1. Liquids and Solids

2. Relative Magnitudes of Forces The types of bonding forces vary in their strength as measured by average bond energy. Covalent bonds (400 kcal/mol) (intramolecular) Hydrogen bonding (12-16 kcal/mol ) (intermolecular) Dipole-dipole interactions (2-0.5 kcal/mol) (intermolecular) London forces (less than 1 kcal/mol) (intermolecular) Strongest Weakest

3. Intra- vs. Inter- intramolecular forces –inside molecules –hold atoms together into molecule intermolecular forces –between molecules –get weaker as phase changes from S – L – G Generally, intermolecular forces are much weaker than intramolecular forces.Generally, intermolecular forces are much weaker than intramolecular forces.

4. Hydrogen Bonding happens between H and N, O, or F very strong type of dipole-dipole attraction –because bond is so polar –because atoms are so small

5. Hydrogen Bonding Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

6. Attractive forces between polar molecules Orientation of Polar Molecules in a Solid Dipole-Dipole Forces

7. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. Fritz London 1900-1954 increase with the size of the molecules. increase with the size of the molecules. in every molecular compound in every molecular compound only important for nonpolar molecules and noble gas atoms only important for nonpolar molecules and noble gas atoms weak, short-lived weak, short-lived caused by formation of temporarycaused by formation of temporary dipole moments

8. London Dispersion Forces

9. London Forces in Hydrocarbons Dispersion forces usually increase with molar mass.

10. Boiling point as a measure of intermolecular attractive forces

11. Practice which has highest boiling pt? – –HF, HCl, or HBr? Identify the most important intermolecular forces : – –BaSO 4 – –H 2 S – –Xe – –C 2 H 6 – –P 4 – –H 2 O – –CsI ionicdipole-dipoleH-bonding London Dispersion

12. Which has stronger IMF’s? CO 2 or OCS – –CO 2 : nonpolar so only LD – –OCS: polar so dipole-dipole PF 3 or PF 5 – –PF 3 : polar so dipole-dipole – –PF 5 : nonpolar so only LD SF 2 or SF 6 –SF 2 : polar so dipole-dipole –SF 6 : nonpolar so only LD SO 3 or SO 2 –SO 3 : nonpolar so LD only –SO 2 : polar so dipole-dipole ? ? ? ?

13. Some Properties of a Liquid  Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals). Strong intermolecular forces High surface tension =

14.  Capillary Action: Spontaneous rising of a liquid in a narrow tube. Adhesion is an attraction between unlike molecules Adhesion Cohesion attracted to glass attracted to each other Cohesion is the intermolecular attraction between like molecules

15. Some Properties of a Liquid  Viscosity: Resistance to flow  High viscosity is an indication of strong intermolecular forces

16. General Classification of Solids Crystalline Solids: Well-ordered, definite arrangement of atoms. (Examples- metals, H2O, diamond) Amorphous: No pattern to the arrangement of particles. (Examples- glass, plastic, wax)

17. Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

18. Types of Crystalline Solids There are four types of crystalline solid: - Molecular (formed from molecules) - usually soft with low melting points and poor conductivity. - Covalent network (formed from atoms) - very hard with very high melting points and poor conductivity. - Ionic (formed form ions) - hard, brittle, high melting points and poor conductivity. - Metallic (formed from metal atoms) - soft or hard, high melting points, good conductivity, malleable and ductile.

19. Bonding in Crystalline Solids Metallic bonds are formed from metal nuclei floating in a sea of electrons. Know these!!

20. Crystalline Solids Covalent Network IonicMetallic Molecular

21. Metal Alloys  Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Znbrass = Cu/Zn

22. Metal Alloys (continued)  Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbonsteel = iron + carbon

23. Phase Changes & Energy Endothermic: melting, evaporating/boiling & sublimation Exothermic: freezing, condensation, & deposition

24. Phase Changes & Energy heat of vaporization: heat of vaporization: the heat energy required to evaporate a given mass of liquid at a constant temperature heat of fusion: heat of fusion: the heat energy required to melt a given mass of solid at a constant temperature Heating Curve The temperature, (average KE), during a phase change (such as boiling) does not change! Any heat added during boiling gives more molecules enough energy to escape the liquid.

25. Phase Changes & Energy Generally, it will take more heat to vaporize a liquid than to melt a solid… (∆H(vap) > ∆H(fusion) ) Why?Generally, it will take more heat to vaporize a liquid than to melt a solid… (∆H(vap) > ∆H(fusion) ) Why? – Every intermolecular bond is broken when vaporizing, but only some of the intermolecular forces break when melting solids. ?

26. Liquefying Gases Gases can be liquefied by increasing pressure at some temperature. - Critical temperature: the highest temperature at which a substance can remain a liquid regardless of the pressure applied. - Critical pressure: the pressure needed at the critical temperature.

27. Vapor Pressure and Boiling Liquids Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the liquid. These molecules move into the gas phase. As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. After some time the pressure of the gas will be constant at the vapor pressure. Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface. Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

28. Vapor Pressure and Boiling Liquids Liquids boil when the external pressure equals the vapor pressure. Temperature of the boiling point increases as pressure increases. Normal boiling point is the boiling point at 760 mmHg (1 atm). A substance with a high vapor pressure is said to be volatile. It readily evaporates.

29. Vapor Pressure and Boiling Liquids Two ways to get a liquid to boil: 1) increase temperature 2) decrease pressure (vacuum pump) Pressure cookers operate at high pressure. - At high pressure the boiling point of water is higher than at 1 atm. - Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required. - Boiling points increase with molecular weight (as long as the intermolecular forces are similar.) - Example: CH 4 (m.w. = 16) < C 2 H 6 (m.w. = 30) < C 3 H 8 (m.w. = 44) [ lowest B.P.] [highest B.P.]

30. Phase Diagrams Shows the relationship between the 3 phases of matter at various temperatures and pressures. Triple Point: All 3 phases of matter at equilibrium. Critical Point: The highest temperature at which the liquid phase can exist.

31. Phase Diagrams of H 2 O and CO 2 Notice the slope of the solid–liquid equilibrium line. This indicates that water expands when it freezes and CO 2 contracts when it freezes. Notice the slope of the solid–liquid equilibrium line. This indicates that water expands when it freezes and CO 2 contracts when it freezes.