The effects of I ntermolecular forces

1. The effects of Intermolecular forces

2. AIM • Matter can exist in one of four different states: solid, liquid, gas or plasma. The solid and liquid state are also known as the condensed states of matter. • There are forces between the molecules of the substance in the condensed states, causing them to attract one another. These forces are also known as intermolecular forces. • In the five experiments that form part of the formal assessment (term two), learners will investigate the effect of the strength of intermolecular forces on different physical properties of matter.

3. Introduction to Intermolecular Forces • Pure metal atoms bond with strong metallic bonds, e.g. Cu(s), Na(s), Li(s), Fe(s) and Mg(s) • The crystal lattice consists of positive atomic cores that attract a sea of delocalized electrons. • Strong electrostatic forces are responsible for the fact that metals (except mercury) all exist in the solid state at room temperature. • Metals can conduct electricity in the solid and liquid state, because free ions can act as carriers of charge. • The smallest particle in the lattice, is the positive core ion. • When atoms of metals and non-metals react, they form ionic bonds. • Electrons are transferred from the metal atom to the non-metal atom and a cation and anion are formed. • Strong electrostatic (coulomb) forces exist between the oppositely charged ions and they form an ionic crystal lattice. • All ionic substances are solids at room temperature. • The substance can only conduct electricity in the liquid or molten state, when ions are free to act as charge carriers. • A liquid that conducts electricity is called an electrolyte. • Ionic solids are not conductors, because the bonds between the ions are very strong and ions are not free to move. Atoms of non-metal elements can also bond to form new compounds. • Non-metal atoms share one or more electron pairs and form covalent bonds. • Covalent bonds are very strong and known as intramolecular- or interatomic forces. • The product is called a simple molecule, e.g. H2O. • It is also possible that a macro molecule can form, e.g. diamond or graphite.

4. Examples of molecules represented Lewis Diagrams Description • The formation of H2, N2 and CO2-molecules are represented with Lewis diagrams.

5. Examples of molecules represented Ball and Stick Diagrams Description • Simple molecules like P4, S8 and Cℓ2 are represented below with ball-and-stick models.

6. Introduction to Intermolecular Forces continued • Molecular substances consist of molecules that are bonded with intermolecular forces to each other. • Types of intermolecular forces: • Dipole-dipole Van der Waals forces and hydrogen bonds • Induced dipole-dipole Van der Waals forces • Dipole induced-dipole Van der Waals forces • Ion-dipole forces • Ion-induced dipole force • Van der Waals forces are the intermolecular forces between neutral molecules (polar or non-polar) in the solid or liquid state.

7. Introduction to Induced Dipole Forces • These forces exist between non-polar molecules and are also known as London or dispersion forces. • When non-polar molecules or atoms approach each other there is a slight change in the charge distribution of the electron cloud of each particle and temporary dipoles are formed (induced). • These weak forces are responsible for the existence of the substance in the liquid or solid state. • The bigger the molecules (atoms) the stronger the induced forces. • Examples: • H2, N2, O2, F2, Cℓ2, Br2, I2, BH3, BeH2, SO3, C12H22O11(sugar), CH4, CCℓ4, oil, C8H18 (petrol), noble gases.

8. Introduction to Dipole-Dipole Forces • Polar bonds between atoms can occur when one atom is more electronegative than the other. • The shared electron pair(s) is drawn closer to the more electronegative atom causing a dipole. • The molecule is polar if it has an asymmetrical shape which causes an uneven distribution of charge on the molecule. • HCℓ is a diatomic molecule that consists of two different atoms. • The Lewis diagram for the molecule is: • EN (Cℓ) = 3,0 • EN (H) = 2,1 • ∆EN = 0,9 The bond between the atoms is polar. The shape of the molecule is linear and asymmetrical. • There is a net dipole moment in the molecule and it is therefore a polar molecule.

9. Introduction to Dipole-Dipole Forces • Polar bonds between atoms can occur when one atom is more electronegative than the other. • The shared electron pair(s) is drawn closer to the more electronegative atom causing a dipole. • The molecule is polar if it has an asymmetrical shape which causes an uneven distribution of charge on the molecule. • HCℓ is a diatomic molecule that consists of two different atoms. • The Lewis diagram for the molecule is: • EN (Cℓ) = 3,0 • EN (H) = 2,1 • ∆EN = 0,9 The bond between the atoms is polar. The shape of the molecule is linear and asymmetrical. • There is a net dipole moment in the molecule and it is therefore a polar molecule. • The Lewis diagram for a sulphur(IV)oxide molecule (SO2), is: • EN (O) = 3,5 • EN (S) = 2,5 • ∆EN = 1,0 The bond between an oxygen and a sulphur atom is polar. The shape of the molecule is angular and the molecule asymmetrical. • There is a net dipole moment in the molecule and it is a polar molecule. • Polar molecules bond with stronger dipole-dipole Van der Waals forces. • The strength of the forces increases as the size of the molecules increases. • Examples • Hydrides of group V, VI and VII (except NH3, H2O and HF). • Chloroform (CH3Cℓ) • Acetone (CH2COCH3).

10. Introduction to Hydrogen Bonds • Hydrogen bonds occur between molecules in which the hydrogen atom is bonded to a small atom with an extremely high electronegativity – oxygen, nitrogen and fluorine. • Hydrogen bonds are specific types of dipole-dipole forces, but is much stronger. • Examples: • NH3, H2O and HF • Organic alcohols e.g. methanol (CH3OH), ethanol (C2H5OH) and glycerol (C3H5(OH)3). • Organic carboxylic acids e.g. methanoic acid (CHOOH) and ethanoic acid (CH3COOH) that is present in vinegar. • Carboxylic acids can form two hydrogen bonds per molecule and the bonds are therefore stronger than in alcohols.

11. Introduction to Dipole-Induced Dipole Bonds • When a non-polar molecule dissolves in water (polar) dipoles are induced in the non-polar molecules. • These intermolecular forces are normally relative weak. • Examples • Non-polar sugar (C12H22O11) is soluble in polar water.

12. Introduction to Ion-induced dipole forces • Ionic substances are normally not soluble in non-polar liquids. • The bigger the molecules of the non-polar substance, the greater the possibility that dipoles can be induced in these molecules. • Examples • Sodium chloride (NaCℓ) is soluble in hexane (C6H14), which is a non-polar substance, because of ion-dipole interaction.

13. Introduction to Ion-dipole forces • When ionic substances dissolve in polar solvents, e.g. H2O(ℓ), the ions become surrounded by the polar water molecules and there is evidence of ion-dipole interaction. • This process is also called dissociation and the ions that are separated during the dissolution process are responsible for the fact that the solution acts as a conductor of electricity. • Solutions that can conduct electricity are also known as electrolytes.

14. Physical properties and intermolecular forces • The state in which a substance exists at a specific temperature depends on how strong the forces between the particles are. • The stronger the intermolecular forces, the greater the possibility that the substance can be in the liquid or solid state • London < Dipole-dipole < hydrogen bonds. • It is also possible that substances can bond with the same intermolecular forces, but that the strength of the forces is different. • Ethane molecules (C2H6) and octane molecules (C8H18) consist of non-polar molecules that bond with induced-dipole forces, however octane is a liquid at room temperature and ethane is a gas. • The intermolecular forces are weaker between smaller molecules. • The density of a substance is also determined by the intermolecular forces – the stronger the forces, the more dense is the substance. • Liquids with weaker intermolecular forces evaporate more easily than ones with stronger forces. These liquids have a high vapour pressure. • The boiling point of a liquid is the temperature at which the vapor pressure is equal to atmospheric pressure. • The melting point is the temperature at which a solid changes into a liquid. • Liquids with weak bonds and high vapor pressures have low boiling points. • Solids that contain weak bonds in the solid state have low melting points. • Viscosity is an indication of the resistance of a liquid to flow. The stronger the intermolecular forces the greater the viscosity.

15. Physical properties and intermolecular forces • Intermolecular forces also play an important role in how soluble a solute is in a solvent. • The general rule for solubility is: like dissolves like. • It implies that substances with intermolecular forces of the same strength are totally miscible – they form homogeneous solutions. • If the forces are of different strengths, the dissolution process cannot take place. • Non-polar solutes normally dissolve easily in non-polar solvents. • Polar and ionic solutes dissolve easily in polar solvents. • Ionic substances dissociate in water – break up into ions, e.g. NaCℓ(s) in H2O(ℓ). • NaCℓ(s) →┴(𝐻_2 𝑂("ℓ" )) Na+(aq) + Cℓ-(aq) • Polar substances ionise in water – form ions when they come in contact with water, e.g. HCℓ(g) in H2O(ℓ) • HCℓ(g) + H2O(ℓ) ⟶ H3O+(aq) + Cℓ-(aq) • It should be understood that there are exceptions to these rules, e.g. sugar (C12H22O11) is a non-polar substance, but is soluble in water. The large sugar molecules bond with strong London forces, that are of the same strength as the dipole-dipole forces between the water molecules. • Dipole induced-dipole interactions between sugar and water molecules explain the dissolution process.

16. Physical properties and intermolecular forces • The meniscus of a liquid can be concave or convex. • This phenomenon can also be explained in terms of intermolecular forces. • Water has a concave meniscus. The adhesion forces between the liquid and the glass are stronger than the cohesion force between the molecules themselves. • Mercury has a convex meniscus. Adhesion forces are weaker than cohesion forces. • Another important physical property of a liquid that can be explained in terms of intermolecular forces, is surface tension. • It can be defined as the resistance that a liquid offers to a force which tries to increase the surface area. • Molecule B in the diagram, experiences equal cohesion forces in ALL directions. • There is no net force acting on molecule B. • Molecule A experiences a net cohesion force downward, resulting in tension on the surface. • Liquids can rise against the walls of thin tubes – this phenomenon is known as capillary action. • Capillary action allows water to be absorbed by the roots and stems of plants.

17. The Experiment Description Diagram 1: Table of substances used in the 5 experiments • In this experiment you are going to compare the effect of intermolecular forces on the following physical properties of matter. • Vapour pressure • Surface tension • Solubility • Boiling point of liquids • Capillary action

18. Keywords