Chem 1212: Chapters 12-14

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Boyle’s Law
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P1V1=P2V2
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Charles’ Law
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V1/T1=V2/T2
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Combined Gas Law
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P1V1/T1=P2V2/T2
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Avogadro’s Law
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At the same temperature and pressure, equal volumes of all gases contain the same number of molecules.
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Standard Molar Volume
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22.4 L — One mole of an ideal gas occupies 22.4 L at 0C and 1 atm of pressure.
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Ideal Gas Equation
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PV=nRT
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Dalton’s Law of Partial Pressures
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The total pressure exerted by a mixture of ideal gases is the sum of the partial pressures of those gases.
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Vapor Pressure
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Pressure exerted by its gaseous molecules in equilibrium with the liquid.
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Kinetic-Molecular Theory
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1. Gases consist of discrete molecules. The individual molecules are very small and far apart relative to their sizes.

2. Gaseous molecules are in continuous, rapid random, straight-line motion with varying velocities.

3. The collisions between gas molecules and with the walls of the container are elastic; the total energy is conserved during a collision.

4. Between collisions, the molecules exert no attractive or repulsive forces on one another; each molecule travels in a straight line with a constant velocity.

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Root Mean Square
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Urms=sq rt(3RT/Mm)

R=8.314 kg m2/s2 K mol

Molar mass, Mm, must be in KG/mol

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Effusion
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Gases exiting through the walls of a container with porous walls.
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Diffusion
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The movement of a substance into a space or the mixing of one substance with another.
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Ratio of Rate of Effusion
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R1/R2=sqrt(M2/M1) OR (D2/D1)
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Deviations from Ideality
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When molecules become close together through increased T or P — intermolecular forces become important.
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Miscible Liquids
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Diffuse into each other; they are soluble in each other and form homogeneous solutions.
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Intermolecular Forces
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Forces between (among) individual particles (atoms, ions, molecules).
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Ion-Ion Interactions
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Force of attraction between two oppositely charged ions is governed by Coulomb’s Law:

F=abs(q1q2)/r2

F=magnitude of force
r=distance in m
q1,q2=charges on bodies (coulombs)

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Intermolecular Interaction Strengths
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I-I>HB>D-D>>LF
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Dipole-Dipole
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Occur between the + end of one polar molecule and the – end of the other.
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Hydrogen Bonding
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Especially strong D-D interaction. Occurs between an H atom in one molecule (attached to an F, O, or N), and an F, O, or N atom in another molecule.
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Dispersion Forces (London Forces)
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Present in all substances. ONLY kind of interactions for nonpolar substances. Result from the attraction of the nucleus of one atom for the electron cloud of another atom (in a different molecule). They increase with increasing polarizability, the ability of an electron cloud to be distorted. Polarizability increases with increasing numbers of electrons and therefore with increasing sizes of molecules.
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Ion-Dipole Interactions
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Ion and Dipole. Weaker than Dipole-Dipole.
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Change in EN
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Nonpolar: 0-0.5
Polar: 0.5-2.0
Ionic: > 2.0
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Viscosity
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Resistance to flow. One measure of the forces of attraction within a liquid. Honey = high viscosity; gasoline = low viscosity.
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Surface Tension
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Measure of the inward forces that must be overcome to expand the surface of a liquid. Molecules on the surface are attracted only toward the interior, while those on the interior are attracted equally in all directions.
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Capillary Action
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Capillary Rise: Adhesive Forces > Cohesive Forces
Capillary Fall: Cohesive Forces > Adhesive Forces
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Evaporation
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Process by which molecules escape from the surface of a liquid (vaporize). Occurs more rapidly as temperature increases because a greater fraction of molecules possess the necessary escape velocity and kinetic energy.
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LeChatelier’s Principle
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A system at equilibrium, or changing toward equilibrium, responds in the way that tends to relieve or “undo” any stress placed on it.
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Vapor Pressure
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Pressure exerted by the vapor of the liquid on its surface at equilibrium in a closed container. Because vapor pressures increase as temperature increases, evaporation occurs more rapidly as T increases.
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Boiling Point
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Temperature at which the vapor pressure of the liquid equals the applied (usually atmospheric) pressure.
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Distillation
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Process by which a mixture or solution is separated into its components on the basis of differences in BPs of the components.
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Specific Heat
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Amount of heat required to raise the T of 1g of a substance 1C with no change in state.

Q=mC(change in T)

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Molar Heat Capacity
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Amount of heat required to raise the temperature of 1 mol of a substance 1C with no change in state.

Q=mol x molar heat capacity x change in T

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Heat of Vaporization
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Amount of heat that must be absorbed to convert 1g of a liquid at its BP to a gas with no change in temperature. Usually J/g.
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Heat of Condensation
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Reverse of Heat of Vaporization. Amount of heat that must be released to liquefy 1g of a gas at its condensation (BP) with no change in T.
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Molar Heat of Vaporization
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The amount of heat that must be absorbed to convert one mole of a liquid at its BP to a gas with no change in T.
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Molar Heat of Condensation
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Reverse of molar heat of vaporization.
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Melting Point
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Temperature at which liquid and solid coexist at equilibrium under a pressure of 1 atm.
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Heat of Fusion
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The amount of heat required to melt 1g of a solid at its melting point with no change in T. Endothermic.
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Heat of Crystallization
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Amount of heat liberated by the crystallization of one gram of liquid at its freezing point. Exothermic.
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Specific Heat of Water
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4.18 J/gC
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Density of Water
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1.00 g/mL
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Specific Heat of Ice
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2.09 J/gC
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Sublimation
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Conversion of a solid directly into vapor. Dry Ice is solid CO2. The white vapors are due to water condensing in the very cold gaseous CO2 near the solid.
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Deposition
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Vapor directly to solid.
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Triple Point
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Point at which three phases of a substance can coexist in equilibrium.
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Critical Temperature
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Temperature above which a gas cannot be liquefied, i.e., the temperature above which the liquid and gas do not exist as distinct phases.
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Amorphous Solids
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Have no well-defined ordered structure. Waxes, asphalt, and glass.
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Critical Pressure
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Pressure required to liquefy a gas at its critical temperature.
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Critical Point
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The combination of critical temperature and critical pressure.
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Metallic Solids
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Metals.
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Ionic Solids
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Consist of positive and negative ions arranged in a definite crystal structure. Attractions between oppositely charged ions are strong.
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Molecular Solids
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Consist of discrete molecules that occupy positions in unit cells. Intermolecular forces are relatively weak. London Forces and D-D.
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Covalent Solids
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Individual atoms are covalently bonded to several other atoms; thus they are very hard with very high melting points. Diamond, graphite, SO2 (sand), SiC.

Molecular less than Metallic less than Ionic less than Covalent

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Two major factors affect dissolution of solutes
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Enthalpy and Entropy
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Crystal Lattice Energy
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Measure of attractive forces among particles of a solid. Increases as the charges on ions increase. The higher the CLE, the stronger the bond.
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Enthalpy
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Change of energy content of solution (ChangeHsoln). If change Hsoln is exothermic (<0), dissolution is favored. -changeH
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Entropy
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Change in disorder (randomness( of the solution. ChangeSmixing.

NaCl –> Na+ + Cl- increasing entropy +changeS

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Molality
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mol solute/kg solvent
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Molarity
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mol solute/L solution
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Mole Fraction
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Xa=(# mols of A)/(# mols of A + # mols of B)
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Percent by Mass
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g solute/100 g solution

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