1. INTERMOLECULAR FORCES LIQUIDS AND SOLIDS Chapter 11

2. STATES OF MATTER

3. Weak attractive forces between molecules Intermolecular forces stronger Strong intermolecular forces The state of a substance depends largely on the balance between the kinetic energies of the particles and the interparticle energies of attraction.

4. INTERMOLECULAR FORCES intramolecular force - covalent bond between atoms in a molecule When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).  Many properties are governed by the strength of the intermolecular forces e.g. boiling point, melting point, vapor pressure, viscosity, etc. intermolecular force - attraction between molecules Strong 431 kJ/mol weak 16 kJ/mol (can also be ionic bond) HCl boils at -85 o C at 1 atm

5. A liquid boils when: A solid melts when:

6. 4 types of intermolecular forces: Ion-Dipole Forces Dipole-Dipole Forces London Dispersion Forces Hydrogen Bonding electrostatic forces Can you arrange them in order of increasing strength? What is a dipole? + - van der Waals forces

7. Ion-Dipole Forces Interaction between an ion and a dipole Strongest of all intermolecular forces. Ion-dipole interactions make it possible for ionic substances to dissolve in polar solvents.

8. Dipole-Dipole Forces Interaction between two dipoles Weaker than ion-dipole forces. Polar molecules need to be close together. Molecules in liquids are free to move  results in both attractive and repulsive forces.  overall and stronger attractive force

9. If two molecules have about the same mass and size, then intermolecular attractions (dipole-dipole forces) increase with increasing polarity.

10. London Dispersion Forces Interaction between nonpolar atoms or molecules, but can also occur between polar molecules Occur only when atoms or molecules are close together Weakest of all intermolecular forces. Instantaneous dipoles formed! While the electrons in the 1s orbital of helium would repel each other (and therefore tend to stay far away from each other), it does occasionally happen that they wind up on the same side of the atom. These dipoles are temporary

11. Polarisability: the ease with which the distortion of the charge distribution occurs. Greater polarisablility  stronger dispersion In general, larger molecules tend to have greater polarisability as they have more electrons which are further away from the nuclei.

12. Exercise Explain why n-pentane has a higher boiling point than neopentane? Both have the molecular formula C 5 H 12.

13. Hydrogen Bonding Boiling pints Special case of dipole-dipole forces nonpolar polar

14. Hydrogen bonds are dipole-dipole interactions between the H-atom in a polar bond (usually H-F, H-O or H-N) and an unshaired e - pair on a nearby small electronegative ion or atom (usually F, O, N) Hydrogen bond

15. Small and electronegative!Why are we specifying F, O and N? Hydrogen only has 1e -  +ve nucleus is rather exposed. Hydrogen can approach the small electronegative atom closely and interact strongly.

16. Exercise Why does ice float on water?

17. Exercise Which of the following molecules can hydrogen bond with itself? CH 2 F 2 NH 3 CH 3 OH H 3 C– C–CH 3 O

18. SUMMARY

19. PROPERTIES OF LIQUIDS Viscosity The resistance of a liquid to flow Which intermolecular forces also increase with molecular weight? - related to the ease with which molecules can move past each other Viscosity increases with molecular weight and decreases with higher temperature.

20. Surface tension A measure of the inward forces that must be overcome in order to expand the surface area of a liquid Surface tension of water at 20 o C: 0.0729 J/m 2 Molecules in the interior are attracted equally in all directions. Surface molecules experience a net inward force. Molecules at the surface can pack more closely together

21. PHASE CHANGES Every phase change is accompanied by a change in energy of the system. (fusion)  H fus  H (change in enthalpy)  H vap  H cond  H freez  H sub  H depos

22.  H fus  H vap  H vap >  H fus  In the transition from liquid to vapour phase, the molecules must essentially sever all intermolecular interactions.  In melting, many of these interactions remain.  H sub =

23. Plot of temperature versus heat added is a heating curve e.g. Ice initially at -25 o C is heated (constant P = 1 atm) While melting, heat added is used to break intermolecular forces While evaporating, heat added is used to break intermolecular forces From the graph determine:  H fus  H vap Recall:  H vap >  H fus  H AB  H CD  H EF ?

24. -30 o C 0oC0oC 50 o C D CB A water ice Exercise Calculate the enthalpy change, ΔH, when 100 g of water at 50 o C is cooled down to ice at -30 o C. Given: Specific heat capacities: Water = 4.18 J/g K, Ice = 2.09 J/g K Δ H fus = 6.01 kJ/mol  H AB = (4.18 J/g K)(100 g)(0 - 50 K) = -20.9 kJ  T = T f - T i  H BC = ΔH freez = -ΔH fus = -(6.01 kJ/mol)(5.55 mol) = -33.4 kJ n = (100 g)/(18.016 g/mol) n = 5.55 mol  H CD = (2.09 J/g K)(100 g)(-30 - 0 K) = -6.27 kJ  H AD =  H AB +  H BC +  H CD = -60.6 kJ  H AB  H CD  H BC =  H freez  H = Cm  T What is the enthalpy change when 100 g of ice at -30 o C is heated up to water at 50 o C?

25. Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid  these molecules move into the gas phase. VAPOUR PRESSURE As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. After some time the pressure of the gas will be constant at the vapour pressure  when liquid and vapour reach dynamic equilibrium. Closed container

26. As the temperature increases, the fraction of molecules that have enough energy to escape increases.

27. The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. Normal boiling point  boiling point at 1 atm. Why does water boil at a higher temperature at the coast than here in Jhb?

28. PHASE DIAGRAMS Phase diagrams display the state of a substance under various pressure and temperature conditions. The solid lines show the conditions P,T conditions under which equilibrium exists between phases.

29. Line AD: solid -liquid interface - Each point along this line is the melting point of the substance at that pressure. Critical point B: above this critical temperature and critical pressure the liquid and vapour are indistinguishable from each other. Line AB: liquid-vapour interface - Each point along this line is the boiling point of the substance at that pressure.

30. Line AC: solid-vapour interface - Each point along this line is the sublimation point of the substance at that pressure. - Note: the substance cannot exist in the liquid state below A Triple point A: the temperature and pressure condition at which all three states are in equilibrium.

31. The slope of the solid–liquid line is negative.  The melting point decreases with increasing pressure. Phase Diagram of Water Note the high critical temperature and critical pressure:  due to the strong van der Waals forces between water molecules. WHY?

32. Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm. CO 2 sublimes at normal pressures. Phase Diagram of Carbon Dioxide

33. BONDING IN SOLIDS Solids can be: crystalline Particles are in highly ordered arrangement amorphous No particular order in the arrangement of particles. or

34. The physical properties of crystalline solids (e.g. m.p., hardness) depend on the arrangement of particles and on attractive forces between particles. There are 4 types of crystalline solids: Molecular Covalent network Ionic Metallic

35. Most substances that are gases or liquids at room temperature form molecular solids at low temperature. Atoms or molecules are held together by intermolecular forces  dipole-dipole forces, London dispersion forces, hydrogen bonds Molecular Solids Because of these weak forces they are soft and have relatively low melting points (<200 o C) Benzene Toluene Phenol m.p./ o C 5 -95 43 b.p./ o C 80 111 182 Explain the m.p.’s and b.p.’s observed below:

36. DiamondGraphite Which one is harder and has the higher m.p.? Explain. Covalent-Network Solids Atoms are held together in large networks or chains by covalent bonds. Covalent bonds much stronger than intermolecular forces  Harder solids and higher melting points than molecular solids.

37. Ionic Solids Ions are held together by ionic bonds  strength of the ionic bond depends on the charges of the ions Ions pack themselves so as to maximize the attractions and minimize repulsions between the ions  Depends on relative size and charge of ions

38. Metallic Solids Consist of entirely metal atoms. Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. Bonding due to valence electrons delocalized throughout the solid. In general, the strength of bonding increases as the no. of electrons available for bonding increases The m.p. for sodium is 97.5 o C and for chromium is 1890 o C. Explain.

39. SUMMARY