Intermolecular Forces

1. Intermolecular Forces Forces Which Determine States and Changes in State

2. IMF vs Intramolecular Intramolecular forces are those which hold a substance together�Ionic, covalent, and metallic bonds Intermolecular forces are those which hold molecules close to one another. They are primarily the reason that substances can be liquid or solid

3. Intermolecular Forces Forces between (rather than within) molecules. dipole-dipole attraction: molecules with dipoles orient themselves so that �+� and �?� ends of the dipoles are close to each other. hydrogen bonds: dipole-dipole attraction in which hydrogen is bound to a highly electronegative atom. (F, O, N)

5. London Dispersion Forces relatively weak forces that exist among noble gas atoms and both nonpolar and polar molecules. (Ar, C8H18, C6H12O6) caused by instantaneous dipole, in which electron distribution becomes asymmetrical. the ease with which electron �cloud� of an atom can be distorted is called polarizability.

7. Some Properties of a Liquid Surface Tension: The resistance to an increase in its surface area. Capillary Action: Spontaneous rising of a liquid in a narrow tube. Cohesion and Adhesion: Forces within a liquid(cohesion) and toward other substances (adhesion) Viscosity: Resistance to flow (molecules with large intermolecular forces).

8. Forces vs. Properties Intermolecular forces cause differences IMF increases surface tension, viscosity, capillary action, melting/boiling points IMF decreases volatility (ability of a liquid to evaporate easily), and thusly also vapor pressure is decreased.

9. Types of Solids Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)]. Amorphous solids: considerable disorder in their structures (glass). Glass is considered a supercooled (or frozen) liquid.

10. Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance. Unit Cell: The smallest repeating unit of the lattice. simple cubic body-centered cubic face-centered cubic

12. Bragg Equation Used for analysis of crystal structures. n? = 2d sin ? d = distance between atoms n = an integer ? = wavelength of the x-rays

14. Types of Crystalline Solids Ionic Solid: contains ions at the points of the lattice that describe the structure of the solid (NaCl). Held together by intramolecular ionic bonds. Molecular Solid: discrete covalently bonded molecules at each of its lattice points (sucrose, ice). Held together by intermolecular forces.

16. Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors. hexagonal closest packed (�aba�) cubic closest packed (�abc�)

18. Bonding Models for Metals Electron Sea Model: A regular array of metals in a �sea� of electrons. Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.

21. Homework 1 p. 489ff 1,2,5,8,13,19,21,26,28,29,33,34 Turn in ONE assignment per lab group--divide your LABOR!!

22. Metal Alloys 1. Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn 2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon 3. Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (carbon) alloys.

25. Semiconductors Conductivity is enhanced by doping with group 3a or group 5a elements.

27. Network Solids Composed of strong directional covalent bonds that are best viewed as a �giant molecule�. Held together by intramolecular covalent bonds. brittle do not conduct heat or electricity well carbon, silicon-based usually graphite, diamond, ceramics, glass

30. Vapor Pressure . . . is the pressure of the vapor present in a container of liquid at equilibrium. . . . is determined principally by the size of the intermolecular forces in the liquid. . . . increases significantly with temperature. Volatile liquids have high vapor pressures.

32. Melting Point Molecules break loose from lattice points and solid changes to liquid. (Temperature is constant as melting occurs.) vapor pressure of solid = vapor pressure of liquid

33. Boiling Point Constant temperature when added energy is used to vaporize the liquid. vapor pressure of liquid = pressure of surrounding atmosphere

35. Energy Change Within Phases As energy is added to a solid substance, its molecules begin vibrating more and more. This additional motion is the increase in temperature of the solid. The heat energy can be calculated as q= m�c �DT where m is the mass in grams, c is the specific heat capacity of the substance in a specific phase, and DT is the change in temperature

36. Energy in Phase Changes When a solid has absorbed enough energy, its particles overcome the forces holding them together and become a liquid. This happens also where a liquid becomes a gas (vapor), and the equation to find the energy necessary is q= m L where L is the heat of fusion(melting) or vaporization.

37. Energy in Heating How much energy to change 50 g of ice at -10oC to steam at 110oC?

38. Phase Diagram Represents phases as a function of temperature and pressure. critical point: critical temperature and pressure (for water, Tc = 374�C and 218 atm). critical temperature: temperature above which the vapor can not be liquefied. critical pressure: pressure required to liquefy AT the critical temperature.

43. Homework 2 p 492ff 39, 40, 61, 62, 68, 69, 71, 83, 88