Gen. Chem. II Exam I review sheet (Ch. 10, 11, 13, 14) Ch. 10 Gases behave physically similarly. KMT (Kinetic Molecular Theory): particles in a gas:

are in constant rapid motion are tiny compared to the distances between them have little to no attraction to each other collisions are elastic Temperature is a measure of average KE of gas particles Properties of gases: V, T, n, P Independent variable / dependent variable 1 atm (typical P on earth) = 760 mmHg = 760 torr = 101.325 kPa Boyle’s Law: P1V1 Charles’s Law: V1/T1 Avogadro’s Law: V1/n1 Ideal gas Law: PV = nRT T in Kelvin. Gas constant, R, match units (unit analysis). Dalton’s Law of Partial Pressures: Ptotal χA = PA Ptotal = P1 + P2 + P3 + ….. Real gases: can stick together take up some volume themselves ____ Diffusion: distance gas1 = √MM2 distance gas2 √MM1 ____ Effusion: rate gas1 = √MM2 (Graham’s Law) ratee gas2 √MM1 Appendix B (p. 1111) has more data on the properties of water

Ch. 11 IMF, electrostatic forces between molecules (NOT chemical bonds). In order of decreasing strength: 1) Ion-ion: in ionic compounds. Strength depends on size of ion charges. 2) a) Hydrogen bonding (special case of dipole-dipole, see part b): H must be directly attached to: N, O, or F. Strength depends on: # H-bonds, electronegativity differences, total # of e–s b) Dipole-dipole: occurs in polar compounds. Strength depends on: total # e–s 3) London Dispersion forces (instantaneous dipole; a temporary uneven distribution of

e– density): occurs in all compounds (because all compounds have e–s!) but it is the only force in non-polar compounds. Strength depends on: total # e–s and surface area.

Properties of liquids: depend on intermolecular forces. Surface tension- skin or elastic film at surface Cohesive forces- between same molecules Adhesive forces- between two different molecules Viscosity- resistance to flow; shape- tangling ability surface area Volatility- tendency to go into g phase at a T below the bp. Solids: Crystalline- very regular array of atoms/molecules/ions Amorphous- very viscous liquids Classification of solids: 1) Atomic (London)- usually Noble Gases 2) Molecular (London, dipole-dipole, H-bonding) individual molecules. 3) Covalent/Network solid- a large molecule- extends infinitely in 3-D graphite diamond SiO2

4) Metallic- metals 5) ionic- ionic compounds Polymorphs- “many forms”- different connectivities of atoms. Unit cells- lattice points must be in same place throughout crystalline lattice.

3–D cubic unit cells (high symmetry, 90o angles) simple cubic (sc) body-centered cubic (bcc) face-centered cubic (fcc)

_ _ r = a/2 r = a √3 r = a √2 4 4 Used to find: atomic radius, density, empirical formula Just as the 2–D lattice points define a unit cell and “cut” the pattern, the atoms can be “cut” when defining unit cells.

(Completely in unit cell = 1; corner = 1/8; face = ½; edge = 1/4 of an atom: (not seen in class)) Heating curves (thermodynamics)

Phase diagrams (applying pressure will always cause the more compact form to be made. The melting point curve on water slopes to the left, meaning applying pressure to solid water forms the denser liquid water. Ice is less dense than water because hydrogen bonding leaves open gaps. Most phase diagrams slope to the right in this region; the solid form is usually the most dense).

Ch. 13 Solute, solvent Nonelectrolytes, weak electrolytes, strong electrolytes 3 things to make a solution: 1) Break solute IMF’s (pull solute apart), ∆H = + 2) Break solvent IMF’s (pull solvent apart), ∆H = + 3) Make solute-solvent interactions, ∆H = – (most likely to

form a solution, although entropy can make ∆H = 0 or small + also form a solution. Entropy drives mixing)

These are mixed systems (potentially mixed IMF’s). “Like dissolves like”, similar IMF’s will dissolve. Hydrophobic, hydrophilic (compounds aren’t sentient!) Henry’s Law, sgas = k Pgas Gases behave as if they were the only ones in the container.

Solubility- amount of solute needed to form a saturated solution in a given quantity of solvent. Unsaturated Saturated Supersaturated Most compounds are more soluble as the T increases. Gases are less soluble with increasing T. Concentration units: M, molarity = moles solute / L solution m, molality = moles solute / kg solvent (independent of volume)

Mass (wt.) % = (mass solute / total mass of solution) x 100 pph ppt x 103 ppm x 106

ppb x 109 mole fraction, χ = moles A / total moles Colligative properties- depends on # of particles, not structure. 1) Vapor P lowering Raoult’s Law: Ptotal χA = PA 2) bp elevation ∆T = m i kb 3) fp depression ∆T = m i kf 4) osmotic P ΠV = nRT Ch. 14 Kinetics (RXN rate) [ ] = concentration of (in M) Rate = k[reactant]x = ∆[ ] / ∆time, units = M/s x = order of reaction, found experimentally; overall order = summation of the individual orders Instantaneous rate at t=0 is the initial rate of the reaction. Rates change as RXNs proceed due to less reactant, side RXNs, …. Orders of RXN 0 order; change [ ], rate is not affected. 1st order; change [ ], rate changes by same amount. 2nd order; double [ ] , rate 4x.

Integrated form of the rate law:

Rate constant, k, found via experiment. Units depend on order of the RXN. (Rate is always M/s) ½ life: 1st order, [ ] independent, ln 0.693 / k = t1/2

RXN profiles: (exothermic shown here) Elementary RXNs: unimolecular, bimolecular, termolecular Mechanisms: can use stoichiometry as exponent because each step is an elementary RXN rds- sets rate forward rate = reverse rate for steady state step Catalysts lower Ea, they don’t change the RXN composition, just change the rate. Enzymes are specialized biological catalysts. Arrhenius equation, slope = –Ea/R