1. 1 Intermolecular Forces, Liquids, and Solids

2. 2 Some Characteristic Properties of the States of Matter Gas: Assumes both the volume and shape of the container Is Compressible Diffusion within a gas occurs rapidly Flows readily Liquid Assumes the shape of the portion of the container it occupies Does not expand to fill the container Is virtually incompressible Diffusion within a gas occurs slowly Flows readily Solid Retains its own shape and volume Is virtually incompressible Diffusion within a solid occurs extremely slowly Does not flow

3. 3 The strength of intermolecular forces vary over a wide range, but are generally weaker than covalent or ionic bonds 16 kJ will overcome the intermolecular attraction leading to vaporization 431 kJ of energy is required to break the covalent bond between HCl

4. 4 Dipole Dipole forces Dipole Dipole forces exist between neutral polar molecules, are effective only when polar molecules are very close together, and are generally weaker then ion-dipole interactions Molecules that are attracting one another spend more time near each other than those molecules which repel. Thus the overall effect is a net attraction For two particles of equal mass and size, the strengths of the intermolecular attractions increase with increasing polarity Substance Dipole moment Bp Propane 0.1 231 Dimethyl ether 1.3 249 Methyl chloride 2.0 249 Acetaldehyde 2.7 293

5. 5 Ion-Dipole Forces Ion-Dipole Forces exist between an ion and the partial charge on the end of a polar molecule The magnitude of the interaction depends upon the charge of the ion, the dipole moment of the polar molecule and the distance from the center of the ion to the midpoint of the dipole Na + Cl -

6. 6 London Dispersion Forces London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment The instantaneous distribution at any given moment can be different from the average distribution producing an instantaneous dipole moment Because electrons repel, the motions of electrons in one atom influence the motions of electrons on its near neighbors

7. 7 London Dispersion Forces London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment The ease with which the charge distribution in a molecule can be distorted by an external force is called its polarizability. The greater the polarizability of a molecule, the more easily its electron cloud can be distorted to give a momentary dipole. In general, larger molecules tend to have greater polarizability because their electrons are farther from the nucleus Because molecular size and molecular mass tend to parallel one another, dispersion forces tend to increase in strength with increasing molecular weight. Halogen Boiling pt (K) Noble gas Boiling point (K) F 2 85.1 He 4.6 Cl 2 238.6 Ne 27.3 Br 2 332.0 Ar 87.5 I 2 457.6 Kr 120.9

8. 8 London Dispersion Forces London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment The shapes of molecules can play a role in the magnitude of dispersion forces.

9. 9 London Dispersion Forces London Dispersion Forces can exist between non-polar atoms and molecules where the movement of electrons create an instantaneous dipole moment London dispersion forces may also operate within polar molecules and contribute to the overall attractive forces between the molecules; some times more than the dipole-dipole forces HCl: bp = 189.5 K Dipole Moment = 1.03 D HBr: bp = 206.2 K Dipole Moment = 0.79 D

10. 10 Hydrogen bonding Hydrogen bonding is a special type of intermoleculer attraction that exists between the hydrogen atom in a polar bond (e.g., H F, H O, or H N) and an unshared electron pair on a nearby electronegative atom (usually an F, O, or N atom on another molecule) Because F, N, and O are so electronegative, a bond between hydrogen and any of these elements is quite polar, with hydrogen at the positive end.

11. 11 Metallic solids consist of metal atoms. Bonding is due to valence electrons that are delocalized throughout the entire solid. We can visualize these electrons as an array of positive ions immersed in a sea of delocalized electrons.

12. 12 Intermolecular Forces, Liquids, and Solids Metallic bond examples: Cu, Fe, or alloy

13. 13 Intermolecular Forces, Liquids, and Solids Properties of Liquids: Viscosity and Surface Tension The resistance of liquids to flow is called their viscosity. The greater the viscosity, the Viscosity decreases with increasing temperature more slowly the liquid flows. Viscosity decreases with increasing temperature. Viscosity is related to the ease with which individual molecules of a liquid can move with respect to one another. It thus depends on the attractive forces between molecules and and their structural character.

14. 14 Intermolecular Forces, Liquids, and Solids Properties of Liquids: Viscosity and Surface Tension Molecules in the interior are attracted equally in all directions, whereas those at the surface experience a net inward force This net inward force pulls molecules from the surface into interior, thereby reducing the surface area Surface tension is therefore defined as the energy required to increase the surface of a liquid by a unit amount Forces that bind like molecules to one another are called cohesive forces. Forces that bind molecules to a surface are called adhesive forces

15. 15 Changes of State Whenever a change of state involves going to a less ordered state, energy must be supplied in order to overcome intermolecular forces Melting is also called fusion. The enthalpy of change associated with melting is called heat of fusion. (e.g. 6.01 kJ/mol for water) The heat needed for vaporization is called the heat of vaporization (e.g. 40.67 kJ/mol for water) ENDOTHERMIC EXOTHERMIC

16. 16 Changes of State Enthalpy and Temperature Changes Accompanying Heating

17. 17 Calculate the enthalpy change associated with converting 1.00 mol of ice at -25°C to water vapor at 125 °C and 1 atm. The heat capacities per gram (specific heats) of ice, water, and steam are 2.09 J/g-°C, 4.18 J/g-°C, and 1.84 J/g-°C, respectively. The heat of fusion of ice is 6.01 kJ/mol, and the heat of vaporization of water is 40.67 kJ/mol Heat required to bring the ice from -25 °C to 0 °C : 1.00 mol H 2 O 18.0 g. 1.00 mol H 2 O 2.09 J g -°C 25 °C= 940 J Heat of fusion: 1.00 mol H 2 O 6.01 kJ mol H 2 O = 6.01 kJ Heat required to bring liquid water from 0 °C to 100 °C : 1.00 mol H 2 O 18.0 g. 1.00 mol H 2 O 4.18 J g -°C 100 °C=7520 J Heat of vaporization 1.00 mol H 2 O 40.67 kJ mol H 2 O = 40.67 kJ Heat required to raise the temperature of the vapor to 125 °C 18.0 g. 1.00 mol H 2 O 1.84 J g -°C 25°C=830 J  H t =.940 kJ + 6.01 kJ + 7.52 kJ + 40.67 kJ +.830 kJ = 55.97 kJ 1.00 mol H 2 O

18. 18 Changes of State Vapor Pressure: pressure exerted by evaporating liquid in the space above that liquid once dynamic equilibrium has been established The vapor pressure of a substance increases as the temperature increases Substances with a high vapor pressure evaporate more rapidly than thosewith low vapor pressures. Theses substances are more volatile

19. 19 Changes of State Vapor Pressure and Boiling Point: A liquid boils when its vapor pressure equals the external pressure acting on the surface of the liquid The higher the outside pressure, the higher the boiling point

20. 20 A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the difference states of matter. Line A-B is the vapor curve of the liquid The line A-C represents the variation in vapor pressure during sublimation. Line A-D represents the change in melting point of the solid under increased pressure

21. 21 A phase diagram is a graphical way to summarize the conditions under which equilibria exist between the difference states of matter. Note that the gas phase is the stable phase under condition of low pressure and high temperature Note that the solid phase is the stable phase under condition of high pressure and low temperature

22. 22 Why are the fusion-freezing lines different for H 2 O and CO 2 ? Why does CO 2 sublime under normal conditions? How is freeze-dry food prepared?

23. 23 AS THE TEMPERATURE IS INCREASED AS THE TEMPERATURE IS INCREASED : At point 6-A, the H 2 O exists entirely as a solid. When the temperature reaches point 4-B, the solid begins to melt, and an equilibrium condition occurs between the solid and the liquid. At a yet high temperature, point 7-C, the solid has been converted entirely to a liquid. When point 8-D is encountered, vapor forms, and a liquid-vapor equilibrium is achieved. Upon further heating, to point 9-E, the H 2 O is converted entirely to the vapor phase. 6 4 7 9 8 7 53215321

24. 24 A crystalline solid is a solid whose atoms, ions, or molecules are ordered in well-defined arrangements. Amorphous solids have no structure. A unit cell is the repeating unit that makes up the crystalline solid. The array of repeating points is called the crystal lattice

25. 25 The lattices of crystalline solids can be described in terms of seven basic types. The following three types are the simplest of the seven. Coordination number = # of nearest neighbors # of atoms in unit cell

26. 26 The Crystalline (Body-Centered) Lattice of Sodium Chloride Na + Cl -

27. 27 The total cation-anion ratio in each unit cell must be the same as that for the entire crystal. Therefore a unit cell of sodium chloride must have equal number of Na + and Cl - Na + Cl - (¼ Na + per edge ) (12 edges) = 3 Na + (1 Na + per center) (1 center ) = 1 Na + Na + ( 1/8 Cl - per corner ) (8 corners) = 1 Cl - ( 1/2 Cl - per face) (6 faces ) = 3 Cl - Cl -

28. 28 Aluminum crystallizes in a face- centered arrangement a.how many atoms are in the unit cell? b.what is the coordination number? c.each atom has diameter of 2.86 Å, what is the length of a side? d. what is the density? °

29. 29 The Close Packing of Spheres: A Study of Metallic Solids Lets make an assumption using Methane as a example that many molecules may be approximated as being roughly spherical

30. 30 The coordination number of close packing structures = 12 The coordination number of body-centered cubic structures = 8 The coordination number of primitive cubic structures = 6 The Close Packing of Spheres: A Study of Metallic Solids hexagonal close pack cubic close pack = face-centered cubic

31. 31 Using X-Ray Diffraction to Determine the Structure of Crystalline Solids

32. 32 Covalent Network Solids are Held together in large networks or chains by covalent bonds. These solids are much harder and have higher melting points than molecular solids Molecular solids consist of atoms or molecules held together by dipole-dipole or London dispersion forces or hydrogen bonds

33. 33 Ionic solids consist of ions held together by ionic bonds. The strength of the bonds depend upon the charges of the ions. Melting point increases as bond strength increases The structure of the solid is dependent upon the charges and relative sizes of the ions. The coordination number increases as the ratio of the cation radius to the anion radius increases. NaCl

34. 34 Atomic solids – (london) eg. He, Ar etc. atoms exist at lattice pts molecular solids (London, dipole, H-Bond) CH 4, NH 3 molecules exist at lattice pts Ionic solids (electrostatic attractions) NaCl cations & anions exist at lattice pts Metallic solids- (Cations held by a sea of electrons) Cu, Fe cations exist at lattice pts Network solids – (3 dimensional covalent bonds) diamond (C), quartz (SiO 2 ) atoms exist at lattice pts