CHEM 112 Final Review: Equilibrium, Acids and Bases, Electrochemistry and Kinetics

Terms in this set (31)

1. Write the equilibrium constant expression for aA + bB <=> cC + dDK = [C]^c[D]^d / ([A]^a[B]^b). Pure solids and pure liquids are left out of the expression.
2. Reaction quotient Q vs K: which way does the reaction shift?If Q < K, the reaction proceeds forward toward products. If Q > K, it shifts in reverse toward reactants. If Q = K, the system is at equilibrium.
3. Le Chatelier's principle: what happens if a reactant is added at equilibrium?The system shifts toward products to consume some of the added reactant. K itself does not change.
4. Le Chatelier: effect of decreasing the volume of a gas-phase equilibriumPressure increases, so the equilibrium shifts toward the side with fewer moles of gas. If both sides have equal moles of gas, there is no shift.
5. What is the only change that alters the value of K?Temperature. For an exothermic reaction, raising the temperature decreases K; for an endothermic reaction, raising the temperature increases K.
6. How does a catalyst affect an equilibrium?It speeds up the forward and reverse reactions equally, so equilibrium is reached faster, but the position of equilibrium and K are unchanged.
7. Relationship between Kp and KcKp = Kc(RT)^(delta n), where delta n = moles of gaseous products minus moles of gaseous reactants, R = 0.08206 L·atm/(mol·K) and T is in kelvin.
8. Bronsted-Lowry definitions of acid and baseAn acid is a proton (H+) donor and a base is a proton acceptor. After the transfer, the acid becomes its conjugate base and the base becomes its conjugate acid.
9. Identify the conjugate pairs in NH3 + H2O <=> NH4+ + OH-NH3 (base) and NH4+ (its conjugate acid); H2O (acid) and OH- (its conjugate base).
10. pH, pOH and Kw at 25 degrees CpH = -log[H3O+], pOH = -log[OH-], Kw = [H3O+][OH-] = 1.0 x 10^-14, so pH + pOH = 14.00.
11. Mini-problem: pH of 0.010 M HClHCl is a strong acid and ionizes completely, so [H3O+] = 0.010 M and pH = -log(0.010) = 2.00.
12. Weak acid pH shortcut and when it is valid[H3O+] is approximately sqrt(Ka x C). It is valid when the percent ionization is under about 5 percent (x much smaller than C).
13. Mini-problem: pH of 0.10 M weak acid with Ka = 1.0 x 10^-5[H3O+] = sqrt(1.0 x 10^-5 x 0.10) = 1.0 x 10^-3 M, so pH = 3.00. Ionization is 1 percent, so the approximation holds.
14. Relationship between Ka and Kb for a conjugate acid-base pairKa x Kb = Kw (1.0 x 10^-14 at 25 degrees C), so pKa + pKb = 14. The stronger the acid, the weaker its conjugate base.
15. Henderson-Hasselbalch equationpH = pKa + log([A-]/[HA]), where [A-] is the conjugate base concentration and [HA] is the weak acid concentration. Used for buffer calculations.
16. What is a buffer and when is it most effective?A solution of a weak acid and its conjugate base (or weak base and conjugate acid) that resists pH change when small amounts of acid or base are added. It works best when pH is within about 1 unit of the pKa.
17. Weak acid titrated with strong base: pH at half-equivalence and equivalenceAt half-equivalence, [HA] = [A-] so pH = pKa. At the equivalence point the pH is above 7 because the conjugate base A- reacts with water to form OH-.
18. Mini-problem: molar solubility of AgCl if Ksp = 1.8 x 10^-10AgCl(s) <=> Ag+ + Cl-, so Ksp = s^2 and s = sqrt(1.8 x 10^-10), about 1.3 x 10^-5 M.
19. Common ion effectAdding an ion already present in an equilibrium shifts it away from that ion. For example, AgCl is less soluble in NaCl solution than in pure water.
20. Oxidation vs reduction, and where each occurs in a galvanic cellOxidation is loss of electrons and occurs at the anode. Reduction is gain of electrons and occurs at the cathode. Electrons flow through the wire from anode to cathode.
21. How do you calculate the standard cell potential?E cell = E(cathode) - E(anode), using standard reduction potentials for both. A positive E cell means the reaction is spontaneous as written.
22. Link between delta G, E cell and Kdelta G = -nFE cell and delta G standard = -RT ln K, where n is moles of electrons transferred and F = 96485 C/mol. Positive E means negative delta G and K greater than 1.
23. Nernst equationE = E standard - (RT/nF) ln Q. At 25 degrees C this becomes E = E standard - (0.0592 V/n) log Q.
24. Galvanic vs electrolytic cellA galvanic (voltaic) cell uses a spontaneous redox reaction to produce electricity. An electrolytic cell uses an external power source to drive a nonspontaneous reaction. Oxidation is at the anode in both.
25. General form of a rate lawRate = k[A]^m[B]^n, where k is the rate constant and m and n are reaction orders. Orders must be determined experimentally, not from the balanced equation's coefficients.
26. Mini-problem: doubling [A] quadruples the rate. What is the order in A?Second order, since 2^m = 4 gives m = 2.
27. First-order reactions: integrated rate law and half-lifeln[A] = ln[A]0 - kt, and t(1/2) = 0.693/k. The half-life does not depend on starting concentration.
28. Arrhenius equationk = A e^(-Ea/RT), where A is the frequency factor, Ea the activation energy, R = 8.314 J/(mol·K) and T the kelvin temperature. Higher T or lower Ea gives a larger k.
29. What is the rate-determining step?The slowest elementary step in a reaction mechanism. It limits the overall rate, so the rate law follows from that step.
30. Gibbs free energy and spontaneitydelta G = delta H - T delta S. If delta H is negative and delta S is positive, the process is spontaneous at all temperatures; if delta H is positive and delta S negative, it is never spontaneous.
31. Second law of thermodynamicsFor any spontaneous process, the total entropy of the universe (system plus surroundings) increases.