Homogeneous reactions e.g (when reactants and products are in same phases) H 2 (g)+ Cl 2 (g) 2HCl (g) or 2NO(g) +O 2 (g) 2NO 2 (g)

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1 1. CHEMICAL KINETICS: The branch of chemistry that deals with the study of rate of reaction and their mechanisms. Significance to know: the time required to complete the reaction, understand mechanism of reaction, to know optimum conditions for a maximum yield of the product. Example of fast reaction: NaOH + HCl NaCl + H 2 O Example of slow reaction :rusting of iron Oxidation : Fe(s) Fe 2+ (aq) + 2e- Reduction : O 2 (g) + 4H + (aq) + 4e- 2H 2 O(l) Atmosphere oxidation: 2Fe 2+ (aq) + 2H 2 O(l) + ½ O 2 (g) Fe 2 O 3 (s) 4H + (aq) Homogeneous reactions e.g (when reactants and products are in same phases) H 2 (g)+ Cl 2 (g) 2HCl (g) or 2NO(g) +O 2 (g) 2NO 2 (g) Hetrogeneous reactions: e.g (when reactants and products are in different phases) 2Mg (s) + O 2 2MgO (s) 2. RATE OF REACTION : Change in concentration of reactant (or product) / time taken for the change. Rate of reaction: -ΔC Δt Where C is concentration, t is time, Δ is the change from initial point to final. Note: concentration of reactant decreases w.r.t time, therefore ve, while of products increase with time. 2A+ B C Av. Rate of reaction = -dc/dt = -1/2d(rate of disappearance of A)/dt = -½ d[a]/dt = -db/dt = -d(rate of disappearance of B)dt = -d[b]/dt = +dc/dt = +(rate of appearance of C) = +d[c]/dt Rate = d(concentration )/dt. Unit of rate of reaction is (mol L -1 s -1 ) aa+ bb products for a gas phase reaction : C= P/RT where partial pressure = mol fraction (moles of gas/ total moles) x P t (total pressure) Law of mass action Rate [A] a x [B] b Rate = k [A] a [B] b where k is rate constant or velocity constant or specific reaction rate a and b are stoichiometric coefficient. 3. AVERAGE RATE: The rate of reaction measured over a long time interval is called average rate of reaction. It is equal to (x 2 -x 1 )/(t 2 -t 1 ) in concentration time graph. Average rate of reaction decreases w.r.t time as concentration of reactants decreases. 4. INSTANTANEOUS RATE: It is the rate of reaction at a particular moment of time. It is equal to dx/dt. Average rate of reaction approaches instantaneous rate when Δt 0 i.e. becomes infinitesimally small. Instantaneous rate of reaction by slope of tangent of a concentration time graph. 1

2 Slope is ve for reactants as reactants decreases wrt time while +ve for products. 5.RATE LAW OR RATE EQUATION: expression is rate C (concentration). where C is concentration of reactants or products. Rate =k [C] order, [] means concentration in moles/ litre. k is known as rate constant. aa+ bb products rate law equation = k[a] p [B] q where p and q are experimentally determined Difference between rate law and law of mass action Law of mass action For a reaction aa+ bb products rate = k [A] a [B] b where k is rate constant or velocity constant a and b are stoichiometric coefficient. For Simple reactions rate law = law of mass action Rate law rate law equation = k[a] p [B] q where p and q are experimentally determined and p+ q = order of a reaction. Complex reactions Rate law may not be equal to law of mass action 6. RATE CONSTANT: k, When the concentration of reactants is unity, then the rate of reaction is known as rate constant. It is also called specific reaction rate. Units of k dx/dt = rate = k [conc] n n is order of reaction k = [conc] 1-n s -1 for zero order moll -1 s -1,1 st order reaction s -1, 2 nd order, Lmol -1 s -1, 3 rd order L 2 mol -2 s -1 or k = (atm) 1-n s -1 Rate of reaction Rate of reaction depends upon the concentration. rate = k [conc] n n is order of reaction Rate of reaction depends upon concentration, T, surface area, reactant states, etc Reaction rate constant It is constant of proportionality in the rate of law equation and equal to rate if concentration of reactants are unity. Rate = k [1][1] k depends upon temperature by Arrhenius equation Units mol L -1 s -1 Units of k = (conc ) 1-n s -1 Except zero order reaction, it goes on decreasing wrt time 7. FACTORS AFFECTING RATE OF REACTION: It remains constant for a given reaction at a given temperature. (a) concentration of reacting species: rate concentration of reactants. Rate of reaction increases with increases in concentration of reactants. (b) nature of reactants: rate of reaction gas> liquids> solids. Ionic reactions are fast as compared to non ionic (molecular) reactions.recations involving cleavage of lesser number of bonds of reactants proceeds faster than cleavage of more number of bonds of reactants. E.g. oxidation of Fe +2 to Fe +3 by acidified KMnO 4 takes place more readily than oxidation of oxalate ion (C 2 O 4 2- ) to CO 2 by same reagents. 2

3 (c) Temperature: rate of reaction Temperature. Increase in 10 0 C temperature, doubles the rate of reaction. Increase in temperature increases the collision frequency and effective collisions. Colliding atoms, ion, molecule must have certain minimum amount of energy (called threshold energy) and proper orientation at the time of collisions. Effective collision is much more important than collision frequency. THRESHOLD ENERGY: the minimum amount of energy which must be associated with colliding species so that their mutual collisions may bring about a chemical reaction is known as threshold energy. (d) Catalyst: Catalyst is a substance which alters the rate of reaction, however its composition and mass remains unchanged at the end of reaction. (i) catalyst may be consumed in one step and is regenerated in any one of the subsequent steps. (ii) catalyst provides a new reaction path with lower energy barrier and thus activation energy gets lowered. As a result, energy barrier and orientation barrier are overcome by more fraction of molecules, thereby increasing the rate of reaction. (e) surface area of reactants in heterogeneous system: in a heterogeneous system, smaller particles react more readily than bigger particles of the same mass; wood chips burn more readily than log of wood of same mass. (f) Photo chemical reactions: the rate of reaction increases by absorption of radiations. Such reactions are known as photo chemical reactions. E.g. H 2 and Cl 2 reacts more readily in UV radiations than in dark. MOLECULARITY of a reaction: Total number of atoms, ions or molecules of the reactants involved in the balanced chemical reaction is termed as its molecularity. It is always in whole number. Molecularity of >3 is rare because chances of simultaneous collisions between more than 3 are rare. It cannot be zero. Always a whole no. PCl 5 PCl 3 + Cl 2 unimolecular 2HI H 2 + I 2 CH 3 COOC 2 H 5 + H 2 O CH 3 COOH + C 2 H 5 OH Molecularity 2, Order 1 8. ORDER OF A REACTION: The sum of the power of the concentration of reactants in the rate law equation. It can be in fraction. It can be zero also. rate = k [A] n [B] m order = m+n order of reaction is experimentally. Cannot be written from balanced chemical reaction, molecularity and order may be same of different. order of reaction can change with T,P, while e.g. of zero order reaction: H 2 + Cl 2 (h ) 2HCl (g) photo chemical reactions 2NH 3 (Mo or W) N 2 + 3H 2 rate = k for zero order reaction. H 2 O 2 slow H 2 O+ O, O+ O fast O 2 slow step is rate determining order =1 Although reaction is 2H 2 O 2 2H 2 O + O 2 2N 2 O 5 4NO 2 +O 2 order =1 k units s -1 (COOH) 2 CO+ CO 2 + H 2 O order =1 2N 2 O 5 4NO 2 + O 2 order=1 Hydrogenation of ethene order = 1 Hydrolysis of ester =1 = inversion of sugar 3

4 2HI H 2 + I 2 2nd order reaction 2NO 2 (g) + F 2 (g) 2NO 2 F (g) order =2 Saponification CH 3 COOC 2 H 5 + NaOH CH 3 COONa + C 2 H 5 OH 2 nd order 2NO + O 2 2NO 2 4KClO 3 3KClO 4 + KCl Zero order reaction: where reaction rate is not affected by concentration of reactants. i.e. rate =k (conc) 0 = k = mol L -1 s -1 H 2 (g) + Cl 2 (g) 2HCl(g) Fractional order : H 2 (g) H 2 (g) para hydrogen to ortho hydrogen at high T rate =k[p H2 ] 1.5 Negative order reaction: 2O 3 3O 2 ½ d(o 3 )dt =k d(o 3 ) 2 /[O 2 ] =k [O 3 ] 2 [O 2 ] -1 order wrt O 3 is 2 while wrt O 2 is -1, overall order 1 2HI H 2 (g) + I 2 (g) rate = k[hi] 2 CH 3 COOH + NaOH CH 3 COONa + H 2 O order wrt CH 3 COOH =1 and wrt NaOH =1 overall order =1+1 =2, units of k = (mol l -1 ) -1 s -1 Rate = k [CH 3 COOH] 1 [NaOH] 1 Molecularity It is the total no. of atoms, ions or molecule which must collide at a time to form products Theoretical determined Molecularity is a whole number Can never be zero. Does not give any clue about mechanism of reaction Order of reaction It is the sum of powers to which concentration terms of the reactants in the rate law equation Experimental determined Order is a whole number or a fractional value. Can be zero Reaction is based on overall reaction. 9. HALF LIFE OF A REACTION: Time required to convert the original concentration of reactant to half. For first order reaction, t 1/2 = / k, where k is rate constant.. Half life is independent of initial concentration for first order reaction for 1 st order reaction. For nth order reaction t 1/2 1/ a n-1 No.o f half life periods = total time / (time of one half periods) Quantity left after n half periods= (½) n x initial quantity of reactants. [A] = [A 0 ] (1/2) n 10.PSEUDO MOLECULAR REACTIONS: A first order reaction which appears to be of higher order is known as pseudo first order reaction. i.e. where molecularity and order does not match. C 12 H 22 O 11 + H 2 O (excess H + ) C 6 H 12 O 6 glucose + C 6 H 12 O 6 fructose molecularity 2, order 1 rate = k[c 12 H 22 O 11 ] Hydrolysis of ester: RCOOR + H 2 O (excess H + ) RCOOH + ROH rate = k[ch 3 COOCH 3 ] When one of the reactants is in large excess as compared to other reactant, second order reaction conforms to 1 st order as slow step determine the rate of reaction Pseudo first order reaction. 4

5 Reaction which is bimolecular but order is one is called pseudo first order reaction. CH 3 COOH + H 2 O (H + ) CH 3 COOH + CH 3 OH 11. ACTIVATION ENERGY: (E a ): It is extra energy over and above the average energy of reactants which must be supplied to the reactant molecules so that collision between reactant molecules is effective and leads to the formation of product molecules. Threshold energy: it is the minimum amount of energy required by a reactant species to break the bond to form products. For effective collision energy barrier must be crossed by the reactant molecules to form products. Also reactant must react with proper orientation. Lowers the activation energy, faster is the rate of reaction. 12. ARRHENIOUS EQUATION: of reaction rate: It gives the relation between reaction rate constant and temperature. k is rate constant, increases wrt Temperature. For every 10 0 C rise in temperature rate of reaction (k) gets doubled. Temperature coefficient = k at t+10 0 C k at t 0 C Arrhenius Equation: k= Ae -Ea/RT where k = rate constant, A = frequency factor, E a = energy of activation (it is the extra amount of energy required by molecule to overcome the energy barrier) term e -Ea/RT also known as Boltzman factor. R = gas constant, T = temperature in Kelvin, A is related with frequency factor or pre exponential factor. Which is related to frequency of collisions and orientation of reacting molecules. ln k = ln A - E a /RT or log k = log A - E a /2.303RT R =8.314 J/gmol K where T in kelvin plot of log k wrt. 1/T is a straight line whose slope is equal to E a /2.303R for two different temperature log(k 2 /k 1 ) = Ea/2.303R (1/T 1-1/T 2 ) T in Kelvin only 13. MECHANISM OF A REACTION: The sequence of elementary processes leading to the overall stoichiometric of a chemical reaction is known as mechanism of a reaction. It is actually very difficult to determine the mechanism of a reaction experimentally, however, the rate equation can be determined experimentally. The slowest step is the rate determining step and helps in determining the mechanism of reaction. NO 2 (g) + CO(g) NO(g) + CO 2 (g) The rate equation below 500K is rate = k[no 2 ] 2 Thus it is a second order reaction and the proposed mechanism is NO 2 + NO 2 NO+ NO 3 (slow step ) NO 3 + CO NO + CO 2 (fast step) NO 2 + CO NO + CO 2 14.ACTIVATED COMPLEX: It is defined as unstable intermediate formed between reacting molecules. It is highly unstable and readily changes into product. 5

6 15. RATE DETERMINING STEP: The slowest step in the reaction mechanism is called rate determining step. 16. COLLISION FREQUENCY: The number of collisions per second per unit volume of the reaction mixture is known as collision frequency (Z). Collision theory state that rate of reaction = -dx/ dt = f x z Where f is the fraction of effective collision (proper orientation between reactants and having sufficient energy), z is the frequency of collision. dx/ dt is the change in concentration wrt time. rate of reaction = PZ AB e -Ea/RT Where: Z AB represents the collision frequency of reactants, A and B represents the fraction of molecules with energies equal to or greater than E a P is called the probability or steric factor 17. CATALYST: Substance which increases or decreases the rate of reaction without getting consumed in the reaction. Reactants on which catalyst operates are called substrates. 2KClO 3 (s) (MnO 2 ) 2KCl(s) + 3O 2 (g) N 2 (g) + 3H 2 (g) (Fe 2 O 3 (s)) 2NH 3 (g) Haber Bosch Process 2SO 2 (g) + O 2 (g) (Pt (s) or V 2 O 5 ) 2SO 3 (g) Contact process 2SO 2 (g) + O 2 (g) (NO(g)) 2SO 3 (g) CH 2 =CH 2 + H 2 (g) (Ni) CH 3 -CH 3 (g) Hydrogenation reaction C 12 H 22 O 11 + H 2 O (H 2 SO 4 ) C 6 H 12 O 6 (glucose) + C 6 H 12 O 6 (fructose) Catalyst which decreases the rate of reaction is ve catalyst or retarder or inhibitor. H 3 PO 4 retards the decomposition of H 2 O 2, alcohol slows down the oxidation of chloroform to phosgene. 18. HETROGENEOUS CATALYST: when reactants and catalyst are in different phase. CH 2 =CH 2 (g) + H 2 (g) (Ni(s)) CH 3 -CH 3 (g) hydrogenation of oils. (1 st order reaction) 19. CHARACTERISTICS OF A CATALYST: (a) unchanged chemically during/ end of the reaction (b)small quantity of catalyst is required (c) do not alter the position of equilibrium in reversible reaction. (d) they cannot make impossible reaction to occur. (e) Every catalyst has its optimum temperature at which it s efficiency is maximum. (f) it decreases the activation energy of the reaction and hence increases the rate of reaction. (g) catalyst are very specific in respect of reactions. E.g. HCOOH (Cu or ZnO) H 2 +CO HCOOH (Al 2 O 3 or TiO 2 ) H 2 O + CO (h) A catalyst does not alter the free energy change (ΔG) of the reaction, A catalyst does not alter the enthalpy change (ΔH) of the reaction. (i) for reversible reaction, it accelerated the forward and backward reaction to same extent. Hence does not disturb the equilibrium. i.e. does not change the equilibrium constant K but simply helps to attain the equilibrium quickly. 6

7 20. CATALYTIC PROMOTER: Those substance which are not catalysts but when present with some catalysts greatly enhances their action. e.g. N 2 (g) + 3H 2 (g) (Fe 2 O 3 (s)+ Mo catalytic promoter) 2NH 3 (g) Haber Bosch Process 21. CATALYST POISON or ANTI CATALYSTS: those substance which decreases or inhibit the activity of catalysts. E.g. oxides of arsenic poison the catalytic power of platinum in the contact process. 2SO 2 (g) + O 2 (g) (Pt (s) or V 2 O 5 ) 2SO 3 (g) contact process Pd/ BaSO 4, S act as poison for Pd catalyst in Rosenmund reaction RCOCl + H 2 Pd, BaSO 4, S RCHO +HCl. 22. AUTO CATALYSIS: some reactions are slow during start, but reaction rate becomes fast when products are formed as one of products act as catalysis. KMnO 4 solution addition with Oxalic acid, disappearance of pink colour is slow at start but rapid after formation of Mn +2 ions. 2 MnO C 2 O H + 2Mn CO 2 + 8H 2 O 23. ENZYMES: Catalyst found in living tissues are called biological catalysts called enzymes. They are highly sensitive to temperature. e.g. curd will not form at low temperature. Catalytic activity of enzyme is due to their capacity to lower the E a. Source Enzymes Saliva Salivary Amylas (ptyalin) and mucin Gastric juice pepsin, rennin and gastric lipase Pancreatic juice trypsin, chymotrypsin, steapsin. C 12 H 22 O 11 sucrose + H 2 O (invertase) C 6 H 12 O 6 (glucose) + C 6 H 12 O 6 (fructose) C 6 H 12 O 6 (glucose) Zymase 2C 2 H 5 OH + 2CO 2 (g) Process /reaction Catalyst Haber s process N 2 (g) + 3H 2 (g) (Fe 2 O 3 (s)) 2NH 3 Iron, Mo (promoter) (g) Contact process 2SO 2 (g) + O 2 (g) (Pt (s) or V 2 O 5 ) Pt, V 2 O 5 2SO 3 (g) Ostwald process for mfg of HNO 3 Pt gauze Deacon process for mfg of Cl 2 CuCl 2 Bosch process for mfg of H 2, CO +H 2 O CO 2 + H 2 Fe 2 O 3 + Cr 2 O 3 Mfg of vegetable ghee Ni. Fischer Tropsch process (synthesis of H/c) from H 2 Co, ThO 2 and CO Synthetic petrol from coal (Bergius process) Iron oxalate Mfg of C 2 H 5 OH from sucrose Invertase, Zymase Mfg of C 2 H 5 OH from starch Diastase, maltase Pthalic acid from naphthalene HgSO 4 Polymerisation of alkene H 3 PO 4 on kieselguhr Methanol from H 2 and CO ZnO and Cr 2 O 3 Dehydrogenation of alcohol to aldehyde Cu 573K Acetic acid from aldehyde by oxidation with air V 2 O 5 Preparation of O 2 from KClO 3 MnO 2 Zeigler Natta catalyst TiCl 4 + (C 2 H 5 ) 3 Al HDPE 7

8 Formulas: 1. Instantaneous rate = dx/ dt, where dx is small change in conc. and dt is the smallest interval of time. x is concentration in moles/lt for liquids for reactions in gas phase PV = nrt n/ V = C = P/RT 2. Average rate =x 2 -x 1 /t 2 -t 1, where is change in concentration and is large interval of time. 3. A + B C + D Rate of disappearance of, where d[a] is small change in conc. of A and dt is small interval of time Rate of disappearance of B = -d[b]/dt, Rate of appearance of C = +d[c]/dt Rate of appearance of D = + d[d]/dt Unit of rate of reaction for 2 nd order reaction = mol L -1 s Order of reaction: If rate law expression for a reaction is Rate = k [A] x [B] y Order of reaction = x + y Remember: Order cannot be determined with a given balanced chemical equation but molecularity can. It can be experimentally determined. 5. Integrated rate law for zero order reaction: R P dx/dt = k[r] 0 Units of k is moll -1 s -1 k= [C 0 ]-[C] t. If we plot a graph between concentration of C v/s time, the graph is a straight line with slope equal to -k and intercept is equal to [C o ] 6. Half- life reaction for a for zero order reaction: t ½ = [C 0 ] 2k 7. Rate law for Ist order reaction: t = {log [Co ]} where k is rate constant or specific reaction rate, k C [Co] is initial molar conc. in moles/ lt, [C] is final molar conc. in mol/lt after time t If we plot a graph between ln[co/c] with time, we get a straight line whose slope k and intercept ln[a / a-x ] where a is initial conc. in mol L -1, x moll -1 have reacted in time t final conc., after time t is (a x) 8. Half- Life for a first order reaction is: t ½ =0.693 / k t 1/2 is independent of initial concentration of the reactants. Half life for nth order reaction t ½ (conc) 1-n 9. Formula to calculate rate constant for first order gas phase reaction of the type A(g) B(g) + C(g) Where: p i is initial pressure of A 8

9 P t is total pressure of gaseous mixture containing A, B, C Remember: P t = p A + p B + p C 10. Arrhenius equation: k = Ae -Ea/RT 11. log(k 2 /k 1 ) = E a [T 2 -T 1 ] 2.303R T 1 T rate = PZ AB e -Ea/RT Where: ZAB represents the collision frequency of reactants, A and B represents the fraction of molecules with energies equal to or greater than E a and P is called the probability or steric factor 13. Av time period = 1/ t 1/2 units sec for 1st order. 14. amount of reactant left after n half lives = [A 0 ] times 2 n 9

C H E M I C N E S C I

C H E M I C N E S C I C H E M I C A L K I N E T S C I 4. Chemical Kinetics Introduction Average and instantaneous Rate of a reaction Express the rate of a reaction in terms of change in concentration Elementary and Complex