1 / 110

Huntington Beach High School AP CHEMISTRY EXAM REVIEW WINTER REVIEW

Huntington Beach High School AP CHEMISTRY EXAM REVIEW WINTER REVIEW. Intermolecular Forces Forces between molecules, between ions, or between molecules and ions. Table 13.1: summary of forces and their relative strengths. Intermolecular Forces Ion-Ion Forces. Na + — Cl - in salt.

trella
Télécharger la présentation

Huntington Beach High School AP CHEMISTRY EXAM REVIEW WINTER REVIEW

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Huntington Beach High SchoolAP CHEMISTRYEXAM REVIEWWINTER REVIEW

  2. Intermolecular Forces Forces between molecules, between ions, or between molecules and ions. Table 13.1:summary of forces and their relative strengths.

  3. Intermolecular ForcesIon-Ion Forces Na+ — Cl- in salt. These are the strongest forces. Solids have high melting temperatures. NaCl, mp = 800 oC MgO, mp = 2800 oC

  4. Attraction Between Ions and Permanent Dipoles Water is highly polar, interacts with positive ions, gives hydrated ions in water.

  5. Attraction Between Ions and Permanent Dipoles Many metal ions are hydrated. It is the reason metal salts dissolve in water. Co(H2O)62+

  6. Attraction Between Ions and Permanent Dipoles Attraction between ions and dipole depends on ion charge and ion-dipole distance. -1922 kJ/mol -405 kJ/mol -263 kJ/mol See Example 13.1, page 588.

  7. Dipole-Dipole Forces Such forces bind molecules having permanent dipoles to one another.

  8. Dipole-Dipole Forces Influence of dipole-dipole forces is seen in the boiling points of simple molecules. Compd Mol. Wt. Boil Point N2 28 -196 oC CO 28 -192 oC Br2 160 59 oC ICl 162 97 oC

  9. Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. Hydrogen bonding in HF

  10. Hydrogen Bonding A special form of dipole-dipole attraction, which enhances dipole-dipole attractions. H-bonding is strongest when H is joined to N, O, or F

  11. H-Bond Between Ethanol and Water - H-bond + -

  12. Hydrogen Bonding in H2O H-bonding is very strong in water: • the O—H bond is very polar • there are 2 lone pairs on the O atom • Accounts for many of water’s unique properties

  13. FORCES WITH INDUCED DIPOLES • Non-polar molecules condense to (l) and (s) by “induced dipoles.” • Example: I2 dissolving in alcohol. • Alcohol INDUCES a temporary dipole in I2 .

  14. FORCES INVOLVING INDUCED DIPOLES Dipoles can form in nonpolar molecules. Electrons in one molecule REPEL the electrons in another and create TEMPORARY DIPOLES (van der Waals or London Forces)

  15. FORCES INVOLVING INDUCED DIPOLES The size of the dipole depends on the tendency to be distorted. Higher molec. weight ---> larger dipoles. Molecule Boiling Point (oC) CH4 (methane) - 161.5 C2H6 (ethane) - 88.6 C3H8 (propane) - 42.1 C4H10 (butane) - 0.5

  16. Boiling Points of Hydrocarbons C4H10 Linear relation between bp & molar mass C3H8 C2H6 CH4

  17. LiquidsSection 13.3 • molecules are in constant motion • almost incompressible • do not fill the container • appreciable intermolecular forces • molecules close together

  18. LiquidsSection 13.3 evaporation---> <---condensation

  19. LiquidsSection 13.3 To evaporate, molecules must have sufficient energy to break IM forces. Breaking IM forces requires energy. The process of evaporation is endothermic.

  20. LiquidsSection 13.3 When molecules of liquid evaporate, they exert VAPOR PRESSURE EQUILIBRIUM VAPOR PRESSURE = pressure of the vapor over a liquidin a closed container when rate of evaporation = rate of condensation. See Fig. 13.18

  21. Boiling Liquids Liquid boils when its vapor pressure equals atmospheric pressure.

  22. Boiling Point at Lower Pressure When pressure is lowered, the vapor pressure can equal the external pressure at a lower temperature.

  23. Liquids T of boiling when P = 1 atm is the NORMAL BOILING POINT VP of a molecule depends on IM forces. Note IM forces in molecules below.

  24. Liquids HEAT OF VAPORIZATION = heat needed (at constant P) to vaporize the liquid. LIQ + heat ---> VAP Compd. DHvap (kJ/mol) IM Force H2O 40.7 (100 oC) H-bonds SO2 26.8 (-47 oC) dipole Xe 12.6 (-107 oC) induced dipole

  25. Types of SolidsTable 13.6 TYPE EXAMPLE FORCE Ionic NaCl, CaF2, ZnS Ion-ion Metallic Na, Fe Metallic Molecular Ice, I2 Dipole Ind. dipole Network Diamond Extended Graphite covalent

  26. Network Solids Diamond Graphite

  27. Properties of Solids 1. Molecules, atoms or ions locked in a CRYSTAL LATTICE 2. Particles are CLOSE together 3. STRONG IM forces 4. Highly ordered, rigid, incompressible ZnS, zinc sulfide

  28. Crystal Lattices Regular 3-D arrangements of equivalent LATTICE POINTS. The lattice points define UNIT CELLS, the smallest repeating internal unit that has the symmetry characteristic of the solid. 7 basic crystal systems -- we are only concerned with CUBIC.

  29. Cubic Unit Cells

  30. Cubic Unit CellsFigure 13.28 • simple cubic (SC) • body centered cubic (BCC) • face centered cubic (FCC)

  31. Simple Ionic Compounds Lattices of many simple ionic solids are built by taking a SC or FCC lattice of ions of one type and placing ions of opposite charge in the holes in the lattice. EXAMPLE: CsCl has a SC lattice of Cs+ ions with Cl- in the center.

  32. Simple Ionic Compounds 1 unit cell has 1 Cl- ion plus (8 corners)(1/8 Cs+ per corner) = 1 net Cs+ ion.

  33. TRANSITIONS BETWEEN PHASESSection 13.7, Figure 13.37

  34. TRANSITIONS BETWEEN PHASESSection 13.7 Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line. At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example. INTERPRET PHASE DIAGRAMS.

  35. TRANSITIONS BETWEEN PHASES As P and T increase, you reach CRITICAL T and P Above critical T no liquid exists no matter how high the pressure.

  36. Critical T and P COMPD Tc(oC) Pc(atm) H2O 374 218 CO2 31 73 CH4 -82 46 Freon-12 112 41(CCl2F2) Tc and Pc depend on IM forces.

  37. Solid-Vapor Equilibrium At P < 4.58 mmHg and T < 0.0098 ° C solid H2O can go directly to vapor. This process is called SUBLIMATION This is how a frost-free refrigerator works.

  38. Some Definitions A solution is a HOMOGENEOUS mixture of 2 or more substances in one phase. One substance is the SOLVENT and the others are the SOLUTES.

  39. Definitions Solutions can be unsaturated or saturated. A saturated solution contains the maximum quantity of solute that dissolves at that temperature. Unsaturated solutions have less than the maximum solute.

  40. Definitions SUPERSATURATED SOLUTIONS contain more than is possible (at the ambient temp-erature) and are unstable. Prepare them at higher temperature and allow to cool.

  41. Colligative Properties The properties of the solvent in a solution are changed. • Vapor pressure decreases • Melting point decreases • Boiling point increases • Osmosis is possible (osmotic pressure)

  42. Colligative Properties These changes are called COLLIGATIVE PROPERTIES. They depend only on the NUMBER of solute particles relative to solvent particles, not on the KIND of solute particles. Applies to IDEAL SOLUTIONS. Properties depend only on the concentration of solute.

  43. Concentration Units Need concentration units to tell us the number of solute particles per solvent particle. The unit “molarity” does not do this!

  44. Concentration Units MOLE FRACTION, X For a mixture of A, B, and C

  45. Concentration Units MOLALITY, m

  46. Calculating Concentrations Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H2O. Calculate X, and m of glycol. 250. g H2O = 13.9 mol X glycol = 0.0672

  47. Calculating Concentrations Dissolve 62.1 g (1.00 mol) of ethylene glycol in 250. g of H2O. Calculate X, m, and % of glycol. Calculate molality

  48. Dissolving Gases & Henry’s Law Gas solubility (M) = kH • Pgas kH for O2 = 1.66 x 10-6 M/mmHg When Pgas drops, solubility drops. Note M is used for gases since each gas molecule is independent of others.

  49. Understanding Colligative Properties Colligative properties depend on the LIQUID-VAPOR EQUILIBRIUM of a solution.

  50. Understanding Colligative Properties To understand colligative properties, study the LIQUID-VAPOR EQUILIBRIUM for a solution.

More Related