KINETICS

1. MATERIALS SCIENCE & ENGINEERING Part of A Learner’s Guide AN INTRODUCTORY E-BOOK Anandh Subramaniam & Kantesh Balani Materials Science and Engineering (MSE) Indian Institute of Technology, Kanpur- 208016 Email:anandh@iitk.ac.in, URL:home.iitk.ac.in/~anandh http://home.iitk.ac.in/~anandh/E-book.htm KINETICS "Thermodynamics Warrants Kinetics Prohibits"

2. Kinetics • In the chapter on Equilibrium we had seen that the thermodynamic feasibility of processes is dictated by ‘Thermodynamic Potentials’ (e.g. Gibbs Free Energy, at constant T, P, Ni). • If (say) the Gibbs Free Energy for a process is negative then the process CAN take place spontaneously. • However, IF the process WILL actually take place (and if it will take place- how long will it take to occur?) → will be determined by the ‘Kinetics of the process’. • Deeper the ‘metastable energy well’, higher will be activation energy$ required to pull the system out of the ‘well’ and take it to the equilibrium state (or some other metastable state). • To give an example: the window pane glass is in a metastable state* → there is a tendency for it to crystallize and to lower the Gibbs Free Energy of the system. However, at room temperature the crystallization is very slow and the glass pane can remain amorphous** for hundreds of years. Case of ‘Thermodynamics warrants, Kinetics delays’ • For a given process to occur heat and mass transfer may have to take place and this would take time → hence in ‘kinetics’ we deal with time and rates (1/t) $ defined in an upcoming slide (must have read it in school as well) * All glasses are considered to be metastable and there exists at least one crystalline state with a lower G. ** The terms glass and amorphous material are more often used synonymously

3. Some basics • A homogenous reaction is one which involves only one phase. E.g. a reaction involving only gaseous phase • In a heterogeneous reaction more than one phase is present. Let us consider a homogenous balanced chemical reaction, occurring in a single step: Rate of consumption of a reactant is proportional to the stoichiometric coefficient in a balanced reaction • nA→ number of moles of A present at time t • J → Rate of conversion • r → rate of reaction (= J/V) Rate of Conversion (J) is defined as: A, B are being consumed and hence dn/dt for these species is negative and J is positive J depends on system size and is an extensive quantity, the conversion rate per unit volume (J/V) is the reaction rate is an intensive quantity The reaction rate (r) is a function of P, T and the concentration of species

4. In many of the reactions the volume is constant. If the volume is constant during the reaction: • nA→ number of moles of A present at time t • J → Rate of conversion • r → Rate of reaction (= J/V) • [A] → Molar concentration of A (= cA) For many reactions it is seen that the rate can be related to the concentration of species by a reaction of the form: Partial orders: Integer or half integer • The rate constant k is a function of T and P The pressure dependence is usually small • The exponents: ,  are usually integers or half integers: (1, ½, 3/2, …) and are the partial orders (i.e. the reaction has got an order  wrt to A and  wrt B •  +  = n is the order of the reaction (overall order) • Units of k → [concentration]1n [t]1 Order wrt A Order wrt B Rate Constant = f(T,P) Molar Concentrations

5. Some basics • Gas phase reaction • Overall Order as expected = 2 • Gas phase reaction • Partial and Overall Order as expected • Overall order = 3 The above are cases which are ‘intuitively’ easy to correlate with the concept of order. But, reactions may have partial order as below or even the cocept of order of a reaction may not apply! • Gas phase reaction • Partial overall order • Gas phase reaction • Catalyst NO appears in the order while reactant SO2 does not • Gas phase reaction • Concept of order does not apply

6. Looking at some of the examples in the previous slide it is clear that the exponents in the rate law can be different from the numerical coefficients in the balanced chemical reaction equation. • Rate laws are to be determined from measurement of reaction rates and cannot be determined from reaction stoichiometry. • Additionally, the use of concentrations in the rate equations is valid only for ideal systems.

7. Rate constants depend strogely on temperature (usually increasing rapidly with increasing T). • For many reactions in solution, a thumb rule can be used that near room temperature ‘k’ doubles or triples every 10C increase in temperature. • In 1889 Arrhenius noted that for many reactions k = f(T) fits an exponential function. Arrhenius equation Activation EnergyAffected by catalyst • A → pre-exponential factor [units of k] • Q → activation energy [J/mole] • R → gas constant T in Kelvin Frequency factor A is a term which includes factors like the frequency of collisions and their orientation. It varies slightly with temperature, although not much. It is often taken as constant across small temperature ranges.

8. Fraction of species having energyhigher than Q (statistical result) ln (Rate) → 0 K

9. Reactants Products A + BC AB + C A + BC (ABC)* AB + C Activated complex

10. Activated complex (ABC)* Preferable to use G ΔH Energy A + BC AB + C Configuration

11. The average thermal energy is insufficient to surmount the activation barrier (~ 1eV) • The average thermal energy of any mode reaches 1eV at ~ 12000 K • But reactions occur at much lower temperatures • Fraction of species with energies above the activation barrier make it possible • Lost species by reaction are made up by making up the distribution • Rate  fraction of species with sufficient energy Rate  vibrational frequency (determines the final step) • Rate  n