling point is: butane (C₄H₁₀) < pentane (C₅H₁₂) < hexane (C₆H₁₄).

Guided Practice (With Solutions)

Question: Draw the Lewis diagram for carbon tetrachloride (CCl₄). Is the molecule polar or nonpolar? Explain.