Chapter 10

1. Chapter 10 • Liquids and Solids and Phase Changes

2. gas liquid solid Kinetic-Molecular Description of Liquids and Solids Addition or removal of energy can convert one state of matter into another state. cool cool heat heat

3. Kinetic-Molecular Description of Liquids & Solids • Strengths of interactions among particles and • Degree of ordering of particles • Gases < Liquids < Solids • Miscibleliquids diffuse into one another • Miscible is another word for soluble. • Examples: • water/alcohol • gasoline/motor oil

4. Kinetic-Molecular Description of Liquids & Solids • Immiscible liquids do not diffuse into each other • (they are insoluble in each other or relatively insoluble in each other) • For example: • water/oil • water/cyclohexane

5. Intermolecular Attractions and Phase Changes • Ion-Ion interactions • The force of attraction between two oppositely charged ions is determined by Coulomb’s law

6. Intermolecular Attractions and Phase Changes • The energy of attraction between two ions is given by: Generally it is the magnitudeof the charge and not the number of ions with that charge that are considered when evaluating this strength of attraction.

7. Intermolecular Attractions and Phase Changes • Coulomb’s Law and the Attraction Energy determines: • melting points of ionic compounds • boiling points of ionic compounds • the solubility of ionic compounds • the heat of solvolysis of ionic compounds

8. Intermolecular Attractions and Phase Changes • Example: Arrange the following ionic compounds in the expected order of increasing melting and boiling points. • NaF, CaO, CaF2 • Consider the relative charges on the ions. (the greater the charge, the greater the attraction) • Consider the relative size of the ions (this equals separation of charge).

9. Intermolecular Attractions and Phase Changes • Dipole-dipole interactions • These are weak electrostatic attractions between the polarized molecules.

10. Intermolecular Attractions and Phase Changes • Hydrogen bonding • This is a relatively strong attraction between a H attached to a N, F, or O to another N, F, or O in the solution.

11. Intermolecular Attractions and Phase Changes Affects of Hydrogen Bonding

12. Intermolecular Attractions and Phase Changes • London Forces • These are very weakattractive forces caused by induced dipole moments. • London Forces are also known as dispersion forces. Isolated Ar atom Temporary induced dipole moments

13. The Liquid State • Viscosity • Viscosity is resistance to flow. • dependent upon intermolecular forces • consider water and molasses • temperature dependent • motor oil is rated on this property • W is the winter rating for viscosity • Ostwald viscometer measures this property • Viscosity is measured in centipoise

14. The Liquid State • Surface Tension • Surface tension is the measure of the unequal attractions that occur at the surface of a liquid. • Molecules at surface are attracted unevenly. • There is a stronger attraction to the neighboring molecules within the liquid than there is for molecules at the surface (air or other floating objects). • Water bugs • Soap or camphor boats • Floating needles • Soap bubbles

15. The Liquid State • Capillary Action • Capillary action is the ability of a liquid to rise (or fall) in a glass tube. • cohesive forces– hold liquids together adhesive forces– forces between a liquid and another surface • capillary rise implies adhesive > cohesive • capillary fall implies cohesive > adhesive

16. The Liquid State • Capillary Action rise – adhesive, H2O fall – cohesive, Hg

17. The Liquid State • Evaporation • Evaporation is the process in which molecules escape from the surface of a liquid. • The process is temperature dependent.

18. The Liquid State Blowing on a liquid cools it. Covering the liquid slows the cooling. • Evaporation WHY? Why then can you blow on your hands to warm them?

19. The Liquid State • Vapor Pressure • Vapor pressure is the pressure exerted by a liquid’s vapor at its surface at equilibrium. • Vapor Pressure (torr) for 3 Liquids Normal B.P. • 0 °C 20 °C 30 °C • diethyl ether 185 442 647 36°C • ethanol 12 44 74 78 °C • water 5 18 32 100 °C When vapor pressure equals current atmospheric pressure the liquid boils.

20. The Liquid State • Vapor Pressure

21. The Liquid State • Boiling Points & Distillation • Boiling point is the temperature at which the liquid’s vapor pressure is equal to applied pressure. • The boiling points you are familiar with (water at 100 °C) is the boiling point at 1 atm. • Distillation is a method used to separate mixtures of based on their differences in boiling points. • Distillation is another vapor phase phenomenon.

22. The Liquid State • Specific heat– the energy required to raise 1 g of a substance 1 °C • Molar heat capacity– the energy required to raise 1 mole of a substance 1 °C • Energy is generally calculated in joules, J or kJ (1000 J = 1 kJ) • Calories (cal) are an older unit of heat transfer and are defined as the amount of energy required to raise 1 g of water 1 °C.

23. The Liquid State • Changes ofState • Changes of state involve a change in energy. • Changes of state do not involve temperature change. • Heat of vaporization or DHvap is the amount of energy required to change 1 g of liquid to a gas (J/g) • Heat of condensation or DHcondis the amount of energy required to change 1 g of gas to a liquid (J/g) • The energy required for DHvap or DHcond is of exactly the same magnitude, the sign varies depending upon whether heat is absorbed (+) or released (–) At 100 °C

24. The Liquid State • Clausius-Clapeyron Equation This equation allows: 1. determination of the vapor pressure of a liquid at a new T (if DHvap is known) 2. determination of the T a liquid must be heated to in order to get a specified vapor pressure 3. determination of DHvap if vapor pressure is known at 2 T’s

25. The Liquid State • Example : In Denver the normal atmospheric pressure is 630 torr. At what temperature does water boil in Denver?

26. The Liquid State • Example : What is the vapor pressure of water at 50 °C? (Hvap = 40.7 kJ/mol)

27. The Liquid State Boiling Points of Various Substances Gas mol.wt. b.p. (°C) He 4 –269 Ne 20 –246 Ar 40 –186 Kr 84 –153 Xe 131 –107 Rn 222 –62 Few molecular attractions, b.p. related to mol.wt.

28. The Liquid State Boiling Points of Various Substances Compound mol.wt. b.p. (°C) CH4 16 –161 C2H6 30 –88 C3H8 44 –42 n-C4H10 58 –0.6 n-C5H12 72 +36 Few molecular attractions, b.p. related to mol.wt.

29. The Liquid State Boiling Points of Various Substances Compound mol.wt. b.p. (°C) HF 20 19.5 HCl 37 –85 HBr 81 –67 HI 128 –34 Notice how HF is out of order because of hydrogen bonding.

30. The Liquid State Boiling Points of Various Substances Compound mol.wt. b.p. (°C) H2O 18 100 H2S 34 –61 H2Se 81 –42 H2Te 130 –2 Notice how H2O is out of order because of hydrogen bonding.

31. The Liquid State • Example: Arrange the following substances in order of increasing boiling points. • C2H6, NH3, Ar, NaCl, AsH3 Ar < C2H6 < AsH3 < NH3 < NaCl nonpolar very polar ionic Although ammonia is much lower weight than AsH3, arsine, it is higher boiling because of hydrogen bonding.

32. The Solid State • Normal Melting Point • This is defined as the temperature at which the solid melts (liquid and solid in equilibrium) at 1 atm of pressure. • Both the liquid and solid states co-exist at the melting (or freezing) point • The melting point increases as intermolecular attractions increase.

33. Heat Transfer Involving Solids • Heat of Fusion, DHfus • amount of heat required to melt one gram of a solid at its melting point at constant T at 0 °C • Heat of Crystallization • the exact reverse of heat of fusion Important! At the melting point the liquid and the solid are both present and at equilibrium. DG = 0

34. Heat Transfer Involving Solids • Example : Calculate the amount of heat required to convert 150.0 g of ice at -10.0 °C to water at 40.0 °C. (specific heat of ice is 2.09 J/g ·°C; Hfus = 334 J/g·°C; specific heat of water is 4.183 J /g ·°C) ice melting heating combine

35. Sublimation & Vapor Pressure of Solids • Sublimation • Sublimation is the process by which a solid transforms directly to a vapor (gas). • Common examples include dry ice (CO2), camphor, and moth balls.

36. Phase Diagrams (P vs T) • Phase Diagrams only deal with closed systems • Triple point – the point at which the three states of matter can exist simultaneously • Critical Point – the point at which critical temperature and critical pressure meet • Supercritical liquid – liquid existing at temperatures and pressures beyond the critical point • A phase diagram is a convenient method to simultaneously display all the phases of a substance.

37. Phase Diagrams (P vs T) • Phase Diagram for Water

38. Phase Diagrams (P vs T) • Phase Diagram for Carbon Dioxide

39. Amorphous & Crystalline Solids • Amorphous solids do not have a well ordered structure. • Crystalline solids have well defined structures that consist of extended array of repeating units giveX-ray diffraction patterns. • Polymorphic solids can crystallize in one or more form.

40. Structure of Crystals • unit cell - smallest repeating unit of a crystal bricks are repeating units for buildings • 7 basic crystal systems

41. Structure of Crystals • Simple cubic Halite Crystals

42. Structure of Crystals • Simple cubic • each particle at a corner is shared by 8 unit cells • 1 unit cell contains 8(1/8) = 1 particle

43. Structure of Crystals • Body centered cubic (bcc) • 8 corners + 1 particle in center of cell • 1 unit cell contains 8(1/8) + 1 = 2 particles

44. Structure of Crystals • Face centered cubic (fcc)

45. Structure of Crystals • Face centered cubic (fcc) • 8 corners + 6 faces • 1 unit cell contains 8(1/8) + 6(1/2) = 4 particles

46. Bonding in Solids • Molecular Solids • molecules occupy unit cells • low melting points, volatile and insulators (nonconductive)

47. Bonding in Solids • Covalent Solids • atoms that are covalently bonded to one another

48. Bonding in Solids • Ionic Solids • ions occupy the unit cell

49. Bonding in Solids • Metallic Solids • positively charged nuclei surrounded by a sea of electrons • positive ions occupy lattice positions

50. Bonding in Solids • Variations in Melting Points • Molecular Solids • Compound Melting Point (°C) • ice 0 • ammonia -77.7 • benzene, C6H6 5.5 • naphthalene, C10H8 80.6 • benzoic acid, C6H5CO2H 122.4 • These are relatively low melting compounds.