Chapter 11 Intermolecular Forces, Liquids, and Solids

1. Chapter 11Intermolecular Forces, Liquids, and Solids

2. States of Matter The fundamental difference between states of matter is the distance between particles.

3. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

4. The States of Matter • The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: • the kinetic energy of the particles; • the attractive strength between the particles.

5. What are Intermolecular Forces? The intermolecular attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

6. Why care about Intermolecular Forces? They are strong enough to impact physical properties: • boiling and melting points • vapor pressures • viscosities. • READ p. 428

7. Names of Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces.

8. van der Waals Forces 1. Dipole-dipole interactions 2. Hydrogen bonding 3. London dispersion forces And maybe …. 4. Ion – Dipole Interaction

9. #4. Ion-Dipole Interactions • Ion-dipole interactions (a fourth type of force), are important in solutions of ions. • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

10. #1 Dipole-Dipole Interactions • Molecules that have permanent dipoles are attracted to each other. • The positive end of one is attracted to the negative end of the other and vice-versa. • These forces are only important when the molecules are close to each other.

11. Dipole-Dipole Example Attractive D-D Forces Attractive + Repulsive D-D Forces

12. Dipole-Dipole Interactions Turn to page 431 – Look at the molecules above. The more polar the molecule, the greater the strength of the dipole attractions, the higher the boiling point.

13. #2 London Dispersion Forces 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 happen that they occasionally wind up on the same side of the atom.

14. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side. Electron Traffic Jam! - termed Instantaneous Dipole Moment

15. London Dispersion Forces Another helium nearby would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in cloud on helium atom 1. Traffic Pile Up! - termed induced dipole moment

16. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and another induced dipole.

17. London Dispersion Forces • These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called polarizability. • Is the electron cloud “squashable”? • READ p. 429

18. Dispersion Forces Trends • Generally, the greater the polarizability, the more easily the electron cloud can be distorted (squashed) to produce an instant dipole. • Generally, polarizability increases as the number of electrons increases. • Strength of dispersion forces increases with increasing size or molecular mass.

19. Shape Influences London Forces • The shape of the molecule affects strength of dispersion forces: • - long, skinny ones have stronger forces • - short, fat ones have weaker forces • This is due to the increased surface area so molecules can come into greater contact with one another in longer molecules.

20. Factors Affecting London Forces • The strength of dispersion forces tends to increase with increased molecular weight. • Larger atoms have larger electron clouds which are easier to polarize or “squash”.

21. Which Have a Greater Effect?Dipole-Dipole Interactions or Dispersion Forces • If two molecules are of comparable size and shape, • If one molecule is much larger than another, dipole-dipole interactions will likely the dominating force. . disperson forces will likely determine its physical properties.

22. How Do We Explain This? • Read Go Figure! p. 431 • The nonpolar series (SnH4 to CH4) follow the expected trend. • The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

23. #3. Hydrogen Bonding • The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. • We call these interactions hydrogen bonds.

24. Hydrogen Bonding • Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

25. Importance of H Bonding • Stabilize Protein Structures • Responsible for the way that DNA is able to carry genetic code

26. Importance of H Bonding Read p. 433: Be able to express the relationship between H bonds and floating ice!

27. Summarizing Intermolecular Forces

28. Strengths of Forces

29. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

30. Viscosity • Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

31. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

32. Cohesive vs. Adhesive Forces • Forces that bind similar molecules together = cohesive • Forces that bind a substance to a surface = adhesive

33. Water vs. Mercury Turn to page 438

34. Phase Changes

36. Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid.

37. Energy Changes Associated with Changes of State The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.

38. Energy Changes Associated with Changes of State • The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. • The temperature of the substance does not rise during a phase change. • Energy is used to overcome attractive forces

39. Sample Exercise 11.3 p. 441

40. Vapor Pressure • At any temperature some molecules in a liquid have enough energy to escape. • As the temperature rises, the fraction of molecules that have enough energy to escape increases.

41. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

42. Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condenseat the same rate.

43. Vapor Pressure • The boiling point of a liquid is the temperature at which it’s vapor pressure equals atmospheric pressure. • The normal boiling point is the temperature at which its vapor pressure is 760 torr.

44. Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

45. Phase Diagrams • The circled line is the liquid-vapor interface. • It starts at the triple point (T), the point at which all three states are in equilibrium.

46. Phase Diagrams It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

47. Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

48. Phase Diagrams • The circled line in the diagram below is the interface between liquid and solid. • The melting point at each pressure can be found along this line.

49. Phase Diagrams • Below the triple point the substance cannot exist in the liquid state. • Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

50. Phase Diagram of Water • Note the high critical temperature and pressure. • These are due to the strong van der Waals forces between water molecules.