1. Condensed Phases and Intermolecular Forces

2. Fundamentals How do particle diagrams of liquids & solids compare to those of gases?

3. Describe the relative positions and motions of particles in each of the 3 phases.

4. The Question Why do some substances exist as gases at room temperature… …some as liquids …and some as solids?

5. TWO KINDS OF FORCES between Part of the answer has to do with the forces between separate molecules. There are two broad categories of forces you need to be aware of.

6. Intramolecular Forces Intramolecular forces  forces that hold particles together in ionic, covalent, or metallic bonds. Intra means “within.” Intramolecular forces = bonding forces.

7. Intermolecular Forces Inter means “between” or “among.” Inter molecular forces are forces between molecules. Inter molecular forces are weaker than intramolecular forces. There are three types of

8. Three Types of Intermolecular Forces Dispersion Forces Dispersion Forces  These are the weakest. Occur between nonpolar molecules. Dipole-dipole forces Dipole-dipole forces  These are intermediate. Occur between polar molecules. Hydrogen bonds Hydrogen bonds  These are the strongest. Occur between molecules containing an H-F, H-O, or H-N bond.

9. Dispersion Forces & Nonpolar molecules Instantaneous and momentary; fluctuating. Results from motion of electrons. Avg = fig. a. induce If the charge cloud is not symmetrical it will induce an asymmetry in its neighbor’s charge cloud!

10. Nonpolar molecules Nonno Nonpolar means no poles. Can’t tell one end of the molecule from the other. Symmetrical.

11. Examples of Nonpolar Molecules Monatomic gas molecules: He, Ne, Ar, Kr, Xe, Rn Diatomics where both atoms are the same element: H 2, N 2, O 2, Cl 2, F 2, I 2, Br 2 Larger molecules if they are very symmetrical molecules: CH 4, C 2 H 6, C 3 H 8

12. Dispersion Forces and Size Dispersion forces increase with the size of the molecule. The larger the electron cloud, the greater the fluctuations in charge can be. Rn > Xe > Kr > Ar > Ne > He I 2 > Br 2 > Cl 2 > F 2 C 8 H 18 > C 5 H 12 > C 3 H 8 > CH 4

13. Boiling point of N 2 is 77 K. Intermolecular forces are very weak dispersion forces.

14. Dipole-dipole Forces & Polar Molecules Molecule shows permanent separation of charge. It has poles – one end is somewhat negative & one end is somewhat positive.

15. Polar means the molecule has poles, +’ve & -’ve. The geometry and the charge distribution are not symmetrical. Polar Molecules

16. What do you know about charge? Opposites Attract! This time, the situation is permanent! Examples: HI, CH 3 Cl

17. Hydrogen Bonding Occurs between molecules containing an H-F, H-O, or H-N bond. (FON!!!)

18. Hydrogen Bonding Hydrogen bonding is the extreme case of dipole-dipole bonding. electronegative F, O, and N are all small and electronegative. They really attract electrons. H has only 1 electron, so if it’s spending time somewhere else, the proton is almost “naked.” The H end is always positive and the F, O, or N is always negative.

19. The bonding electrons spend more time by the oxygen atom than by the hydrogen atoms. Oxygen end – bit negative Hydrogens – bit positive

20. Hydrogen Bonding in Water

21. Hydrogen bonding is the strongest intermolecular force and influences the physical properties of the substance a great deal.

23. Result of Hydrogen Bonding Increased Boiling Point over predicted value based on molecular size & comparison with similar compounds.

24. Strength of Hydrogen Bonding Fluorine is the most electronegative element, so H-F bonds are the most polar and exhibit the strongest hydrogen bonding. H-F > H-O > H-N

25. Intermolecular Forces vs. Physical Properties If intermolecular forces increase, Boiling point  Melting point  Heat of Fusion  Heat of Vaporization  Evaporation Rate  Vapor Pressure 

26. Intermolecular Forces vs. Temperature Intermolecular forces become more important as the temperature is lowered. EX:  Low temperature – low evaporation rate  High temperature – high evaporation rate

28. Indicate the kind of intermolecular forces for each molecule below. NH 3 Ar N 2 HCl HF Ne O 2 HBr CH 3 NH 2 Hydrogen bonding Dispersion forces Dipole-dipole forces Hydrogen bonding Dispersion Dipole-dipole Hydrogen bonding

29. Forces INTERMOLECULAR Dispersion Dipole-Dipole Hydrogen Bonding INTRAMOLECULAR Covalent Ionic Metallic

30. Intermolecular forces determine phase. “Competition” between strength of intermolecular forces and kinetic energy determines phase. If intermolecular forces are strong, substance will be a solid or liquid at room temperature. Particles want to clump together. If intermolecular forces are weak, substance will be a gas at room temperature. Particles spread apart.

31. It’s a balancing act! Intermolecular Forces Kinetic Energy This substance = a gas at room temperature.

32. Intermolecular Forces vs. Kinetic Energy Intermolecular Forces Kinetic Energy This substance = a condensed phase.

33. Why Temperature Changes Affect Phase Since temperature is a measure of avg kinetic energy, changing temperature changes phase. Changing temperature changes the average kinetic energy

34. Changing the temperature IntermolecularForces Kinetic Energy

35. Intermolecular Forces 5% to 15% the strength of the intramolecular or bonding forces Account for phase at room temperature. Account for phase at room temperature. Strong intermolecular forces  condensed phase. Weak intermolecular forces  gas phase

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