Unit 13 – Intermolecular Forces/Solids, Liquids and Solutions

1. Unit 13 – Intermolecular Forces/Solids, Liquids and Solutions

2. IQ #1 • What is bond polarity? • What determines the type of bond that exists between two atoms? • List the 3 major bonds types and the difference in electronegativity that exists between the atoms.

3. Definition of IMF • Attractive forces between molecules. • Much weaker than chemical bonds within molecules. • a.k.a. van der Waals forces

4. Attractions between molecules • They are what make solid and liquid molecular compounds possible. • The weakest are called van der Waal’s forces - there are two kinds: 1. Dispersion forces weakest of all, caused by motion of e- increases as # e- increases - halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A

5. Dispersion Forces • London Dispersion Forces View animation online.

6. Dipole interactions • Occurs when polar molecules are attracted to each other. • Dipole interactionhappens in water • Figure 8.25, page 240 • positive region of one molecule attracts the negative region of another molecule.

7. δ+δ- δ+ δ- H F H F Dipole interactions • Occur when polar moleculesare attracted to each other. • Slightly stronger than dispersion forces. • Opposites attract, but not completely hooked like in ionic solids.

8. d+d- d+d- d+d- d+d- d+d- d+d- d+d- Dipole Interactions d+d-

9. + - • Dipole-Dipole Forces View animation online.

10. 3. Hydrogen bonding • …is the attractive force caused by hydrogen bonded to N, O, F, or Cl • N, O, F, and Cl are veryelectronegative, so this is a very strong dipole. • The hydrogen partially share with the lone pair in the molecule next to it. • This is the strongestof the intermolecular forces.

11. Hydrogen bonding defined: • When a hydrogen atom is: a) covalently bonded to a highly electronegative atom, AND b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom. • The hydrogen is left very electron deficient, thus it shares with something nearby • Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!

12. d- d+ O d+ H d+ d- H H O H d+ Hydrogen Bonding

13. H O O H H O H H H H O H H H H O O O H H H Hydrogen bonding

14. Hydrogen Bonding

16. Types of IMF

17. Determining IMF • NF3 • polar = dispersion, dipole-dipole • CH4 • nonpolar = dispersion • HF • H-F bond = dispersion, dipole-dipole, hydrogen bonding

18. 0 Non – Polar Covalent 0.5 Polar Covalent Ionic character 1.7 Ionic 4.0 Examples: 1. Explain, in terms of intermolecular forces, why (a) the boiling point of O2 (-183oC) is higher than that of N2 (-196oC). (b) the boiling point of NO is higher than either N2 or O2. (a) Both O2 & N2 are non-polar molecules  it is based on molar mass. As molar mass increases, so does the dispersion force resulting in stronger bonds in turn a higher boiling pt. O2 has a higher BP, because it has a greater molar mass in turn a greater dispersion force. (b) Both O2 & N2 are non-polar molecules, but NO is a polar molecule.  NO has stronger intermolecular forces in turn a higher boiling pt.

19. •• •• C •• •• O S •• N H H H IQ #2 What types of intermolecular forces are present in H2? CCl4? OCS? NH3? H2= dispersion CCl4 = dispersion OCS = dispersion, dipole-dipole dipole-dipole, hydrogen bonding NH3 =

20. Attractions and properties • Why are some chemicals gases, some liquids, some solids? • Depends on the type of IMF! • Table 8.4, page 244

21. Kinetic Molecular Theory • KMT • Particles of matter are always in motion. • The kinetic energy (speed) of these particles increases as temperature increases.

22. Forces and Phases • Substances with very little intermolecular attraction exist as gases. • Substances with strong intermolecular attraction exist as liquids. • Substances with very strongintermolecular (or ionic) attraction exist as solids.

23. Phase Differences Solid – definite volume and shape; particles packed in fixed positions; particles are not free to move Liquid– definite volume but indefinite shape; particles close together but not in fixed positions; particles are free to move Gas– neither definite volume nor definite shape; particles are at great distances from one another; particles are free to move

24. Three Phases of Matter

25. Liquid Properties • Surface Tension • attractive force between particles in a liquid that minimizes surface area.

26. water mercury Liquid Properties • Capillary Action • attractive force between the surface of a liquid and the surface of a solid.

27. Applications: 1) Blood tests (finger) 2) Plants: absorb subsurface ______ with tiny tubes in the _____. This can lift water about a maximum of __ ft. or ___ cm. Plants taller than one foot must use _______ (_____________). 3) Paper: ______________ 4) Sponges, towels, diapers, etc: ____________ water roots 1 30 xylem Active transport Cellulose fibers Cotton fibers

28. Definition: The _________ of a liquid to flow. Examples of viscous liquids: Cause? The more the molecules ______ each other, the ______ the viscosity. - Effect of temperature: As the temperature increases, the viscosity _________. Examples: Fudge, syrup, motor oil (summer: ______ viscosity vs. winter: ______ viscosity) Viscosity resistance Malasses, oil, & honey attract higher decreases high low

29. http://www.emporia.edu/ The Solid State 1. Types of Solids a) Crystalline: A solid in which the particles are arranged in an orderly, ____ repeating pattern. Example: _____ Seven types of crystals: cubic, orthorhombic, tetragonal, monoclinic, triclinic, hexagonal, rhombohedral.  b) Amorphous: Without ______. A non-crystalline solid whose particles are in a ______ arrangement. Example: _______ 3-D NaCl shape random glass

30. IMF m.p. Phase Changes • Melting Point • equal to freezing point • Which has a higher m.p.? • polar or nonpolar? • covalent or ionic? polar ionic

31. Phase Changes

32. Phase Changes • Evaporation • molecules at the surface gain enough energy to overcome IMF. • Volatility • measure of evaporation rate • depends on temp & IMF.

33. Phase Changes • Equilibrium • trapped molecules reach a balance between evaporation & condensation

34. IMF v.p. Phase Changes • Vapor Pressure • pressure of vapor above a liquid at equilibrium p.478 v.p. • depends on temp & IMF • directly related to volatility temp v.p. temp

35. Patm b.p. IMF b.p. Phase Changes • Boiling Point • temp at which v.p. of liquid equals external pressure. • depends on Patm & IMF • Normal B.P. = b.p. at 1 atm

36. Think About It! Example: Which substance would have a higher vapor pressure at 25°C: O ║ or H2O? H3C—C—CH3 (acetone) Hydrogen Bonding Dipole-Dipole

37. Effect of Pressure on Boiling Point

38. Think About It! 1) If you place a glass of water in a bell jar and turn on the vacuum pump, what will happen to boiling point? 2) Can you cook an egg faster if you turn up the flame under a pan of boiling water? Explain. Patm , B.P.  No, the temperature remains constant at the boiling point. High energy molecules escape, which cools the liquid. Thus, continuing to heat the water just maintains the temperature.

39. More time. The atmospheric pressure on Mt. Everest is only 240 mmHg. , the water boils at 70 oC, and the food would take longer to cook at the lower temperature. 3) Does it take more or less time to boil an egg on Mt. Everest or here in Fullerton? Explain.  4) Does food cook faster in a pan with a lid on it? Explain. 5) How does a pressure cooker work? Yes. The lid traps the high energy molecules, which keeps the heat from escaping. The pressure cooker increases the pressure, which increases the boiling point of water to ~ 150 oC –200 oC. , more heat-faster cooking.

40. IQ #3 • What is primarily responsible in determining the state of a compound or element? Explain. • Define: surafce tension, capillary action, and viscosity. • Explain in terms of intermolecular forces why: (a) ICl has a higher melting point that Br2. (b) C2H6 has a higher boiling point than CH4

41. IQ #3 cont. • Define: Volatility, Boiling Point, Vapor Pressure, Melting Point. • What relationship does IMF have with all of these? • Will increasing the elevation lower or raise your boiling point?

42. Solution Chemistry- Definitions Solution - homogeneous mixture Solute- substance being dissolved Solvent-present in greater amount

43. Definitions Solute - KMnO4 Solvent - H2O

44. Concentrated vs. Dilute

45. Solvents Solvents at the hardware store

46. Solvation Solvation – the process of dissolving solute particles are surrounded by solvent particles First... solute particles are separated and pulled into solution Then...

47. Dissolution of sodium Chloride

48. NONPOLAR NONPOLAR POLAR POLAR Solvation “Like Dissolves Like”

49. Solvation • Soap/Detergent • polar “head” with long nonpolar “tail” • dissolves nonpolar grease in polar water

50. Solubility • Solubility • maximum grams of solute that will dissolve in 100 g of solvent at a given temperature • varies with temp