Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
| At the same temperature and pressure, equal volumes of all gases contain the same number of molecules. |
|
|
Term
|
Definition
| 22.4 L -- One mole of an ideal gas occupies 22.4 L at 0C and 1 atm of pressure. |
|
|
Term
|
Definition
|
|
Term
| Dalton's Law of Partial Pressures |
|
Definition
| The total pressure exerted by a mixture of ideal gases is the sum of the partial pressures of those gases. |
|
|
Term
|
Definition
| Pressure exerted by its gaseous molecules in equilibrium with the liquid. |
|
|
Term
|
Definition
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. |
|
|
Term
|
Definition
Urms=sq rt(3RT/Mm)
R=8.314 kg m2/s2 K mol
Molar mass, Mm, must be in KG/mol |
|
|
Term
|
Definition
| Gases exiting through the walls of a container with porous walls. |
|
|
Term
|
Definition
| The movement of a substance into a space or the mixing of one substance with another. |
|
|
Term
| Ratio of Rate of Effusion |
|
Definition
| R1/R2=sqrt(M2/M1) OR (D2/D1) |
|
|
Term
|
Definition
| When molecules become close together through increased T or P -- intermolecular forces become important. |
|
|
Term
|
Definition
| Diffuse into each other; they are soluble in each other and form homogeneous solutions. |
|
|
Term
|
Definition
| Forces between (among) individual particles (atoms, ions, molecules). |
|
|
Term
|
Definition
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) |
|
|
Term
| Intermolecular Interaction Strengths |
|
Definition
|
|
Term
|
Definition
| Occur between the + end of one polar molecule and the - end of the other. |
|
|
Term
|
Definition
| 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. |
|
|
Term
| Dispersion Forces (London Forces) |
|
Definition
| 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. |
|
|
Term
|
Definition
| Ion and Dipole. Weaker than Dipole-Dipole. |
|
|
Term
|
Definition
Nonpolar: 0-0.5
Polar: 0.5-2.0
Ionic: > 2.0 |
|
|
Term
|
Definition
| Resistance to flow. One measure of the forces of attraction within a liquid. Honey = high viscosity; gasoline = low viscosity. |
|
|
Term
|
Definition
| 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. |
|
|
Term
|
Definition
Capillary Rise: Adhesive Forces > Cohesive Forces
Capillary Fall: Cohesive Forces > Adhesive Forces |
|
|
Term
|
Definition
| 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. |
|
|
Term
|
Definition
| A system at equilibrium, or changing toward equilibrium, responds in the way that tends to relieve or "undo" any stress placed on it. |
|
|
Term
|
Definition
| 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. |
|
|
Term
|
Definition
| Temperature at which the vapor pressure of the liquid equals the applied (usually atmospheric) pressure. |
|
|
Term
|
Definition
| Process by which a mixture or solution is separated into its components on the basis of differences in BPs of the components. |
|
|
Term
|
Definition
Amount of heat required to raise the T of 1g of a substance 1C with no change in state.
Q=mC(change in T) |
|
|
Term
|
Definition
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 |
|
|
Term
|
Definition
| 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. |
|
|
Term
|
Definition
| 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. |
|
|
Term
| Molar Heat of Vaporization |
|
Definition
| 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. |
|
|
Term
| Molar Heat of Condensation |
|
Definition
| Reverse of molar heat of vaporization. |
|
|
Term
|
Definition
| Temperature at which liquid and solid coexist at equilibrium under a pressure of 1 atm. |
|
|
Term
|
Definition
| The amount of heat required to melt 1g of a solid at its melting point with no change in T. Endothermic. |
|
|
Term
|
Definition
| Amount of heat liberated by the crystallization of one gram of liquid at its freezing point. Exothermic. |
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
| 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. |
|
|
Term
|
Definition
|
|
Term
|
Definition
| Point at which three phases of a substance can coexist in equilibrium. |
|
|
Term
|
Definition
| Temperature above which a gas cannot be liquefied, i.e., the temperature above which the liquid and gas do not exist as distinct phases. |
|
|
Term
|
Definition
| Have no well-defined ordered structure. Waxes, asphalt, and glass. |
|
|
Term
|
Definition
| Pressure required to liquefy a gas at its critical temperature. |
|
|
Term
|
Definition
| The combination of critical temperature and critical pressure. |
|
|
Term
|
Definition
|
|
Term
|
Definition
| Consist of positive and negative ions arranged in a definite crystal structure. Attractions between oppositely charged ions are strong. |
|
|
Term
|
Definition
| Consist of discrete molecules that occupy positions in unit cells. Intermolecular forces are relatively weak. London Forces and D-D. |
|
|
Term
|
Definition
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 |
|
|
Term
| Two major factors affect dissolution of solutes |
|
Definition
|
|
Term
|
Definition
| Measure of attractive forces among particles of a solid. Increases as the charges on ions increase. The higher the CLE, the stronger the bond. |
|
|
Term
|
Definition
| Change of energy content of solution (ChangeHsoln). If change Hsoln is exothermic (<0), dissolution is favored. -changeH |
|
|
Term
|
Definition
Change in disorder (randomness( of the solution. ChangeSmixing.
NaCl --> Na+ + Cl- increasing entropy +changeS |
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
| Xa=(# mols of A)/(# mols of A + # mols of B) |
|
|
Term
|
Definition
|
|
