Chemistry 100
description
Transcript of Chemistry 100
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Chemistry 100
Chapter 14 - Chemical Kinetics
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The Connection Between Chemical Reactions and Time
Not all chemical reaction proceed instantaneously!!!
2 A B combination reaction Can we quantify the length of time it
takes for this (or any) chemical reaction to occur?
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Practical examples of how long!
H2(g) + ½ O2 (g) H2O (l) Þ Very SlowN2O(g) N2 (g) + ½ O2 (g) Þ SLOW
combustion reactions Þ fast process TNT exploding Þ very fast reaction Food spoilage Drug decomposition
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Chemical Kinetics
Chemical kinetics is concerned with determining the speed or rate at which a reaction occurs.
How is the reaction rate affected by temperature? states of reactants? amount of reactants? catalyst? surface area of the reacting species?
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Example
Br2 (aq) + HCOOH (aq) 2 Br - (aq) + 2 H+ (aq) + CO2 (g)
Define the average rate
initialfinal
initial2final22
tt]Br[]Br[
t]Br[
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A Sample
ReactionBr2 (aq) + HCOOH (aq) 2 Br – (aq) + CO2 (g) +
2 H+ (aq) Note – orange colour fades as reaction
proceeds.
t / s [Br2] / M
0 0.012050 0.0101
100 0.00846150 0.00710200 0.00596250 0.00500300 0.00420350 0.00353400 0.00296
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Average Rate Data
Avg. Rate (150 s)
Avg. Rate (100 s)
Avg. Rate (50 s)
3.27 x 10-5 3.54 x 10-5 3.80 x 10-5
2.76 x 10-5 3.00 x 10-5 3.28 x 10-5
2.31 x 10-5 2.50 x 10-5 2.72 x 10-5
1.93 x 10-5 2.10 x 10-5 2.28 x 10-5
1.62 x 10-5 1.76 x 10-5 1.92 x 10-5
1.47 x 10-5 1.60 x 10-5
1.34 x 10-5
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Instantaneous Rates
td]Brd[ 2
It’s best to define an instantaneous ‘speed of reaction’
Plot of the [Br2] vs. time
0.000
0.002
0.004
0.006
0.008
0.010
0.012
0.014
0 100 200 300 400 500
time / s
[Bro
min
e] /
M
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Instantaneous Rate DataTime (s) Rate (M/s)
0 4.19 x 10-5
50 3.52 x 10-5
100 2.95 x 10-5
150 2.48 x 10-5
200 2.08x 10-5
250 1.75 x 10-5
300 1.47 x 10-5
400 1.03 x 10-5
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Plot of the ln [Br2] vs. time
-11.600
-11.400
-11.200
-11.000
-10.800
-10.600
-10.400
-10.200
-10.0000 100 200 300 400 500
time / s
[Bro
min
e] /
M
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The Rate Law
Relates rate of the reaction to the reactant concentrations and rate constant
For a general reaction a A + b B + c C d D + e E
rate = k[A]x[B]y[C]z
The exponents (x,y, and z) are called the reactant orders.
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The Rate Constant
The rate constant relates the ‘speed’ of the chemical reaction to the instantaneous reactant concentration.
k = constant for constant temperature The rate of the reaction is dependent on
reactant concentration RATE CONSTANT IS INDEPENDENT OF THE
REACTANT CONCENTRATION.
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The Reaction Orders
Determine the superscripts (x, y, and z) for a non-elementary chemical reaction by experimentation.
S x + y + z = reaction ordere.g. x = 1; y = 1; z = 0
2nd order reaction (x + y + z = 2)x = 0; y = 0; z = 1 (1st order reaction)
x = 2; y = 0; z = 0 (2nd order)
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Reaction Rates and Reaction Stoichiometry
Look at the reaction
O3(g) + NO(g) NO2(g) + O2(g)
t
]O[+ = t
]NO[+ = t
[NO]- = t
O- = rate 223
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Another Example
ttt
2Cl
11NO
21NOCl
21- rate
2 NOCl (g) 2 NO + 1 Cl2 (g)
WHY? 2 moles of NOCl disappear for every 1 mole Cl2 formed.
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The General Case
a A + b B c C + d D
rate = -1 [A] = -1 [B] = +1 [C] = +1 [D] a t b t c t d t
Why do we define our rate in this way? removes ambiguity
obtain a single rate for the entire equation, not just for the change in a single reactant or product.
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The Isolation Method of Obtaining Rate Laws
a A + b B c C + d Drate = k[A]x[B]y
Fix the concentration of one reactant (say reactant A).
We then perform a series of experiments to examine how changing the [B] affects the initial reaction rate?
rate = (constant) [B]y
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Isolation Method (cont’d)
Now we fix the concentration of reactant B. We then perform another series of experiments
to examine how changing the [A] affects the initial reaction rate?
rate = (constant) [A]x
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Types of Reactions
The rate law gives us information about how the concentration of the reactant varies with time
How much reactant remains after specified period of time?
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First Order Reaction
A product rate = -[A]/ t = k[A]
How does the concentration of the reactant depend on time?
k has units of s-1
ktAAo
ln
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The Half-Life of a First Order Reaction
For a first order reaction, the half-life t1/2 is calculated as follows.
k6930t 21.
/
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Radioactive Decay
Radioactive Samples decay according to first order kinetics.
This is the half-life of samples containing e.g. 14C , 239Pu, 99Tc.
Example 01
1414 NC
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Second Order Reaction
A + B products Rate = k[A][B]A products Rate = k[A]2
Reaction 1 is 1st order in A and B and 2nd order overall
Reaction 2 is 2nd order in A
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The Dependence of Concentration on Time
For a second order process where rate = k[A]2
ktA1
A1
o
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Half-life for a Second Order Reaction.
[A] at t = t½ = ½ [A]0
021
21o0
Ak1 = t or
kt + A1 =
/2A1
][
][][
/
/
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A Pseudo-First Order Reaction
Example hydration of methyl iodideCH3I(aq) + H2O(l) CH3OH(aq) + H+(aq) + I-(aq)
Rate = k [CH3I] [H2O] Since we carry out the reaction in aqueous
solution [H2O] >>>> [CH3I] / [H2O] doesn’t change by a lot
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Since the concentration of H2O is essentially constant
rate = k [CH3I] [constant]
= k` [CH3I] where k` = k [H2O]
Pseudo first order since it appears to be first order, but it is actually a second order process.
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Collision Theory of Kinetics
With few exceptions, the reaction rate increases with increasing temperature.
Chemical reactions take place due to collisions between reactant molecules
i.e. rate µ number of collisions / unit time
A2 + B2 product rate = k[A2][B2]
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The Reaction Profile
How does the energy of the reactants vary during the reaction sequence?
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The Activation Energy The minimum amount of
energy need for initiation of a chemical reaction is the activation energy (Ea).
Colliding reactant molecules possess kinetic energy > the activation energy or Ea.
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The Activated Complex
The species temporarily formed by the reactant molecules – the activated complex.
A small fraction of molecules usually have the required kinetic energy to get to the transition state The concentration of the activated
complex is extremely small.
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The Arrhenius Equation
/RTE - A ln k ln a
• Arrhenius showed how the rate constant depended on temperature.
A is called the frequency factor – an estimate of the number of reactive collisions in the system
Ea is the activation energy
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12
a
1
2
T1 -
T1
RE -
kkln
Reaction possesses a large activation energy Þ small rate constant Slow reaction!!
Measure k at several different temperatures
Activation Energies and the Arrhenius Equation
R = 8.314 J/(K mole) T in Kelvin units!!!
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Arrhenius Equation (cont’d)
The Arrhenius equation is best suited for studying reactions between simple species (atoms, diatomic molecules).
The orientation of the reactants (how they collide) becomes very important when the species get bigger.
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Catalysts
So far, we have considered one way of speeding up a reaction increasing T usually increases k.
Another way is by the use of a catalyst.
A catalyst - a substance that speeds up the rate of the reaction without being consumed in the overall reaction.
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· look at the following two reactions
A+B C rate constant k
A+B C rate constant with catalyst is kc
· NOTE: RATE WITH CATALYST > RATE WITHOUT CATALYST
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Types of Catalyst
We will briefly discuss three types of catalysts. The type of catalyst depends on the phase of the catalyst and the reacting species. Homogeneous Heterogeneous Enzyme
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Homogeneous Catalysis
The catalyst and the reactants are in the same phase
e.g. Oxidation of SO2 (g) to SO3 (g) 2 SO2(g) + O2(g) 2 SO3 (g) SLOW
Presence of NO (g), the following occurs.
NO (g) + O2 (g) NO2 (g)NO2 (g) + 2 SO2 (g) 2 SO3 (g) + NO (g)
FAST
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SO3 (g) is a potent acid rain gasH2O (l) + SO3 (g) H2SO4 (aq)
Note the rate of NO2(g) oxidizing SO2(g) to SO3(g) is faster than the direct oxidation.
NOx(g) are produced from burning fossil fuels such as gasoline, coal, oil!!
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Heterogeneous Catalysis
The catalyst and the reactants are in different phases adsorption the binding of molecules to
the surface to a surface. Adsorption on the surface occurs on
active sites. An active site is a place where reacting
molecules are adsorbed and physically bond to the metal surface.
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The hydrogenation of ethene (C2H4 (g)) to ethane
C2H4 (g) + H2(g) C2H6 (g) Reaction is energetically favourable
rH = -136.98 kJ/mole of ethane. With a finely divided metal such as Ni
(s), Pt (s), or Pd(s), the reaction goes very quickly .
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Common Heterogeneous Catalysts
Two other important heterogeneous catalysis processes petroleum cracking (refining crude oil) catalytic converters very efficient in
reducing exhaust emission when hot; cold is another story!
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Enzyme Catalysis
Enzymes - proteins (M > 10000 g/mol) High degree of specificity (i.e., they will
react with one substance and one substance primarily
Living cell > 3000 different enzymes
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The Lock and Key Hypothesis
Enzymes are folded into fixed configurations.
According to Fischer, active site is rigid.
The substrate’s molecular structure exactly fits the “lock” (hence, the “key”).
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Simplified Model for Enzyme Catalysis
E º enzyme; S º substrate; P º product
E + S ES ES P + E
rate = k [ES] The reaction rate depends directly on
the concentration of the substrate.
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Rate Laws for Multistep Processes
Chemical reactions generally proceed via a large number of elementary steps - the reaction mechanism
The experimentally established rate law must reflect the reaction rate of the slowest elementary step Þ the rate determining step (rds)
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What do we mean by an rds?
A commuter goes through a two step process to get to work in Halifax. (1) highway MacKay Bridge Toll
booth (2) toll booth downtown
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MacKay Bridge Toll Booth
Overall reaction
highway downtown
Situation 1. Þ highway clogged, toll booth is fast.
Situation 2. Þ fast highway, clogged toll booth.
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The Rate Determining Step (rds)
Situation 1 - clogged highway is the slowest step in the commuting process (rds).
Situation 2 - the clogged toll-booth is the slowest step in the commuting process (the rds).
Speed of overall process (highway downtown) depends which step is slowest! Which is the rate-determining step.
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Elementary steps and the Molecularity
Any chemical reaction occurs via a sequence of elementary steps.
Kinetics of the elementary step only depends on the number of reactant molecules in that step! Molecularity the number of reactant
molecules that participate in elementary steps
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The Kinetics of Elementary Steps
Classes of elementary steps
Akrate productsA
B Akrate productsBA
A bimolecular step
A unimolecular step
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For the step
BAkrate productsBA2 2
A termolecular (three molecule) step.