You Ever Ask???

1. You Ever Ask??? • Why is water usually a liquid and not a gas? • Why does liquid water boil at such a high temperature for such a small molecule? • Why does ice float on water? • Why do snowflakes have 6 sides? • Why is I2 a solid whereas Cl2 is a gas? • Why are NaCl crystals little cubes?

2. Chapter 11: Liquids, Solids and Intermolecular Forces

3. Chapter Objectives Chapter Topics • Kinetic-Molecular Description of Liquids & Solids • Intermolecular (IMF) Attractions & Phase Changes • Properties of Liquids • Viscosity, Surface Tension, Capillary Action • Evaporation, Vapor Pressure, • Boiling Points & Distillation, Heat transfer • Phase Changes The following will NOT be covered in class Types of crystals

4. States of Matter • List all the differences between • Solids • Liquids • Gases • Kinetic Energy?

5. States of Matter (KE of Matter) The fundamental difference between states of matter is the distance between particles. Intermolecular attractionsin liquids & solids are strong: KE of molecules << IMF

6. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases (depends on T and P).

7. The States of Matter • The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: • The kinetic energy of the particles • The strength of the attractions between the particles

8. Examples • What is the difference between intermolecularforces and intramolecular forces? • List all the intermolecular forces you are familiar with. • List all the intramolecular forces you are familiar with. • Which forces are stronger?

9. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. Intermolecular forces are weaker than intramolecular forces ( to break 2 O-H in bonds in water: 930 kJ/mol; to vaporize water: 43 kJ/mol) Responsible for the existence of condensed states (liquids, solids)

10. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, surface tension, and viscosities (reflect the strength of the bond). INTRAmolecular forces—the forces that holds atoms together o form molecules INTERmolecular forces: the forces between molecules, ions and molecules-ions.

11. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces (electrostatic forces).

12. IMF Problem For each of the molecules below, Determine the geometry of the molecule Determine the polarity of the molecule List the types of intermolecular force which act between pairs of these molecules. • (a) CH4 • (b) PF3 • (c) CO2 • (d) HCN, • (e) HCOOH (methanoic acid)

13. Types of IMF • Ion – Ion • Van der Waals Forces • Dipole – dipole (for molecules with dipole moments) • Dipole – induced dipole • Dispersion forces (London) • Hydrogen Bond (special case of dipole-dipole (IMF) • London Dispersion Forces (induced dipole-dipole) • Ion – induced dipole • Ion – dipole • Total attraction between molecules may depend on more than one type of intermolecular force.

14. Ion – Ion Forces 08M16VD1

15. Ion-Ion Forces Ion-ion forces: electrostatic forces of attraction between __________________ of ionic compounds. Generally very strong → 250 kJ. (not a true intermolecular force) Ionic compounds: metal and nonmetal or polyatomic anions (NH4+) Coulomb’s law & the attraction energy determine: melting & boiling points of ionic compounds the solubility of ionic compounds

16. IMF: Ionic Solids Ion-ion interactions force of attraction between two oppositely charged ions is determined by Coulomb’s law • Energy of attraction between two ions is given by: D:\Media\Movies\08M17AN1.MOV

17. Ion-Ion Forcesfor comparison of magnitude Na+—Cl- in salt These are the strongest forces. Lead to solids with high melting temperatures. NaCl (lattice energy = 788 kJ/mol), mp = 800 oC MgO (lattice energy = 3890 kJ/mol) , mp = 2800 oC

18. Covalent Bonding Forcesfor comparison of magnitude C=C, 610 kJ/mol C–C, 346 kJ/mol C–H, 413 kJ/mol CN, 887 kJ/mol

19. Ion – Dipole Force

20. Ion-Dipole Interactions • Ion-dipole interactions are an important force in solutions of ions. • The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents. Na+(g) + 6H2O(l)→ [Na(H2O)6]+(aq) ΔHrxn = -405 kJ Hydrated Ions?Coordination Number?

21. Attraction Between Ions and Permanent Dipoles

22. Attraction Between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion charge and ion-dipole distance. Measured by ∆H for Mn+ + H2O → [M(H2O)x]n+ -1922 kJ/mol -405 kJ/mol -263 kJ/mol

23. Dipole – Dipole Force D:\Media\Movies\13M04AN2.MOV

24. Dipole-Dipole Forces • Molecules that have permanent dipoles are attracted to each other. • The positive end of one is attracted to the negative end of the other and vice-versa. • These forces are only important when the molecules are close to each other. • Note the difference between solid and liquid. Liquid liquid Solid

25. Effect of Dipole Moment on BP

26. Dipole-Dipole Interactions The more polar the molecule (higher μ), the higher is its boiling point. • Basic attraction : electrostatic, Coulomb’s Law Examples: HCl, CO, SO2, NF3, etc

27. IMF: Dipole-Dipole • Dipole-dipole are of the order of 5 to 20 kJ/mol. (KE due to temp at 25oC about 4 kJ/mol). Cmpds that have these forces (dipole-dipole) are frequently solids and liquids at room temp. • The stronger the forces, the ______ the melting and boiling points of the compounds.

28. Hydrogen Bond: type of dipole-dipole force 13M07AN2

29. Boiling Points of Simple Hydrogen-Containing Compounds The nonpolar series (SnH4 to CH4) follow the expected trend. The polar series follows the trend from H2Te through H2S, but water is quite an anomaly. EXPLAIN!

30. Intermolecular Forces: H-bond Which of these are capable of forming hydrogen bonds among themselves? • CH3OH • C2H4 • CH3NH2 • HCN • NH4+ • KF • CH3COOH

31. Hydrogen Bonding • The dipole-dipole interactions experienced when H is bonded to N, O, or F (HIGH ELECTRONEGATIVITY) are unusually strong. • Hydrogen nucleus is exposed. • We call these interactions hydrogen bonds.

32. Hydrogen Bonding

33. H-Bonding Between Methanol and Water H-bond - + -

34. Hydrogen Bonding in H2O H-bonding is especially strong (40 kJ/mol) in water because • the O—H bond is very polar • there are 2 lone pairs on the O atom Accounts for many of water’s (and other molecules such as DNA, proteins) unique properties such as anomalous high BP and high viscosity.

35. Hydrogen Bonding in H2O Ice has open lattice-like structure. Ice density is < liquid. And so solid floats on water. Snow flake: www.snowcrystals.com

36. Hydrogen Bonding in H2O Ice has open lattice-like structure. Ice density is < liquid and so solid floats on water. One of the VERY few substances where solid is LESS DENSE than the liquid.

37. Hydrogen Bonding H bonds leads to abnormally high boiling point of water. D:\Media\Movies\13M07AN1.MOV See Screen 13.7

38. Boiling Point of Hydrides in ºC

39. Hydrogen Bonding in Biology H-bonding is especially strong in biological systems — such as proteins and DNA. D:\Media\Movies\09S03AN1.MOV DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O. Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases. —adenine with thymine —guanine with cytosine

40. Double helix of DNA Portion of a DNA chain

41. Base-Pairing through H-Bonds

42. Induced Dipole –Induced Dipole (London Dispersion Forces)

43. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

44. London Dispersion Forces Instantaneous dipole The helium atom becomes polar, with an excess of electrons on the left side and a shortage on the right side. Instantenous dipole forms (for an instant)

45. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

46. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

47. London Dispersion Forces • These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called POLARIZABILITY.

48. Forces Involving Dipole -Induced Dipole • Process of inducing a dipole is polarization • Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.

49. IMF: London Dispersion Forces • Induced Dipoles: the temporary separation of positive and negative charges in a neutral particle due to the proximity of an ion, dipole, or another induced dipole. On average μ = 0. • London Dispersion Forces: attractive forces (electrostatic in origin) that arise as a result of temporary dipoles induced in atoms or molecules (instantaneous dipoles). Weak: 0.1- 5 kJ/mol • Dispersion forces allow non-polar molecules to condense. • Exist in all molecules!!!!! Importance depends on the type of intermolecular forces.

50. IMF: London Dispersion Forces • London Forces very weak only attractive force in nonpolar molecules Ar atom Cluster of Ar atoms