1. Chapter 11 Intermolecular Forces and Liquids and Solids

2. Unit Objectives 1.0 Define key terms and concepts. 1.1 Determine if two liquids are miscible using intermolecular forces. 1.2 Explain the properties of gases, liquids, and solids in terms of density, compressibility, motion of molecules, shape, and volume. 1.3 Explain the structure and properties of water. 1.4 Interpret a phase diagram. 1.5 Explain how a phase change occurs. 1.6 Calculate the Molar Heat of Vaporization for a liquid. 1.7 Identify and explain the properties of different types of crystalline structures.

3. Solids, Liquids, and Gases Solid Highly ordered assembly of particles Has a definite volume and shape Liquid Particles are close together, but free to move randomly Has a definite volume, but takes the shape of the container it is in Gas Particles are relatively far apart and move randomly in all directions at rapid speeds Has no definite volume or shape Easily compressible

4. Kinetic Theory of Molecular Gases A gas consists of small particles (atoms or molecules) that are in constant motion, moving randomly and rapidly in straight paths. The attractive forces between the particles of a gas and be neglected since they are so far apart. The actual volume occupied by gas molecules is extremely small compared to the distance between molecules Energy is conserved in collisions between gas molecules because it is transferred from one molecule to another. The average kinetic energy of gas molecules is proportional to the Kelvin Temperature. Any two gases at the same temperature will have the same kinetic energy.

5. Forces  Intramolecular Forces  Holds atoms together within a molecule  Intermolecular Forces (https://www.youtube.com/watch?v=PVL24HAesnc)https://www.youtube.com/watch?v=PVL24HAesnc  Forces between molecules  When a state change occurs, it’s because intermolecular forces are disrupted

7. Dipole-Dipole Forces  Occurs between two molecules with a dipole moment (polar molecules)

8. Dipole-Dipole Forces

9. Ion-Dipole Forces  The attraction between an ion and a polar molecule  Strength of the bond is dependent on the charge/size of the ion/dipole moment.

10. Dispersion Forces Weakest of intermolecular forces Induced Dipole A dipole can be induced in a molecule if it is placed near an ion or a molecule with a dipole moment. The proximity of the polar molecule or ion distorts the electron distribution of the non-polar molecule. Factors that effect the creation of induced dipoles Ion charge/dipole moment strength Polarizability How easily the electron distribution can be distorted in a molecule

11. Dispersion Forces Instantaneous Dipole Weak, temporary attraction between two non-polar molecules as the result of electron imbalances caused by the constant movement of electrons around the nucleus

12. Hydrogen Bonds The strongest of intermolecular forces Occurs between hydrogen and typically oxygen, fluorine, and nitrogen.

13. Intermolecular Forces

14. Intermolecular Forces and Changes of State

15. Identify the intermolecular forces each of the following chemicals can undergo. NH 3 CH 4 H 2 S CH 3 Br HCl CH 3 CH 2 CH 3 CH 3 CH 2 OH

16. Identify the intermolecular forces at work between the following pairs of chemicals. Are they miscible? Na + and CH 4 H 2 O and HCl H 2 S and CH 3 CH 2 CH 3 H 2 O and NH 3 CH 3 CH 2 CH 3 and CH 4 CH 3 CH 2 OH and CH 3 CH 2 CH 3

17. What Are Your Questions?

18. Properties of Liquids Surface Tension The amount of energy required to stretch or increase the surface of a liquid by a unit of area. Liquids undergoing stronger intermolecular forces have greater surface tension.

19. Properties of Liquids Capillary Action Cohesion The intermolecular attraction between like molecules Adhesion Attraction between unlike molecules If adhesion > cohesion, the liquid will move up the tube. https://www.youtube.com/watch?v=wOOY1szbcX4

20. Properties of Liquids Viscosity Measure of a fluids resistance to flow Viscosity decreases as temperature increases Liquids with stronger intermolecular forces have higher viscosity.

21. Properties of Water “Universal Solvent” Has a high specific heat due to hydrogen bonding. Formation of Ice https://www.youtube.com/watch?v=HVT3Y3_gHGg

22. Crystal Structure Crystalline Solid A very orderly and ridged chemical compound Atoms within the structure occupy specific positions The structure shape tends to repeat Unit Cell Basic repeating structural unit Lattice point Represents an atom, ion or molecule within the structure

24. Bonding in Solids The forces that hold particles together determine properties of the structures, such as the melting point, density, and hardness. Solids are classified according to the type of forces present Ionic Molecular Covalent Metallic

25. Ionic Crystals Held together by ionic bonds. Examples: NaCl, MgO, KNO 3 Are hard, brittle, have high melting points and are poor conductors of heat and electricity. The structure of the crystal is dependent on the cation and anion involved in the bond and their radii.

26. Molecular Crystals Held together by van der Waals Forces or hydrogen bonding. Examples: Ice, SO 2, I 2 Are easier to break apart than ionic or covalent crystals.

27. Covalent Crystals Molecules are held together by covalent bonds. Forms 3-D structures Examples: diamond, graphite, quartz

28. Metallic Crystals Bonding electrons are delocalized across the entire crystal, which gives the metal its strength. As the number of bonding electrons present on the metal increases, so does the metals melting point. Good conductors of heat and electricity.

29. Crystal Structure

30. What Are Your Questions?

31. Phase Changes Converts from one physical state to another Melting When a substance changes from solid to liquid Freezing When a substance changes from liquid to solid Melting Point/Freezing Point The temperature at which the solid and liquid phases exists in equilibrium

32. Phase Changes Molar Heat of Fusion (∆H fus ) The energy required to melt 1 mole of a solid Substance∆H fus (kJ/mole)∆H vap (kJ/mole)Boiling Point (˚C) Argon1.36.3-186 Benzene10.931.080.1 Diethyl Ether6.926.034.6 Ethanol7.6139.378.3 Mercury23.459.0357 Methane0.849.2-164 Water6.0140.79100

33. Phase Changes Vaporization Particles of a liquid have enough kinetic energy to change to a gas Condensation Gas changes to liquid as the concentration of vapors increases Boiling Point The temperature at which the vapor pressure of a liquid is equal to the external pressure.

34. Phase Changes Sublimation When a substance goes from a solid to gas without passing through the liquid state Deposition The reverse process of sublimation Heat of Molar Sublimation (∆H sub ) The heat required to sublime 1 mole of a solid ∆H sub = ∆H fus + ∆H vap

36. Phase Changes Dynamic Equilibrium The rate of a process and the reverse of that process occur at the same rate. Equilibrium Pressure The vapor pressure measured under dynamic equilibrium of condensation and evaporation.

37. Phase Changes Molar Heat of Vaporization (∆H vap ) The energy required to vaporize one mole of a liquid. Measured in kJ/mole Directly related to the strength of the intermolecular forces between molecules of the liquid. Vapor pressure and the temperature of a liquid are related by the Clausius-Clapeyron Equation Where R=8.314J/K x mole C is a constant In P = - ∆H vap RT + C

38. Phase Changes The Clausius-Clapeyron Equation can also be adapted to determine the vapor pressure of a liquid at a different temperature if you know the values of ∆H vap and P at another temperature. Where R=8.314J/K x mole Temperature is measured in K In P 1 In P 2 ∆H vap R T 1 -T 2 T 1 T 2 = In P 1 In P 2 ∆H vap R 1T21T2 = - 1T11T1

39. Calculate the vapor pressure at 85°C if the heat of vaporization of water is 40.7kJ/mole.

40. Carbon disulfide has a boiling point of 46°C and a heat of vaporization of 26.8kJ/mole. What is the vapor pressure of CS 2 at 35°C?

41. Selenium tetrafluoride is a colorless liquid with a vapor pressure of 757mmHg at 105°C and 522mmHg at 95°C. What is it’s heat of vaporization?

42. Phase Diagrams Summarizes the conditions under which a substance exists as a solid, liquid, and gas Triple Point The only temperature at which all three phases can exist in equilibrium Critical Point The point at which the distinction between two phases ceases to exist

43. Phase Diagram

45. What Are Your Questions?