Lesson Notes By Weeks and Term v5 - Grade 11

Lesson Video

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Performance objectives

• Define and differentiate between intramolecular and intermolecular forces.
• Draw Lewis dot diagrams for simple molecules.
• Explain the relationship between intermolecular forces and physical properties.
• Identify types of intermolecular forces in different substances.
• Predict relative boiling and melting points of substances.

Lesson summary

This week, we begin our journey into the fascinating world of Matter and Materials, focusing specifically on the fundamental building blocks of everything around us: molecules and the forces that hold them together. Understanding molecular structure and intermolecular forces is crucial because it explains the properties of the materials we use every day. From the strength of building materials used in houses and bridges to the properties of the fuels that power our cars and the plastics we use for packaging, these concepts are at the heart of understanding the world around us.

Lesson notes

2.1 Intramolecular Forces: The Bonds WITHIN Molecules These forces hold atoms together within a molecule. They are generally much stronger than intermolecular forces.

Covalent Bonds: Formed when atoms share electrons. This usually occurs between nonmetal atoms.

Polar Covalent Bonds:* Unequal sharing of electrons due to differences in electronegativity. This creates partial charges (δ+ and δ-) on the atoms.

Example: Water (H₂O). Oxygen is more electronegative than hydrogen, so the oxygen atom has a partial negative charge (δ-) and each hydrogen atom has a partial positive charge (δ+).

Nonpolar Covalent Bonds:* Equal sharing of electrons. Occurs when atoms have similar electronegativities.

Example: Diatomic molecules like H₂, Cl₂, and molecules like methane (CH₄), where the electronegativity difference between C and H is small.

Ionic Bonds: Formed when electrons are transferred from one atom to another, creating ions (charged particles). This typically occurs between a metal and a nonmetal. The electrostatic attraction between the oppositely charged ions holds them together.

Example: Sodium chloride (NaCl). Sodium (Na) loses an electron to chlorine (Cl), forming Na+ and Cl- ions.

Metallic Bonds: Formed between metal atoms. Metal atoms donate their valence electrons to form a "sea" of delocalized electrons. This "sea" of electrons is attracted to the positively charged metal ions, holding the metal together. This explains why metals are good conductors of electricity and heat. 2.2 Intermolecular Forces: The Attractions BETWEEN Molecules These forces are weaker than intramolecular forces and determine many of the physical properties of substances. They arise from the attraction between the partial charges or temporary charges on molecules.

Van der Waals Forces: A general term for intermolecular forces.

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

F. LDF is the only intermolecular force present in nonpolar molecules.

Example: Methane (CH₄). Even though methane is nonpolar, temporary fluctuations in electron density create temporary dipoles.

Explanation of LDF:* Imagine a molecule of methane. At any given instant, the electrons might be slightly more concentrated on one side of the molecule. This creates a temporary, slight negative charge on that side and a slight positive charge on the other side. This temporary dipole can then induce a dipole in a neighboring methane molecule. The attraction between these temporary dipoles is the London Dispersion Force.

Dipole-Dipole Forces: Present in polar molecules. They arise from the attraction between the positive end of one polar molecule and the negative end of another polar molecule. These forces are stronger than LDF for molecules of similar size.

Example: Propanone (acetone) (CH₃COCH₃). The carbonyl group (C=O) makes the molecule polar.

Hydrogen Bonding: A particularly strong type of dipole-dipole force that occurs when hydrogen is bonded to a highly electronegative atom such as oxygen (O), nitrogen (N), or fluorine (F). The hydrogen atom has a significant partial positive charge and is attracted to the lone pair of electrons on the electronegative atom of another molecule. Hydrogen bonds are much stronger than typical dipole-dipole forces.

Example: Water (H₂O), ammonia (NH₃), hydrogen fluoride (HF). The hydrogen bonds between water molecules are responsible for many of water's unique properties, such as its high surface tension and high boiling point. 2.3 Molecular Shape and Polarity The shape of a molecule can determine whether it is polar or nonpolar. To determine the shape, we first need to draw the Lewis structure.

Lewis Structures (Electron Dot Diagrams): A diagram that shows the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule.

Rules for drawing Lewis structures:* Calculate the total number of valence electrons. Draw the skeletal structure of the molecule, with the least electronegative atom in the center (usually). Hydrogen is always terminal. Distribute the remaining electrons as lone pairs, starting with the most electronegative atoms. If necessary, form multiple bonds to satisfy the octet rule (or duet rule for hydrogen).

Example: Carbon Dioxide (CO₂)* Carbon has 4 valence electrons, and each oxygen has 6, so the total is 4 + 2(6) =

1

6. The skeletal structure is O-C-

O. Distribute the remaining electrons: O:C:O Satisfy the octet rule by forming double bonds: O=C=

O. Each atom now has 8 electrons around it.

Molecular Shape: The three-dimensional arrangement of atoms in a molecule. VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the shape of molecules based on the idea that electron pairs (both bonding and nonbonding) repel each other and want to be as far apart as possible.