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Types of Chemical Reactions Rate of a Reaction Factors Affecting Rate of A B. Sc. II-Sem Rate of formation of a product is positive d[C] Rate of formation of C dt Chemical Kinetics d[D] Rate of formation of D The branch of physical chemistry which deals with the rate at dt which the chemical reactions occur, the mechanism by which the In terms of stoichiometric coefficient rate may be chemical reactions take place and the influence of various factors expressed as dx 1 d[A] 1 d[B] 1 d[C] 1 d[D] such as concentration, temperature, pressure, catalyst etc., on the reaction rates is called the chemical kinetics. dt a dt b dt c dt d dt The rate of reaction is always positive. Types of chemical reactions The rate of chemical reaction decreases as the reaction On the basis of reaction rates, the chemical reactions have proceeds. been classified into the following three types, Unit of conc. Unit of rate of a reaction = =mole L–1 time –1 (1) Very fast or instantaneous reactions: These reactions Unit of time occur at a very fast rate generally these reactions involve ionic In term of gaseous reaction the unit is atm time-1 and species and known as ionic reactions. It is almost impossible to Rate in atm time-1= Rate in mole L1time 1 RT determine the rates of these reactions. Examples (i) AgNO NaCl AgCl NaNO (Precipitation reaction) 3 3 Conc. of product (PPt.) (ii) HCl NaOH NaCl H 2 O (Neutralization reaction) (acid) (base) (Salt) (2) Moderate reaction: These reactions proceed with a measurable rates at normal temperature and it is these reactions (mole/lit.) Conc. are studied in chemical kinetics. Mostly these reactions are Conc. of reactant molecular in nature. Time (Sec) Examples Change of concentration (i) Decomposition of H 2O2 : 2H 2O2 2H 2O O2 with increase in time (ii) Decomposition of N 2O5 : 2N 2O5 2N 2O4 O2 (3) Very slow reactions: These reactions are extremely Factors affecting rate of a reaction slow and take months together to show any measurable change. The rate of a chemical reaction depends on the following things Examples (1) Nature of reactants (i) Rusting of iron: Fe2O3 xH 2O Fe2O3 . xH 2O Hydrated ferric oxide (Rust) (i) Physical state of reactants: This has considerable effect over rate of reaction. Room temperatu re (ii) 2H 2 O2 2H 2O Gaseous satae Liquid state Solid state Rate of a reaction Decreasing rate of reaction The rate (speed or velocity) of a reaction is the change in (ii) Physical size of the reactants: Among the solids, rate concentration in per unit time. increases with decrease in particle size of the solid. x dx x x (iii) Chemical nature of the reactants 2 1 or t dt t2 t1 (a) Reactions involving polar and ionic substances including the proton transfer reactions are usually very fast. On where x or dx is the concentration change, i.e., (x x ) 2 1 the other hand, the reaction in which bonds is rearranged, or in the time interval t or dt, i.e., (t t ). 2 1 electrons transferred are slow. –1 Concentration is generally expressed in active mass, i.e., mole L (b) Oxidation-reduction reactions, which involve The rate measured over a long time interval is called transfer of electrons, are also slow as compared to the ionic average rate and the rate measured for an infinitesimally small substance. time interval is called instantaneous rate and (c) Substitution reactions are relatively much slower. Instantaneous rate (Average rate)t0 For the reaction aA bB cC dD (2) Effect of temperature: The rate of chemical reaction generally increases on increasing the temperature. The rate of a Rate of disappearance of a reactant is negative o d[A] reaction becomes almost double or tripled for every 10 C rise in Rate of disappearance of A dt temperature. d[B] Temperature coefficient of a reaction is defined as the ratio Rate of disappearance of B dt of rate constants at two temperatures differing by (generally 25°C and 35°C) 10°C. o k The value of rate constant depends on, nature of k at (t 10 C) 35 o C Temperatur e coefficient o reactant, temperature and catalyst. It is independent of k at t C k o 25 C concentration of the reactants. (3) Concentration of reactants: The rate of a chemical n1 1n litre 1 mol 1 reaction is directly proportional to the concentration of the Unit of rate constant sec sec reactants means rate of reaction decreases with decrease in mol litre concentration. Where n order of reaction. (4) Presence of catalyst: The function of a catalyst is to Rate law: Molecularity and Order of a reaction lower down the activation energy. The greater the decrease in the Molecularity is the sum of the number of molecules of activation energy caused by the catalyst, higher will be the reactants involved in the balanced chemical equation. Molecularity reaction rate. Reaction path of a complete reaction has no significance and overall kinetics of Without catalyst the reaction depends upon the rate determining step. Slowest step is the rate-determining step. This was proposed by Van't Hoff. Ea Example : NH 4 NO 2 N 2 2H 2 O (Unimolecular) Ea Reaction path with catalyst NO O3 NO 2 O 2 (Bimolecular) Reactants 2NO O 2NO (Trimolecular) Energy of Reaction 2 2 The total number of molecules or atoms whose concentration Products determine the rate of reaction is known as order of reaction. Energy Potential A catalyst changes the reaction path Order of reaction = Sum of exponents of the conc. terms in rate law (5) Effect of sunlight : There are many chemical reactions For the reaction xA yB Products whose rate are influenced by radiations particularly by ultraviolet The rate law is Rate [A]x[B[y and visible light. Such reactions are called photochemical reactions. For example, Photosynthesis, Photography, Blue Then the overall order of reaction. n x y printing, Photochemical synthesis of compounds etc. where x and y are the orders with respect to individual reactants. The radiant energy initiates the chemical reaction by If reaction is in the form of reaction mechanism then the supplying the necessary activation energy required for the order is determined by the slowest step of mechanism. reaction. 2A 3B A2 B3 Law of mass action and Rate constant A B AB(fast) The rate at which a substance reacts is directly proportional AB B AB (slow) (Rate determining step) to its active mass and the rate at which a reaction proceeds is 2 3 proportional to the product of the active masses of the reacting AB3 A A2 B3 (fast) substances. (Here, the overall order of reaction is equal to two.) For a reaction, aA bB product Molecularity of a reaction is derived from the mechanism of dx a b dx a b the given reaction. Molecularity can not be greater than three because Rate [A] [B] ; k[A] [B] dt dt more than three molecules may not mutually collide with each other. Where k is called rate constant or velocity constant. Molecularity of a reaction can't be zero, negative or dx fractional. order of a reaction may be zero, negative, positive or in When [A] [B] 1 mol / litre , then k dt fraction and greater than three. Infinite and imaginary values are Thus, rate constant k is also called specific reaction rate. not possible. When one of the reactants is present in the large excess, the second order reaction conforms to the first order and is known as pesudo unimolecular reaction. (Table 11.1) Table : 11.1 Order and molecularity of some reaction S. Chemical equation Molecularity Rate law Order w.r.t. No. First Second reactant Overall reactant 1. aA bB product a + b dx a b a + b k[A]a[B]b dt 2. aA bB product a + b dx 2 0 2 zero, if B is in 2 k[A] [B] excess dt 3. Pt, 2 1* ----- 1 2H 2O2 2H 2O O2 dx (Bimolecular) k[H 2O2] dt 4. H 2 1* Zero, if H2O is in 1 CH COOC H H O dx 3 2 5 2 (Bimolecular) k[CH 3COOC 2 H5 ] excess dt CH 3COOH C2 H5OH 5. H 2 dx 1* Zero, if H2O is in 1 C12 H 22 O11 H 2 O (Bimolecular) k[C12 H 22O11 ] excess Sucrose dt C6 H12O6 C6 H12O6 Glucose Fructose 6. 2 1* Zero, if OH– does 1 (CH 3 )3 CCl OH dx (Bimolecular) k[(CH 3 )3 CCl] not take part in dt (CH 3 )3 COH Cl slow step 7. 2 dx 1 1 2 CH 3Cl OH CH OH Cl 3 (Bimolecular) k[CH 3Cl][OH ] dt 8. 1 1 ---- 1 C6 H5 N 2Cl C6 H5Cl N 2 dx (Unimolecular k[C6 H5 N 2Cl] dt ) 9. 1 1.5 ---- 1.5 CH 3CHO CH 4 CO dx 3 / 2 (Unimolecular k[CH 3CHO] dt ) 10. 5 1 1 2 H 2O2 2I 2H 2H 2O I2 dx k[H 2O2 ][I ] (H+is medium) dt 11. 2O3 3O2 2 dx 2 1 -1 with respect to 1 (Bimolecular) k[O3 ] [O2 ] O2 dt *Pseudo-unimolecular reactions. Table : 11.2 Rate constant and other parameters of different order reactions Order Rate constant Unit of rate Effect on rate by changing (Half-life period) constant conc. to m times T50= 0 x conc. time–1 No change a k0 (mol L–1 s–1) t 2k0 1 time–1 (s–1) m times 0.693 2.303 a k1t k1 log10 , C C0e t a x k1 k1t 2.303 (a x1) N N0e , k1 log10 (t2 t1) (a x 2 ) 2 1 1 1 x conc–1 time–1 m2 times 1 k2 (for the case (mol L–1) s–1 t (a x) a ta(a x) k 2 a L mol–1 s–1 when each reactant has equal concentration) 2.303 b(a x) k2 log10 (for the case when t(a b) a(b x) both reactants have different concentration) 3 1 1 1 conc–2 time–1 m3 times 3 k3 2 2 (mol L–1)–2 s–1 2 2t (a x) a 2k3a L2 mol–2 s–1 n 1 1 1 conc(1–n) time–1 mn times 2n1 1 kn n1 n1 ; n 2 (mol L–1)(1–n) s–1 n1 (n 1)t (a x) (a) (n 1)kn(a) L(n–1) mol(1–n) s–1 1 1 1 Methods for determination of order of a reaction k (For second order reactions) t a a x (1) Integration method (Hit and Trial method) (i) The method can be used with various sets of a, x and t 1 1 1 k (For third order reactions) 2 2 with integrated rate equations.
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