Chemical Kinetics Notes

CHEMICAL KINETICS CHEMICAL KINETICS Chemical Kinetics: (Study of reaction rates and their mechanisms. ) Rate of Reaction : The change in concentration of reactant or product per unit time .(SI unit molL–1s– 1 ). Let a reaction : aA + bB  cC + dD. Average rate of reaction: The rate of reaction measured over a long period of time. Rav =  1 A   1 B 1 C 1 D    a t b t c t d t Instantaneous rate of reaction: The rate of reaction at a particular instant of time i.e.(T approaches zero) Rinst =  1 dA   1 dB 1 dC  1 dD    a dt b bt c dt d dt Factors affecting rate of reaction: 1. Rate of reaction  Conc. of reactants 2.  Temperature 3.  Surface area of reactants 4. 1  Bond dissociation energy of reactants 5. 6. Increases in the presence of catalyst Depends on pressure in case of gases. Rate law or rate equation: The mathematical expression which denotes the experimentally observed rate of a reaction in terms of the concentration of the reactants which influences the rate of reaction. Rate = k[A]x[B]y k = rate constant and x  a , y  b. Order of a reaction : The sum of powers of the concentration of the reactants in the rate law expression. Order = x + y eq. Rate = k[A]1/2 [B]3/2 so order = * (k) 1 3  =2 2 2 Order of a reaction can be 0, 1, 2, 3 and even a fraction . Rate constant : It is defined as the rate of reaction, when the conc. of each of the reactants is unity. The unit of rate constant for nth order = (molL–1)l–ns–1 Reaction n Units of rate constant zero order 0 mol L–1s–1 1st Order 1 s–1 nd 2 Order 2 mol–1L s–1 CHEMICAL KINETICS Molecularity of a reaction : The number of reacting species which collide simultaneously to bring about a chemical reaction. Molecularity = a + b eg. NH4NO2 N2 + 2H2O 2HI  H2 + I2 2NO + O2  2NO2 (Unimolecular reaction) (Bimolecular reaction) (Tri or termolecular) Elementary reactions : The reactions taking place in one step Complex reactions : When a sequence of elementary reactions gives us a products, the reactions are called complex reactions. * Molecularity is applicable only for elementary reactions. For complex reactions it has no meaning * For complex reactions, order is given by the slowest step (i.e. rate determining step) and generally, molecularity of the slowest step is same as the order of the over all reaction. Integerted rate equations: Zero order reactions: Let a reaction : So. Rate Thus =  dR  k Ro dt Ro R k eg. 2NH3(g) R  P. (Product) t Pt 1130 K [R]o = Initial conc. of reactant [R] = Conc. of reactant after trums N2(g) + 3H2(g) Rate = k[NH3]º = k First order reaction: R P eg.1 So. Rate = Thus k  dR  kR dt Ro 2.303 log t R  C2H4(g) + H2(g) C2H6(g) Rate = k[C2H4] eg.2 All natural and artificial radioactive decay of unstable nuclei 226 88 Ra eg.3 22  24 He  86 Rn Decomposition of N2O5 and N2O. Half life of reaction (t1/2) : The time in which the concentration of a reactant is reduced to one half of its initial concentration. CHEMICAL KINETICS eg. For a Zero Order reaction t1/2: = eg.1 Hydrolysis of cane sugar. Ro 2k 2.303 0.693 eg. For a first order reaction : t1/2= log2 = k k Psendo First Order reaction : A reaction which behaves as a first order reaction even when there is more than one concentration involved. It is because the other concentration is in large excess or present as a catalyst for which the concentration does not change. H+ C12H22O11 + H2O C6H12O6 (excess) eg.2 t=0 t Thus. + C6H12O6 Glucose Fructose Hydrolysis of ester CH3COOC2H5 + H2O  CH3COOH + C2H5OH 0.01 mol 10 mol 0 mol 0 mol 0 mol 9.9 mol 0.01 mol 0.01 mol Rate = k'[H2O][CH3COOC2H5] [H2O] taken as constant as it does not altered much. Rate = k[CH3COOC2H5] Where k = k' [H2O] k  Ae Ea / RT Arrhenius equation : * The temperature dependence of the rate of chemical reaction can be accurately explained by Arrhenius equation A = Arrhenius factor or frequency factor R = Gas constant , Ea = Activation energy *It has been found that for a chemical reaction with rise in temperature by 10º, the rate constant is nearly doubled. Reaction H2(g) + I2(g)  2HI H H + I H I H I H I H + I I P.E C A H2 + I2 B Reaction Co-ordinate Diagram showing plot of P.E v/s reaction co-ordinate. CHEMICAL KINETICS The energy required to form this intermediate, called activated complex (C) is known as Activation Energy (Ea). Fraction of molecules t t +10 *Peak of the curve correspond to the most probable K.E. Kinetic energy * If rate constants of a reaction at T1 and T2 are K1 and K2, log K2 Ea  1 1     K1 2.303R  T1 T2  Collision Theory : By Max Trautz and william Lewis. According to it (a) For a chemical reaction to occur, there must be collisions between the reactants molecules. (b) Only a certain fraction of the total number of collisions are effective in forming the products. (c) For effective collisions, the molecules should have sufficient energy as well as orientation. A + B  Products. Rate  Z AB e Ea / RT ZAB = Collision frequency : No. of collisions per second per unit volume of the reaction mixture. Effective collistion : The collisions in which molecules collide with suffcient K.E. (Called threshold energy) and proper orientation, so as to facilitate breaking of bonds between reacting species and formation of new bonds to form products are called as effective collisions. Threshold Energy = Ea + energy possesed by reacting species. * To account for effective collisions another factor 'P' called steric factor was introduced. Rate = PZAB e–Ea/RT . H Improper H H H _ C –Br + OH Orientation H _ C –Br H Proper Orientation HO H OH H C No products H Br H – HO – C – H + Br H