Ch. 11: Liquids, Solids, and Intermolecular Forces

1. Ch. 11: Liquids, Solids, and Intermolecular Forces Dr. Namphol Sinkaset Chem 200: General Chemistry I

2. I. Chapter Outline • Introduction • Intermolecular Forces • Vaporization and Vapor Pressure • Energies of Phase Changes • Phase Diagrams

3. I. Electrostatic Forces • Every molecule in a sample of matter experiences two types of electrostatic forces. • Intramolecular forces: the forces that exist within the molecule (bonding). These forces determine chemical reactivity. • Intermolecular forces: the forces that exist between molecules. These forces determine physical properties.

4. I. Solid, Liquid, or Gas? • Whether a substance exists as a solid, liquid, or gas depends on the relationship between the intermolecular attractions and the kinetic energy of the molecules. • It’s a battle – which dominates? The KE or the IM attractions? • Recall that the average KE of a sample is related to its temperature.

5. I. KE vs. IM Forces • Gas: the kinetic energy of the molecules is much greater than the intermolecular attractions. • Liquid: the kinetic energy of the molecules is moderately greater than the intermolecular attractions. • Solid: the kinetic energy of the molecules is less than the intermolecular attractions.

6. II. Intermolecular Forces • IM forces originate from interactions between charges, partial charges, and temporary charges on molecules. • IM forces are relatively weak because of smaller charges and the distance between molecules.

7. II. Types of IM Forces • There are different kinds of IM forces, each with a different level of strength. • Dispersion force • Dipole-dipole force • *Hydrogen “bonding” • Ion-dipole force

8. II. Dispersion Force • Dispersion force (London force) is present in all molecules and atoms and results from changes in e- locations.

9. II. Instantaneous Dipoles • Charge separation in one creates charge separation in the neighbors.

10. II. Dispersion Force Strength • The ease with which e-’s can move in response to an external charge is known as polarizability. • Large atoms with large electron clouds tend to have stronger dispersion forces. • Larger molecules tend to have stronger dispersion forces.

11. II. Noble Gas Boiling Points

12. II. Dispersion Force and Size • Molecular size is not the only factor…

13. II. Dispersion Force and Shape • Shape influences how the molecules interact with one another…

14. II. Dipole-Dipole Force • Occurs in polar molecules which have permanent dipoles, so attraction is always present.

15. II. Effect of Dipole-Dipole Force • Polar molecules have dispersion forces and dipole-dipole forces. • Effects can be seen in boiling and melting points.

16. II. “Like Dissolves Like” • Polar liquids are miscible with other polar liquids, but not with nonpolar liquids. • Can be explained with intermolecular forces.

17. II. Hydrogen “Bonding” • This IM force is a misnomer since it’s not an actual bond. • Occurs between molecules in which H is bonded to a highly electronegative element (N, O, F), leading to high partial positive and partial negative charges. • It’s a “super” dipole-dipole force.

18. II. H “Bonding” in Ethanol & Water

19. II. Effect of H “Bonding” • Hydrogen “bonding” is a very strong intermolecular force. • Without hydrogen “bonding” life as we know it could not exist!

20. II. Ion-Dipole Force • This type has already been discussed. • Example: NaCl(s) dissolved in water.

21. II. Summary of IM Forces

22. III. Vaporization and IM Forces • From experience, we know that water evaporates in an open container. • What factors influence rate of vaporization?

23. III. Vaporization Variables • Temperature • Surface area • IM forces

24. III. Heat of Vaporization • The energy needed to vaporize 1 mole of a liquid to gas is the heat of vaporization. • Can be thought of the energy needed to overcome IM forces of the liquid.

25. III. Dynamic Equilibrium • In an open flask, a liquid will eventually evaporate away. • What about a closed flask?

26. III. Dynamic Equilibrium • As evaporation occurs, headspace fills with gas molecules. • Gas molecules condense back to liquid phase. • Eventually, rates become equal. • Pressure of gas at dynamic equilibrium is called the vapor pressure.

27. III. Dynamic Equilibrium • Systems at dynamic equilibrium will seek to return to dynamic equilibrium when disturbed.

28. III. Vapor Pressure and Temp. • Vapor pressure depends on temperature and IM forces. • Why?

29. III. Clausius-Clapeyron Equation • The nonlinear relationship between vapor pressure and temperature can be written in a linear form:

30. III. Clausius-Clapeyron Equation, 2-point Form • If you have two sets of pressure, temperature data for a liquid, the more convenient 2-point form of the Clausius-Clapeyron equation can be used.

31. III. Boiling Point • When temperature is increased, the vapor pressure increases due to the higher number of molecules that can break away and enter the gas phase. • What if all molecules have the necessary thermal energy? • At this point, vapor pressure equals the external pressure, and the boiling point is reached.

32. III. Boiling Point • At the boiling point, those aren’t air bubbles!

33. IV. Energies of Phase Changes • The enthalpies involved in a phase change depends on the amount of substance and the substance itself. • We look at a heating curve for 1.00 moles of H2O at 1.00 atm pressure. • Note that there are sloping regions and flat regions in the curve. (Why?)

34. IV. Heating Curve for H2O

35. IV. Heating Curve, Segment 1 • At this stage, we are heating ice from -25 °C to 0 °C, increasing KE (vibrational motions). • The heat required depends on the specific heat capacity of ice.

36. IV. Heating Curve, Segment 2 • Here, the temperature stays the same, so the average KE stays the same. • Thus, the PE must be increasing. • The heat gained is a factor of the heat of fusion, the heat needed to melt 1 mole of solid.

37. IV. Heating Curve, Segment 3 • During this stage, water is being heated from 0 °C to 100 °C; again, KE is increasing. • The heat gained depends on the specific heat capacity of water.

38. IV. Heating Curve, Segment 4 • Again, the temperature stays the same, so the average KE stays the same. • PE must be increasing. • The heat gained is a factor of the heat of vaporization.

39. IV. Heating Curve, Segment 5 • During this stage, steam is heated from 100 °C to 125 °C; average KE is increasing. • The heat gained depends on the specific heat capacity of steam.

40. V. Phase Diagrams • The relationship between pressure, temperature, and the three phases can be summarized in a phase diagram. • A phase diagram allows the prediction of how a substance will respond to changes in pressure and/or temperature.

41. V. Phase Diagram of Water

42. V. I2 and CO2 Phase Diagrams