Chemical Reaction Engineering

1. Chemical Reaction Engineering Lecture (1) Week 1

2. Chemical Engineering Reaction • CourseLecturer : • Prof. Dr. Mona E. Ossman & • Course Assistant • Eng AmiraElgindi & • Eng Dina Sobhy

3. Course Objectives • This course is designed to give the students an ability to exploit and interpret chemical reaction kinetics and to apply the basic principles to design various types of reactor systems.

4. Course Description • In this course a general methodology will be developed for analysis and design of a variety of systems for which engineering of reactions is needed.

5. Recommended Text Books • H. Scott Fogler 2006. Elements of Chemical Reaction Engineering, 4th Ed, Prentice Hall Inc. • Missen, R.W., 1999. Introduction to Chemical Reaction Engineering and Kinetics, John Wiley sons.

6. Course Assessment • Assignments, attendance, • Presentations and lab 10% • Quizzes (4 bi weekly) 10% • Mini project 10% • Midterm 20% • Final 50% • Total 100%

7. Introduction • The study of chemical reaction engineering (CRE) combines the study of chemical kinetics with the reactors in which the reactions occur. • Chemical kinetics is the study of chemical reaction rates and reaction mechanisms. • Chemical kinetics and reactor design are at the heart of producing almost all industrial chemicals

8. Introduction • Chemical Identity • A chemical species is said to have reacted when it has lost its chemical identity. • The identity of a chemical species is determined by the kind, number, and configuration of that species’ atoms.

9. The Rate of Reaction, • The reaction rate is the rate at which a species looses its chemical identity per unit volume. • The rate of reaction tells us how fast a number of moles of one chemical species are being consumed to form another chemical species. • A given number of molecules (e.g., mole) of a particular chemical species have reacted or disappeared when the molecules have lost their chemical identity. • The rate of a reaction (mol/dm3.s) can be expressed as either the rate of Disappearance: -rA or as the rate of Formation (Generation): rA

10. The Rate of Reaction, Consider the isomerizationAB rA = the rate of formation of species A per unit volume -rA = the rate of a disappearance of species A per unit volume rB = the rate of formation of species B per unit volume

11. The Rate of Reaction, • Example: AB If Species B is being formed at a rate of 0.4 moles per decimeter cubed per second, ie, rB = 0.4 mole/dm3.s • Then A is disappearing at the same rate: -rA= 0.4 mole/dm3.s • The rate of formation (generation of A) is rA= - 0.4 mole/dm3.s

12. Relative rates of reaction (stoichiometrec coefficients) • We see that for every mole of A that is consumed, c/a moles of C appear. In other words,

13. Similarly, the relationship between the sates of formation of C and D is • The relationship can be expressed directly from the stoichiometry of the reaction,

14. For example, • the reaction equation for the well-known Haber process, used industrially to produce ammonia, is: • N2 + 3 H2 2 NH3 • N2 has a stochiometric coefficient of 1, H2 has a coefficient of 3, and NH3 has a coefficient of 2. • We could determine the rate of this reaction in any one of three ways, by monitoring the changing concentration of N2, H2, or NH3. Say we monitor N2, and obtain a rate of -d[N2]/dt = x mol dm-3 s-1. • Since for every mole of N2 that reacts, we lose three moles of H2, if we had monitored H2 instead of N2 we would have obtained a rate -d[H2]/dt = 3x mol dm-3 s-1.

15. Similarly, monitoring the concentration of NH3 would yield a rate of 2x mol dm-3 s-1. • Clearly, the same reaction cannot have three different rates, so we appear to have a problem. The solution is actually very simple: the reaction rate is defined as the rate of change of the concentration of a reactant or product divided by its stochiometric coefficient. • For the above reaction, the rate (r) is therefore

16. Heterogeneous reactions involve more than one phase. In heterogeneous reaction systems, the rate of reaction is usually expressed in measures other than volume, such as reaction surface area or catalyst weight. • For a gas-solid catalytic reaction, the gas molecules must interact with the solid catalyst surface for the reaction to take place. • The dimensions of this heterogeneous reaction ( catalytic reaction) rate, rA’; (prime). are the number of moles of A reacted per unit time per unit mass of catalyst (mol/s. g catalyst).

17. The Rate of Reaction, Consider species j: • rj is the rate of formation of species j per unit volume [e.g. mol/dm3.s] • rj is function of concentration, temperature, pressure, and the type of catalyst (if any) • rj is independent of the type of reaction system i.e. the reactor (batch reactor, plug flow reactor, etc.) • rj is an algebraic equation, not a differential equation • NOTE: dCA/dt is not the rate of reaction

18. Parameters Affecting Rate of Reaction: The Rate Law • Rate of reaction depends on a number of parameters, the most important of which are • usually (1)The nature of the species involved in the reaction; • Many examples of types of very fast reactions involve ions in solution, At the other extreme, very slow reactions may involve heterogeneous reactions. (2) Concentrations of species; • and usually increases as concentration of reactants increases. (3) Temperature; • and usually increases nearly exponentially as temperature increases.

19. (4) Catalytic activity; • Many reactions proceed much faster in the presence of a substance which is itself not a product of the reaction. This is the phenomenon of catalysis, and many life processes and industrial processes depend on it. (5) Nature of contact of reactants; • Thus the titration of an acid with a base occurs much faster if the acid and base are stirred together than if the base is simply allowed to “dribble” into the acid solution. (6) Wave-length of incident radiation. • Some reactions occur much faster if the reacting system is exposed to incident radiation of an appropriate frequency thus , a mixture of hydrogen and chlorine can be kept in the dark, and the reaction to form hydrogen chloride is very slow; however, if the mixture is exposed to ordinary light, reaction occurs with explosive rapidity. Such reactions are generally called photochemical reactions.