Transcript Kinetics - Pre

Kinetics
Introduction
• Rate of Reaction – the rate at which reactants
are consumed and products are formed.
• Chemical Kinetics – the study of the rates of
reactions.
Factors that affect Reaction Rates
• Nature of the Reactants – surface area and
color
• Concentration
• Temperature
• Presence of a Catalyst
Rate = k[A]x[B]y
• Zero order – the concentration of the reactant
does not affect the rate of production of
product. The concentration of reactant is
doubled, the rate of production of product not
affected.
• First order – the concentration of the reactant
does affect the rate of production of product.
The concentration of reactant is doubled, the
rate of production of product double.
• Second order – the concentration of the
reactant does affect the rate of production of
product. The concentration of reactant is
doubled, the rate of production of product
quadruple.
• Overall order = sum of all orders
Example 1
Expt #
1
2
3
4
[CH3CHO]
0.10
0.20
0.30
0.40
Rate (M/sec)
0.020
0.080
0.180
0.320
• Expt #
1
2
3
4
5
Example 2
[CO]
5.1 x 10-4
5.1 x 10-4
5.1 x 10-4
1.0 x 10-3
1.5 x 10-3
[NO2]
0.35 x 10-4
0.70 x 10-4
0.18 x 10-4
0.35 x 10-4
0.35 x 10-4
Rate (M/hr)
3.4 x 10-8
6.8 x 10-8
1.7 x 10-8
6.8 x 10-8
1.02x 10-7
• Expt
1
2
3
4
2A + B  C + D
[A]
0.10
0.20
0.30
0.20
[B]
0.05
0.05
0.05
0.15
Rate (M/sec)
6.0 x 10-3
1.2 x 10-2
1.8 x 10-2
1.1 x 10-1
• Expt
1
2
3
4
2A + B + C  D + E
[A]
0.20
0.40
0.20
0.20
[B]
0.20
0.30
0.30
0.40
[C]
0.20
0.20
0.20
0.60
Rate (M/min)
2.4 x 10-6
9.6 x 10-6
2.4 x 10-6
7.2 x 10-6
C2H4 + O3  2CH2O + ½ O2
• Expt
1
2
3
[O3]
5.0 x 10-8
1.5 x 10-7
1.0 x 10-7
[C2H4]
1.0 x 10-8
1.0 x 10-8
2.0 x 10-8
Rate (M/sec)
1.0 x 10-12
3.0 x 10-12
4.0 x 10-12
First Order Reactions
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ln([A]0/[A]) = akt
A0 = Initial Concentration
A = Concentration at some point
a = Coefficient of A
k = Rate Constant
t = Time
Half Life of a First Order Rxn
• t1/2 = ln2/ak
• Compound A decomposes to form B and C in a
reaction that is first order with respect to A
and first order overall. At 25C, the specific
rate constant for the reaction is 4.50 sec-1.
What is the half life of A?
AB+C
• The reaction 2N2O5  2N2O4 + O2 obeys the
rate law: rate = k[N2O5], in which the specific
rate constant is 0.00840 sec-1 at a certain
temperature. If an initial molarity is 4.2M,
what would the molarity of it be after 1
minute?
• How long would it take for it to decrease to
0.5M?
Second Order Rxn
• 1/[A] - 1/[A]0 = akt
• t1/2 = 1/(ak[A]0)
• Compound A reacts to form C and D in a
reaction that was found to be second order
overall. The rate constant is 0.622 M-1 min-1.
What is the half life of A when 4.10 x 10-2 M A
is reacted.
• The gas phase decomposition of NOBr is
second order in [NOBr] with k = 0.810 M-1sec-1
at 10C. We start with 4.00 x 10-3M NOBr in a
flask at 10C. How many seconds does it take
to reduce the concentration of NOBr to
2.5 x 10-3 M?
2NOBr  2 NO + Br2
• Consider the reaction in the previous problem.
If we start with 6.2 M NOBr, what
concentration of NOBr will remain after 5.00
minutes of reaction?
Rate Determining Step
• A reaction can never proceed faster than its
slowest step.
• Intermediates – Substances that are produced
by one reaction and later consumed by
another reaction.
• Catalyst – Substances that start a reaction and
later produced by another reaction
Example 1
• NO2 + NO2  N2O4 (fast)
• N2O4 + CO  NO + CO2 + NO2 (slow)
• What are the intermediates?
• What is the rate law for this reaction?
Example 2
• NO + Br2  NOBr2 (slow)
• NOBr2 + NO  2NOBr (fast)
• What are the intermediates?
• What is the rate law for this reaction?
2B  C + D
Expt
[B]
Rate (M/sec)
1
0.40
6.0 x 10-3
2
0.80
1.2 x 10-2
3
1.20
1.8 x 10-2
1.
2.
3.
4.
5.
What is the order with respect to B?
What is the rate law?
What is k?
If the initial concentration of B is 4.2, how long would it take for it to reach 0.5M?
Which of the following mechanisms correctly justifies your rate law? Justiify
1. B  C + E (Slow)
B + E  D (Fast)
2. A + B  C + E (Fast)
B + E  A + D (Slow)
• How to get a date to the Dance?