Chapter 11 Liquids and Solids

1. Chapter 11Liquids and Solids

2. Characteristic Properties of Gases, Liquids, and Solids • Intermolecular forces are the attractions that hold molecules together in the liquid and solid states.

3. Kinetic Molecular Theory • What was kinetic molecular theory? • Intermolecular forces: • Do not change with temperature • However, kinetic energy does • So…

4. Example • At room temperature, chlorine is a gas, bromine is a liquid, and iodine is a solid. Arrange the molecules in order of increasing intermolecular forces.

5. Phase Changes • Intermolecular forces determine phase, but temperature and pressure can influence phase too. • Definitions: • Evaporation is the process by which molecules escape from the liquid to the gas phase. • Condensation is the process by which molecules go from the gas phase to the liquid phase.

6. Vapor Pressure • The vapor pressure is the partial pressure of the gas when the rate of evaporation equals the rate of condensation.

7. Dynamic Equilibrium • A state of dynamic equilibrium is one in which the two opposing changes occur at equal rates, so no net change is apparent. At constant temperature

8. Factors that Affect Vapor Pressure • As temperature increases, the vapor pressure of a liquid increases. • The stronger the intermolecular forces, the lower the vapor pressure of the liquid at any temperature.

9. Vapor Pressure Curves Which substance has the weakest intermolecular forces? (a) diethyl ether (b) ethanol (c) water

10. Boiling Point • The boiling point of a liquid is the temperature at which the vapor pressure is equal to the external pressure. • The normal boiling point of a liquid is the temperature at which its equilibrium vapor pressure is equal to 1 atmosphere. • At the boiling point, bubbles filled with vapor form below the surface of the liquid.

11. Other Vaporization Properties • Enthalpy of vaporization (DHvap) is the enthalpy change that accompanies the conversion of one mole of a substance from a liquid to a gas at constant temperature. • The critical temperature is the maximum temperature at which a substance can exist in the liquid state. • The critical pressure is the minimum pressure needed to maintain the liquid state up to the critical temperature.

12. Vaporization and Intermolecular Forces • As the strength of intermolecular forces increase: • vapor pressure of the liquid decreases; • boiling point increases; • enthalpy of vaporization increases; • critical temperature increases.

13. Example Calculation • Butane boils at a temperature of -0.6°C and has a ΔHvap= 22.3 kJ/mol. How much energy is necessary to boil 150 g of butane?

14. Liquid-Solid Equilibrium • The changes of a substance from liquid to solid (freezing) and from solid to liquid (melting or fusion) are also opposing changes that involve a dynamic equilibrium.

15. Definitions • The melting point of a substance is the temperature at which the solid and liquid phases are in equilibrium when the pressure is one atmosphere. • There is very little effect of pressure on the melting point of a solid. • The enthalpy of fusion (DHfus) is the enthalpy change that accompanies the change of one mole of solid into liquid at constant temperature.

16. Heating Curves • A heating curve is a graph of temperature of a sample versus heat added.

17. Only One Phase is Present • When only one phase is present (up to A, B to C, D and after), then q = mCsDT.

18. Phase Transitions • During a phase transition (A to B, C to D), the temperature remains constant and q = DH of the transition.

19. Example: Heating Curve • In the heating curve below, identify the phase transition between A and B.

20. Example: Heating Curve • From the heating curve below, determine which phase (solid, liquid, or gas) has the largest specific heat.

21. Solid-Gas Equilibrium • Sublimation is the direct conversion of a substance from the solid to the gas phase. • Deposition is the reverse of the sublimation process. • Enthalpy of sublimation (DHsub) is the enthalpy change for the conversion of one mole of substance from solid to gas. • DHsub = DHfus + DHvap

22. Enthalpy Diagram for Phase Changes

23. A Phase Diagram • A phase diagram is a graph of pressure versus temperature that shows the region of stability for each phase.

24. Triple Point • There is a unique combination of pressure and temperature, called the triple point (T), at which all three phases (solid, liquid, gas) are at equilibrium.

25. Melting Point and Pressure • The melting point of a substance changes very little with pressure. • The effect of pressure on the melting point of a substance depends on the relative density of the two phases.

26. Melting Point and Pressure • If the solid is denser than the liquid (which is the more common case), the melting point increases with increasing pressure. • If the liquid is denser than the solid (as in H2O), the melting point decreases with pressure.

27. Intermolecular Attractions • Electrostatic forces account for all types of intermolecular attractions. There are three types of attractions: • Dipole-dipole attractions • London dispersion forces • Hydrogen bonding

28. Dipole-Dipole Attractions • Dipole-dipole attractions result from electronic forces between molecular dipoles:

29. London Dispersion Forces • London dispersion forces arise from the attractions between instantaneous dipoles and induced dipoles.

30. Dispersion Forces and Periodic Trends • Polarizability is the ease with which a charge distorts the electron cloud in a molecule. • Polarizability generally increases with the number of electrons in the molecule. • For related series of molecules, London dispersion forces increase going down any group in the periodic table.

31. Boiling Points of Some Nonpolar Substances Substance Molar Mass Boiling Point (C) CH4 16 -184 SiH4 32 -112 GeH4 77 -90 SnH4 123 -52 F2 38 -188 Cl2 71 -35 Br2 160 59 I2 254 184

32. Hydrogen Bonding • The unexpectedly high boiling points of water, ammonia, and hydrogen fluoride requires another kind of intermolecular force.

33. Hydrogen Bonding • Hydrogen bonding occurs between a hydrogen atom bonded to N, O, or F, and a lone pair of electrons on a second N, O, or F. • Hydrogen bonds are sometimes shown as dotted lines.

34. Structure of Solid Water • Hydrogen bonding causes ice to have a lower density than liquid water.

35. Example: Intermolecular Forces • Identify the kind of intermolecular forces, and predict which substance in each pair has the stronger forces of attraction. (a) BF3, BBr3 (b) C2H5OH, C2H5Cl

36. Liquids: Surface Tension • Surface tension is the energy needed to increase the surface area of a liquid. • Surface tension results from intermolecular interactions.

37. Liquids: Capillary Action • Capillary action causes water to rise in a small diameter glass tube. • Capillary action is the result of a competition between: • cohesion: the attraction of molecules for other molecules of the same substance. • adhesion: the attraction of molecules for other molecules of a different substance.

38. Capillary Action • Water rises because adhesion is stronger than cohesion. • Mercury is lowered because cohesion is stronger than adhesion.

39. Liquids: Viscosity • Viscosity is the resistance of a fluid to flow. • The stronger the intermolecular forces of attraction, the greater the viscosity. • Other factors contribute to viscosity as well, like structure, size, and shape of molecules.

40. Solids • A crystalline solid: the units that make up the solid are arranged in a very regular, repeating pattern. • Ionic compounds, metals, and solids of small molecules are usually crystalline. • An amorphous solid lacks the long range order of a crystalline solid. • Most plastics are amorphous solids. (they are polymers)

41. Crystalline Solids • Crystalline solids can be classified by the nature of the forces that hold the units together in a regular arrangement. • These forces are usually referred to as crystal forces.

42. Molecular Solids • Molecular solids consist of atoms or small molecules held together by van der Waals forces and/or hydrogen bonding. • Because these crystal forces are fairly weak, molecular solids are generally soft and low-melting. • Examples are CO, Ar, I2, and most organic molecules.

43. Covalent Network Solids • In a covalent network solid, all of the atoms in a crystal are held together by covalent bonds. • Solids of this kind are high melting and often very hard because strong covalent bonds hold the atoms together. • Some examples of covalent network solids are diamond (C), boron nitride (BN), and silicon dioxide (SiO2).

44. Allotropes • Allotropes are two or more molecular or crystalline forms of an element in the same physical state. • O2 (oxygen) and O3 (ozone) are examples of gas-phase allotropes. • Many elements have two or more allotropes in the solid phase: C, S, P, Sn, among others.

45. Allotropes of Carbon • Graphite and diamond are allotropes of carbon that have different covalent network structures.

46. Ionic Solids • An ionic solid consists of oppositely charged ions, held together by strong electrostatic interactions. • Ionic solids are high melting and usually brittle – they tend to shatter under impact. • Binary compounds made up of a metal and a nonmetal are in this category.

47. Metallic Solids • Metallic solids are formed from metal atoms, and are characterized by high thermal and electrical conductivity, metallic luster, and malleability. • A special kind of bonding, metallic bonding, is needed to account for these unique properties.

48. Metallic Bonding • The electron sea model for metallic bonding views the solid as metal ions in a “sea” of electrons formed from the valence shell electrons. • The electrons are very mobile and adequately account for the conductivity and malleability of metals. • Another model for metallic bonding will be discussed in Chapter 20.

49. Properties of Solids - Summary

50. The Bragg Equation • The distances between layers of atoms in a crystal, as measured by x ray diffraction, are given by the Bragg equation: where l = wavelength of x rays, d = distance between layers of atoms, q = angle of x ray diffraction, and n is a whole number called the order.