Chpt 10 - Condensed Phases

1. Chpt 10 - Condensed Phases • Condensed phases • Intermolecular forces • Special bonding - molecular solids, network solids, metallic • Phase diagrams & Heating curves • HW: Chpt 10 - pg. 487-496, #s 12, 14-16, 19-21, 24, 26, 31, 32, 34, 40, 44 Due Fri Dec. 4

2. States of Matter Differences? What do the phases look like? What makes the state of matter at a given temperature? Intermolecular forces

3. Intermolecular Forces • Intramolecular forces (chemical bonds) - forces that hold atoms together within a molecule • Intermolecular forces - forces between molecules - aggregate or bulk material - Is it a solid, liquid or gas? • dipole-dipole force (~1% of strength of a bond) • Hydrogen bonding H and N,O,F bond • London dispersion forces

4. Dipole-Dipole Force • Dipole moment – molecules with polar bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge. • Molecules with dipole moments can attract each other electrostatically. They line up so that the positive and negative ends are close to each other.

5. Hydrogen Bonding Very strong dipole-dipole force between H and N,O,F (most electro-negative elements) (a) Polar water molecule (b) hydrogen bonding between water molecules - blue dotted lines

6. Hydrogen bonding graph of covalent hydrides • Why are these interaction forces happening? • Especially polar X-H bond • Small size of N,O, and F allow close approach of dipoles

7. London Dispersion Forces • Weakest of the intermolecular forces • Important for atoms & non-polar molecules • As the motion of these atoms and molecules slows (low T) the interaction becomes apparent. • Halogens Trend!!! • Occurs in all molecules even polar ones

8. London Dispersion Forces - How? Moving e- make a momentary nonsymmetric e- distribution, which produces a temporary dipole. This then can induce a similar dipole in a neighboring atom or molecule. Becomes significant for large atoms with large # of electrons. Termed polarizability of an electron cloud.

9. Characteristics Intermolecular Forces • In general, the stronger the intermolecular forces, the higher the melting and boiling points. • Decrease rapidly with increasing intermolecular distance especially for London dispersion • Nonpolar solids (I2 and CO2) sublimate

10. Dry Ice Sublimation at RTemp

11. Liquids characteristics • Low compressibility, lack of rigidity, and high density compared with gases. • Surface tension – resistance of a liquid to an increase in its surface area: • Liquids with large intermolecular forces tend to have high surface tensions. H2O droplets Playing with Hg video YouTube http://www.youtube.com/watch?v=31CE2BYicyU&feature=fvw • Capillary action – spontaneous rising of a liquid in a narrow tube: YouTube video water special http://www.youtube.com/watch?v=CT4pURpXkbY&feature=related • Cohesive forces – intermolecular forces among the molecules of the liquid. • Adhesive forces – forces between the liquid molecules and their container.

12. Liquid - Cohesive or adhesive? • Which force dominates alongside the glass tube – cohesive or adhesive forces? adhesive forces “Like attract like” determines which will dominate

13. Cohesive vs. Adhesive meniscus graphic Water (polar) interaction with glass surface (polar) and mercury (non-polar) with glass surface (polar)

14. Liquids characteristics - cont • Viscosity – measure of a liquid’s resistance to flow: • Liquids with large intermolecular forces or molecular complexity tend to be highly viscous.

15. Solids • Amorphous solids • Non-uniform structure • glasses • waxes • Crystalline solids • Uniform lattice structure (regular arrangement of atoms) • Unit Cell - smallest repeating unit of the lattice

16. Cubic Unit cell and lattices X-ray diffraction (crystallography) used to determine arrangement of atoms n = integer lambda = wavelength of the X rays d = distance between the atoms theta = angle of incidence and reflection

17. Bragg Diffraction graphic Bragg equation

18. Types of Crystalline solids • Ionic Solids – ions at the points of the lattice that describes the structure of the solid. • Molecular Solids – discrete covalently bonded molecules at each of its lattice points. • Atomic Solids – atoms at the lattice points that describe the structure of the solid.

19. Lattice of crystalline solids

20. Structure and bonding in Metals • Closest Packing: • Assumes that metal atoms are uniform, hard spheres. • Spheres are packed in layers. Like oranges in grocery store display abab packing - 3rd directly over 1st layer - called hexagonal closest pack (hcp)

21. Structure and bonding in Metals (con’t) abca packing - 3rd layer not directly over 1st, 4th layer is over 1st - cubic closest pack (ccp) or face centered cubic (fcc) see next slide

22. Face Centered Cubic (FCC)

23. Metallic Bonding Nearest Neighbors • The Indicated Sphere Has 12 Nearest Neighbors Each sphere in closest packed (both fcp and hcp) has 12 equivalent nearest neighbors. What about bcc ? simple cubic ?

24. Unit cell atoms fcc and hcp 8 x 1/8 spheres and 6 x 1/2 spheres = 4 total atoms in unit cell What about bcc? Or simple cubic? What does that say about density of metals?

25. Metallic Bonding • Sea of electrons - regular array of cations surrounded by its valence electrons

26. Metallic bonding MO model • Band Model (MO Model) - combinations of atomic orbitals. Virtual continuum of levels, called bands. Many semiconductor applications

27. Metal alloys • Metals melted together to make a solution (homogeneous solid!!) - 2 types • Substitutional Alloy – some of the host metal atoms are replaced by other metal atoms of similar size. • Interstitial Alloy – some of the holes in the closest packed metal structure are occupied by small atoms.

28. Metal alloys graphics Which is a substitutional alloy? Which is an interstitial alloy?

29. Network atomic solids 2 main allotropes of carbon (3rd is buckyballs). What is hybridization on each C atom in these two structures?

30. Graphite - sp2 hybridization p-orbitals and Pi system in graphite for 1 layer (sheet). Graphite layers slide by each other because of e- repulsion. Large difference between diamond and graphite is type of bonding

31. Carbon Atoms in Graphite

32. Types and Properties of Solids - Table

33. Vapor pressure graphic Not equilibrium (pressure increasing) Equilibrium (pressure constant) Not closed --> no Pvap just Patm

34. Vapor pressure rate diagram Why does rate of condensation increase? While the rate of evaporation remain essentially constant ?

35. Surface molecule interactions

36. Vapor Pressure definition • Pressure of the vapor present at equilibrium. • The system is at equilibrium when no net change occurs in the amount of liquid or vapor because the two opposite processes exactly balance each other. • The boiling point of the liquid is when the Pvap = Patm • Normal boiling point of liquid is at 1 atm.

37. Vapor pressure trends • Liquids in which the intermolecular forces are strong have relatively low vapor pressures. • Vapor pressure increases significantly with temperature.

38. Vapor pressure of various liquids

39. Pvap rationale Temp vs. KE plot T2 > T1, which means on average more molecules have sufficient energy to overcome liquid intermolecular forces (more evaporate --> rate faster)

40. Pvap - Clausius-Clapeyron equation Plots of In(Pvap) vs. (b) 1/T

41. Clausius–Clapeyron Equation Pvap = vapor pressure ΔHvap = enthalpy of vaporization R = 8.3145 J/K·mol T = temperature (in kelvin) For calculation: to undo ln use ex

42. Vapor pressure calc. problem The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kJ/mol. Calculate the vapor pressure of water at 65°C. 194 torr

43. Heating curve for water Temp changing use Q = c x m x T Temp not changing use Hvap liquid <--> gas Hfus solid <--> liquid Why is Hvap > Hfus ?

44. Phase Diagrams (P,T) • A convenient way of representing the phases of a substance as a function of temperature and pressure: • Triple point • Critical point • Phase equilibrium lines Phase diagram for CO2

45. Phase Diagram for Water What is different about phase diagram for water from most other substances?

46. Concept check As intermolecular forces increase, what happens to each of the following? Why? • Boiling point • Viscosity • Surface tension • Enthalpy of fusion • Freezing point • Vapor pressure • Heat of vaporization