2. Define: polar, nonpolar, dipole-dipole forces, ion-dipole forces, Hydrogen “bonding”, and London dispersion forces; sublimation, condensation, evaporation, freezing point, boiling point, vapor pressure, viscosity, surface tension, delta H of fusion, vaporization, and sublimation. Distinguish between intermolecular and intramolecular attractions Put a list of compounds in order of increasing melting point, boiling point, and vapor pressure Use and label the parts of a phase diagram Understand the expansion of ice due to hydrogen bonding Use the Clausius-Clapeyron equation

3. Three factors determine whether a substance is a solid, a liquid, or a gas: 1.The attractive intermolecular forces between particles that tend to draw the particles together. 2.Temperature: The kinetic energies of the particles (atoms, molecules, or ions) that make up a substance. Kinetic energy tends to keep the particles moving apart. 3.Pressure: pressure is increased or decreased as the volume of a closed container changes Solid, Liquid, or Gas

4. There are several types of attractive intermolecular forces: 1.Hydrogen bonding 2.Dipole-dipole forces 3.Induced-dipole forces 4.London dispersion forces Types of Attractive Forces All of the intermolecular forces that hold a liquid together are called cohesive forces.

5. Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

6. Hydrogen Bonding in DNA TA Thymine hydrogen bonds to Adenine

7. Hydrogen Bonding in DNA CG Cytosine hydrogen bonds to Guanine

8. Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching). Dipole-Dipole Forces

9. An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases. Ion-Dipole Forces

10. Induced dipole forces result when an ion or a dipole induces a dipole in an atom or a molecule with no dipole. These are weak forces. Induced-Dipole Forces Ion-Induced Dipole Forces An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

11. A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species. Dipole-Induced Dipole Forces

12. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. Fritz London 1900-1954 London forces increase with the size of the molecules.

13. London Dispersion Forces

14. London Forces in Hydrocarbons

15. Boiling point as a measure of intermolecular attractive forces

16. Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Hydrogen bonding (12-16 kcal/mol ) Dipole-dipole interactions (2-0.5 kcal/mol) London forces (less than 1 kcal/mol) Strongest Weakest Ion-dipole interactions Ionic bonds Ion induced dipole interactions Induced Dipole-dipole interactions

18. Intermolecular Quiz Identify the most predominant intermolecular force in the following molecules: 1)CF 2 Cl 2 2)CaCl 2 3)C 6 H 12 O 6 4)PCl 5 5)Na + in H 2 O 6)Cl - in hexane (C 6 H 14 ) 1)Dipole-Dipole 2)Ionic 3)Hydrogen Bonding 4)London Dispersion 5)Ion-dipole 6)Ion-induced Dipole

19. Intermolecular Quiz Determine which of the pairs of molecules would have the: 1)Highest Boiling Point a)CF 2 Cl 2 or H 2 O b)SO 3 orSO 2 2)Highest Vapor Pressure a)KCl orHCl b)H 2 Oor C 6 H 12 O 6

20. This bottle contains both liquid bromine [Br 2 (l), the darker phase at the bottom of the bottle] and gaseous bromine [Br 2 (g), the lighter phase above the liquid]. The circles show microscopic views of both liquid bromine and gaseous bromine. A liquid is a state of matter in which a sample of matter: is made up of very small particles (atoms, molecules, and/or ions). flows and can change its shape. is not easily compressible and maintains a relatively fixed volume. The particles that make up a liquid: are close together with no regular arrangement, vibrate, move about, and slide past each other. What Is a Liquid?

21. More Properties of a Liquid  Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).  Capillary Action: Spontaneous rising of a liquid in a narrow tube.

22. Even More Properties of a Liquid  Viscosity: Resistance to flow  High viscosity is an indication of strong indication of strong intermolecular forces intermolecular forces

23. Microscopic view of a liquid. Microscopic view after evaporation. When a liquid is heated sufficiently or when the pressure on the liquid is decreased sufficiently, the forces of attraction between molecules do not prevent them from moving apart, and the liquid evaporates to a gas. Example: The sweat on the outside of a cold glass evaporates when the glass warms. Example: Gaseous carbon dioxide is produced when the valve on a CO 2 fire extinguisher is opened and the pressure is reduced. Evaporation Evaporation is the change of a liquid to a gas.

24. Condensation Condensation is the change from a vapor to a condensed state (solid or liquid). When a gas is cooled sufficiently or, in many cases, when the pressure on the gas is increased sufficiently, the forces of attraction between molecules prevent them from moving apart, and the gas condenses to either a liquid or a solid. Example: Water vapor condenses and forms liquid water (sweat) on the outside of a cold glass or can. Example: Liquid carbon dioxide forms at the high pressure inside a CO 2 fire extinguisher. Microscopic view of a gas. Microscopic view after condensation.

25. When a solid or a liquid evaporates to a gas in a closed container, the molecules cannot escape. Some of the gas molecules will eventually strike the condensed phase and condense back into it. When the rate of condensation of the gas becomes equal to the rate of evaporation of the liquid or solid, the amount of gas, liquid and/or solid no longer changes. The gas in the container is in equilibrium with the liquid or solid. Vapor Pressure Revealed

26. Types of Molecules: the types of molecules that make up a solid or liquid determine its vapor pressure. If the intermolecular forces between molecules are: Factors That Affect Vapor Pressure substance vapor pressure at 25 o C diethyl ether0.7 atm bromine0.3 atm ethyl alcohol0.08 atm water0.03 atm Surface Area: the surface area of the solid or liquid in contact with the gas has no effect on the vapor pressure.

27. The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid (or solid) The vapor pressure of a liquid varies with its temperature, as the following graph shows for water. The line on the graph shows the boiling temperature for water. Vapor Pressure As the temperature of a liquid or solid increases its vapor pressure also increases. Conversely, vapor pressure decreases as the temperature decreases.

28. Temperature Dependence of Vapor Pressures The vapor pressure above the liquid varies exponentially with changes in the temperature. The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

29. Clausius – Clapeyron Equation A straight line plot results when ln P vs. 1/T is plotted and has a slope of  H vap /R. Clausius – Clapeyron equation is true for any two pairs of points. Write the equation for each and combine to get:

30. Using the Clausius – Clapeyron Equation Boiling point - the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere. Normal boiling point - the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm). E.g. Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of 190.0 mmHg at 25.0°C. E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it has a vapor pressure of 100.0 mmHg. What is the heat of vaporization? 335 K 33.0 kJ/mol

31. Solids

32. Types of Solids  Amorphous solids: considerable disorder in their structures (glass).

33. Types of Solids  Crystalline Solids: highly regular arrangement of their components

34. Metal Alloys  Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn

35. Metal Alloys (continued)  Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon

36. Network Atomic Solids Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of covalently bonded carbon atoms Graphite, a network of covalently bonded carbon atoms

37. Molecular Solids Strong covalent forces within molecules Weak covalent forces between molecules Sulfur, S 8 Phosphorus, P 4

38. Phase Transitions Melting: change of a solid to a liquid. Freezing: change a liquid to a solid. Vaporization: change of a liquid to a gas. Condensation: change of a gas to a liquid. Sublimation : Change of solid to gas Deposition: Change of a gas to a solid. H 2 O(s)  H 2 O(l) H 2 O(l)  H 2 O(s) H 2 O(l)  H 2 O(g) H 2 O(g)  H 2 O(l) H 2 O(s)  H 2 O(g) H 2 O(g)  H 2 O(s)

39. Water phase changes Temperature remains constant during a phase change. Energy

40. Energy of Heat and Phase Change Heat of vaporization: heat needed for the vaporization of a liquid. H 2 O(l)  H 2 O(g)  H = 40.7 kJ Heat of fusion: heat needed for the melting of a solid. H 2 O(s)  H 2 O(l)  H = 6.02 kJ Temperature does not change during the change from one phase to another.

41. E.g. Start with a solution consisting of 50.0 g of H 2 O(s) and 50.0 g of H 2 O(l) at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water.

42. Phase Diagrams Triple point- Temp. and press. where all three phases co-exist in equilibrium. Critical temp.- Temp. where substance must always be gas, no matter what pressure. Critical pressure- vapor pressure at critical temp. Critical point- point where system is at its critical pressure and temp.

43. Phase changes by Name

44. Water

45. Carbon dioxide

46. Carbon

47. Sulfur