Chemical Kinetics Problems and Questions (Chapter 18 in the textbook)

  1. Gasoline is highly flammable, but it doesn’t burst into flame just because there is oxygen in the air. Explain what must occur for the gasoline molecules to react with oxygen molecules.
  2. Define collision theory.
  3. What is an activated complex and at what point during a reaction does it exist?
  4. Draw a potential energy diagram for an endothermic reaction including activation energy (Ea) and the change in enthalpy (H).
  5. List two examples of rates that were not in the notes.
  6. Considering collision theory, why does increasing temperature increase the rate of a reaction?
  7. Considering collision theory, why does increasing concentration increase the rate of a reaction?
  8. What is a catalyst and how does it work?
  9. What is an inhibitor and how does it work?
  10. What pollutants are removed from tailpipe emissions by a catalytic converter? What compounds are emitted instead?
  11. What are the basic ingredients of gunpowder?
  12. What did Alfred Nobel invent? Why did it lead him to initiate the Nobel Prize awards?
  13. How/why does the presence of oxygen in an explosive increase the rate of reaction (compared to combustible materials that contain little or no oxygen)?
  14. Why do high explosives always contain nitrogen? Discuss two reasons why its presence can result in an explosion.

Reversible Reactions and LeChatelier’s Principle

  1. In a typical “forward” reaction, products are formed? What can form in a “reverse” reaction?
  2. The concentrations of the reactants and products are NOT the same in a system in chemical equilibrium. What is the same?
  3. What does it mean to say that the reactants are favored in a system in chemical equilibrium?
  4. According to LéChâtelier, what happens if a stress is applied to a system in chemical equilibrium?
  5. Considering probability, why does the addition of a reactant favor the production of product?
  6. Determine how the following changes would affect this equilibrium (reactants favored or products favored):

PCl5(g) + heat ↔ PCl3(g) + Cl2(g).

(a) temperature increased, (b) PCl5 removed, (c) pressure reduced, (d) Cl2 added, (e) system is placed into a refrigerator, (f) gases compressed, (g) phosphorus trichloride added.

Equilibrium Constant (Keq)

  1. If Keq for a system in chemical equilibrium is less than 1, what side of the equilibrium is favored?
  2. Consider the reaction: CH4(g) + H2O(g) ↔ CO(g) + 3H2(g).

What is Keq and what side of the equilibrium is favored given the following concentrations?

[CH4] = 0.0370 M, [H2O] = 0.827 M, [CO] = 2.51, [H2] = 1.56 M

  1. Consider the reaction: 2NO(g) + O2(g)↔ 2NO2(g)

What is Keq and what side of the equilibrium is favored given the following concentrations? [NO] = 0.0062 M, [O2] =0.000831 M, [NO2] = 0.073 M

  1. Consider the reaction: CH3OH(g) ↔ CO(g) + 2H2(g)

What is Keq and what side of the equilibrium is favored given the following concentrations? [CH3OH] = 0.818 M, [CO] = 1.402 M, [H2] = 0.602 M

  1. The equilibrium constant of the following reaction for the decomposition of phosgene at 25oC is 4.282 x 10-2.

COCl2(g) ↔ CO(g) + Cl2(g). If the concentration of CO is 5.90 x 10-3 M and Cl2 is 5.90 x 10-3 M, what is the concentration of COCl2? (Hint: you have to do some algebra! You can do it!!)

Answers: (21) 311, products, (22) 1.7x105, products, (23) 0.622, reactants (24) 0.000813 M

Entropy and Spontaneous Processes

G = (n ·Gof)products - (n ·Gof)reactants

  1. Define entropy.
  2. Which processes tend to be spontaneous: (a) endothermic or exothermic reactions? (b) increased entropy or decreased entropy?
  3. Consider the chemical reaction:

H3C6H5O7 (aq) + 3 NaHCO3(s) + energy → 3 CO2(g) + 3 H2O(l) + Na3C6H5O7(aq)

Give two reasons why this reaction would tend to be spontaneous and one reason why it would not (think about enthalpy and entropy).

  1. Consider the chemical reaction:

C (s) + O2 (g) → CO2(g) + 394 kJ

Considering both enthalpy and entropy, would you expect this reaction to spontaneous or non-spontaneous? Explain why.

Gibbs Free Energy Calculations (G0)

  1. What three factors determine if free energy is released in a spontaneous process?
  2. Explain how increasing entropy can be used to do work. Hint: think about expanding gases!

NOTE: for the following problems, it is necessary to break aqueous compounds into their individual ions. For example, the Gf0value for CaCl2(aq) cannot be found on the table, but values for both Ca2+ (aq) and 2Cl-(aq) can be found.

  1. Calculate G0 for the following reaction (at 25oC) and indicate if the reaction is spontaneous or non-spontaneous.

4Al(s) + 3O2(g) → 2Al2O3(s)

  1. Calculate G0 for the following reaction (at 25oC) and indicate if the reaction is spontaneous or non-spontaneous.

ZnCl2(aq) + H2(g) → Zn(s) + 2HCl(aq)

  1. Calculate G0 for the following reaction (at 25oC) and indicate if the reaction is spontaneous or non-spontaneous.

CaCO3(s) → CaO(s) + CO2(g)

  1. Calculate G0 for the following reaction (at 25oC) and indicate if the reaction is spontaneous or non-spontaneous.

MgCO3(s) + 2NaCl(aq) → MgCl2(aq) + Na2CO3(aq)

Answers: (31) -3164.6 kJ, spontaneous, (32) +147.1 kJ, non-spontaneous, (33) +130.4 kJ,non-spontaneous, (34) +29.5 kJ

There will be extra-credit solubility constant (Ksp) problems on the test. See Chapter 18 for examples.