Intermolecular Forces

1. Intermolecular Forces (a) Particles in solid (b) Particles in liquid (c) Particles in gas

2. Intermolecular Forces Objectives • Distinguish various properties of liquids and solids. • Relate different intermolecular forces to explain observations in lab and nature. • Use heating curves to calculate energy transfers of substances across various phase changes (melting, freezing, boiling, condensing, sublimation, deposition). • Use phase diagrams to relate pressure and temperature changes to physical states.

3. Properties of Solids, Liquids, and Gases Property Solid Liquid Gas Shape Has definite shape Takes the shape of Takes the shape the container of its container Volume Has a definite volume Has a definite volume Fills the volume of High densities High densities the container Low densities Bonding Ionic, Metallic, Covalent Covalent Covalent Arrangement of Fixed, very close Random, close Random, far apart, Particles Crystalline or amorphous Collisions Interactions BetweenVery strong forces: Strong forces: Essentially none Particles (i.e. Melting point, (i.e. Boiling point, malleability, ductility, Surface Tension, conductivity…) Viscosity, Vapor pressure…)

4. Liquid Surface Tension • molecules minimize • their surface area (“skin”) • molecules at surface • interact only with molecules • in the interior of liquid • surface molecules • subjected to inward force, • so surface is under tension • surface tension increases • with increasing intermolecular • forces H2O(l) Water Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 31

5. Liquid Viscosity • resistance of a liquid to flow • greatest in substances with • strong intermolecular forces, • which hinder flow • longer molecules higher • viscosity than shorter ones H2O(l) Water Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 31

6. Acetone

7. Gasoline (Hexane)

8. Corn Syrup

9. Motor Oil

10. Molasses

11. - + + - - + - + - + + - - + + - Repulsion Attraction Intermolecular ForcesDispersion (London Force) • (a) Interaction of any two atoms or molecules. Electrons unevenly distributed. Creates instantaneous (temporary dipole). Polarization increases with size. • (b) interaction of many dipoles. WEAKEST forces! - + - + Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 442

12. - + + - - + - + - + + - - + + - Repulsion Attraction Intermolecular Forces Dipole-Dipole • (a) Interaction of two polar molecules. Polar molecules have permanent dipoles from electronegativity difference. Higher melting and boiling points due to stronger IM forces. • (b) interaction of many dipoles in a liquid. - + - + Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 442

13. H H O O H H Intermolecular Forces Hydrogen Bonding • Strong intermolecular forces of attraction between molecules containing fluorine, oxygen, or nitrogen bonded to hydrogen • Results from large electronegativity difference and small atomic size of hydrogen Hydrogen bonds Chemical Bonds Chemical Bonds Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 442

14. Molecular Structure of Ice Hydrogen bonding Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 455

15. Water boiling on a stove

16. Hot vs. Cold Tea Low temperature (iced tea) Many molecules have an intermediate kinetic energy High temperature (hot tea) Few molecules have a very high kinetic energy Percent of molecules Kinetic energy

17. Evaporation (Vaporization) • Molecules must have sufficient energy to break IM bonds. • Molecules at the surface break away and become gas (“volatility”). • Only molecules with enough KE escape. • Breaking IM bonds absorbs energy. Evaporation is endothermic. • Rate of evaporation increases with increasing surface area, increasing temperature, and weaker IM forces

18. Condensation • Forming IM bonds from gas to liquid • Condensation is exothermic because energy is released. • Dynamic equilibrium: rate of vaporization equals rate of condensation (gas molecules above liquid becomes constant). • Vapor pressure: partial pressure of gas in dynamic equilibrium with liquid • Vapor pressure increases with increasing temperature and weaker IM forces

19. Boiling point: temperatureat which the vapor pressure of a liquid is equal to the pressure above it Microscopic view of a liquid near its surface Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 446

20. Formation of a bubble is opposed by the pressure of the atmosphere Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 452

21. Effect of Pressure on Boiling Point

22. Heating Curves • Temperature Change • change in KE (molecular motion) • depends on heat capacity • Phase Change • change in PE (molecular arrangement) • temperature remains constant • Heat Capacity • energy required to raise the temp of 1 gram of a substance by 1°C (q = mC∆T) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

23. Heating Curve for Water vaporization E gas D 100 condensation C liquid melting Temperature (oC) B 0 A freezing solid Heat added LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 487

24. Gas - KE Boiling - PE  Liquid - KE  Melting - PE  Solid - KE  Heating Curves 140 120 100 80 60 40 Temperature (oC) 20 0 -20 -40 -60 -80 -100 Energy

25. Heating Curves • Heat of Fusion (Hfus) • energy required to melt 1 gram of a substance at its melting point. Breaking intermolecular forces in the solid. Water = 6.02 kJ/mol • Heat of Vaporization (Hvap) • energy required to vaporize 1 gram of a substance at its boiling point. • usually larger than Hfus (requires complete separation of molecules) • higher temperatures = lower (Hvap) Water = 40.7 kJ/mol (@ B.P.) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

26. Calculating Energy Changes - Heating Curve for Water 140 DH = mol xDHvap DH = mol xDHfus 120 100 80 Heat = mass xDt x Cp, gas 60 40 Temperature (oC) 20 Heat = mass xDt x Cp, liquid 0 -20 -40 -60 Heat = mass xDt x Cp, solid -80 -100 Energy

27. Specific Heat Capacities Tro's "Introductory Chemistry", Chapter 3

28. States of Matter

29. Heating Curves • Heat of Vaporization (Hvap) • energy required to vaporize 1 gram of a substance at its boiling point. • usually larger than Hfus (requires complete separation of molecules) • higher temperatures = lower (Hvap) • EX: sweating, steam burns, the drinking bird Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

30. Energy Changes Accompanying Phase Changes Gas Vaporization Condensation Sublimation Deposition Energy of system Liquid Melting Freezing Solid Brown, LeMay, Bursten, Chemistry2000, page 405

31. Phase Diagrams • Show the phases of a substance at different temps and pressures. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

32. Intermolecular Forces • Determine the kinds of IM forces that are present in the following: Kr, N2, CO, HF, NCl3, NH3, SiH4, CCl4 • Which of the following has the highest boiling point. Why? CH4, CH3CH3, CH3CH2CH3, CH3CH2CH2CH3 • A flask containing a mixture of NH3(g) and CH4(g) is cooled. At -33.30C a liquid begins to form in the flask. What is the liquid?

33. GeH4 Boiling Points of Covalent Hydrides H2O 100 0 H2Te Boiling point (oC) H2Se SnH4 H2S -100 SiH4 CH4 -200 50 100 150 Molecular mass

34. Melting Points H -259.2 He -269.7 Mg 650 1 1 Symbol Melting point oC Li 180.5 Be 1283 B 2027 C 4100 N -210.1 O -218.8 F -219.6 Ne -248.6 2 2 > 3000 oC 2000 - 3000 oC Na 98 Mg 650 Al 660 Si 1423 P 44.2 S 119 Cl -101 Ar -189.6 3 3 K 63.2 Ca 850 Sc 1423 Ti 1677 V 1917 Cr 1900 Mn 1244 Fe 1539 Co 1495 Ni 1455 Cu 1083 Zn 420 Ga 29.78 Ge 960 As 817 Se 217.4 Br -7.2 Kr -157.2 4 4 Rb 38.8 Sr 770 Y 1500 Zr 1852 Nb 2487 Mo 2610 Tc 2127 Ru 2427 Rh 1966 Pd 1550 Ag 961 Cd 321 In 156.2 Sn 231.9 Sb 630.5 Te 450 I 113.6 Xe -111.9 5 5 Cs 28.6 Ba 710 La 920 Hf 2222 Ta 2997 W 3380 Re 3180 Os 2727 Ir 2454 Pt 1769 Au 1063 Hg -38.9 Tl 303.6 Pb 327.4 Bi 271.3 Po 254 At Rn -71 6 6 Ralph A. Burns, Fundamentals of Chemistry , 1999, page 1999

35. Densities of Elements H 0.071 He 0.126 1 1 Li 0.53 Be 1.8 B 2.5 C 2.26 N 0.81 O 1.14 F 1.11 Ne 1.204 2 2 Na 0.97 Mg 1.74 Al 2.70 Si 2.4 P 1.82w S 2.07 Cl 1.557 Ar 1.402 3 3 K 0.86 Ca 1.55 Sc (2.5) Ti 4.5 V 5.96 Cr 7.1 Mn 7.4 Fe 7.86 Co 8.9 Ni 8.90 Cu 8.92 Zn 7.14 Ga 5.91 Ge 5.36 As 5,7 Se 4.7 Br 3.119 Kr 2.6 4 4 Rb 1.53 Sr 2.6 Y 5.51 Zr 6.4 Nb 8.4 Mo 10.2 Tc 11.5 Ru 12.5 Rh 12.5 Pd 12.0 Ag 10.5 Cd 8.6 In 7.3 Sn 7.3 Sb 6.7 Te 6.1 I 4.93 Xe 3.06 5 5 Cs 1.90 Ba 3.5 La 6.7 Hf 13.1 Ta 16.6 W 19.3 Re 21.4 Os 22.48 Ir 22.4 Pt 21.45 Au 19.3 Hg 13.55 Tl 11.85 Pb 11.34 Bi 9.8 Po 9.4 At --- Rn 4.4 6 6 8.0 – 11.9 g/cm3 12.0 – 17.9 g/cm3 > 18.0 g/cm3 Mg 1.74 Symbol Density in g/cm3C, for gases, in g/L W