chemical kinetics

1. ChemE 2200 – Chemical Kinetics Lecture 1
Today:
Chemical Kinetics vs. Chemical Thermodynamics.
Reaction-Coordinate Energy-Level Diagrams.
Defining Question:
“How are Chemical Kinetics and Chemical
Thermodynamics of reactions represented graphically?”
Reading for Today’s Lecture:
McQuarrie & Simon, Chp 28.1.
Reading for Kinetics Lecture 2:
McQuarrie & Simon, Chp 28.2-28.4.

2. Career Guidance Mondays
12:25-1:15 p.m.
165 Olin

3. ChemE 2200 Final Exam
Thursday, May 16, 9:00 - 11:30 a.m.
The Final Exam will be ‘Prelim 3’
and will cover Chemical Kinetics only.

4. ChemE 2200 - Physical Chemistry II for Engineers
Part 3 - Chemical Kinetics
Reaction-Coordinate Energy-Level Diagrams
Thermodynamics vs. Kinetics
Rate Equations from Experimental Data
Method of Initial Rates
Saturation Methods
Method of Half Lives
Rate Constants and Activation Energies from Experimental Data
Arrhenius Plots
Rate Equations from Mechanisms of Elementary Reactions
Reactive Intermediates - the usual suspects
Steady-State Approximation
Pre-equilibrium Approximation
Rate-Limiting Step
Special Classes of Reactions
Chain Reactions
Photo-initiated Reactions
Polymerization
Homogeneous Catalysis - Enzymes
Heterogeneous Catalysis - Solid Acids and Transition Metals
Autocatalysis

5. ChemE 2200 - Physical Chemistry II for Engineers
Ancillary Skills
Mathematical Modeling & Graphical Modeling
How to translate a chemical and physical description
into equations and graphs.
Approximation
How to identify dominant effects and estimate
the consequences of neglecting secondary effects.
Evaluation
How to test assumptions and assess predictions.
Numerical Methods
Numerical Integration
Statistical methods - linear regression and least-squares fits
Analytical Concepts
Rate constants and time constants

6. The Rate Equation
Reactor Design & Analysis:
A reactor
A  P
A
P
Mass balance on reactor:
At steady state,
rate in - rate out + rate of formation - rate of consumption = 0
= 0 for A (reaction is irreversible)
For A,
FA, in - FA, out - rAVreactor = 0
rate equation: rA = (concentration, T)
The Two Goals of Chemical Kinetics:
1. Predict the reaction rate; obtain the rate equation.
Optimize reactor design.
Optimize reactor performance.
2. Devise the reaction mechanism; describe reactions at the molecular level.
The core of chemical engineering
Improve product selectivity.
Increase reaction rate.

7. Kinetics and Thermodynamics
Thermodynamics: The energy of a chemical system.
Predicts the equilibrium configuration at a given T.
The equilibrium configuration is independent of path.
The equilibrium configuration is a state function.
Kinetics: The rate a chemical system changes configuration.
The rate depends on the path.

8. Kinetics and Thermodynamics
Two stable states:
Example 1: The energy states of your textbook.
vertical horizontal
H
W
P.E. = mg(½H) P.E. = mg(½W)
energy
mg(½H)
vertical
mg(½W)
horizontal
The horizontal state
is lower energy than
the vertical state.
Why doesn’t the
book spontaneously
go to the lower
energy state?
Consider a path from
the vertical state to
the horizontal state.
unstable intermediate
P.E. = mg(½(H 2+W 2)½)
mg(½(H 2+W 2)½)
reaction coordinate
a reaction-coordinate
energy-level diagram
thermodynamics
kinetics

9. Kinetics and Thermodynamics
A third stable state:
Example 1,cont’d: The energy states of your textbook.
vertical on binding
H
W
P.E. = mg(½H) P.E. = mg(½D)
unstable intermediate
P.E. = mg(½(H 2+D 2)½)
energy
mg(½(H 2+W 2)½)
mg(½H)
vertical
mg(½W)
horizontal
reaction coordinate
D
D
mg(½D)
on binding
mg(½(H 2+D 2)½) Kinetics favors
transition to the
horizontal state.
Thermodynamics
favors the horizontal
state at equilibrium.

10. Kinetics and Thermodynamics
An alternative reaction-coordinate energy-level diagram.
Example 1,cont’d: The energy states of your textbook.
energy
vertical
horizontal
reaction coordinate
on binding
Initial state may
‘react’ to the right
or to the left.
‘On binding’ state
may ‘react’ to
horizontal state.

11. Kinetics and Thermodynamics
Example 2: H2 + ½O2  H2O(g)
energy
reaction coordinate
H2 + ½O2
DGrxn = -229 kJ/mol = -2.4 eV/molecule Reaction is highly exothermic.
Half life is ~ age of the universe. Reaction is extraordinarily slow.
0 kJ/mol
-229 kJ/mol
H2O
Reaction conversion is ~100%.
Thermodynamics
Kinetics
H2 + •O•
+249 kJ/mol
catalyst: -Pt-Pt-Pt-Pt-Pt-
H2 + O2
H H O H O
H2O
A catalyst can provide
an alternative path with
a lower energy barrier.
A catalyst can accelerate
the reaction rate to the
same thermodynamic state.

12. Kinetics and Thermodynamics
Example 3: CH2=CH2 + ½O2 
energy
reaction coordinate
CH2=CH2 + ½O2
DGrxn = -81 kJ/mol
+68 kJ/mol
-1246 kJ/mol
CH2=CH2 + •O•
+317 kJ/mol
catalyst: -Ag-Ag-Ag-Ag-Ag-
O
A catalyst can provide
an alternative path with
a lower energy barrier.
A catalyst can kinetically
favor the desired product.
O
CH2
CH2
undesired reaction: CH2=CH2 + 3O2  2CO2 + 2H2O DGrxn = -1314 kJ/mol
CH2 CH2
O
-13 kJ/mol
2CO2+ 2H2O
CH2 CH2
O
CH2=CH2 + O2
CH2CH2 CH2CH2

13. Kinetics and Thermodynamics - Most Probable Reaction Mechanism?
Example 4: CO + 3H2  CH4 + H2O
energy
+442 kJ/mol
+903 kJ/mol
+157 kJ/mol
reaction coordinate
CH4 + H2O
-137 kJ/mol
-280 kJ/mol
CO + 3H2
C + O + 3H2
C + H2O + 2H2
CH2 + H2O + H2
+269 kJ/mol
CO + 2H + 2H2
-103 kJ/mol
H2CO + 2H2
+618 kJ/mol
H2C + O + 2H2