Chemistry Chapter (4) Chemical Equilibrium

Points To Remember

   Homogeneous Reaction: The reaction in which all the components of the system (reactants and products) are in the same physical state , are called homogeneous reactions.

        Example:   H_2(g)+ I_{2(g) <--------------------->}   2H1_(g)

·       Heterogeneous Reaction: The reaction in which all the components of the system (reactants and products) are in different physical state, are called heterogeneous reactions.

        Example: CaCO_3(s) <--------------------->  CaO_(s)+ CO_2(g)

·       Irreversible Reaction: A reaction is which reactants are completely converted into product and the reaction is fast but cannot be reversed is called irreversible or unidirectional reaction.

       Example: HCL + NaCl <---------------------> NaCl + H₂O

·       Reversible Reaction: If in a chemical reaction, reactants slowly change into products but not completely and the products are also capable of converting into reactants, is called a reversible reaction.

     Example N_2(g) + 3H_2(g) <---------------------> 2NH_3(g)

 This reactions is represented by    (<---------------------> )

·       Equilibrium State of Chemical Reaction: Chemical equilibrium is the state when rates of two opposing reaction (i.e. forward and reverse) are the same and concentration of reactant and products  do not change with time.

·       Dynamic Chemical Equilibrium: Chemical equilibrium is the state when two opposing reactions proceed simultaneously at identical rates. It is an equilibrium because concentration do not change. It is dynamic because opposing reactions never cease.

  ·       Characteristics of Chemical Equilibrium: 
      Chemical equilibrium is dynamic in nature.

  ·       It can be approached from both sides.

  ·       The equilibrium readjusts with the changing conditions and spontaneously goes back to the original state when the disturbing factors are removed.

  ·       Equilibrium is not affected by the presence of a catalyst. The catalyst only hastens the approach of the equilibrium state.

  ·       Change in temperature, pressure or concentration favours one of the reactions and thus shifts the equilibrium point in one direction.

  ·       There is no apparent change in properties with time.

  ·       Active Mass: Molar concentration or the concentration expressed in mole per dm³ is also known as active mass.

  ·       Rate Expression: The equation which show the relationship between rate of a chemical reaction and the mole concentration of the reactant is known as rate expression.

  ·       Law of Mass Action: At constant temperature, rate of chemical reaction is directly proportion to the product of molar concentration of reacting species.

  ·       Kc (equilibrium constant in terms of concentration):

       Kc = [product] Equilibrium /[Reactant] Equilibrium

  ·       It is independent of initial concentration of reactants but depends upon      temperature.

  ·       Kp (equilibrium constant in terms of partial pressure): For a general reversible gaseous reaction

         n₁A_(g) + n₂ B_(g) <----------->  n₃  C_(g) + n₄  D_(g) 

KP = (P_c)ⁿ₃  (P_D)ⁿ₄ / (P)ⁿ₁  (P)ⁿ₂   

  ·       Law of Chemical Equilibrium: At the equilibrium point, the ratio of the product of concentration of the substance produced and the product of the concentrations of the reactants with exponents of their respective coefficients is constant.

  ·       Relation between Kc & Kp:

    Kp = Kc when there is no charge in volume; e.g.

        H₂ + I₂  <--------->2HI

    Kp > Kc when reaction occurs with increase in volume; e.g.

           2NH_{3 <------------>} N₂ +3H₂

    Kp < Kc when there is decrease in volume on product side; e.g.

           2SO₂ + O₂                           2SO₃

                           Kp = Kc  (RT)^Δn 

  ·      

Whe

   R = gas constant, T = absolute temperature

  ·       Equilibrium Constants and their Units: Both Kp and Kc will be dimension i.e. without units when Δn = 0

  ·        If the total number of moles of reactants is not equal to that of the products, then the units of equilibrium constant change depending on the reaction involved. Consider the following examples;

       H₂ + I_{2 <---------->} 2HI;     Kc = [HI]² / [H₂] [I₂] (no units of Kc)

       PCl_{5  <======>} PCl₃ + Cl₂;  Kc : [PCl₃] [Cl₂]/ [PCl₅] (unit of kc will be mole litre^-1)

       N₂ + 3H₂ <=======>2NH₃; Kc = [NH₃]² / [N₂][H₂]³ ( unit will be mole^-2 liter²)

   Similarly, for the reaction;

                     2SO_2(g) + O_{2(g)  <=====>} 2SO_3(g)

                     Kp = P² _SO3 / P²_SO2  P_o2 = (Pressure)²/ (Pressure)² (Pressure)                                      i.e. atmosphere^-1

   Thus, the unit of equilibrium constant change depending on the reaction involved.

·       Is Equilibrium Constant Really Constant? To answer this we need to think about changes in four factors that are often used to bring about chemical changes.

INFLUENCES ON EQUILIBRIUM CONSTANT

CHANGE	EFFCT OF EQUILIBRIUM CONSTANT
Temperature	Change
Concentration	No charge
Pressure	No Charge
Catalyst	No charge (but equilibrium reached    more quickly)

·       Temperature: Equilibrium constant change when temperature changes; but provided the temperature does not change, an equilibrium constant really is constant. For example, in the reaction;

             2SO_2(g)+ O_2(g)                          2SO_3(g)

At 298k, the equilibrium constant is always Kp = 4×10²⁴ atm^-1

·       Concentration: If concentration of the reactants or products in a reaction at equilibrium is changed, the proportions of reactants and products adjust themselves in such a way that Kc does not change (provided the temperature does not change). For example , if we add alcohol to the equilibrium;

                Alcohol + Acid                       Ester + water

          For which Kc = [easter][water] / [alcohol][acid]

Then some of extra alcohol reacts to make more ester and water. As a result, the increase in [alcohol] is just balanced by the increase in [ester] and [water].

·       Pressure: Equilibrium constants do not charge when pressure changes. For example, if sulphur dioxide reaction above is performed at 1atm or 10atm then Kp keeps the same value, provided the temperature is 298k. However , just as they do when concentrations are changed.

·       Catalyst: We known that at equilibrium there is no overall change in the proportion of the reactant and products. However this is not to say that chemical life has come to an end equilibrium. The reactions between the chemicals are still taking place i.e. reactant and products come to a state of dynamic equilibrium.

At equilibrium, the reactants are changing into products at the same rate as the products are changing back to give reactants. If a catalyst is added to the reaction mixture, then the rates of the reactions increase and equilibrium is achieve more quickly. However, the proportion of the reactants and products at equilibrium do not change. For this reason, a catalyst has no effect on the value of an equilibrium constant.

·       How does Temperature Affect an Equilibrium Reaction? For an endothermic reaction , an increase in temperature favours the products.

      For endothermic reaction, an increase in temperature avours the reactants.

   The effect of temperature on an equilibrium can be explained by the principle proposed by the French chemist Henri Le chatelier in 1885.

·       Le Chatelier’s Principle: “If a constraint is placed on an equilibrium mixture, then the equilibrium will shift so as to oppose the constraint”.

     At present, we are thinking of ‘a constraint’ as a change in temperature . Thus, the principle says that if the temperature is raised then the equilibrium will shift so as to reduce the temperature.

                     In an exothermic reaction, the charge from reactant to products gives out heat; the reverse change, from products to reactants absorbs heat.

·       Exothermic Reaction:

            Reactants favoured at    ^{Heat released}          Products favoured at

            High temperatures       _{Heat absorbed}          low temperature

Therefore, if the temperature is increased, the reaction will shift to the left. As the proportion of reactants increase more heat is absorbed which results in the temperature being reduced.

·       Endothermic Reaction:

          Reactants favoured          ^{Heat absorbed}        Products favoured at

          At low temperatures     ^{Heat released}         high temperature

In this cause, if temperature is increased, then the equilibrium will shift to the right, in favour of products.

·       APPLICATIONS OF CHATELIER’S PRINCIPLE TO CHEMICAL EQUILIBRIA

·       Effect of change of temperature:

     Synthesis of Nitric oxide (Birekeland and Eyde process);

                N_2(g) + O_2(g)                     2NO_(g) ΔH = 43.2 Kcals

·       As the forward reaction is endothermic, the increase in temperature will favour forward reaction.

·       Effect of Change of pressure: The reaction in which there is no change in the number of moles of reactants and products (Δn=0), are not affected by change in pressure or volume. e.g.;

              H_2(g) + I_2(g)                        2HI_(g)

               N₂ + O_2(g)                         2NO_(g)

    Pressure would have no effect on above equilibria.

  ·        If pressure is increased, the equilibrium will shift in that direction in which total number of moles is decreased. e.g.

     N_2(g) + 3H_2(g)                     2NH_3(g)

       1    +  3 = 4 moles           2 moles

    (It is noted that increase in pressure does not affect the volume of solids and liquid to an appreciable extent, the number of moles per unit volume also remains practically constant, Hence the equilibrium is not affected by change of pressure in reactions involving solid or liquid system).

 ·       Effect of Change of Concentration: If concentration of the reactant is increased or product is removed, the equilibriumwill shift in the forward direction so as to favour the reaction which proceeds with a decrease in the number of moles.

     For e.g.         N_2(g) + 3H_2(g)   <=====> 2NH_3(g)

                          2SO_2(g)  +   O_{2(g) <======>} 2SO_3(g)

    The forward reaction in both cases are exothermic and takes place with a decrease in volume. thus an increase on reactant side will result in good yield of products.

Sher Afzal Ranais

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