1. Chemical Reaction Engineering Lecture (1) Week 2

2. 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

3. 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.

4. 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/dm 3.s) can be expressed as either the rate of Disappearance: -r A or as the rate of Formation (Generation): r A

5. The Rate of Reaction, Consider the isomerization A  B r A = the rate of formation of species A per unit volume -r A = the rate of a disappearance of species A per unit volume r B = the rate of formation of species B per unit volume

6. Example: A  B If Species B is being formed at a rate of 0.4 moles per decimeter cubed per second, ie, r B = 0.4 mole/dm 3.s – Then A is disappearing at the same rate: -r A = 0.4 mole/dm 3.s – The rate of formation (generation of A) is r A = - 0.4 mole/dm 3.s The Rate of Reaction,

7. 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,

8. 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,

9. For example, the reaction equation for the well-known Haber process, used industrially to produce ammonia, is: N 2 + 3 H 2 2 NH 3 N 2 has a stochiometric coefficient of 1, H 2 has a coefficient of 3, and NH 3 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 N 2, H 2, or NH 3. Say we monitor N 2, and obtain a rate of - d[N 2 ]/dt = x mol dm -3 s -1. Since for every mole of N 2 that reacts, we lose three moles of H 2, if we had monitored H 2 instead of N 2 we would have obtained a rate -d[H 2 ]/dt = 3x mol dm -3 s -1.

10. Similarly, monitoring the concentration of NH 3 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

11. Basic Definitions 1- Types of reactions Homogenous reactions It is one that involves only one phase. Reversible reactions It can proceed in either direction, depending on the concentrations of reactants and products relative to the corresponding equilibrium concentrations. Heterogeneous reactions It involves more than one phase, and the reaction usually occurs at the interface between the phases. Irreversible reactions It is one that proceeds in only one direction and continues in that direction until the reactants are exhausted. It behaves as if no equilibrium condition exists.

12. Basic Definitions 2. Molecularity of a reaction It is the number of atoms, ions. or molecule involved in a reaction step. The terms uni- molecular, bimolecular and ter-molecular refer to reactions involving, respectively. one. two, or three atoms (or molecules) interacting in any one reaction step.

13. Basic Definitions 3. Elementary reactions It is a reaction that involves: – single step, – The stoichiometrec coefficient in this reaction is identical to the power in the rate law. Example: in the following reaction

14. Consider species j: r j is the rate of formation of species j per unit volume [e.g. mol/dm 3.s] r j is function of concentration, temperature, pressure, and the type of catalyst (if any) r j is independent of the type of reaction system i.e. the reactor (batch reactor, plug flow reactor, etc.) r j is an algebraic equation, not a differential equation NOTE: dC A /dt is not the rate of reaction The Rate of Reaction,

15. 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.

16. (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.

17. Rate law and Reaction order The limiting reactant is usually chosen as our basis for calculation. The rate of disappearance of A, -r A, depends on temperature and composition. For many reactions, the rate of reaction can be written as the product of a reaction rate constant k A, and a function of the concentrations of the various species involved in the reaction: -r A = [k A (T)] [fn (C A, C B, …)] The algebraic equation that relates –r A, to the species concentrations is called the kinetic expression or rate law. The specific rate of reaction (also called the rate constant). k A

18. Rate law and Reaction order The dependence of the reaction rate. –r A, on the concentrations of the species present. fn(C A ), is almost without exception determined by experimental observation. The rate law is the product of concentrations of the individual reacting species, each of which is raised to a power. for example -r A = k A C A α C B β The exponents of the concentrations in the above Equation lead to the concept of reaction order.

19. Rate law and Reaction order The order of a reaction refers to the powers to which the concentrations are raised in the kinetic rate law. the reaction in the above equation is α order with respect to reactant A. and β order with respect to reactant B. The overall order of the reaction, n, is n = α + β The units of –r A, are always in terms of concentration per unit time while the units of the specific reaction rate, k A, will vary with the order of the reaction.