Download Ch 13 kinetics

Chemistry 201B Dr. G Baxley Spring 2015
A. Reaction Rates (section 13.2)
How do chemists calculate reaction rates, and what units are typically used?
B. Five important factors which affect rates of reactions:
1. ___________________________________________________________
2. ___________________________________________________________
3. ___________________________________________________________
4. ___________________________________________________________
5. ___________________________________________________________
C. Calculating Reaction Rates
For X + Y → Z there are two ways of measuring rate:
There are three types of rate calculations:
Average: ______________________________________________________________________
Instantaneous:
________________________________________________________________
Initial: special in kinetics _________________________________________________________
DATA: with t in sec and [ ] in M
At t = 0 min, [A] = 1.00 M (100 red spheres) and [B] = 0 M
At t = 20 min, [A] = 0.54 M
and [B] = 0.46 M
At t = 40 min, [A] = 0.30 M
and [B] = 0.70 M
Calculate average rate of [A] or [B] from 0 to 20 min
Chapter 13
C. Calculating Reaction Rates
For the following reaction, C4H9Cl + H2O → C4H9OH + HCl
What happens to concentration of C4H9Cl over time?
What happens to rate of this reaction over time?
D. Reaction Rates and Stoichiometry
What does the reaction equation tell a chemist?
How many mol of H2O will be formed for every 2 mol of H2 consumed? 2H2 + O2 → 2H2O
For reaction of 2H2 + O2 → 2H2O, draw a graph showing relative rates of disappearance of H2 and O2.
For the gaseous reaction: 3H2 + N2 → 2 NH3, What is the reaction rate with respect to N2 if
the rate with respect to H2 is 0.15 M · s–1?
the rate with respect to NH3 is 0.15 M · s–1?
E. Rate Laws
Rate law: __________________________________________________________________________________
General form of rate law: _________________________________
Rate constant: ______________________________________________________________________________
The only way to determine rates laws and rate constants and exponents is with __________________________
The overall order of reaction is _________________________________________________________________
F. Rate Laws-Method of Initial Rates
Run several experiments ______________________________________________________________________
Measure ___________________________________________________________________________________
Compare ___________________________________________________________________________________
Example 1: Determine the rate law for the reaction of A + B → C
Example 2: For reaction of A + B2 → AB2 Calculate the rate law and the full value of the rate constant
Concept questions:
Given rate law of rate = k[A]3[B], what happens to a rate if both [A] and [B] are triple their initial value?
What has a greater affect on the rate, increasing [A] or [B]?
G. Rate Laws-Concentration w/ Time
Goal: convert rate law into a convenient equation for conc and time.
1st order:
2nd order
Zero order
Summary: 2 methods for determining rate laws
1.
_____________________________________________________________
2.
_____________________________________________________________
3. Can also look at ______________________________________________________________
H. Half-Life
The half-life is _______________________________________________________________________________
Does the concentration change the by the same molarity during each
1st order half-life? ______________________
Will a sample be completely gone after two half-lives? _____________
Does each 1st order half-life have the same length of time? _____________________
For a second order reaction, the half-life depends upon:
_________________________________________
Does each 2nd order half-life have the same length of time? _____________________
Examples:
a) If a reaction of 2A → B is 1st order and has a rate constant of 0.150 s–1, how much time will it take for [A] to
equal 0.0100 M if [A]0 = 0.0500 M? What is the value of [A] at 27.0 seconds?
b) If a reaction of 2A → B is 2nd order and has a rate constant of 0.150 M–1 s–1, how much time will it take for [A]
to equal 0.0100 M if [A]0 = 0.0500 M?
c) A first order reaction takes 400.0 s to decrease from 0.200 M to 0.00820 M. What is the half-life of this reaction?
How long will it take for the same reaction to decrease from 1.00 M to 0.125 M?
I: Temperature and Reaction Rates How does Temp affect reactions?
What portion of a rate law changes when temperature changes?
I: Temperature and Reaction Rates
A rate constant must depend on ________________________________________________________________
I: Temperature and Reaction Rates: The Orientation Factor
For the reaction Cl + NOCl → NO + Cl2 Draw the 2 reactants colliding. Are all collisions effective?
I: Temperature and Reaction Rates: Activation Energy
Arrhenius: molecules must possess ______________________________________________________________
Why?
bonds in reactants ___________________________________________________________________________
Bond breakage_________________________________ energy.
Activation energy, Ea, is the ___________________________________________________________________ .
For the 1st order reaction: CH3NC → CH3CN, draw a reaction coordinate diagram (PE vs. reaction time) for this
exothermic reaction.
PE
Reaction progress
I: Temperature and Reaction Rates: Activation Energy
The activation energy is the ____________________________________________________________________
The rate of any reaction depends on _____________________________________________________________
K: Reaction Mechanisms
More than a balanced chemical equation, a reaction mechanism ________________________________________ .
Provides a detailed picture of how a reaction occurs.
Elementary step:
Any process that occurs ____________________________________________________________________
Makes either ____________________________________________________________________________
the rate law for an ____________________________________ is determined from the equation as written
Intermediate: ______________________________________________________________________________ .
Molecularity: _______________________________________________________________________________
–
__________________________________: one molecule in the elementary step,
–
__________________________________::
–
__________________________________: three molecules in the elementary step
–
__________________________________: are not common in the elementary step
two molecules in the elementary step
K: Reaction Mechanisms
In a multistep process, one of the steps will be slower than all others.
The overall reaction cannot occur faster than this slowest, _____________________________________________ .
For the reaction:
NO2 (g) + CO (g) → NO (g) + CO2 (g)
•
The experimentally determined rate law is
__________________________________
•
Is CO necessary for the reaction? ________________
•
Does the rate depend on [CO]? _________________
•
So, the reaction must occur in ___________________
A proposed mechanism for this reaction is
Step 1: NO2 + NO2 → NO3 + NO (slow)
Step 2: NO3 + CO → NO2 + CO2 (fast)
NO3 is an ___________________________ and is consumed in the second step.
CO can’t be involved in the _________________________________________________________
Mechanisms with an Initial Fast Step
• For the reaction:
2NO(g) + Br2(g) → 2NOBr(g)
The experimentally determined rate law is
Rate = k[NO]2[Br2]
Are termolecular reactions common? __________________
Chemists job: propose mechanism that explain data and are reasonable in terms of statistics.
•
•
•
Rate law of rate determining step is
_____________________________________________
The rate law can’t have an intermediate.
Show [NOBr2] in terms of NOBr and Br2 assuming an equilibrium in step 1 (rates are equal)
Rate = k1[NO][Br2] and rate = k–1[NOBr2]
since rates are equal,
k1[NO][Br2] = k–1[NOBr2]
Rearrange: k1/k–1 [NO][Br2] = [NOBr2]
Substitute for NOBr2 in the slow step: Rate = k[NO][NOBr2] becomes
Rate = k×k1/k–1 [NO]2Br2]
L: Catalysts
A catalyst _________________________________________________________________________________ .
Is not ______________________________________________________________________________________
Hydrogen peroxide decomposes very slowly:
2H2O2(aq) → 2H2O(l) + O2(g)
In the presence of the bromide ion, the decomposition occurs rapidly:
2Br – + H2O2 + 2H+ → Br2 + 2H2O
Br2 + H2O2 → 2Br – + 2H+ + O2
Catalysts operate by:
1
2
3