Chapter 9 Liquids and Solids

1. Chapter 9Liquids and Solids

2. Intro Vocabulary • Gas: no definite shape or volume • Remember kinetic theory of gases • Liquid: definite volume – no definite shape • Some attraction between molecules or atoms • Solid: definite shape and volume • Strong intermolecular bonding 1. Molecules are much closer together in liquids and solids than in gases • In gases, molecules are separated by ten or more molecular diameters • In liquids and solids, the molecules are in contact with each other 2. Intermolecular forces play a major role in the behavior of liquids and solids, whereas they are negligible in gases

3. Phase transitions • Melting/Freezing • Vaporization/Condensation • Sublimation/Deposition • Note: Energy is absorbed or released during a phase change even though there is no temperature increase/decrease.

5. Phase Diagrams • A. Heating/Cooling Curves definitions • 1. Conversion of a solid to a liquid is:_______________ • 2. Conversion of a liquid to a solid is:_______________ • 3. The freezing point = melting point • 4. Energy needed to melt a given quantity of solid is • called the ___________________________________.

6. Examples: • Example: How much energy is required to melt 100.0 grams of ice? The heat of fusion is 6.01 kJ/mole. • Example: How much energy in kJ is required to heat 100.0 grams of liquid water from zero to 100°C, and then vaporize all of it? ∆Hvap= 40.79 kJ/mole

7. 9.1 Liquid - Vapor Equilibrium A. Vaporization (evaporation) process in an open container - evaporation will continue until all the liquid is gone - the energy required for vaporization comes from the surroundings and system - vaporization leaves the remaining liquid cooler - evaporation will occur below the boiling point of a substance - evaporation below the boiling point is slower than at the boiling point

8. B. Enthalpy of vaporization • Definition – the amount of energy change that occurs during the vaporization of 1 mole of a substance q = m c ∆T q = n ∆Hvap

9. C. Vapor Pressure – the pressure of the gas above a liquid in a closed container; dependent on temperature 1. Closed container vs. open container • In an open container the system includes the surroundings and the liquid will evaporate • In a closed system the liquid will evaporate and begin to condense when equilibrium is established between the liquid and gas

10. 2. Dynamic Equilibrium • When the rate at which the liquid vaporizes is equal to the rate at which the vapor condenses • The liquid level in the container does not change • Molecules are constantly moving between phases with no net change

11. 3. Pressure and Volume • As long as some liquid remains when equilibrium is established, the equilibrium vapor pressure will be the same regardless of the volume of the container

12. 33% humidity = 7.9 mm Hg Use PV = nRT moles H20 = 13.9 mol 1.75 L

13. E. Vapor Pressure Curves and Temperature 1. Relationships • Vapor pressure of liquid increases as temperature increases

14. 2. General Graph • What does this graph tell you about the relative attraction between molecules for substances a-e?

15. Boiling Point 1. Definition: • The boiling point is the temperature at which the vapor pressure equals atmospheric pressure 2. Normal Boiling Point: • The boiling point at exactly 1 atm of pressure 3. Dependency on pressure • At a certain temperature, large bubbles form throughout the liquid; i.e., the liquid boils • The temperature at which a liquid boils depends on the pressure above it

16. Dependency on pressure (continued) • Variation of atmospheric pressure will change the boiling point • At high elevation, atmospheric pressure is lower, so the boiling point is lower • To elevate the boiling point and allow food to cook more quickly, a pressure cooker can be used

17. Figure 9.3 - Boiling

18. C. Phase Diagrams 1. definition: A pressure vs. temperature graph depicting phases of a substance; specific for each substance 3. Triple Point – the point where all phases can exist in equilibrium

20. Figure 9.7

21. Figure 9.5 – A Phase Diagram

22. 4. Critical Temperature and Pressure • Critical temperature (Tc)- the temperature above which only vapor can exist b. Critical pressure (Pc)- the pressure above which on vapor can exist c. Critical point – the meeting point on the graph of the Tc and Pc

23. Example 9.2 (refer back to diagram on pg 7) • What phase is present at a temperature of 75°C and a pressure of 1 atm? • What phases of matter are present at point F? • What would happen if you started at point F and increased T with no change in P?

24. Permanent Gases • Permanent gases are substances with critical temperatures below 25 °C. • Usually stored in cylinders • Only vapor is present • Pressure drops as the gas is released

25. Condensable Gases • Condensable gases have critical temperatures above 25 °C. • Pressure starts to drop after all the liquid is gone

26. 9.3 Intermolecular Forces • Many gases, most liquids and many solids are molecular • Molecules are the structural units of such matter • Properties of molecular substances include • They are nonconductors of electricity when pure • They are insoluble in water but soluble in nonpolar solvents such as CCl4 or benzene • They have low melting points • These properties depend on the intermolecular forces between the molecules of the substance

27. A. 3 types of intermolecular forces • Dispersion • Dipole-dipole • Hydrogen bonding

28. Dispersion Forces • Definition – a force of attraction between molecules that is caused by temporary dipoles dipole – a molecule with a positive a negative end

29. Dispersion Forces 2. All molecules have some dispersion forces acting between them temporary dipoles form as a result of the natural movement of e-s in the e- cloud creating areas of positive and negative charge

30. Dispersion Forces 3. Strength of dispersion forces - all molecules have dispersion forces - strength increases with increasing # of e-s

31. Dispersion Forces 4. Dispersion forces increase as molar mass increases - directly proportional - Why? As molar mass increases the # of e-s increases - higher Dispersion forces = higher boiling and melting points because molecules tend to “stick together”

32. Figure 9.8 and Table 9.2

33. Example 9.3 Account for the fact that chlorine is a gas, bromine is a volatile liquid, and iodine is a volatile solid at room temperature.

34. C. Dipole-Dipole Forces • Definition and example - a force of attraction between molecules that is caused by permanent dipoles - CO = polar bond resulting in permanent dipole

35. 2. Higher bp and mp than expected because of the D-D forces the molecules “stick together” and require much more energy for the phase change

36. Figure 9.9

37. Table 9.3

38. Example 9.4

39. D. Hydrogen Bonds 1.Unusually strong type of dipole force • H attached to an N or O or F • The H from one molecule is strongly attracted to the negative end of the dipole of another • The strong dipole forms from the large difference in electronegativities of H and (N, O, or F) 2.Hydrogen bonds are the strongest intermolecular force • unusually high boiling points (H2O vs. CH4) • Small size of H allows the unshared pair from the negative end of the dipole to approach the H closely

40. 4. examples of hydrogen bonding • H2O, NH3, HF

41. 5. Unusual properties of Water • Because of H-bonding: • High specific heat • High boiling point • Liquid phase more dense than solid phase = ice floats

42. Figure 9.10

43. Example 9.6

44. What types of intermolecular forces are present in the following substances? Rank these substances in order of increasing bp. N2 HF SiCl4 CH3Cl NH3

45. Covalent vs. Intermolecular Forces • Three types of intermolecular force • Dispersion • Dipole • Hydrogen bond • All three intermolecular forces are weak relative to the strength of a covalent bond • Attractive energy in ice is 50 kJ/mol • Covalent bond in water is 928 kJ/mol

46. Covalent, Ionic, Metallic, and Amorphous • Network covalent solids • Continuous network of covalent bonds • Crystal is one large molecule • Ionic solids • Oppositely-charged ions held together by strong electrical forces • Non-conductors and water soluble • Metallic solids • Structural units are +1, +2 and +3 metals with associated electrons • Conductors, ductile and malleable

47. Figure 9.11

48. Network Covalent Solids • Characteristics • High melting points, often above 1000 °C • Covalent bonds must be broken to melt the substance • Examples • Graphite and diamond: allotropes • Diamond is three-dimensional and tetrahedral • Graphite is two-dimensional and planar

49. Figure 9.12

50. Figure 9.13