Ch 11 States of Matter and Intermolecular Forces

1. Ch 11 States of Matter and Intermolecular Forces

2. Chapter 11 Preview • Intramolecular forces (bonds) govern molecular properties. • Intermolecular forces are between molecules and determine the macroscopic physical properties of liquids and solids. • This chapter: • describes changes from one state of matter to another. • explores the types of intermolecular forces that underlie these and other physical properties of substances.

3. Review:Intramolecular Forces • Ionic • Covalent • Metallic

4. Ionic bond- electron donated and accepted

5. Covalent- sharing electrons

6. Metallic-sea of electrons

7. Ionic Bonds as“Intermolecular” Forces • There are no molecules in an ionic solid, and therefore there can’t be any intermolecular forces. • These forces increase: • as the charges on the ions increase. • as the ionic radii (size) decrease.

8. Interionic Forces of Attraction Mg2+ and O2– have much stronger forces of attraction for one another than do Na+ and Cl–. Melting point of MgO is about 2800 oC. Melting point of NaCl is about 801 oC.

9. Molecular Forces Compared

10. States of Matter Compared Intermolecular forces are very important. Intermolecular forces are of little significance; why? Intermolecular forces must be considered.

11. Intermolecular Forces • Hydrogen bonding • Dipole-dipole • London Forces

12. Cohesion • Attraction for each other • Water • Mercury • Boiling point varies based on cohesion

13. Adhesion • A liquids attraction for solid particles • Water’s attraction for glass etc.

14. WATER MERCURY

15. Meniscus Formation What conclusion can we draw about the cohesive forces in mercury? Water wets the glass (adhesive forces) and its attraction for glass forms a concave-up surface.

16. Plant root- capillary action

17. Surface Tension • The forces present in liquids • A molecule inside the liquid experiences cohesive forces with other molecules in all directions. • A molecule at the surface experiences only downward cohesive forces

18. Surface Tension There is no force above the molecule on the top of the surface.

19. Pool ball floating on mercury

20. Adhesive and Cohesive Forces The liquid spreads, because adhesive forces are comparable in strength to cohesive forces. The liquid “beads up.” Which forces are stronger, adhesive or cohesive?

21. Cohesive vs. Adhesive 1 • Water on plastic • Water on metal • Water on glass 2 3

22. Intermolecular Forces • Hydrogen bonding • Dipole-dipole • London Forces

23. Hydrogen bonding • A hydrogen is attracted to a highly electronegative atom like O, N, or Cl.

24. Hydrogen Bonding in Water

25. Hydrogen Bonding in Ice Hydrogen bonding arranges the water molecules into an open hexagonal pattern. “Hexagonal” is reflected in the crystal structure. “Open” means reduced density of the solid (vs. liquid).

26. Hydrogen Bonding in Acetic Acid Hydrogen bonding occurs between molecules.

27. Intermolecular Hydrogen Bonds Intermolecular hydrogen bonds give proteins their secondary shape, forcing the protein molecules into particular orientations, like a folded sheet …

28. Intramolecular Hydrogen Bonds … while intramolecular hydrogen bonds can cause proteins to take a helical shape.

29. In which of these substances is hydrogen bonding an important intermolecular force: N2, HI, HF, CH3CHO, and CH3OH? Explain.

30. In which of these substances is hydrogen bonding an important intermolecular force: N2, HI, CH3CHO, and CH3OH? Explain. CH3CHO and CH3OH because of the attraction between the H of one molecule and the O of another. HI would not have hydrogen bonding b/c iodine is not highly electronegative.

31. Dipole–Dipole Forces • A polar molecule has a positively charged “end” (δ+) and a negatively charged “end” (δ–). • When molecules come close to one another, repulsions occur between like-charged regions of dipoles. Opposite charges tend to attract one another.

32. Dipole Forces • The more polar a molecule, the more pronounced is the effect of dipole–dipole forces on physical properties.

33. Opposites attract! Dipole–Dipole Interactions

34. London Forces-aka Dispersion Forces • At first no dipole, like Argon. • But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end -. The other end will be temporarily short of electrons and so becomes +. • An instant later the electrons may have moved up to the other end, reversing the polarity of the molecule.

35. Induced London Forces What would happen if we mixed HCl with the element argon, which has no dipole? • The electrons on an argon atom are distributed homogeneously around the nucleus of the atom. But these electrons are in constant motion. When an argon atom comes close to a polar HCl molecule, the electrons can shift to one side of the nucleus to produce a very small dipole moment that lasts for only an instant.

36. Dispersion Forces Illustrated (1) At a given instant, electron density, even in a nonpolar molecule like this one, is not perfectly uniform.

37. Dispersion Forces Illustrated (2) … the other end of the molecule is slightly (+). The region of (momentary) higher electron density attains a small (–) charge … When another nonpolar molecule approaches …

38. Dispersion Forces Illustrated (3) … this molecule induces a tiny dipole moment … … in this molecule. Opposite charges ________.

39. Types of forces • Click here for an Animation of the forces

40. Molecular Shape and Polarizability … can have greater separation of charge along its length. Stronger forces of attraction, meaning … Long skinny molecule … … higher boiling point. … giving weaker dispersion forces and a lower boiling point. In the compact isomer, less possible separation of charge …

41. Arrange the following substances in the expected order of increasing boiling point: Carbon tetrabromide, CBr4; Butane, CH3CH2CH2CH3; Fluorine, F2; Acetaldehyde, CH3CHO.

42. Arrange the following substances in the expected order of increasing boiling point: Carbon tetrabromide, CBr4; Butane, CH3CH2CH2CH3; Fluorine, F2; Acetaldehyde, CH3CHO. Answer: F2, CBr4, CH3CH2CH2CH3, CH3CHO

43. vaporization LiquidVapor condensation Vapor Pressure • The vapor pressure of a liquid is the partial pressure exerted by the vapor when it is in dynamic equilibrium with the liquid at a constant temperature.

44. Liquid–Vapor Equilibrium … until equilibrium is attained. More vapor forms; rate of condensation of that vapor increases …

45. Phase Diagrams A—D, solid-liquid equilibrium. A—C, liquid-vapor equilibrium. A phase diagram is a graphical representation of the conditions of temperature and pressure under which a substance exists as a solid, liquid, a gas, or some combination of these in equilibrium. A—B, solid-vapor equilibrium. Triple point

46. Phase Diagram for CO2 Note that at 1 atm, only the solid and vapor phases of CO2 exist.

47. Phase Diagram for H2O

48. Supercritical Fluid • Above the critical temperature and pressure, only one phase exists…a combination of liquid and gas. • Properties are in between those of liquids and gases. • They act as solvents and dissolve well. • They diffuse well like gases. • CO2 and H2O- environmentally friendly

49. The Critical Point At Tc, the densities of liquid and vapor are equal; a single phase. At room temperature there is relatively little vapor, and its density is low. At higher temperature, there is more vapor, and its density increases … … while the density of the liquid decreases; molecular motion increases.