SUBJECT Physical Sciences WEEK 8 TOPIC Physical state and density in terms of intermolecular

forces. Lesson 3

LESSON SUMMARY FOR: DATE STARTED: DATE COMPLETED:

LESSON OBJECTIVES

At the end of the lesson learners should be able to:

• Illustrate the proposition that intermolecular forces increase with increasing molecular size.

• Explain density of material in terms of the number of molecules in a unit volume, e.g. compare gases, liquids and solids TEACHING and LEARNING ACTIVITIES

1. TEACHING METHOD/S USED IN THIS LESSON:

Question and answer, explanation, practical experiment and demonstration.

2. LESSON DEVELOPMENT

2.1. Introduction [10 min]

Educator marks the homework given in lesson 2. Ensure that learners understand the difference between interatomic- and intermolecular forces as well as the different types of intermolecular forces.

ANSWERS TO HOMEWORK ACTIVITY

1. Covalent bonds are the sharing of one or more electron pairs between atoms in a molecule. Van der Waals forces are the forces that bond molecules in the liquid and solid phase.

2. Van der Waals forces increase in strength as the molecular mass of the molecules increase. CH4 molecules are bonded with weaker forces than C7H16 molecules. 3. Dipole-dipole forces (interaction) are the forces between polar molecules in the liquid or solid phase. In molecules, where a highly electronegative atom is bonded to

one or more H atoms, the intermolecular forces are stronger than dipole-dipole forces. These forces are called hydrogen bonds. 4.

a. Sublimates – Iodine changes phase from a solid directly to a gas. I2(s) ⟶ I(g) Diatomic refers to elements that exist as molecules e.g. N2; O2 etc.

b. I2(s) and Cℓ2(g) are non-polar molecules. c. Induced-dipole Van der Waals forces.

d. The I2(s) molecule has a molar mass of 253,8 g∙mol-1. Cℓ2(g) has a molar mass of 71 g∙mol-1. Both substances have induced-induced dipole forces (London forces) between the molecules. The strength of the intermolecular forces increases as the molar mass of the molecules increase. London forces are strong enough to bond I2 molecules in the solid phase, but not strong enough to bond the smaller Cℓ2 molecules in the solid phase.

5. Elements like N, O and F because they are highly electronegative. 6. Intermolecular forces.

7. The O-H bonds in the H2O molecule are very polar and the O atom takes on a significant negative charge. The proton of the H atom can get very close to the negative end of another H2O molecule. The electrostatic attraction between them is unusually strong. The H atom becomes a bridge between the two electronegative atoms in the adjacent molecules.

8.

a. Covalent bonds (Interatomic forces).

b. C12H22O11 is in the solid phase and C4H10 is in the gas phase.

c. The molecules in both substances are bonded with weak London forces. The London forces between the bigger sugar molecules are stronger than in the smaller butane molecules.

Solid Liquid Gas Ice, liquid water and

water vapour 9.

a. London forces or induced-dipole-induced-dipole forces. b. Hydrogen bonds

c. Hydrogen bonds and dipole-dipole forces. d. Dipole-dipole forces

e. Hydrogen bonds

f. Ionic bonds (Coulomb or electrostatic forces). 10. a. Dipole-dipole forces b. Dipole-dipole forces c. London forces d. London forces e. London forces f. Hydrogen bonds g. Hydrogen bonds PRE-KNOWLEDGE

A basic understanding of the following: Interatomic- and intermolecular forces. The kinetic particle model of matter.

2.2. Main Body (Lesson presentation) [20 min]

• The effect of molecular size on the strength of intermolecular forces

 The size of molecules (determined by atomic mass) plays an important role in the strength of intermolecular forces. Alkanes like methane (CH4) and octane (C8H18) consist of H-atoms and C- atoms, but methane is in the gaseous phase at room temperature, while octane is a liquid. Both these molecules are non-polar and bonded with similar intermolecular forces, but the bonds in octane (molar mass: 114 g∙mol-1_{) are stronger than in methane (molar mass: 16 g∙mol}-1_).

 The following elements and molecules can also be used as examples to explain the influence of atomic and molecular size on the strength of intermolecular forces: He(g), O2(g), C8H18(ℓ) (petrol) and C23H48(s) (wax).

• Density

 The common statements that mercury is “heavier” than water or that iron is “heavier” than aluminium are actually not correct. It is not the mass of the substances that is compared, but the mass per unit volume, which is known as the density.

 The density of a substance is independent of the amount and size of the sample and can therefore

be used as an aid to distinguish one pure substance from another. Stronger intermolecular forces will result in substances having greater densities.  According to the kinetic molecular theory, particles in a gas are far apart with no regular motion. Particles in a liquid are close together with no regular

arrangement. Particles in a solid are close together, usually in a regular pattern.

 It implies that the number of particles per unit volume will decrease from the solid to the liquid state. (Refer to diagram).

Method

• Add 3 – 5 cm3 _{of water and xylene (or CS}

2)

to test tubes 1 and 2.

• Add 3 -5 cm3 _{of water and ethanol to test}

tube 3.

• Add 3 -5 cm3 _{of ethanol to test tube 4 and}

5.

• Add a few KMnO4 crystals to the test tubes

1 and 4, and a few I2(s) crystals to test

tubes 2 and 5.

from the gas to the solid state.

 Note: The density of most substances known, decreases as the temperature rises. However, the density of water increases as the temperature is raised from 0 0_C to 4 0_{C. (Refer to the lesson: Macroscopic properties of the three phases of water related to their microscopic structure)}

LEARNER ACTIVITY(20 min) EXPERIMENT 1– GROUP ACTIVITY