SI Chem 178
Ch 14/19 Exam Review
Leader: Emily
th
April 15 and 16th, 2013
Conceptual Concepts: Chapter 14
The speed at which a chemical reaction occurs is called ____________ ____________.
What are the “Big 4” things that affect reaction rates? Indicate what is required in each
situation to INCREASE reactant rate.
Why would the average rate of a reaction decrease over time? (Consider what’s occurring
in the reaction vessel.)
What are the units for reaction rate? ________________ or ____________________.
Reaction rates are always
+
- or either +/How can you express the disappearance of “A” and appearance of “B” if you were given
their concentration values at particular times? (aAbB) Does stoichiometry matter here?
Delta () signifies ______________ minus _________________.
Describe how you approach determining rate law based on experimental data and initial
rates. What is the equation for this type of rate law determination? Does stoichiometry
matter here?
With your generalized rate law equation from above, how do you determine rate order?
- - - Hint _____________ + ____________ = rate order
Rate order, in turn affect the ___________ _____________, symbolized as k.
T/F Concentration affects k
T/F Concentration affects rate
Generally, reaction rates ______________ as temp increases.
[A]^m
m=?

[2A]^___________

2^_____= 8
Helpful Tables:
Order
Rate =
Rate Order
0
Derived rate equation
Units of k
Slope of
linear
graph
____ versus t
gives a linear
graph
t ½ Equation
1
2
Copy the given Energy Graph: Label it, write out elementary equations, and consider the
reverse reaction.
As Ea increases, reaction rate _________________ .
What is the Arrhenius equation? Write out its basic form and its manipulated form for
comparing k at 2 temperatures.
Write out the Arrhenius plot (lnk vs 1/T)
How do you determine the molecularity of an elementary reaction?
T/F elementary rate laws relate directly to molecularity.
The rate determining step of a reaction is generally the ___________________.
Homogeneous catalysts are in the _______________ _________________ as the
reactants that they are interacting with.
T/F An example of a heterogeneous catalyst could be that of a solid interacting with
gaseous reactants.
_________________ are biological catalysts.
Conceptual Concepts: Ch 19
T/F A spontaneous process requires outside help to begin or proceed.
T/F Spontaneous processes occur in both directions (they can be reversed.)
T/F Thermodynamics tells us about the speed of a reaction.
Spontaneity often depends on experimental ____________ and ______________.
Isothermal means that T = _____or that the reaction occurred at a
_________________________temperature. In what cases is this applicable?
T/F In an irreversible process, it’s possible to return a system and its surroundings to their
original states via the same path of change they took in the first place.
Reversibility generally relates to the concept of __________________ between system
and surroundings.
Entropy describes ________________, therefore a negative value describes:
Entropy and enthalpy are ___________ functions, this means that: (circle one)
The path of their change matters OR
Only their final and initial values matter
At a constant T, S=
What are the units of S and T?
A microstate describes 1 snapshot of
________________________________________________________________, and how
does this relate to entropy?
What are things that increase microstates?
What are the three motions a real molecule can have? Can an atom have all of these?
Things that increase S…
At absolute 0, there is/are ___________ microstate(s) and the entropy is _________ .
(Consider a crystalline substance.)
S =
The (  ) sign indicates: _________________ __________________.
Are Ssurr and Ssys always equal and opposite in magnitude? What are the equations
for each at standard state?
G =
(generalized form)
G= ______________________________ or __________________________________
How can we relate G to G in equilibrium situations?
How can this be solved for K? (Equilibrium constant)
Ch 14 Problems
14.22 The following table gives data about the disappearance of HCl in this reaction:
CH3OH(aq) + HCl(aq)  CH3Cl(aq) + H2O(l)
Time (min)
[HCl] (M)
0.0
1.85
54.0
1.58
107.0
1.36
215.0
1.02
430.0
0.580
a. Calculate the avg rate of reaction between 54 and 107 minutes, 215 and 230 minutes,
and then find the average rate of reaction over the course of the entire time (0-430 min)
Sample Exercise 14.3 p.563
2O3(g)3O2(g)
1. Write out how their rates of disappearance/appearance relate
2. If O2 appears at a rate of 6x10^-5M/s, ([O2]/ t) what is the
rate that O3 disappears?
14.37 Consider the gas-phase reaction of 2NO2(g) + Br2(g)  2NOBr(g) Use the
following table to find: 1. Rate law 2. Rate constant (k) 3. How is the rate of appearance
of NOBr related to the appearance of Br2? 3. What is the rate of disappearance for Br2
when [NO]=0.075M and [Br2]=0.25M
Experiment #
[NO] (M)
[Br2] (M)
Initial Rate
(M/s)
1
0.10
0.20
24
2
0.25
0.20
150
3
0.10
0.50
60
4
0.35
0.50
735
14.45 The first order rate constant for the decomposition of SOCl2 at 660K is 4.5x10^-2
s^-1. If its initial pressure is at 450 torr, what is its pressure after 1 minute? At what time
will its pressure reach 1/10 of its original value? What about ½ its value?
14.64 NO decomposes to N2 and O2 with a rate constant of 0.0796 M-1s-1 at 737*C and
0.0815 M-1s-1 at 947*C. Calculate its activation energy.
What is the molecularity of the following? Write a rate law for each.
Cl2(g)  2Cl(g)
OCl-(aq) + H2O(l)  HOCl(aq) +OH-(aq)
A+A+AB+C
X+Y+ZD+E+F
Given: Step 1: NO(g) + NO(g)  N2O2(g) (fast step, and at equilibrium) (k1, k-1)
Step 2: N2O2(g) Br2(g)  2NOBr(g) (slow step) (k2)
Which step determines the rate law? What intermediates are observed? Write the rate law
into terms that would be experimentally observed.
Ch 19 Problems
Exercise 19.1 p787 Which of the following would you predict to be spontaneous, nonspontaneous, or at equilibrium:
a. 40*C water gets hotter as a 150*C piece of metal is added to it.
b. Water at room temperature decomposes into hydrogen and oxygen gas.
c. Benzene vapor condenses to liquid benzene at 80.1*C, its normal boiling point.
Exercise 19.2 p791 Mercury is a liquid at room temperature. Its freezing point is -38.9*C
and its Hfusion= 2.29kJ/mol. What is the entropy change of the system when 50g of
Hg(l) freezes at -38.9*C?
Exercise 19.5 Calculate S*univ based on the following data:
N2(g) + 3H2(g) 2NH3
At 298K, the S* (in J/mol x K) values for the three chemicals are as follows: 191.5,
130.6, and 192.5 respectively. The H* values (kJ/mol) are 0, 0 and -46.19
1. Calculate S*sys Does the sign make sense?
2. Calculate S*surr for the same reaction. Which equation must we use? What standard
values do we need?
3. What is S*univ?
Sample Exercise 19.12: N2(g) + 3H2(g)  2NH3(g) G*=33.3 kJ/mol
Calculate K.
19.69 For a particular reaction, H= -32 kJ and S=-98 J/K
a. At what temperature will the reaction have G=0? What does this mean?
b. If T is increased from that in part a., will the reaction be spontaneous or nonspontaneous?
19.79 2NO2(g)  N2O4(g) G* (kJ/mol) are 51.84 and 98.28 respectively
a. Calculate G*
b. Calculate G if the pressure of NO2 is 0.40atm and the pressure of N2O4 is 1.60atm.