12.5 Van der Waals Forces

1. June 12, 2009 – Class 39 and 40 Overview • 12.5 Van der Waals Forces • Dispersion (London) forces, polarizability, dipole-dipole interactions, permanent, instantaneous and induced dipoles, molecular shape and polarizability, predicting relative boiling points from relative strength of intermolecular forces. • 12.6 Hydrogen Bonding • Properties of hydrogen bonds (strong vs weak hydrogen bonds), hydrogen bonding in water and living matter, intermolecular and intramolecular hydrogen bonding, effect on viscosity.

2. June 12, 2009 – Class 39 and 40 Overview • 12.1 Intermolecular Forces and Some Properties of Liquids • 12.2 Vaporization of Liquids: Vapour Pressure • Vaporization(evaporation), enthalpy of vaporization, vapor pressure curves, boiling and the boiling point, normal boiling point, measuring vapor pressure data, use of the Clausius-Clapeyron equation. • 12.3 Some Properties of Solids • Melting, melting (freezing) point, heat of fusion, sublimation, deposition.

3. Intermolecular Forces and Some Properties of Liquids • Molecules at the surface interact with other surface molecules and with molecules directly below the surface. • Molecules in the interior experience intermolecular interactions with neighboring molecules in all directions. • Intramolecular Forces: attraction within a molecule. A covalent bond is an example of an intramolecular force. • Intermolecular Forces: an attraction between molecules. • Responsible for condenses phases (solids and liquids). • Intermolecular forces explain many properties of liquids.

4. Intermolecular forces • There are three primary types of intermolecular forces: • (London) Dispersion Forces 2. Dipole-Dipole Forces (van der Waal's Forces) • Hydrogen Bonding • Note: polarizability: the ease with which a particle’s electron cloud can be distorted.

5. Dispersion (London) forces • Dispersion (London) forces: Intermolecular forces associated with instantaneous and induced dipoles. • Electron location is expressed in terms of probabilities only • For non-polar molecules it is most probable that the electrons will be evenly distributed in the atomic or molecular orbitals. • It is possible for the electrons in one molecule to flicker into an arrangement that results in partial positive (d+) and partial negative (d-) charges. • When this occurs the molecules acquire an instantaneous (it lasts less than 10-16 seconds!) dipole. • A molecule that has acquired an instantaneous dipole can then induce a dipole in another molecule.

6. Dispersion (London) forces • Normal Condition: A non-polar molecule has a symmetrical charge distribution • b) Instantaneous Condition: A displacement of the electronic charge produces an instantaneous dipole with a charge separation represented as d+ and d-. • c) Induced Dipole: The instantaneous dipole on the left induces a charge separation in the molecule on the right. The result is a dipole-dipole interaction. • The two dipoles, in the two molecules, will attract each other, and the result is that the potential energy of the two is lowered.

7. Dispersion (London) forces • Dispersion forces are present in both polar and non-polar molecules, however, it is the main type of intermolecular force between non-polar molecules. • There are two factors that effect the magnitude of dispersion forces: 1. Number of electrons: Larger molecules with more electrons more easily undergo vibrations that lead to uneven distribution of charge. Molecules with more electrons will have stronger dispersion forces.

8. Dispersion (London) forces • There are two factors that effect the magnitude of dispersion forces: 2. Surface area: Molecules with a larger surface area offer a greater opportunity for a molecule to induce a dipole in a nearby molecule. Spherical molecules with the same number of electrons as more branched molecules will have weaker dispersion forces.

9. Dispersion (London) forces • Example: F2 has 18 electrons and Cl2 has 34 electrons. The dispersion forces for Cl2 are stronger than those of F2. • Example: Cl2 and C4H10 each have 34 electrons. C4H10 has a larger, more complex shape therefore it has stronger dispersion forces than Cl2. • When a substance melts or boils, the intermolecular forces are overcome. • Molecules with greater dispersion forces will have higher boiling points because more energy is required to overcome the attraction between molecules.

10. Dipole-Dipole Forces • Dipole-Dipole Forces: intermolecular attractions associated with molecules with permanent dipoles. • Recall that the bond dipoles for molecules do not always cancel. When the bond dipoles do not cancel the resulting molecule is polar (it has a permanent dipole). • This leads to polar molecules trying to line up with the positive of one dipole directed toward the negative end of neighbouring dipoles.

11. Dipole-Dipole Forces Dipole-dipole forces involve the displacement of electrons in bonds, rather than the displacement of all the electrons in a molecule, as in dispersion forces. Dipole-dipole forces add to dispersion forces (which are present for all molecules). Dipole-dipole forces also affect physical properties such as melting and boiling point.

12. Dipole-Dipole Forces Example: The boiling point of N2(l) is -195.81 oC. The boiling point of O2(l) is -182.96 oC. If only dispersion forces are considered, one would predict the boiling point of NO(l) to be in between that of O2(l) and N2(l) . The actual boiling point of NO(l) is -151.76 oC, much higher than either that of O2(l) or N2(l) . NO is a polar molecule, hence it has both dispersion forces and dipole-dipole forces present. It has a higher boiling point than O2(l) or N2(l) because extra energy is required to overcome the dipole-dipole interactions.

13. Dispersion (London) forces Problem: Which would you expect to have the higher boiling point? • C4H10 or (CH3)2CO • C3H8, CO2, CH3CN

14. Hydrogen Bonding • Hydrogen Bonding: an intermolecular force of attraction in which an H-atom covalently bonded to one highly electronegative atom is simultaneously attracted to another highly electronegative atom of the same or a nearby molecule. • The only atoms that are electronegative enough to participate in hydrogen bonding are fluorine, nitrogen and oxygen. • Hydrogen bonding, like dispersion forces and dipole-dipole forces, also affects physical properties, like melting and boiling points. • Notice the anomalies in the following graph of boiling points for hydrides of Groups 4A, 5A, 6A and 7A.

15. Hydrogen Bonding The values for NH3, H2O and HF are unusually high compared to those of other member of their groups!

16. Hydrogen Bonding • Consider the molecule HF as our example. • Fluorine is highly electronegative, and pulls the bonding pair of electrons closer to itself, leaving the hydrogen nucleus unshielded with a partial positive charge. • The hydrogen nucleus is then attracted to the lone pair of electrons on another highly electronegative atom. • Recall that the other highly electronegative atom will have a partial negative charge, acquired by attracting electrons involved in its molecular bonds.

17. Hydrogen Bonding Covalent Bond Hydrogen Bond • The result for gaseous HF is that it often forms cyclic (HF)6.

18. Hydrogen Bonding N-H Bond Example: • Urea is CO(NH2)2 and it undergoes extensive hydrogen bonding due to its four hydrogen atoms bonded directly to nitrogen, and also H-bonded to both nitrogen and oxygen. Several of these interaction are depicted below.

19. Hydrogen Bonding The molecules can still move around because they have enough kinetic energy to break the hydrogen bonds, which then reform with another H2O molecule. O-H Bond Example: • In liquid H2O each hydrogen atom is bonded to at least four other H2O molecules. • A few of these bonds are illustrated below.

20. Hydrogen Bonding This results in the molecules being organized into a crystalline arrangement, which is less dense than the arrangement of the H2O molecules in the liquid. • Hydrogen bonding in H2O also explains why ice floats! • To float a solid must be less dense than the liquid. • In the liquid there is enough kinetic energy to overcome some of the hydrogen bonds, but in the solid the kinetic energy is no longer sufficient to overcome the hydrogen bonds.

21. Comparing Melting and Boling Points • Classify your compound as ionic or molecular. Ionic compounds have greater intramolecular forces than do molecular compounds, hence they have higher melting and boiling points. • For your molecular compounds, list the intermolecular forces present. The more intermolecular forces, the higher the melting or boiling point for a compound.

22. Comparing Melting and Boling Points

23. Intermolecular Forces and Some Properties of Liquids Despite being more dense than water, a needle is supported on the surface due to surface tension. At 20 oC the surface tension of water is 7.28 x 10-2 J/m2. • Surface Tension (g): the energy or work required to extend the surface of a liquid, measured in J/m2. • Increasing temperature leads to more molecular motion, and less work is required to extend a surface. Therefore, surface tension decreases when temperature increases.

24. Intermolecular Forces and Some Properties of Liquids • Problem: What intermolecular forces account for the following surface tension data? substance (formula) surface tension / J m-2 at 20oC (a) pentane (C5H12) 1.6 x 10-2 (b) diethyl ether (C2H5OC2H5) 1.7 x 10-2 (c) ethanol (C2H5OH) 2.3 x 10-2 (d) butanol (C4H9OH) 2.5 x 10-2 (e) water (H2O) 7.3 x 10-2 (f) mercury (Hg) 48 x 10-2

25. Intermolecular Forces and Some Properties of Liquids (a) Water spreads into a thin film on clean glass; adhesive forces are greater than cohesive, and work done by a collapsing drop supplies energy for spreading. (b) Glass coated in grease or oil; drops of water stand on glass. Here, cohesive forces are greater than adhesive. • Cohesive Forces: intermolecular forces between like molecules, such as within a drop of liquid. • Adhesive Forces: intermolecular forces between unlike molecules, such as molecules of a liquid and of a surface with which it is in contact.

26. Intermolecular Forces and Some Properties of Liquids Adhesive forces lead to a concave meniscus for water (g = 7.28 x 10-2 J/m2). Cohesive Forces in mercury consist of strong metallic bonds (g = 47.2 x 10-2 J/m2) . It forms a convex menisucus. Adhesive forces (as evidenced by the level of the meniscus) are magnified in tubes of small diameter, known as capillary tubes.

27. Intermolecular Forces and Some Properties of Liquids Stronger intermolecular forces lead to slower flow and greater viscosity for a substance. Viscosity can be determined by measuring the velocity of a ball dropping through a liquid. • Viscosity: refers to a liquid’s resistance to flow. Its magnitude depends on intermolecular forces of attraction and in some cases, on molecular sizes and shapes. • When liquid flows, one portion of the liquid moves with respect to neighboring portions. Cohesive forces are responsible for the fast (ex. water or ethanol) or slow (ex. honey or motor oil) rate of flow.

28. Vaporization of Liquids: Vapour Pressure • Vaporization (evaporation): the passage of molecules from the liquid to the gaseous state. Intermolecular forces must be overcome in order for a sample to vaporize. Vaporization occurs more readily with: • Increased temperature: more molecules have sufficient kinetic energy to overcome intermolecular forces. • Increased surface area: a greater proportion of the molecules are at the surface of the liquid. • Decreased strength of intermolecular forces: less kinetic energy is required to overcome the forces of attraction.

29. Vaporization of Liquids Recall from Chapter 7 • Enthalpy of vaporization (DHvap): quantity of heat that must be absorbed in order to vaporize a liquid at constant temperature. • Enthalpy of condensation (DHcondensation) is equal in magnitude but opposite in sign to DHvap.

30. Vaporization of Liquids Problem: How much energy is released when 0.459 g of H2O (g) is cooled from 136 oC to 85 oC?

31. Vaporization of Liquids: Vapour Pressure Liquid vaporizes in a closed container. Condensation begins. Rate of vaporization > rate of condensation. Equilibrium. Rate of vaporization = rate of condensation. • Vapor pressure: the pressure exerted by a vapor when it is in dynamic equilibrium with its liquid at a fixed temperature.

32. Vaporization of Liquids: Vapour Pressure • Liquids with high vapor pressure at room temperature are volatile. • Liquids with high vapor pressure at room temperature are nonvolatile. • Vapor pressure increases with temperature. This is illustrated by vapor pressure curves. • Vapor pressure curves: a graph of vapor pressure as a function of temperature.

33. Vapour Pressure Curves • Diethyl ether, C4H10O • Benzene, C6H6 • Water, H2O • Toluene, C7H8 • Aniline, C6H7N • The normal boiling points are the temperatures at the intersection of the dashed line (P = 760 mmHg).

34. Vapour Pressure Curves

35. Vaporization of Liquids: Vapour Pressure • Boiling: a process in which vaporization occurs through a liquid. It occurs when the vapor pressure of a liquid is equal to barometric pressure. • Normal boiling point: the temperature at which the vapor pressure of a liquid is 1 atm. It is the temperature at which the liquid boils in a container open to the atmosphere at 1 atm. • Boiling points vary with barometric pressure. • At an altitude of 1609 m (Denver, Colorado), barometric pressure is 630 mmHg and the boiling point of water is 95 oC. It takes longer to cook food under these conditions!

36. Vaporization of Liquids: Vapour Pressure Vapor pressure plotted as ln P vs. 1/T (P measured in mmHg and T in Kelvin) – note that the relationships are linear.

37. Vaporization of Liquids: Vapour Pressure Where P1 and T1 are one pair of pressure and temperature conditions P2 and T2 are the other pair of pressure and temperature conditions. DHvap is the heat of vaporization R is the gas constant in units of 8.314 J mol−1K−1 • Clausius-Clapeyron equation: can be used to predict the temperature at a certain pressure, given the temperature at another pressure, or vice versa. Alternatively, if the corresponding temperature and pressure is known at two points, the enthalpy of vaporization can be determined.

38. Vaporization of Liquids: Clausius-Clapeyron Equation Problem: The normal boiling point of isooctane is 99.2 oC, and its enthalpy of vaporization is 35.76 kJ/mol C8H18. Calculate the vapor pressure of isooctane at 25 oC. Problem: The vapor pressure of methanol is 100 mmHg at 21.2 oC, and is 121 mmHg at 25 oC. What is the enthalpy of vaporization for methanol?

39. Some Properties of Solids • Melting: the transition of a solid to a liquid that occurs at the melting point. The melting point and freezing point of a substance are identical. • Heat of fusion (DHfus)

40. Some Properties of Solids Sublimation of solid iodine, and the deposition of the vapor to solid on the cooler walls of the flask. • Sublimation: the passage of molecules from the solid to the gaseous state. • Deposition: the passage of molecules from the gaseous to the solid state. • Enthalpy of sublimation: