Notes of Chemical Equilibrium

SLOs (Topics)	Sub -Topics	Chemical Equilibrium
Reversible reaction  and Dynamic equlibrium	7.1.1	All chemical reactions do not proceed to same extent. Some reactions proceed to completion after sometime, but there are many reactions, which are never completed. On the basis of extent reactions, chemical reactions are classified into: Ø  Irreversible Reactions Ø  Reversible Reactions Irreversible Reactions: “Those reactions that proceed to complete in a definite direction are called Irreversible Reactions”. OR Those reactions in which the reactants react to form product which do not change back are known as “Irreversible Reactions”. Reversible Reaction: There are some reactions in which the products again combine to form the reactants. This reaction therefore precedes in two directions i.e. forward & backward directions. “Such reactions which proceed to both backward & forward directions and are never completed are called Reversible Reactions”. e.g. (a)        2HI                              H2 + I2 (b)        N2 + 3H2                                            2NH3­ (c)        CH3COOH + C2H5OH                          CH3COOC2H5 + H2O The double arrow indicates that the reaction is reversible and both the forward and backward reactions can occur simultaneously.	K
	7.1.2	Equilibrium State: “A reversible reaction is said to be in a state of Equilibrium, when the rate of its forward reaction equals to the rate of its backward reaction and the concentration of various constituents remains unaltered”. Explanation: In a reversible reaction the changes, forward and reverse occur simultaneously. Under these circumstances, a reaction might come to some kind of “Balance” in which the forward and reverse reactions occur at the same rate. For example consider A + B                          C + D             In the beginning, forward reaction predominates, but as soon as C & D are formed the reverse reaction builds up until equilibrium position is reached, where the forward as well as the reverse change proceeds with the same rate i.e. At Equilibrium state.                                                                                                       Rate of forward Reaction       =          Rate of backward Reaction	K
	7.1.3	Law of Mass Action: In 1864, Goldberg and Wage studied the effect of concentration on reversible reactions in equilibrium and put forward a law, which is know as “Law of Mass Action”                                According to this law, “The rate of a chemical reactions is directly proportional to the product of active masses (or molar concentrations) of the reactants”.             The number of moles of a substance in 1dm3 is called its “Molar Concentration” or “Active Mass” and is denoted by square brackets. Expression of   Kc: We consider the following reaction,                      Forward mA + nB                     xC + yD         Backward      According to the law of mass action or --------- (1) Similarly,                --------- (2) or At equilibrium state, Rate of forward = Rate of backward reaction Where ‘Kc’ is called the “Chemical Equilibrium Constant” and the equation is called the chemical equilibrium constant expression, for the general reversible reaction, where x,y,m & n represents the moles of species and are called co-efficient of chemical equations.             In the case of gaseous equilibrium, a partial pressure is sued instead of concentration because at a given temperature, partial pressure of a gas is proportional to its concentration. In this case the equilibrium constant is expressed as KP instead of KC. e.g. for the following gaseous equilibrium; A(g) + B(g)                    C(g) + D(g) Where PA, PB, PC & PD are the partial pressures of gases A,B,C, & D respectively.	U
	7.1.4	Kp      =        [PC] [PD]                                   [PA][PB] Kc      =       [C] [D]                                   [A][B] Kp            =      Kc  x  (RT)∆n Kn      =        [NC] [ND]                                   [NA][NB] Kp            =      Kn  x  (RT)∆n                                  V∆n	U
	7.1.5	N2      +        3H2                        2NH3 t=0           a mol                 b mol                    Nil t=eq.        (a-x)mol           (b-3x)mol            2x mol Acc .to Law of mass action             K =     [NH3]2                          [N2][H2]3       =            (2x/v)2                 ( a-x/v)(b-3x/v)3            H2     +        I2                             2HI t=0      a mol              a mol                    Nil t=eq.    (a-x) mol        (a-x) mol                   2x Acc .to Law of mass action      K =   [HI]2                  [H2][I2]             =        (2x/v)2                 ( a-x/v)(a-x/v)                         PCl5                     PCl3         + Cl2            t=0           a mol                       Nil                    Nil            t=eq.         (a-x)                          x                     x Acc .to Law of mass action       K =      [PCl3][Cl2]                       [PCl5]            =      (x/v) (x/v)                     ( a-x/v)                         2HI                                   H2 +      I2              t=0                        a mol                                Nil             Nil t=eq.                     (a-2x) mol                        x mol        x mol Acc .to Law of mass action K =   [I2][H2]            [HI]2            =      (x/v) (x/v)                     ( a-x/v)2	U
	7.2.1	Le -Chatelier’s Principle: It is a general principle that gives qualitatively the influence of change in temperature, pressure or concentration on system in equilibrium. This principle was first enunciated by a French Chemist                      Henn Le-Chatelier in 1884. According to this principle, “If a system in equilibrium is subjected to a stress, the equilibrium shifts in a direction to minimize or undo the effect of the stress”. Where “Stress” means change in concentration, temperature or pressure. If one of the factors involved in a chemical equilibrium is altered, the equilibrium shifts towards right or left in order to restore the balance of equilibrium.	K
	7.2.2	Effect of Concentration Change: If concentration of reactant or concentrations of or reactants are increase at equilibrium state then equilibrium will shift to words forward direction   . If concentration of product or concentrations of products are increase at equilibrium state then equilibrium will shift to words Backward direction   . If concentration of product or concentrations of  products are removed at equilibrium state then equilibrium will shift to words Backward direction   .   Effect of Temperature Change:          PCl5                             PCl3  +      Cl2     ∆H=Negative    If we increase the temperature of above reaction at equilibrium state the equilibrium state will shift backward direction.    If we decrease the temperature of above reaction at equilibrium state the equilibrium state will shift forward  direction.     Effect of Pressure Change:    N2      +        3H2                        2NH3 If we increase the pressure of above reaction at equilibrium state the equilibrium state will shift forward direction.    If we decrease the pressure of above reaction at equilibrium state the equilibrium state will shift backward direction. Effect of Catalyst : The catalyst does not effect on equilibrium state	U
	7.2.3		U
	7.2.4	Kc for a reaction is 0.0194 and the calculated ratio of the   concentration of the reactants is 0.0116.  Predict the direction of the reaction.	U
	7.3.1	Applications of Le-Chatelier’s Principle: 1.         The Haber’s Process:             Ammonia can be prepared by the Haber’s process as;             N2 + 3HH2Hhhhaa H2                    2NH­3   (DH = -46.2) i.e. The reaction between nitrogen and hydrogen to produce ammonia is accompanied by decrease in volume and it is exothermic. Effect of Concentration:     According to Le-Chatelier’s Principle, the addition of more N2 or H2 or both will move the reaction to the right, thus more NH3 gas will be produced till the equilibrium state is reached again. Similarly addition of NH3 at the equilibrium state will move the reaction to the left. Thus, a part of NH3 gas will decompose into N2 & H2 gases in order to reach the equilibrium state again. Effect of Temperature: It is an exothermic reaction i.e. heat is evolved during the reaction. According to Le-Chatelier’s principle, the low temperature shifts the equilibrium to the right direction  i.e. Effect of Pressure Change:    N2      +        3H2                        2NH3 If we increase the pressure of above reaction at equilibrium state the equilibrium state will shift forward direction.    If we decrease the pressure of above reaction at equilibrium state the equilibrium state will shift backward direction.	U
	7.3.2	Same as   SLOs # 7.2.2	U
	7.4.1		U
	7.4.2		U
	7.4.3
	7.5.1	Common Ion Effect: (i)     CH3COONa                     CH3COO- + Na+          (ii)     CH3COOH(aq)                   CH3COO- + H+ If sodium acetate is added to an acetic acid solution, it ionizes into acetate and sodium ions as CH3COONa                         CH3COO- + Na+ Therefore acetate ion concentration is increased and equilibrium will shift. Since there are more CH3COO- ions, so according to Le-Chatelier’s principle the rate of reaction will increase toward acetic acid. Some of the excess acetate ions unite with H+ ions to form molecular acetic acid, and hence the degree of ionization of acetic acid will reduce. The acetate ion is common to both acetic and sodium acetate. The effect of the acetate ion on the acetic acid and solution is called “Common Ion Effect”.
	7.5.2