Chapter 11 : Liquids, Solids, and Intermolecular Forces

1. Chapter 11 : Liquids, Solids, and Intermolecular Forces • Outline • Climbing Geckos !?! • Solids, Liquids, and Gases • Intermolecular forces • Surface tension, viscosity and capillary action • Water • Crystalline Solids

2. Climbing Geckos A gecko can adhere to a glass ceiling with a single toe ! Given that an average, full grown gecko weighs 100g calculate the force of adhesion it must apply to the glass to do so ? Climbing Geckos

3. Solids, Liquids, and Gases Solids, Liquids, and Gases

4. Solids • Compared to gases, solids are dense • Have a definite shape and volume Liquids • Compared to gases, solids are dense • Have an indefinite shape and volume Solids, Liquids, and Gases

5. Phase Changes Liquid propane is a good example of a pressure-induced phase change. Solids, Liquids, and Gases

6. Phase Changes Remember : A phase change is a physical change. If this picture represents H2O in the liquid state…. Solids, Liquids, and Gases

7. When a molecule in a liquid state undergoes a phase change to a gas, it must break all the intermolecular forces acting upon it. Solids, Liquids, and Gases

8. Intermolecular Forces Intramolecular forces are interactions of atoms within a molecule. Intermolecular forces are interactions between molecules. Among other things, intermolecular forces are responsible for condensation of gases, and freezing of liquids into solids. Examples of intermolecular forces in liquids; • Surface tension • Viscosity • Vapour pressure Intermolecular Forces

9. δ- δ- δ+ δ+ Coulomb’s Law We’ve seen this before when we discussed bonding. In the case of bonding, q = 1.6 x 10-23 C. But we know that molecules can have partial charges if they have polar bonds. The interactions between these partial charges are the basis of intermolecular forces. Intermolecular Forces

10. Because we are not talking about full charges, q < 1.6 x 10-23 C. What does this tell you about the strength of intermolecular forces compared to bonding forces ? Intermolecular forces also act over much larger distances compared to electron-proton interactions. Intermolecular Forces

11. Types of IM Forces There are four categories of IM forces that we will consider. • Dispersion Forces • Dipole-Dipole Forces • Hydrogen Bonding • Ion-Dipole Forces Intermolecular Forces

12. Dispersion (aka London) Forces He is a noble gas and exists as a free atom in the gas state at STP. If there are no bonds, there can be no permanent dipoles. Therefore no partial charges. If partial charges lead to IM forces and IM forces are needed to condense gases into liquids, does that mean you can’t liquefy Helium ? Intermolecular Forces

13. At any instant, the electrons of He (or any other atom) may not be symmetrically distributed about the nucleus. At such an instant, the left side of He will have a slight negative charge whereas the right side will have a slight positive charge. Intermolecular Forces

14. Instantaneous dipoles create induced dipoles in nearby atoms. The forces that pull molecules together due to instantaneous and induced dipoles are called dispersion or London forces. The magnitude of dispersion forces depends on the ability of an atom’s (or molecule’s) electron cloud to be distorted. nucleus Intermolecular Forces

15. Simple Example Match the following noble gases and their boiling points. He 4.2 K Ar 165 K Xe 87 K All molecules are comprised of electrons therefore ALL molecules have dispersion forces. Intermolecular Forces

16. So is it as simple as bigger atoms (or molecules) have larger dispersion forces ? The elongated shape of n-pentane means there is more area of the molecule available to interact with other molecules. Intermolecular Forces

17. Intermolecular Forces

18. Dipole-Dipole Forces Recall our discussion on polar covalent bonds. These partial charges (resulting from the permanent dipole moment) cause the molecules to order themselves. This is an electrostatic interaction. The larger the dipole moment, the larger the partial charges and the greater the dipole-dipole interaction. Intermolecular Forces

19. Which of the following molecules would be expected to have the larger dipole-dipole forces ? Intermolecular Forces

20. Intermolecular Forces

21. Example Which of the following molecules will have dipole-dipole forces ? • CO2 • CH2Cl2 • CH4 Intermolecular Forces

22. Hydrogen Bonding We expect b.p. to increase with increasing molecular mass. What’s going on with HF, NH3, and H2O ? Intermolecular Forces

23. H F When hydrogen is covalently bonded to a very electronegative atom an extremely polar bond forms. Consequently, the hydrogen atom is simultaneously attracted to the electronegative atom (O,N,F) of a neighbouring molecule. An example of a hydrogen bonding network. In the gas state, HF exists in a cyclic “supramolecular” structure consisting of six HF molecules linked with h-bonds. Intermolecular Forces

24. Hydrogen Bonding in Water Crystal structure of ice Four h-bonds per molecules “Open” structure Structure of liquid water Some h-bonds broken “Denser” structure Intermolecular Forces

25. Why Ice Floats Intermolecular Forces

26. Intramolecular and Intermolecular H-Bonding Intermolecular h-bonding occurs between molecules. Intramolecular h-bonding occurs between groups on the same molecules. Intermolecular Forces

27. How Strong are Intermolecular Forces ? Ionic bond Covalent bond Ion-Dipole Hydrogen bonding Perm. Dipole – Perm. Dipole Induced dipole forces Intermolecular Forces

28. Examples 2. Arrange the following in order of increasing boiling point: He, Ne, O3, O2, Cl2, (CH3)2C=O 3. Arrange the following in order of increasing boiling point: C6H6 ; C6H5Cl ; C6H5Br ; C6H5OH Intermolecular Forces

29. Surface Tension Molecules at the surface of a liquid Molecules in the interiour of a liquid Because of the increased number of interactions, molecules in the bulk of solution are at a lower energy state than those on the surface. Surface Tension, Viscosity, Capillary Action

30. To increase the area of the interface, molecules most be moved from the interiour of the liquid to the interface. This requires work be done…NOT energetically favourable. Surface tension is the work (or energy) required to increase the area of a liquid Mathematically : Why are raindrops spherical and not cubic in shape ? Surface Tension, Viscosity, Capillary Action

31. Another example. How does a steel nail float on water ? Surface Tension, Viscosity, Capillary Action

32. Wetting Cohesive forces are intermolecular forces between molecules. Adhesive forces are intermolecular forces between molecules. When a liquid spreads into a film on a surface we say the liquid wets the surface. This will happen if the A.F. > C.F. Surface Tension, Viscosity, Capillary Action

33. Meniscus formation is another example of wetting phenomena. Surface Tension, Viscosity, Capillary Action

34. Viscosity Viscosity describes resistance to flow. Liquids such as honey and ethylene glycol have high viscosity while ethanol and water have low resistance to flow. Strong intermolecular forces cause high viscosities. Unit of viscosity is the Poise, (P), = g cm-1 s-1 Surface Tension, Viscosity, Capillary Action

35. As temperature increases, the intermolecular forces are less pronounced and the viscosity decreases. But shape matters too ! Surface Tension, Viscosity, Capillary Action

36. Vaporization and Vapor Pressure • molecules in the liquid are constantly in motion • Avg. KE is proportional to temperature • But the KE is distributed over a range of values Surface Tension, Viscosity, Capillary Action

37. if these molecules are at the surface, they can escape from the liquid and become a vapor • At a given temperature, the larger the surface area, the faster the rate of evaporation • As the temperature increases, a greater fraction of water molecules have this minimum energy  faster rate of evaporation • The stronger the IMFs, the slower the rate of vaporization

38. What happens to the molecules that have escaped ? • some molecules of the vapor will lose energy through molecular collisions • the result will be that some of the molecules will get captured back into the liquid when they collide with it • also some may stick and gather together to form droplets of liquid • particularly on surrounding surfaces

39. Evaporation vs. Condensation • vaporization and condensation are opposite processes • in an open container, the vapor molecules generally spread out faster than they can condense. Net result is a loss of liquid. • however, in a closed container, the vapor is not allowed to spread out indefinitely • the net result in a closed container is that at some time the rates of vaporization and condensation will be equal

40. Energetics of Vaporization • when the high energy molecules are lost from the liquid, it lowers the average kinetic energy • if energy is not drawn back into the liquid, its temperature will decrease – therefore, • vaporization requires input of energy to overcome the attractions between molecules

41. Heat of Vaporization • the amount of heat energy required to vaporize one mole of the liquid is called the Heat of Vaporization, DHvap • sometimes called the enthalpy of vaporization • always endothermic, therefore DHvap is + • somewhat temperature dependent

42. Dynamic Equilibrium Dynamic Equilibrium is when the rate of vaporization equals the rate of condensation

43. the pressure exerted by the vapor when it is in dynamic equilibrium with its liquid is called the vapor pressure • the weaker the attractive forces between the molecules, the more molecules will be in the vapor • the higher the vapor pressure, the more volatile the liquid

44. Boiling Point • when the temperature of a liquid reaches a point where its vapor pressure is the same as the external pressure, vapor bubbles can form anywhere in the liquid • this phenomenon is what is called boiling and the temperature required to have the vapor pressure = external pressure is the boiling point At the summit of Mt. Everest, water boils at 78oC. Why ?

45. Heating Curve of a Liquid • as you heat a liquid, its temperature increases linearly until it reaches the boiling point • q = mass xCsxDT • once the temperature reaches the boiling point, all the added heat goes into boiling the liquid – the temperature stays constant

46. the graph of vapor pressure vs. temperature is an exponential growth curve

48. Melting = Fusion • as a solid is heated, its temperature rises and the molecules vibrate more vigorously • once the temperature reaches the melting point, the molecules have sufficient energy to overcome some of the attractions that hold them in position and the solid melts (or fuses) • the opposite of melting is freezing

49. Energetics of Melting • when the high energy molecules are lost from the solid, it lowers the average kinetic energy • if energy is not drawn back into the solid its temperature will decrease • and freezing is an exothermic process • melting requires input of energy to overcome the attractions between molecules • the amount of heat energy required to melt one mole of the solid is called the • sometimes called the enthalpy of fusion

50. Sublimation and Deposition • Similar to the process in a liquid, molecules in a solid can break free from the surface and become a gas – this process is called sublimation • the capturing of vapor molecules into a solid is called deposition • the solid and vapor phases exist in dynamic equilibrium in a closed container • at temperatures below the melting point • therefore,molecular solids have a vapor pressure