Chemical Equilibrium I PDF

Chemical Equilibrium : Introduction :  Chemical Reaction : Spontaneous conversion of reactants into products under specific conditions.  Incomplete reactions : Reaction mixture consists of both reactants and products- all reactants not completely converted to products.  State o...

Chemical Equilibrium : Introduction :  Chemical Reaction : Spontaneous conversion of reactants into products under specific conditions.  Incomplete reactions : Reaction mixture consists of both reactants and products- all reactants not completely converted to products.  State of Equilibrium : Two processes occur simultaneously and a state reached when concentrations of both reactants and products become constant.  Reversible Reactions : Simultaneous forward and backward reactions.  State of Chemical Equilibrium : Rates of both reactions equal- concentrations of reactants and products constant. 717 K H2 (g) + I2 (g) 2HI (g)  Eg. PCl5 (s) PCl3 (s) + Cl2 (g) CaCO3 (s) CaO (s) + CO2 (g) Introduction :  Study of state of equilibrium- very important in study of chemical processes- helps us to know how far the reaction proceeds.  Equilibrium in Physical Processes (Physical Equilibria) :  State of equilibrium not only attained in chemical processes but also several physical processes.  Eg. Ice-Water equil. (Melting of ice); Water (l)- Water (g) equil. (boiling of water). Solid Liquid  Gas Three types : Liquid Solid Gas  Solid- Liquid Equilibrium :  Conversion of solid-liquid= melting- takes place at m.p.  Reverse process (conversion of liquid solid=freezing; occurs at freezing point).  For specific solid- liquid system, m.p.= freezing point.  M.p. (freezing point) : Temperature at which solid and liquid phases of a pure substance are at equilibrium at one atm pr. (Physical Equilibria) :  At m.p. of a solid, both solid and liq. phases exist in equilibrium. m.p. Solid Liquid f.p.  Dynamic equilibrium- Two phases simultaneously and continuously change into each other.  Rates are equal- hence, concentration of both phases remain constant as long as equilibrium exists.  At temperatures above 273 K, equil. shifts towards the right, i.e. ice spontaneously melts into water.  At temperatures below 273 K, equil. shifts towards left- water changes into ice. (Physical Equilibria) :  Liquid- Vapour Equilibrium :  Liquid placed in open container- vapour molecules disperse into atmosphere- equilibrium cannot be attained.  Liquid placed in closed container- after certain time, rate of condensation= rate of evaporation- state of dynamic equilibrium reached- concentrations become constant- constant vapour pressure attained- saturation vapour pressure/vapour pressure of liquid at given temperature.  Vapour pressure can be measured with the help of manometer. Liquid Vapour (Physical Equilibria) :  Solid- Vapour Equilibrium :  Subliming solids- equilibrium established between solid and vapour. I2 (solid) I2 (vapour)  Eg. NH4Cl(solid) NH4Cl (vapour) Camphor(solid) Camphor(vapour)  General Characteristics of Physical Equilibria :  Solid- liquid equilibrium : two phases exist together at equil. at a particular temperature- m.p./f.p.  Masses of both phases remain constant if no heat exchange takes place with surroundings.  Liquid-gas equil. : Pressure of gas above liquid remains constant if temperature is constant. (Physical Equilibria) :  General Characteristics :  Equilibrium attained only in closed system- no exchange of matter with surroundings.  Equilibrium always dynamic- forward and reverse processes occur simultaneously at same rate.  At equilibrium- all measurable properties of system remain constant as concentration of substance in two phases remains constant.  When equilibrium is attained, there exists a mathematical expression involving concentration of reactants which attains a constant value at a particular temperature (Table-7.1).  The magnitude of the constant value of concentration related expression indicates the extent to which reaction proceeds before reaching equilibrium. Chemical Equilibrium :  Irreversible and Reversible Reactions :  Irreversible reactions : Chemical reactions proceeding mainly in one direction (forward). Reactions go to almost completion. No reaction truly goes to completion- reactions apparently proceed in one direction. Actually both reactions occur- rate of reverse much smaller than forward. At equilibrium, conc. of reactants almost negligible compared to products. Eg. 2Na (s) + 2H2O (l) 2NaOH (aq) + H2 (g) AgNO3 (aq) + NaCl (aq) NaNO3 (aq) + AgCl (s) BaCl2 (aq) + Na2SO4 (aq) NaCl (aq) + BaSO4 (s) Chemical Equilibrium :  Causes of irreversibility : Any of the products escapes as gas. Any of the products gets precipitated as a solid.  Reversible reactions : Reactions proceed simultaneously in both directions (forward/reverse). Eg. A+B C+D N2(g) + 3H2 (g) 2NH3 (g) N2O4 (g) 2NO2 (g) CH3COOH (l) + C2H5OH (l) CH3COOC2H5 (l) + H2O (l) Reversible reaction reaches equilibrium in closed container. Reversible reaction never goes to completion. Chemical Equilibrium :  General reversible reaction : A+B C+D  At equilibrium, Rate of forward reaction = Rate of backward reaction.  Eg. 717 K H2 (g) + I2 (g) 2HI (g)  Attainment of equilibrium : Constancy of observable properties- eg. colour of reaction mixture (formation of HI), pressure, concentration of reactants/products, specific rotation, etc. Chemical Equilibrium :  Dynamic Nature :  At equilibrium, conc. of reactants/products, other measurable properties become constant- no change with time.  Apparently both reactions seem to stop.  Actually both proceed at the same rate- dynamic nature of equilibrium.  Kinetic Molecular Model : A+B C+D Chemical Equilibrium :  Characteristics :  Constancy of observable properties : All attain constant values at equil. and do not change with time. Eg. Concentration, colour of mixture, specific rotation, etc.  Attainment of equilibrium from either side : reaction can be started from any side- reactants or products- equilibrium state is not changed. 717 K Eg. H2 (g) + I2 (g) 2HI (g) Chemical Equilibrium :  Dynamic nature : Forward and backward reactions do not stop at equil. but occur continuously no matter how long the equil. state is maintained.  Equilibrium cannot be attained in open vessel : Some products may escape and backward reaction may not occur- hence, no equilibrium obtained.  Equilibrium point or state remains unchanged in presence of catalyst : Catalyst catalyses both forward and backward reactions to same extent- hence, equilibrium attained faster but same state. Chemical Equilibrium :  Law of Mass Action : Guldberg and Waage.  Helps in detailed mathematical studies of chemical equilibrium.  At constant temperature, the rate of a chemical reaction is directly proportional to the product of active masses of reacting species, with each active mass raised to the power of the stoichiometric coefficient of that species in the chemical reaction. Simple reaction : A + B Products Law of mass action : Rate of reaction ∝ [A] [B] = k [A] [B]; k= rate constant, [A]/[B]= active masses of reactants A and B. Chemical Equilibrium :  For a more general reaction : aA + bB + cC Products  Applying Law of Mass Action : Rate of reaction ∝ [A]a [B]b [C]c = k [A]a [B]b [C]c  Active Mass :  For a substance in solution/gas phase, active mass expressed as molar concentration.  Units for active mass :  Concentration unit : moles/litre- [A].  Partial pressure unit : For gaseous substances- atm. Chemical Equilibrium :  Law of Chemical Equilibrium :  Very helpful in predicting the extent of a reaction. A+B C+D  Law of mass action : Rate of forward reaction ∝ [A] [B]= k1 [A] [B] Rate of backward reaction ∝ [C] [D]= k2 [C] [D]  At equilibrium, rates are equal. At equilibrium, k1 [A] [B]= k2 [C] [D] k1/k2= [C] [D]/[A] [B].  At a particular temperature, k1, k2= constant. k1/k2= constant= K (equilibrium constant). K= k1/k2= [C] [D]/[A] [B] aA + bB cC + dD  For general reaction : K c= [C] [A] a c [D] [B]b d - Law of chemical equilibrium (Kc= K) Chemical Equilibrium :  Kp : Applicable when all reactants/ products in a reversible chemical reaction are in gaseous phase. Eg. aA (g) + bB (g) cC (g) + dD (g)  p , p , p , p = equil. partial pressures. A B C D  Law of equilibrium :  K = p c. p d p C D pAa. pBb N (g) + 3H (g) 2 2 2NH3 (g)  Eg.  K = p2 /p. p 3 p NH3 N2 H2 Chemical Equilibrium :  Relationship between Kc and Kp : aA (g) + bB (g) cC (g) + dD (g)  Kc= [C]c. [D]d ; Kp= pcC. pdD [A]a. [B]b paA. pbB  For an ideal gas, pV= nRT  p= nRT/V= cRT; c (=n/V)= molar conc. of gas.  pA= cART = [A]RT; pB= cBRT = [B]RT…  Kp= [C]RTc [D]RTd [A]RTa [B]RTb  =[C]c [D]d/[A]a [B]b. (RT)(c+d)-(a+b)  = Kc (RT)Δn; Kp= Kc (RT)Δn Chemical Equilibrium :  Homogeneous and Heterogeneous Equilibria :  Homogeneous Equilibrium : All reactants and products present N2(g) + 3H2 (g) in 2NH the 3 (g)same phase.  Eg. N2O4 (g) 2NO2 (g) CH3COOH (l) + C2H5OH (l) CH3COOC2H5 (l) + H2O (l)  Heterogeneous Equilibrium : Reactants and products present H2O(s) in different H2O (l) phases.  Eg. CaCO3 (s) CaO (s) + CO2 (g) Chemical Equilibrium :  Units of Kc and Kp :  Kc, Kp depend upon stoichiometric coefficients of reactants and products- hence, units different for different reactions.  Sum of stoich. coefficients of reactants= products :  Kc , Kp- dimensionless (no units).  Eg. 717 K H2 (g) + I2 (g) 2HI (g)  Sum of st. coefficients of reactants ≠ products :  Kc= (mol L-1)Δn; Kp= (atm)Δn  Eg. N2 (g) + 3H2 (g) 2NH3 (g)  Kc= (mol L-1)Δn; Kp= (atm)Δn  Δn= 2- (1+3)= -2.  Kc= (mol L-1)-2= mol-2 L-3.  Kp= atm-2. Chemical Equilibrium :  Characteristics of Equilibrium Constant :  Value of Kc/ Kp = constant at particular temperature- on changing temperature, k1 and k2 affected differently.  Value of K does not depend upon initial conc. of reactants.  Value gets reversed on reversing mode of representation of equilibrium, provided the same substances are involved.  Value of K depends on the units of active masses of reactants and products- Kp= Kc (RT)Δn.  Value of K does not change by presence of catalyst.  Value changes on changing stoichiometric coefficients of reactants/products- Eg. NH3 formation (halved).  Value of K predicts the extent to which a chemical reaction proceeds in either direction- K greater if forward is faster (more reactants converted into products); smaller if backward is faster. Chemical Equilibrium :  Applications of Law of Chemical Equilibrium / Law of Mass Action:  Law of mass action and chemical equilibrium very useful to predict the effects of changes in temperature, pressure, concentration or inert gas on state of equilibrium.  Dissociation of N2O4 : 2NO2 N2O4  Initial no. of moles of N2O4= a/V litre closed vessel; moles of NO2= 0 mole.  Moles dissociated till equil.= x moles; moles of NO2 formed= 2x moles.  At equil., no. of moles of N2O4 present= a-x moles; NO2= 2x moles.  [N2O4]= a-x/V; [NO2]= 2x/V.  Law of Mass Action/Law of Equilibrium :  Kc= [NO2]2/[N2O4]…..= 4x2/(a-x)V. Chemical Equilibrium :  Effect of changing pressure :  Change in pr. changes ‘V’- hence, changes value of K.  Increase in pr.- decrease in V- increase in ‘K’.  Increase in ‘K’ not possible as K is constant at constant temperature- hence, x decreases (lesser dissociation of N2O4).  Increase in pressure- lesser dissociation of N2O4; lowering of pr.- higher dissociation.  Effect of change in concentration :  Higher [N2O4]=a, lower ‘K’- hence, x should increase- thus higher dissociation.  Higher [NO2], Kc increases- hence, [N2O4] must increase- conversion of NO2 to N2O4 favoured- backward reaction favoured- suppresses dissociation of N2O4. Chemical Equilibrium :  Synthesis of Ammonia (Haber’s Process) : N2(g) + 3H2 (g) 2NH3 (g) Initial : a b 0 Equil. : (a-x) (b-3x) 2x (x moles of N2 react)  At equil. : [N2] = a-x/V; [H2]= b-3x/V; [NH3]= 2x/V.  Kc= [NH3]2/[N2] [H2]3. = 4x2V2/(a-x) (b-3x)3  Effect of pressure : Increase in pr.- decrease in ‘V’- decrease in Kc- hence, ‘x’ increases- more N2 combining with more H2 to form more NH3- increase in pr. favours formation of NH3.  Effect of Concentration : Increase in ‘a’ or ‘b’ decreases Kc- hence, ‘x’ increases- increase in [N2]/[H2] favours formation of NH3. Chemical Equilibrium :  Synthesis of Sulphur Trioxide (Sulphuric Acid) (Contact Process) : 2SO2 + O2 2SO3 Initially : a b 0 Equil. : (a-2x) (b-x) 2x (x moles O 2 react)  At equil., [SO2]= a-2x/V; [O2]= b-x/V; [SO3]=2x/V.  Kc= [SO3]2/[SO2]2 [O2] = 4x3.V/(a-2x)2 (b-x)  Effect of Pressure : Increase in pr.- decrease in ‘V’- hence, ‘x’ increases to keep Kc constant- hence, more O2 combine with more SO2 to form more SO3.  Increase in pressure favours formation of SO3.  Effect of Concentration :  Increases in ‘a’ or ‘b’- decreases Kc- hence, ‘x’ increases- increase in concentration of SO2/ O2 favours formation of SO3. Chemical Equilibrium :  Hydrolysis of Simple Esters +: H CH3COOC2H5 + H2O C2H5OH + CH3COOH Initial : a b 0 0 Equil. : (a-x) (b-x) x x  [CH3COOC2H5]= a-x/V; [H2O]= b-x/V; [C2H5OH]= x/V; [CH3COOH]= x/V.  Kc= [C2H5OH] [CH3COOH]/[CH3COOC2H5] [H2O] = x2/(a-x) (b-x)  Effect of change in pressure : No effect- no ‘V’ term in Kc expression.  Change in concentration : If a increases, Kc decreases- hence, x has to increase. Hence, more hydrolysis of ester- favours formation of alcohol and acetic acid. Hydrolysis of ester- unaffected by pressure but favoured by increase in conc. of ethyl acetate. Chemical Equilibrium :  Henry Louis Le Chatelier’s Principle :  Predicts the effect of changing conditions of a system at equilibrium.  A system attains equil. under a specific set of conditions.  Any change in conditions- disturbs system- equil. is disturbed- a net reaction takes place until equil. is re- established- new equil. state different from the old.  If a system at equilibrium is subjected to a change which displaces it from the equilibrium, a net reaction will occur in a direction that opposes the change.  Can be applied to any physical/ chemical system at equilibrium- by changing conc., temperature, pressure.  For a chemical system at equilibrium :  When a chemical system at equilibrium is subjected to a change in conc., pressure or  Chemical Equilibrium :  Effect of Change in Concentration : When conc. of a substance is increased, equil. shifts in the direction which neutralizes the effect of the change in conc. A+B C+D Eg.  Conc. of reactant A is increased : equil. shifts in the forward direction so as to use up the excess ‘A’- hence, equil. shifts towards right.  Conc. of ‘A’ is decreased : Equil. shifts towards left so that more ‘A’ is formed.  Conc. of product C is increased: equil. shifts towards backward direction so as to use up the excess ’C’- hence, equil. shifts towards left.  Conc. of ‘C’ is decreased : equil. shifts towards right (forward direction) so more ‘C’ can be formed. Chemical Equilibrium :  Eg. Decomposition of limestone in a kiln :  Continuous removal (escape) of CO2 gas favours forward reaction and continuous decomposition of limestone.  Formation of ammonia (Haber’s Process) :  Once equil. is reached, conc. of NH3 is constant- no more NH3 formed.  Continuous formation of NH3 can be achieved by increasing conc. of N2/ H2 or decreasing conc. of NH3.  Continuous formation of NH3 achieved in Haber’s Process by removal of NH3 from reaction site by liquefaction.  Transport of oxygen by Hb to tissues :  Removal of CO2 from tissues by blood and lungs from body : Chemical Equilibrium :  Effect of Change of Pressure (Pressure of a gaseous system depends on total no. of moles of substances) :  When pressure is increased or decreased on a system containing gaseous substances in equil., then the equil. shifts in the direction which tends to decrease/ increase the pressure and hence, to decrease/ increase the no. of moles. aA + bB cC + dD  When reaction proceeds with decrease in no. of moles (a+b)> (c+d) : eg. Haber’s Process.  Pressure of system decreases with forward reaction- hence, if pr. of system is increased, ‘V’ will decrease- greater no. of moles/ unit vol.  Hence, equil. shifts towards forward direction which involves decreases in no. of moles.  When reaction proceeds with decrease in no. of moles- more products formed by increasing pr. of system. Chemical Equilibrium :  When reaction proceeds with increase in no. of moles (c+d) > (a+b) : Eg. Decomposition of PCl5- more products can be formed by reducing pr. of system.  When reaction proceeds with no change in no. of moles (a+b)= (c+d) :  System remains unaffected by change in pr. Eg. 717 K H2 (g) + I2 (g) 2HI (g) Chemical Equilibrium :  Effect of Change of Temperature :  When the temperature of a system at equilibrium is increased, the equil. shifts in the direction in which heat is absorbed.  For exothermic reaction : Equil. shifts in backward direction (towards left).  Hence, greater yield of products obtained by lowering the temperature. Eg. Haber’s Process : Higher yield of NH3 possible only at low temperatures. N2(g) + 3H2 (g) 2NH3 (g); H= -92.4 kJ  For endothermic reaction : Increase in temperature favours forward reaction-more products formed. Eg. N2 (g) + O2 (g) 2NO(g); ΔH= 180.7 kJ  Higher yield of NO obtained at high temperatures. Chemical Equilibrium :  Effect of Catalyst : Does not affect state of equilibrium or value of equilibrium constant.  Helps to attain equilibrium faster.  Effect of Addition of an Inert Gas :  At constant volume :  Conc. of neither reactant, nor product changes on addition of inert gas.  No change in state of equilibrium or system.  At constant pressure :  Total volume of reaction mixture increases.  Equil. affected if reaction involves change in volume.  Eg. Dissociation of PCl5 : Kc= x2/(a-x)V.  Equil. shifts in forward direction on addition of inert gas. Chemical Equilibrium :  Applications of Le- Chatelier’s Principle to Industrial Processes :  Principle of great significance in predicting optimum conditions for obtaining max. yield of products in industrial processes.  Synthesis of Ammonia2NH (Haber’s (g); Process) : N2(g) + 3H2 (g) 3 H= -92.4 kJ  /-22.0 kcal  Reaction exothermic.  Proceeds with decrease in no. of moles.  Max. yield obtained at low temperature/ high pressure.  Catalyst- finely divided iron with Mo- at low temperatures, activity of catalyst decreases- hence, reaction carried out at 450°C/ 200 atm. Chemical Equilibrium :  Manufacture of Sulphuric Acid (Contact Process) :2SO2 + O2 2SO3 + 42 kcal  Reaction exothermic.  Proceeds with decrease in no. of moles.  Max. yield obtained at low temperature/ high pressure.  400- 450°C/ 2- 3 atm.  Catalyst used- V2O5.