SI Chem 178

Ch 14/19 Exam Review

Leader: Emily

April 15th 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 = / Derived rate equation / Slope of linear graph / ____ versus t gives a linear graph
Rate Order / Units of k / t ½ Equation
0
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 (lnkvs 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, non-spontaneous, 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*sysDoes 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 non-spontaneous?

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.