Chapter 11 Intermolecular Forces, Liquids, and Solids

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1 Chapter 11 Intermolecular Forces, Liquids, and Solids

2 Dissolution of an ionic compound

3 States of Matter The fundamental difference between states of matter is the distance between particles.

4 States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

5 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.

6 At room temperature and pressure, He, Ar, Kr, Xe, and Rn are monatomic gases; H 2, N 2, O 2, F 2, and Cl 2 are diatomic gases; Hg and Br 2 are liquids. All other elements are solids.

7 Bonding Forces

8 Interaction Energy Intermolecular Forces E(kJ/mol)

9 Intermolecular Forces Ion-ion Interactions Ion-dipole Interactions Dipole-dipole Interactions Hydrogen bonds Dispersion Forces (London Forces) Dipole-induced Dipole interactions

10 Interaction Energy Intermolecular Forces E(kJ/mol)

11 What kind of intermolecular forces must be overcome during the following phase changes? CO 2 (s) Sublimes H 2 O(s) Melts CH 3 Cl(l) boils NH 3 (l) vaporizes Are these processes endothermic or exothermic processes?

12 Dipole-Dipole Interactions The more polar the molecule, the higher its boiling point.

13 Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds which are easier to polarize.

14 Factors Affecting London Forces The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). This is due to the increased surface area in n-pentane.

15 Which Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

16 How Do We Explain This? The nonpolar series (SnH 4 to CH 4 ) follow the expected trend. The polar series follows the trend from H 2 Te through H 2 S, but water is quite an anomaly.

17 Coulomb s law states that the energy (E) of the interaction between two ions is directly proportional to the product of the charges of the two ions (Q 1 and Q 2 ) and inversely proportional to the distance (d) between them. E (Q 1 Q 2 ) d

18 Coulombs Law indicates the increases in the charges of ions will cause an increase in the force of attraction between a cation and an anion. Increases in the distance between ions will decrease the force of attraction between them.

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20 Lattice energies are proportional to the Coulombic interaction energies, but also depend on the specific arrangement of the ions. Ionic compounds are generally solids at room temperature. They are solids because intermolecular attraction between ions is the strongest type of intermolecular interaction. The cations and the anions forming an ionic compound often are arranged in well-defined crystal lattice.

21 The lattice energy (U)* of an ionic compound is the energy released when one mole of the ionic compound forms from its free ions in the gas phase. M + (g) + X - (g) ---> MX(s) U = k(q 1 Q 2 ) d *Chapter 8, page 304

22 Lattice Energies of Common Ionic Compounds Compound U(kJ/mol) LiF LiCl -864 NaCl -790 KCl -720 KBr -691 MgCl MgO -3791

23 Determine which salt has the greater lattice energy. A. MgO and NaF B. MgO and MgS

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25 Born-Haber Cycle Is a method for estimating the Lattice Energies for ionic compounds. This is done because Lattice Energies are difficult to measure, and we use Hess s Law to treat the energies associated with it to a series of steps with known energy values. Our starting place for the cycle is the overall reaction and the known reaction enthalpy. Na(s) + ½ Cl(g) NaCl(s) H 0 R = -411kJ

26 Electron Affinity Electron affinity is the energy change occurring when one mole of electrons combines with one mole of atoms or ion in the gas phase. Step 4 in diagram on the last slide. Cl(g) + e - (g) ---> Cl - (g)

27 Successive Electron Affinities (EA) may be defined, e.g. O(g) + e - O - (g) O - (g) + e - O 2- (g) H EA1 = kj H EA2 = kj O(g) + 2e - O 2- (g) Overall H EA = kj

28 Na + (g) + e - (g) ---> Na(g) Na(g) ---> Na(s) Cl - (g) ---> Cl(g) + e - (g) - H IE1 - H sub - H EA Cl(g) ---> 1/2Cl 2 (g) -1/2 H BE Na(s) + 1/2Cl 2 (g) ---> NaCl(s) H f Na + (g) + Cl - (g) ---> NaCl(s) U U = H f - 1/2 H BE - H EA - H sub - H IE1

29 Problem Calculate the lattice energy of potassium chloride from the following data: Ionization energy of K = 425 kj/mol Electron Affinity of Cl for 1 e - = -349 kj/mol Energy to vaporize K = 89 kj/mol Cl 2 bond energy = 240 kj/mol Energy change for the reaction: K(s) + ½ Cl 2 (g) KCl (s) is -438 kj/mol

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32 Summarizing Intermolecular Forces

33 Intermolecular Forces Affect Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution.

34 Viscosity Resistance of a liquid to flow is called viscosity. It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

35 Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

36 Phase Changes

37 Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid. Note the heat of sublimation is the sum of the heats of vaporization and fusion.

38 Energy Changes Associated with Changes of State The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during a phase change.

39 Vapor Pressure At any temperature some molecules in a liquid have enough energy to escape. As the temperature rises, the fraction of molecules that have enough energy to escape increases.

40 Vapor Pressure As more molecules escape the liquid, the pressure they exert increases. At equilibrium, the rate of molecules leaving the liquid phase is equal to those returning.

41 The boiling point of a liquid is the temperature at which it s vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure equals 760 torr. Vapor Pressure

42 The vapor pressure curves represent how P 0 varies with T, but is related to the strength of the intermolecular forces by the enthalpy of vaporization for the liquid ( H 0 vap) For an Ideal solution we can use the Ideal gas law to derive the functional form for these curves. This function is called the Clausius- Clapeyron equation ln(p 1 /P 2 )= -( H 0 /R)[T 2-1 T 1-1 ] Therefore, plots of P vs. 1/T can be used to determine H vap

43 Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

44 Phase Diagrams The circled line is the liquid-vapor interface. It starts at the triple point (T), the point at which all three states are in equilibrium.

45 Phase Diagrams It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

46 Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

47 Phase Diagrams The circled line in the diagram below is the interface between liquid and solid. The melting point at each pressure can be found along this line.

48 Phase Diagrams Below the triple point the substance cannot exist in the liquid state. Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

49 Phase Diagram of Water Note the high critical temperature and critical pressure. These are due to the strong van der Waals forces between water molecules.

50 Phase Diagram of Water The slope of the solidliquid line is negative. This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

51 Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures.

52 Phase Diagram of Carbon Dioxide The low critical temperature and critical pressure for CO 2 make supercritical CO 2 a good solvent for extracting nonpolar substances (like caffeine)

53 Solids We can think of solids as falling into two groups: crystalline, in which particles are in highly ordered arrangement.

54 Solids We can think of solids as falling into two groups: amorphous, in which there is no particular order in the arrangement of particles.

55 Attractions in Ionic Crystals In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.

56 Crystalline Solids Because of the ordered in a crystal, we can focus on the repeating pattern of arrangement called the unit cell.

57 Crystalline Solids There are several types of basic arrangements in crystals, like the ones depicted above.

58 Crystalline Solids We can determine the empirical formula of an ionic solid by determining how many ions of each element fall within the unit cell.

59 Ionic Solids What are the empirical formulas for these compounds? (a) Green: chlorine; Gray: cesium (b) Yellow: sulfur; Gray: zinc (c) Gray: calcium; Blue: fluorine (a) (b) (c) CsCl ZnS CaF 2

60 Types of Bonding in Crystalline Solids

61 Covalent-Network and Molecular Solids Diamonds are an example of a covalentnetwork solid, in which atoms are covalently bonded to each other. They tend to be hard and have high melting points.

62 Covalent-Network and Molecular Solids Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces. They tend to be softer and have lower melting points.

63 Metallic Solids Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. In metals valence electrons are delocalized throughout the solid.