1. Chapter 11 Intermolecular Forces, Liquids, and Solids CHEMISTRY The Central Science 9th Edition Summer 2005 Instructor: Dr. Michael Curry

2. The Behavior of Gases The ideal gas law (PV=nRT) Used to describe how ideal gases behave. However, does not explain why gases behave in such a fashion. The density of a gas can be calculated using the Ideal gas law by multiplying it times it molar mass. In the real world, gases deviate significantly from ideal behavior.

3. Kinetic molecular theory developed by Rudolf Clasius (ca. 1857) to explain gas behavior (PV=nRT only describes the behavior of gases). Why does gases expand when heated? Why do pressure increase as gases are compressed? Sometimes called the “theory of moving molecules”. Basically, kinetic molecular theory gives us an understanding of pressure and temperature on the molecular level. Kinetic Molecular Theory

4. Gases consist of a large number of molecules in constant random motion. Volume of individual molecules negligible compared to volume of container. Intermolecular forces (attractive and repulsive) between gas molecules are negligible. Collisions between molecules are perfectly elastic at constant temperature (no energy is lost). The average kinetic energy of the molecules are proportional to the absolute temperature (i.e., all molecules possess the same average kinetic energy at any given temperature.) Kinetic Molecular Theory Assumptions

5. Questions?

6. Physical properties of substances understood in terms of kinetic molecular theory: Liquids are almost incompressible, assume the shape but not the volume of container: –Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other. Solids are incompressible and have a definite shape and volume: –Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other. A Molecular Comparison of Liquids and Solids

7. Comparison of Liquids and Solids Cont.

8. Converting a gas into a liquid or solid requires the molecules to get closer to each other: Cooling or compressing will result in molecules with smaller distances between them. Converting a solid into a liquid or gas requires the molecules to move further apart: Heating or reducing pressure will result in molecules with larger distances between them. Converting Liquids to Solids and Gases

9. Structure of Liquids, Solids, and Gases

10. These are the forces holding solids and liquids together are called intermolecular forces. Intermolecular Forces

11. The attraction between molecules is an intermolecular force. Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl). When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). However, when a substance condenses, intermolecular forces are formed. What are Intermolecular Forces

12. Vaporizing HCL Vapor phase Liquid phase Heating Liquid Solutions Formation intermolecular forces intermolecular forces are broken

13. Inter- vs. Intramolecular Forces The covalent bond holding a molecule together is an intramolecular force.

14. Properties Reflecting Molecular Force Strengths Boiling and melting points reflect the strengths of intermolecular forces. High boiling points indicate strong attractive forces between molecules. -For example, HCl boils at 85 o C at room temperature due to its weak attractive forces. Melting points increase with increasing attractive forces (i.e., molecules become harder to separate).

15. Types of Molecular Forces There are four types of molecular forces: Ion-dipole Forces Dipole-dipole Forces London Dispersion Forces Hydrogen Bonding Forces The lateral three forces are general called van der Waals forces (developed by Johannes van der Waals) and exist between neutral molecules The ion-dipole forces exist between ions and polar molecules.

16. Water (H 2 O)

17. Molecular Polarity

18. Interaction between an ion and a dipole. Dipole is a polar molecule (e.g. water). Strongest of all intermolecular forces. Ion-dipole Forces

19. Dipole-dipole forces exist between neutral polar molecules. Only effective when polar molecules are close together. These forces are weaker than ion-dipole forces. There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble (free flow in liquids) If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity. Dipole-dipole Forces

20. Dipole-dipole Forces Schematic

21. Dipole-dipole Forces Trend

22. Weakest of all intermolecular forces. Primary property that cause nonpolar substances to condense to liquids and to freeze into solids at low temperatures. Form when electrons occupy positions around the nucleus in two adjacent atoms causing an temporary dipole. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). London Dispersion Forces

23. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). For an instant, the electron clouds become distorted. In that instant a dipole is formed (called an instantaneous dipole). Formation of London Dispersion Forces

24. Dispersion forces are present in all molecules whether polar or nonpolar The larger the molecule (the greater the number of electrons) the more polarizable. London dispersion forces increase as molecular weight increases. London dispersion forces depend on the shape of the molecule. Example: neopentane (gas at 25 o C), n-pentane (liquid at 25 o C) Properties Effecting London Dispersion Forces

25. The Effect of Molecular Shape Higher boiling point indicating stronger dispersion forces. Cylindrical shape of n-petane allows more contact.

26. Trends in London Dispersion Forces Notice that as the molecular weight increases the boiling points of the halogen increases, indicating greater London dispersion forces between atoms.

27. Many elements form compounds with hydrogen - referred to as “Hydrides”. Plotting the boiling points of Group 4 elements show an increase as you go down the group. The increasing boiling point results from stronger dispersion forces because an increase in electron density. Evidence for Hydrogen Bonding Forces Reproduced from http://www.chemguide.co.uk/atoms/bonding/hbond.html

28. Evidence for Hydrogen Bonding Forces Cont. Repeating this for hydrides of elements in Groups 5, 6, and 7, this trend is deviated from with the first element. The hydrides H 2 O, NH 3, and HF have abnormally high boiling points. Reproduced from http://www.chemguide.co.uk/atoms/bonding/hbond.html

29. By experiments: boiling points of compounds with H- F, H-O, and H-N bonds are abnormally high. In the case of NH 3, H 2 O, and HF, additional intermolecular forces must be present which increases the amount of heat energy needed to separate the atoms. These additional intermolecular forces are called hydrogen bonds. Hydrogen Bonding

30. Notice that the hydrogen is attached to the most electronegative elements. Thus, causing the hydrogen to acquire a significant amount of positive charge. The Origin of Hydrogen Bonding

31. Hydrogen Bonding Schematic

32. Hydrogen bonds are responsible for: Ice Floating –Ice is ordered with an open structure to optimize H-bonding. –Therefore, ice is less dense than water. –In water the H-O bond length is 1.0 Å. –The O…H hydrogen bond length is 1.8 Å. –Each  + H points towards a lone pair on O. Hydrogen Bonding in H 2 O

33. Comparing Intermolecular Forces Dispersion forces are found in all substances. Their strength depends on molecular shapes and weights. Dipole-dipole forces add to the effect of dispersion forces. They are found only in polar substances. H-bonding is a special case of dipole-dipole interactions. Strongest of the intermolecular forces involving neutral species. Most important for hydride compounds (NH 3, H 2 O, etc.). Ion-dipole forces are interactions between ionic and polar molecules. Ion-dipole are stronger than H-bonds. Covalent bonds are stronger than any of these reactions.

34. Intermolecular Forces Chart

35. Viscosity Surface Tension Properties of Liquids

36. Viscosity is the resistance of a liquid to flow. A liquid flows by sliding molecules over each other. Viscosity depends on: The attractive forces between molecules: –Stronger the intermolecular forces, the higher the viscosity. The temperature: –Higher temperatures tend to decrease the viscosity. The tendency of molecules to become entangled: –tangled molecules increases the viscosity Viscosity in Liquids

37. Trends in the Viscosities of Hydrocarbons Notice that as molecular weight increases, viscosity increases.

38. Water tends to “bead up” on waxy surfaces. Beading is due to an imbalance of intermolecular forces at the surface of the liquid. The Inward force causes the molecules to pack closer at the surface (forming a skin). A measure of inward forces that must be overcome to expand the surface area of a liquid is termed surface tension. Surface Tension

39. Surface tension is the amount of energy needed to increase the surface area of a liquid by 1 unit area. Surface tension is affected by intermolecular forces. Strong intermolecular forces results in a higher surface tension. –Water has high surface tension (H-bonding). –Hg has even higher surface tension (metallic bonds). –Metallic bonds form when by metal to metal bonding. –Cohesive and adhesive forces are intermolecular forces. Capillary Action: When a narrow glass tube is placed in water, the meniscus pulls the water up the tube. Surface Tension in Liquids

40. Phase changes are changes of state (i.e., matter in one state is converted into another state) Sublimation: solid  gas. Vaporization: liquid  gas. Melting or fusion: solid  liquid. Deposition: gas  solid. Condensation: gas  liquid. Freezing: liquid  solid. Phase Changes

41. Energy Trend in Phase Changes

42. Sublimation:  H sub > 0 (endothermic). Vaporization:  H vap > 0 (endothermic). Heat of Vaporization Melting or Fusion:  H fus > 0 (endothermic). Heat of fusion Deposition:  H dep < 0 (exothermic). Condensation:  H con < 0 (exothermic). Freezing:  H fre < 0 (exothermic). Energy Changes Accompanying Phase Changes

43. Generally heat of fusion (enthalpy of fusion) is less than heat of vaporization: –it takes more energy to completely separate molecules, than partially separate them. All phase changes are possible under the right conditions Phase Changes

44. Heating Curves

45. Gases liquefied by increasing pressure at some temperature. Critical temperature: the minimum temperature for liquefaction of a gas using pressure. Critical pressure: pressure required for liquefaction. Critical Temperature and Pressure

46. Explaining Vapor Pressure on the Molecular Level Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid. These molecules move into the gas phase. As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid. After some time the pressure of the gas will be constant at the vapor pressure. This is called an equilibrium vapor pressure. Vapor Pressure

47. Explaining Vapor Pressure on the Molecular Level Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface. Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium. If equilibrium is never established then vapors continues to form. Eventually, the liquid evaporates to dryness. Volatile substances evaporate rapidly. The Higher the temperature, the faster the liquid evaporates. Vapor Pressure Cont.

48. Volatility, Vapor Pressure, and Temperature Vapor Pressure

49. Phase diagram: A plot of pressure vs. temperature showing all equilibria between phases. Given a temperature and pressure, phase diagrams tell us which phase will exist. Any temperature and pressure combination not on a curve represents a single phase. Phase Diagrams

50. Triple point: temperature and pressure at which all three phases are in equilibrium. Vapor-pressure curve: generally as pressure increases, temperature increases. Critical point: critical temperature and pressure for the gas. Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. Normal melting point: melting point at 1 atm. Features of a Phase Diagram

51. Phase Diagram Example

52. Phase Diagrams of H 2 O and CO 2 Notice that the line separating the liquid and gas phases end rather than continuing to infinite pressure and temperature. (?)

53. Water: The melting point curve slopes to the left because ice is less dense than water. Triple point occurs at 0.0098  C and 4.58 mmHg. Normal melting (freezing) point is 0  C. Normal boiling point is 100  C. Critical point is 374  C and 218 atm. Carbon Dioxide: Triple point occurs at -56.4  C and 5.11 atm. Normal sublimation point is -78.5  C. (At 1 atm CO 2 sublimes it does not melt.) Critical point occurs at 31.1  C and 73 atm. Reading a Phase Diagrams

54. Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions. Crystals have an ordered, repeated structure. The smallest repeating unit in a crystal is a unit cell. Unit cell is the smallest unit with all the symmetry of the entire crystal. Three-dimensional stacking of unit cells is the crystal lattice. Amorphous solid: no definite arrangement of molecules, atoms, or ions (i.e., lack well-defined structures or shapes). Amorphous solids vary in their melting points. Structures of Solids

55. Unit Cells Structures of Solids The smallest repeating unit that shows the symmetry of the pattern is called the unit cell.

56. Primitive Cubic Body-centered Cubic (BCC) Face-centered Cubic (BCC) Hexagonal Close Pack (HCP) Rhombohedral Cubic Diamond Types of Unit Cells

57. Three common types of unit cell. Primitive cubic, atoms at the corners of a simple cube –each atom shared by 8 unit cells; Body-centered cubic (bcc), atoms at the corners of a cube plus one in the center of the body of the cube, –corner atoms shared by 8 unit cells, center atom completely enclosed in one unit cell; Face-centered cubic (fcc), atoms at the corners of a cube plus one atom in the center of each face of the cube, –corner atoms shared by 8 unit cells, face atoms shared by 2 unit cells. Common Types of Unit Cells

58. Unit Cells

60. Table showing Atom Fractions in Unit Cells

61. Determine the number of NaCl ions in a unit cell? Problem Solving

62. The Crystal Structure of Sodium Chloride Two equivalent ways of defining unit cell: –Cl - (larger) ions at the corners of the cell, or –Na + (smaller) ions at the corners of the cell. The cation to anion ratio in a unit cell is the same for the crystal. In NaCl each unit cell contains same number of Na + and Cl - ions. Note the unit cell for CaCl 2 needs twice as many Cl - ions as Ca 2+ ions. Structures of Solids

63. Crystal Structure of NaCl

64. Closer View of a NaCl Unit Cell

65. Crystalline solids have structures that maximize intermolecular forces between particles. These particles can be represent by spheres Atoms and ions are spheres. Molecular crystals are formed by close packing of the molecules. We rationalize maximum intermolecular force in a crystal by the close packing of spheres. The spaces between spheres are called interstitial. A crystal is built up by placing close packed layers of spheres on top of each other. Structures of Crystalline Solids

66. A crystal is built up by placing close packed layers of spheres on top of each other. There are two choices for the third layer of spheres: Third layer eclipses the first (ABAB arrangement). This is called hexagonal close packing (hcp); Third layer is in a different position relative to the first (ABCABC arrangement). This is called cubic close packing (ccp). Packing of Spheres

67. There are four types of solid: Molecular (formed from molecules) - usually soft with low melting points and poor conductivity. Covalent network (formed from atoms) - very hard with very high melting points and poor conductivity. Ions (formed form ions) - hard, brittle, high melting points and poor conductivity. Metallic (formed from metal atoms) - soft or hard, high melting points, good conductivity, malleable and ductile. Bonding in Solids

68. Types of Crystalline Solids

70. Metallic solids usually have metal atoms in hcp, fcc or bcc arrangements. Coordination number for each atom is either 8 or 12. Problem: the bonding is too strong for London dispersion and there are not enough electrons for covalent bonds. Resolution: the bonding is due to valence electrons that are delocalized through the entire solid. Metals conduct because the electrons are delocalized and are mobile. Metallic Bonds

71. End of Chapter 11 Intermolecular Forces, Liquids and Solids