Homework Problems

1. Homework Problems Chapter 11 Homework Problems: 2, 14, 16, 18, 22, 28, 32, 36, 46, 56, 60, 64, 70, 75, 84, 103

2. CHAPTER 11 Intermolecular Forces and the Physical Properties of Liquids and Solids

3. Solids, Liquids, Gases The three common states of matter are solids, liquids, and gases. Their properties can be summarized as follows stateproperties solid definite volume and definite shape liquid definite volume but indefinite shape gas indefinite volume and indefinite shape

4. Compressibility Compressibility refers to the decrease in the volume occupied by a substance when the pressure applied on the substance increases. Because the particles making up solids and liquids are in close contact with one another, these two phases are not easily compressed (to a first approximation solids and liquids are incompressible). Since gases are mostly empty space, they are highly compressible. This also means that solids and liquids will be much higher density than gases.

5. Effect of Temperature and Pressure on Phase The phase of a substance depends on both temperature and pressure. Generally speaking, substances go from solid to liquid to gas as temperature increases. It is often possible to condense a gas into a liquid or solid by increasing pressure. We will discuss this further later in the chapter.

6. Intermolecular Forces Intermolecular forces are the forces that exists between different molecules or particles. We are more concerned with long range attractive forces and will ignore short range repulsive forces. Ion-ion - The attractive force acting between cations and anions. These are strong, and are found in substances where ionic bonding occurs..

7. Dipole-Dipole Forces Dipole-dipole - The attractive force acting between polar molecules. The attraction is between the partial positive charge (+) on one molecule and the partial negative charge (-) on a different molecule. Generally speaking, the larger the partial positive and negative charges the stronger the dipole-dipole attraction.

8. Dipole-Dipole Forces and Boiling Point When molecules have strong intermolecular attractive forces it takes more energy to overcome those attractive forces. One way of seeing this is in the boiling point for a substance. Generally speaking, the stronger the dipole-dipole attraction between molecules the higher the boiling point, parti-cularly for substances with approximately the same molecular mass.

9. Hydrogen Bonding Hydrogen bonding - A particularly strong form of dipole-dipole attractive force. It is the attractive force that exists between a hydrogen atom bonded to an N, O, or F atom and lone pair electrons on a different N, O, or F atom (often an atom in a different molecule).

10. Evidence For Hydrogen Bonding One effect of hydrogen bonding is to raise the boiling point of a liquid. This occurs because it requires more energy (and so a higher temperature) to break apart strong attractive forces between molecules than it does to break apart weak attractive forces between molecules. substanceboiling pointhydrogen bonding? H2O 100.0 C yes H2S - 60.7 C no H2Se - 41.5 C no H2Te - 4.4 C no

11. Boiling Points for Binary Hydrogen Compounds

12. London Dispersion Forces London dispersion forces - The attractive force that is due to the formation of instantaneous dipoles in a molecule. These instantaneous dipoles arise from the random motion of the electrons in the molecule.

13. Strength of London Dispersion Forces London dispersion forces are present in all molecules, but are the only intermolecular force present in nonpolar molecules. The strength of London dispersion forces is approximately proportional to the number of electrons, and so to the size of the molecule. Therefore, as a general rule, the larger the molecule the stronger the London dispersion forces. substanceboiling point He - 268.6 C Ne - 245.9 C Ar - 185.7 C Kr - 152.3 C Xe - 107.1 C

14. Induced Dipole and Polarizability London dispersion forces are closely related to a second property of molecules, called polarizability. The polarizability of a molecule refers to the extent to which the electron cloud distribution is distorted when a positive or negative charge is brought close to the molecule. In general, large molecules, and molecules containing atoms whose valence electrons are far away from the atomic nucleus are more polarizable than other molecules. The more polarizable a molecule the stronger the London dispersion forces it experiences.

15. Ion-Dipole Forces Ion-dipole - The attractive force between an ion and a polar molecule. Responsible for the dissolution of some ionic substances in polar liquids such as water. In general, the smaller the ion and the larger the charge the stronger the ion-dipole attractive force. Solvation - The close association of solvent molecules with solute molecules or ions. Hydration - Solvation when the solvent is water.

16. Summary of the Types of Intermolecular Forces Ion – ion. Forces between cations and anions. Dipole – dipole. Forces between molecules with a permanent dipole moment. This category includes hydrogen bonding, a particularly strong type of dipole – dipole force. London dispersion forces. Due to random movement of electrons. All particles have this type of force, but it is most important in molecules with no permanent dipole moment. Mixed forces: Ion – dipole is the most important, but this also includes ion – induced dipole and dipole – induced dipole. It is responsible for the solubility of some ionic compounds in water. We will discuss these and related forces in detail when discussing solution formation.

17. Properties of Liquids There are several properties of liquids that are related to the strength of the intermolecular forces acting between molecules. Surface tension - Resistance of a liquid to spreading out. Viscosity - Resistance of a liquid to flow. Generally speaking, the stronger the intermolecular forces, the larger the values for viscosity and surface tension.

18. Surface Tension Surface tension is the resistance of a liquid to spreading out. A molecule within the liquid feels forces from all of the surrounding molecules. However, at the surface of the liquid a molecule is attracted back to the liquid by the molecules below. This attractive force tends to cause liquids to arrange themselves to minimize their surface area (the reason small droplets of water form beads.

19. Cohesion and Adhesion Cohesion refers to the attractive force between molecules of the same type. A second term, adhesion, refers to the attractive force between different types of molecules, such as between molecules of a liquid and those making up the surface of the container. In a narrow tube the competition between these forces affects the shape of the liquid meniscus, the curved surface of the liquid.

20. Capillary Action When adhesive forces are stronger than cohesive forces, a liquid will be drawn up into a narrow diameter tube, by a process called capillary action. Note that the narrower the dia-meter of the tube the higher the level of the liquid drawn into the tube. For liquids like mercury, where adhesive forces are larger than cohesive forces, a reverse process occurs.

21. Viscosity Viscosity is the resistance of a liquid to flow. Liquids with strong intermolecular attractive forces will flow more slowly than those where such forces are small. The device at right is a viscometer. It can be used to measure the viscosity of a liquid.

22. Vapor Pressure Vapor pressure is defined as the equilibrium partial pressure of vapor above a pure solid or gas. The vapor pressure of substance increases as temperature increases. The normal boiling point for a substance corresponds to the temperature at which the liquid and vapor pressures are at equilibrium and the vapor pressure is 1.00 atm.

23. Clausius-Clapeyron Equation Based on experiment, the vapor pressure above a liquid is found to obey a simple equation, called the Clausius-Clapeyron equation ln(p) = - Hvap + C RT Based on this equation we expect a plot of ln(p) vs 1/T to give a straight line with slope m = - Hvap/R. The above equation may be used to find a second useful expression ln(p2/p1) = - (Hvap/R) [ (1/T2) – (1/T1) ] where p1 is the vapor pressure at temperature T1 p2 is the vapor pressure at temperature T2

24. Clausius-Clapeyron Equation (Example) slope = - Hvap/R

25. Example: The normal boiling point for water occurs at T = 100.0 C. The enthalpy of vaporization for water is Hvap = 40.67 kJ/mol. Based on this information estimate the vapor pressure of water at T = 20.0 C.

26. Example: The normal boiling point for water occurs at T = 100.0 C. The enthalpy of vaporization for water is Hvap = 40.67 kJ/mol. Based on this information estimate the vapor pressure of water at T = 20.0 C. Recall that one form of the Clausius-Clapyron equation is: ln(p2/p1)= - Hvap1 _ 1 R T2 T1

27. Example: The normal boiling point for water occurs at T = 100.0 C. The enthalpy of vaporization for water is Hvap = 40.67 kJ/mol. Based on this information estimate the vapor pressure of water at T = 20.0 C. Let T1 = 100.0 C = 373. K ; p1 = 1.00 atm T2 = 20.0 C = 293. K ln(p2/p1) = - 40670. J/mol11 = - 3.581 (8.314 J/mol.K) 293. K 373. K So (p2/p1) = e-3.581 = 0.0279 p2 = (0.0279) p1 = (0.0279) (1.00 atm) = 0.0279 atm = 21. torr

28. Solids Solids can be divided into two general categories Crystalline solid - Has a regular arrangement of the particles making up the solid (a crystal structure). Four main types exist: ionic, molecular, covalent, and metallic solids. Amorphous solid - Does not have a regular arrangement of the parti-cles making up the solid (no regular crystal struct-ure).

29. Crystal Structure For crystalline solids, the crystal structure of a solid substance refers to the arrangement of the particles making up the solid. This is often given in terms of the unit cell, the smallest part of the crystal that can be used to construct the entire crystal.

30. Examples of Crystal Structures Examples of different types of crystal structures are given below. Note that information about crystal structures is determined using techniques such as X-ray diffraction or electron diffraction.

31. Types of Crystalline Solids Crystalline solids can be classified into four groups. 1) Ionic – Composed of cations and anions; held together by ionic bonding. 2) Covalent – Composed of atoms where every atom is attached to other atoms in the solid through a network of covalent bonds. 3) Molecular – Composed of molecules; held together by weak intermolecular forces (dipole-dipole, hydrogen bonding, or dispersion) 4) Metallic – Composed of metal atoms, held together by metallic bonding.

32. Ionic Solids An ionic solid is composed of cations and anions. Ionic solids are held together by the strong electrostatic force of attraction that exists between particles of opposite charge. Examples: NaCl, CaCO3, Al2N3, FeCl3. Properties Hard and brittle High melting point High boiling point Poor conductors of heat and electricity in solid state Good conductors of electri-city when dissolved in water NaCl, Tfus = 801 C, Tvap = 1413 C

33. Covalent Solids A covalent solid (sometimes called a network covalent solid) is a “supermolecule” in which every atom is connected to every other atom through a network of covalent bonds. Because all of the atoms are connected by covalent bonds, these substances are generally among the hardest substances known. Exam-ples: C (diamond), C(graphite), SiO2 (quartz). Properties Extremely hard Very high melting point Very high boiling point Poor conductors of heat and electricity in solid state Insoluble in water diamond C , Tfus > 3550 C, Tvap = 4827 C

34. Molecular Solids A molecular solid is composed of molecules. Molecular solids are held together by the weak van der Waals attractive forces (dipole-dipole and London dispersion forces) that exist between molecules. Examples: H2O, Ar, CS2, C10H8 (naphthalene), C6H12O6 (sugar). Properties Soft Low melting point Low boiling point Poor conductors of heat and electricity in solid state Poor conductors of electri-city when dissolved in water CS2, Tfus = - 111 C, Tvap = 46 C

35. Metallic Solids A metallic atomic solid represents the solid form for metals. Metallic solids can be thought of as metal cations immersed in a sea of loosely held valence electrons. Examples: K, Fe, Cu, Na, Pb. Properties Can be hard or soft Low to high melting point Low to high boiling point Good conductors of heat and electricity in solid state Insoluble in water Na, Tfus = 98 C, Tvap = 892 C

36. Phase Transitions The conversion of a substance from one phase to another phase is called a phase transition. Transitions can be caused both by adding heat and by removing heat from a substance. adding heat (H > 0)removing heat (H < 0) s   fusion (melting)   s freezing   g vaporization g   condensation s  g sublimation g  s deposition Recall that the enthalpy change for a phase transition is usually reported at the normal transition temperature, that is, the temperature at which the phase transition occurs when p = 1.00 atm. Since enthalpy is a state function: Hfreez = - Hfus Hcond = - Hvap Hdep = - Hsub

37. Relationships Among Phase Transitions Based on Hess’ law, we would expect Hsub = Hfus + Hvap

38. Thermodynamics of Phase Transitions We can study the thermodynamics of phase transitions by finding the heating curve for a substance. This is simply a plot of temperature vs. amount of heat added, under conditions where the heat is added slowly enough to maintain equilibrium. Experimentally we ex-pect to see two regions in the heating curve. Normally the temperature of the substance will increase as heat is added. However, at the temperature where a phase transition occurs the added heat will be used to carry out the transi-tion, and so temperature will remain constant until the phase transition is complete.

39. Sample Heating Curve

40. Phase Diagram A phase diagram is a diagram indicating which phase or phases are present at equilibrium as a function of pressure and temperature. H2O

41. Important Features in a Phase Diagram Phase boundaries - Indicate where two phases can exist simul-taneously at equilibrium. Triple point - Indicates a point where three phases can exist simultaneously at equilibrium. Normal melting point - Solid-liquid equilibrium at p = 1.00 atm. Normal boiling point - Liquid-gas equilibrium at p = 1.00 atm. Normal sublimation point - Solid-gas equilibrium at p = 1.00 atm. (Note that substances will have either a normal melting and normal boiling point, or a normal sublimation point, but not both.) Critical point - Point below which a gas will undergo a phase transition (g   or g  s) when compressed reversibly at constant temperature. Above the critical point no such phase transition occurs. In this region of the phase diagram a supercritical fluid is present.

42. Phase Diagram For CO2 CO2

43. End of Chapter 11 “Gibbs is perhaps the most brilliant person most people have never heard of. Modest to the point of near-invisibility, he passed virtually the whole of his life, apart from three years spent studying in Europe, within a three-block area bounded by his house and the Yale campus in New Haven, Connecticut. For his first ten years at Yale he didn't even bother to draw a salary. (He had independent means.) From 1871, when he joined the university as a professor, to his death in 1903, his courses attracted an average of slightly over one student a semester.” - Bill Bryson A Short History of Nearly Everything “Of all chemical bonds, hydrogen bonds are the weakest, the most important, the least understood, and the hardest to measure.” - John Emsley, “Science Watch” (2000)