Title: Chemistry 100 Chapter 14 - Chemical Kinetics Chemistry 10
1 Chemistry 100
Chapter 14 - Chemical Kinetics
2 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
3 Practical examples of how long!
H2(g) ½ O2 (g) H2O (l) Þ Very Slow
N2O(g) N2 (g) ½ O2 (g) Þ SLOW
combustion reactions fast process
TNT exploding very fast reaction
4 Chemical Kinetics
Chemical kinetics is concerned with determining the speed or rate at which a reaction occurs.
How is the reaction rate affected by
states of reactants
amount of reactants
surface area of the reacting species
Br2 (aq) HCOOH (aq) 2 Br - (aq) 2 H (aq) CO2 (g)
Define the average rate
6 A Sample
Br2 (aq) HCOOH (aq) 2 Br (aq) CO2 (g) 2 H (aq)
7 Average Rate Data 8 Instantaneous Rates
Its best to define an instantaneous speed of reaction
9 Instantaneous Rate Data 10 (No Transcript) 11 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.
12 Reaction Rates and Reaction Stoichiometry
Look at the reaction
O3(g) NO(g) NO2(g) O2(g)
13 Another Example
2 NOCl (g) 2 NO 1 Cl2 (g)
WHY 2 moles of NOCl disappear for every 1 mole Cl2 formed. 14 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
obtain a single rate for the entire equation not just for the change in a single reactant or product.
15 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
The exponents (xy and z) are called the reactant orders. 16 The Reaction Orders
Determine the superscripts (x y and z) for a non-elementary chemical reaction by experimentation.
S x y z reaction order
e.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)
17 The Isolation Method of Obtaining Rate Laws
a A b B c C d D
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) By
18 Isolation Method (contd)
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) Ax
19 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
20 First Order Reaction
rate -DA/D t kA
How does the concentration of the reactant depend on time
k has units of s-1
21 The Half-Life of a First Order Reaction
For a first order reaction the half-life t1/2 is calculated as follows.
22 Radioactive Decay
Radioactive Samples decay according to first order kinetics.
This is the half-life of samples containing e.g. 14C 239Pu 99Tc.
23 Second Order Reaction
A B products Rate kAB
A products Rate kA2
Reaction 1 is 1st order in A and B and 2nd order overall
Reaction 2 is 2nd order in A
24 The Dependence of Concentration on Time
For a second order process where rate kA2
25 Half-life for a Second Order Reaction.
A at t t½ ½ A0
26 A Pseudo-First Order Reaction
Example hydration of methyl iodide
CH3I(aq) H2O(l) CH3OH(aq) H(aq) I-(aq)
Rate k CH3I H2O
Since we carry out the reaction in aqueous solution
H2O gtgtgtgt CH3I / H2O doesnt change by a lot
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.
28 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 kA2B2
29 The Reaction Profile
How does the energy of the reactants vary during the reaction sequence
30 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 gt the activation energy or Ea.
31 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.
32 The Arrhenius Equation
Arrhenius showed how the rate constant depended on temperature.
Ea is the activation energy A is called the frequency factor an estimate of the number of reactive collisions in the system 33
Activation Energies and the Arrhenius Equation
Reaction possesses a large activation energy small rate constant
Measure k at several different temperatures
R 8.314 J/(K mole) T in Kelvin units!!! 34 Arrhenius Equation (contd)
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.
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.
look at the following two reactions
AB C rate constant k
AB C rate constant with catalyst is kc
NOTE RATE WITH CATALYST gt RATE WITHOUT CATALYST
37 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.
38 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
SO3 (g) is a potent acid rain gas
H2O (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!!
40 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.
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 .
42 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!
43 Enzyme Catalysis
Enzymes - proteins (M gt 10000 g/mol)
High degree of specificity (i.e. they will react with one substance and one substance primarily
Living cell gt 3000 different enzymes
44 The Lock and Key Hypothesis
Enzymes are folded into fixed configurations.
According to Fischer active site is rigid.
The substrates molecular structure exactly fits the lock (hence the key).
45 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.
46 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)
47 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
48 MacKay Bridge Toll Booth
Situation 1. highway clogged toll booth is fast.
Situation 2. fast highway clogged toll booth.
49 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.
50 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
51 The Kinetics of Elementary Steps
Classes of elementary steps
A unimolecular step A bimolecular step 52
For the step
A termolecular (three molecule) step.
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