SI Final Exam Review
Supplemental Instruction
Iowa State University / Leader: / Ryan Gale
Course: / Chem 178
Instructor: / Burnett/Raker
Date: / 5/5/13

1.)  Which of the following interactions does not appear in the formation of solutions?

  1. Solute-Solute
  2. Hydrogen bonds
  3. Solvent-solute
  4. Solvent-solvent

2.)  A solution contains 14.6 g methanol in 184 g water. Calculate the molality of methanol.

  1. 2.48 molal
  2. 2.48 molar
  3. 0.700 molal
  4. 0.00248 molal

3.)  How many grams of ethylene glycol (C2H6O2) must be added to 1.00 kg of water to produce a solution that freezes at -5.00°C?

  1. 23.1 g
  2. 167 g
  3. 48.4 g
  4. -23.1 g

4.)  Which of the following are termed colligative properties?

  1. Freezing Point Depression
  2. Boiling Point Elevation
  3. Melting Point Elevation
  4. Osmotic Pressure

5.)  8A + 4B + C ↔ 5D + 2E

Calculate the equilibrium constant.

  1. K = A8B4CD5E2
  2. K = A8+B4+CD5+E2
  3. K = D5E2A8B4C
  4. K = 5D+2E8A+4B+C

6.)  For N2(g) + 3H2(g) ↔ 2 NH3(g), Kp = 4.34*10-3 at 300°C. What is the value of Kp for the reverse reaction?

  1. Kp doesn’t change for reverse processes
  1. 230
  1. 5.31 * 104
  1. 8.68 * 10-3

7.)  A sample of NOBr decomposes according to the equation:

2 NOBr(g) ↔ 2 NO(g) + Br2(g)

An equilibrium mixture in a 5.00 L vessel at 100 C contains 3.22 g NOBr, 3.08 g NO, and 4.19 g Br2. Calculate Kc

  1. Kc = 0.0184
  1. Kc = 0.0648
  1. Kc = 14.9
  1. Kc = 0.0165

8.)  How do the following changes affect an equilibrium reaction of the following:

2 CH3CH3(g) + 5 O2(g) ↔ 6 H2O(l) + 2 CO2(g)

Increasing [O2], decreasing volume of container, increase temperature

  1. Shift right, shift right, shift left
  2. Shift right, shift right, shift right
  3. Shift left, shift right, shift right
  4. Shift right, shift left, shift right

9.)  At the start of a certain reaction, only reactants are present, products have not been formed yet. What is the value of Q at this time?

  1. Q > 0
  2. Q = 0
  3. Q < 0
  4. Q = K

10.)  Calculate the concentration of OH-(aq) in a solution in which [H+] = 100*[OH-]

  1. 0.010 M
  2. 10-8 M
  3. 10-6 M
  4. 0.001 M

11.)  A 25.0 mL sample of H3PO3 solution titrated with 0.102 M NaOH requires 23.3 mL of NaOH to neutralize both acidic protons. What is the molarity of the solution?

  1. 0.0317 M
  2. 0.095 M
  3. 0.000095 M
  4. 0.000020 M

12.)  Calculate the concentration of C6H5COONa that must be present in a 0.20 M solution of benzoic acid to produce a pH of 4.00. Ka = 6.3 * 10-5

  1. [CH3COONa] = 0.00010 M
  2. [CH3COONa] = 0.13 M
  3. [CH3COONa] = 7.7 M
  4. [CH3COONa] = 0.36 M

13.)  Calculate the pH at the equivalence point when 40.0 mL of 0.100 M NH3 is titrated with 0.100 M HCl. Kb = 1.8*10-5 for NH3

  1. 7.00
  2. 9.50
  3. 5.28
  4. 2.18

14.)  It is found that 1.1*10-2 g of SrF2 dissolves per 100 mL of aqueous solution at 25°C. Calculate Ksp for this solution.

  1. 4.43*10-9
  2. 2.7*10-9
  3. 6.65*10-10
  4. 1.0*10-7

15.)  The decomposition of a certain insecticide in water at 12°C follows first-order kinetics with k = 1.45 yr-1. How long does it take for the concentration to reach ¼ of its initial concentration?

  1. 0.856 yrs
  2. 0.568 yrs
  3. 0.956 yrs
  4. 0.582 yrs

16.)  Consider the gas-phase reaction between nitric acid and bromine at 273 °C:

2 NO(g) + Br2(g) → 2 NOBr(g)

The following data for the initial rate of appearance of NOBr were obtained. Calculate the rate constant.

  1. 1200 M-1s-1
  2. 12000 M-2s-2
  3. 6000 M-1s-2
  4. Cannot calculate with the given data

17.)  Is a spontaneous process reversible?

  1. Yes, always
  2. No, only if ΔG < 0
  3. No, always
  4. Yes, if ΔSuniv < 0

18.)  Calculate ΔS° for the following reactions:

Be(OH)2(s) → BeO(s) + H2O(g)

S°(J/K): 50.21 13.77 188.83

  1. 253 J/K
  2. 152 J/K
  3. -225 J/K
  4. -253 J/K

19.)  Predict the sign of ΔH°, ΔS°, and the magnitude of K for the following reaction:

2 Mg(s) + O2(g) ↔ 2 MgO(s)

  1. Negative, negative, and large
  2. Positive, negative, and small
  3. Positive, negative, and large
  4. Negative, positive, and small

20.)  Consider the reaction

PbCO3(s) ↔ PbO(s) + CO2(s)

ΔH°(kJ) -699.1 -217.3 -393.5

ΔS°(J/K) 131.0 68.70 213.6

Calculate the equilibrium pressure of CO2 in the system at 400 °C

  1. 2.41kJ
  2. 0.24kJ
  3. 2.41*10-4 kJ
  4. 24.1 kJ

21.)  For the redox reaction:

Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO4-(aq)

Write the oxidation half-reaction for the reaction above.

  1. 2 H2O(l) + 2 Mn2+(aq) → 4 MnO4-(aq) + 8 H+(aq)
  1. 2 H2O(l) + 2 Mn2+(aq) → 4 MnO4-(aq) + 8 H+(aq) + 2 e-
  1. 2 H2O(l) + 2 Mn2+(aq) → 4 MnO4-(aq) + 8 H+(aq) + 5e-
  1. 4 H2O(l) + Mn2+(aq) → MnO4-(aq) + 8 H+(aq) + 5e-

22.)  A voltaic cell utilizes the following reaction and operates at 298 K:

3 Ce4+(aq) + Cr(s) → 3 Ce3+(aq) + Cr3+(aq)

What is the emf of this cell under standard conditions?

  1. 0.87 V
  2. -2.35 V
  3. -0.87 V
  4. 2.35 V

23.)  Elemental calcium is produced by the electrolysis of molten CaCl2. What mass of calcium can be produced by this process if 7.5*103 A is flowing for 48 hour period?

  1. 168 g
  2. 2.69*105 g
  3. 4.78*10-3 g
  4. 269 g

24.)  What particle is released during the decay of sodium-24 into magnesium-24.

  1. Alpha particle
  2. Beta particle
  3. Gamma particle
  4. Neutron

25.)  Cobalt-60 which undergoes beta decay, has a half-life of 5.26 yrs. How many beta particles are emitted in 600 s by a 3.75 mg sample of 60Co?

Formulas:

Tf = kfm kf = 1.86 °C/m for H2O(l) pH = -log[H+] Ka*Kb = Kw

For first order processes, half-life, t1/2 = 0.693k ln[A] = -kt + ln[A0] Rate = kN

Δ(G,H,S,)° = G,H,S°products- G,H,S°reactants

ΔG = ΔG° + RTlnQ ΔG° = ΔH°-TΔS° E°cell = E°red(cathode) – E°red(anode)