Title: Chemistry 100 Chapter 14 - Chemical Kinetics Chemistry 10
1Chemistry 100
- Chapter 14 - Chemical Kinetics
2The 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
3Practical 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
- Food spoilage
- Drug decomposition
4Chemical 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
5Example
- Br2 (aq) HCOOH (aq) 2 Br - (aq) 2 H (aq)
CO2 (g) - Define the average rate
6A Sample
- Reaction
- Br2 (aq) HCOOH (aq) 2 Br (aq) CO2 (g) 2
H (aq)
7Average Rate Data
8Instantaneous Rates
- Its best to define an instantaneous speed of
reaction
9Instantaneous Rate Data
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11The 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.
12Reaction Rates and Reaction Stoichiometry
- Look at the reaction
- O3(g) NO(g) NO2(g) O2(g)
13Another Example
- 2 NOCl (g) 2 NO 1 Cl2 (g)
WHY 2 moles of NOCl disappear for every 1 mole
Cl2 formed.
14The 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.
15The 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 kAxByCz
The exponents (xy and z) are called the
reactant orders.
16The 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)
17The Isolation Method of Obtaining Rate Laws
- a A b B c C d D
- rate kAxBy
- 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
18Isolation 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
19Types 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
20First Order Reaction
- A product
- rate -DA/D t kA
- How does the concentration of the reactant
depend on time -
- k has units of s-1
21The Half-Life of a First Order Reaction
- For a first order reaction the half-life t1/2 is
calculated as follows.
22Radioactive Decay
- Radioactive Samples decay according to first
order kinetics. - This is the half-life of samples containing e.g.
14C 239Pu 99Tc. - Example
23Second 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
24The Dependence of Concentration on Time
- For a second order process where rate kA2
25Half-life for a Second Order Reaction.
26A 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
27- 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.
28Collision 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
29The Reaction Profile
- How does the energy of the reactants vary during
the reaction sequence
30The 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.
31The 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.
32The 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 - Slow reaction!!
- Measure k at several different temperatures
R 8.314 J/(K mole) T in Kelvin units!!!
34Arrhenius 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.
35Catalysts
- 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.
36- 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
37Types 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
38Homogeneous 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
39- 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!!
40Heterogeneous 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.
41- 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 .
42Common 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!
43Enzyme 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
44The 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).
45Simplified 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.
46Rate 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)
47What 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
48MacKay Bridge Toll Booth
- Overall reaction
- highway downtown
- Situation 1. highway clogged toll booth is
fast. - Situation 2. fast highway clogged toll booth.
49The 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.
50Elementary 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
51The Kinetics of Elementary Steps
- Classes of elementary steps
A unimolecular step
A bimolecular step
52A termolecular (three molecule) step.