Intermolecular Forces, Liquids and Solids

1. Intermolecular Forces, Liquids and Solids CHAPTER 11CHEM 160

2. A Molecular Comparison of Liquids and Solids • Physical properties of substances understood in terms of kinetic molecular theory: • Gases are highly compressible, assumes shape and volume of container: • Gas molecules are far apart and do not interact much with each other. • Liquids are almost incompressible, assume the shape but not the volume of container: • Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other.

3. A Molecular Comparison of Liquids and Solids • Solids are incompressible and have a definite shape and volume: • Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.

4. A Molecular Comparison of Liquids and Solids

5. A Molecular Comparison of Liquids and Solids

6. A Molecular Comparison of Liquids and Solids • Converting a gas into a liquid or solid requires the molecules to get closer to each other: • cool or compress. • Converting a solid into a liquid or gas requires the molecules to move further apart: • heat or reduce pressure. • The forces holding solids and liquids together are called intermolecular forces.

7. Intermolecular Forces • The covalent bond holding a molecule together is an intramolecular forces. • The attraction between molecules is an intermolecular force. • Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl). • When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

8. Intermolecular Forces

9. Intermolecular Forces • Ion-Dipole Forces • Interaction between an ion and a dipole (e.g. water). • Strongest of all intermolecular forces.

10. Intermolecular Forces • Dipole-Dipole Forces • Dipole-dipole forces exist between neutral polar molecules. • Polar molecules need to be close together. • Weaker than ion-dipole forces. • There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. • If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.

11. Intermolecular Forces Dipole-Dipole Forces

12. Intermolecular Forces Dipole-Dipole Forces

13. Intermolecular Forces • London Dispersion Forces • Weakest of all intermolecular forces. • It is possible for two adjacent neutral molecules to affect each other. • The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). • For an instant, the electron clouds become distorted. • In that instant a dipole is formed (called an instantaneous dipole).

14. Intermolecular Forces • London Dispersion Forces • One instantaneous dipole can induce another instantaneous dipole in an adjacent molecule (or atom). • The forces between instantaneous dipoles are called London dispersion forces.

15. Intermolecular Forces • London Dispersion Forces • Polarizability is the ease with which an electron cloud can be deformed. • The larger the molecule (the greater the number of electrons) the more polarizable. • London dispersion forces increase as molecular weight increases. • London dispersion forces exist between all molecules. • London dispersion forces depend on the shape of the molecule.

16. Intermolecular Forces • London Dispersion Forces • The greater the surface area available for contact, the greater the dispersion forces. • London dispersion forces between spherical molecules are lower than between sausage-like molecules.

17. Intermolecular Forces London Dispersion Forces

18. Intermolecular Forces London Dispersion Forces

19. Intermolecular Forces • Hydrogen Bonding • Special case of dipole-dipole forces. • By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. • Intermolecular forces are abnormally strong.

20. Intermolecular Forces • Hydrogen Bonding • H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N). • Electrons in the H-X (X = electronegative element) lie much closer to X than H. • H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X. • Therefore, H-bonds are strong.

21. Hydrogen Bonding

22. Hydrogen Bonding

23. Intermolecular Forces • Hydrogen Bonding • Hydrogen bonds are responsible for: • Ice Floating • Solids are usually more closely packed than liquids; • Therefore, solids are more dense than liquids. • Ice is ordered with an open structure to optimize H-bonding. • Therefore, ice is less dense than water. • In water the H-O bond length is 1.0 Å. • The O…H hydrogen bond length is 1.8 Å. • Ice has waters arranged in an open, regular hexagon. • Each + H points towards a lone pair on O.

24. Intermolecular Forces Hydrogen Bonding

25. Intermolecular Forces

26. Some Properties of Liquids • Viscosity • Viscosity is the resistance of a liquid to flow. • A liquid flows by sliding molecules over each other. • The stronger the intermolecular forces, the higher the viscosity. • Surface Tension • Bulk molecules (those in the liquid) are equally attracted to their neighbors.

27. Some Properties of Liquids Viscosity

28. Surface Tension

29. Some Properties of Liquids • Surface Tension • Surface molecules are only attracted inwards towards the bulk molecules. • Therefore, surface molecules are packed more closely than bulk molecules. • Surface tension is the amount of energy required to increase the surface area of a liquid. • Cohesive forcesbind molecules to each other. • Adhesive forcesbind molecules to a surface.

30. Some Properties of Liquids • Surface Tension • Meniscusis the shape of the liquid surface. • If adhesive forces are greater than cohesive forces, the liquid surface is attracted to its container more than the bulk molecules. Therefore, the meniscus is U-shaped (e.g. water in glass). • If cohesive forces are greater than adhesive forces, the meniscus is curved downwards. • Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube.

31. Phase Changes • Surface molecules are only attracted inwards towards the bulk molecules. • Sublimation: solid  gas. • Vaporization: liquid  gas. • Melting or fusion: solid  liquid. • Deposition: gas  solid. • Condensation: gas  liquid. • Freezing: liquid  solid.

32. Phase Changes

33. Phase Changes • Energy Changes Accompanying Phase Changes • Sublimation: Hsub > 0 (endothermic). • Vaporization: Hvap > 0 (endothermic). • Melting or Fusion: Hfus > 0 (endothermic). • Deposition: Hdep < 0 (exothermic). • Condensation: Hcon < 0 (exothermic). • Freezing: Hfre < 0 (exothermic).

34. Phase Changes • Energy Changes Accompanying Phase Changes • Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: • it takes more energy to completely separate molecules, than partially separate them.

35. Phase Changes

36. Phase Changes • Energy Changes Accompanying Phase Changes • All phase changes are possible under the right conditions. • The sequence • heat solid  melt  heat liquid  boil  heat gas • is endothermic. • The sequence • cool gas  condense  cool liquid  freeze  cool solid • is exothermic.

37. Phase Changes • Heating Curves • Plot of temperature change versus heat added is a heating curve. • During a phase change, adding heat causes no temperature change. • These points are used to calculate Hfus and Hvap. • Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. • Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces.

39. Phase Changes • Critical Temperature and Pressure • Gases liquefied by increasing pressure at some temperature. • Critical temperature: the minimum temperature for liquefaction of a gas using pressure. • Critical pressure: pressure required for liquefaction.

40. Phase Changes Critical Temperature and Pressure