CHEM 212

1. CHEM 212 Chapter 4 Chemical Equilibrium By Dr. A. Al-Saadi

2. Definition of Chemical Equilibrium In a chemical process, chemical equilibrium is the state in which the concentrations of the reactants and products have NO net change over time. A chemical reaction is in a dynamical equilibrium when it is occurring in forward and reverse directions with the rate being the same in both directions. For a chemical reaction of the form: aA + bB + … … + yY + zZ when the equilibrium condition is reached and at fixed T, we have: This ration is known as the equilibrium constant (KC). For 1st order reactions, the rate is given by: and

3. Definition of Chemical Equilibrium At equilibrium with fixed T, there is no net change in concentrations and both forward and backward rates are equal: Remember, this is only true for 1st order reactions. From previous chapter we showed that at an equilibrium condition, G and A tend to go to minimum, while S tends to go to maximum values. In terms of chemical reaction, this is correct when the forward and backward processes are occurring at equal rates.

4. Relationship between ΔGº and KP ΔGº = –RT lnKP

5. ln Relationship between ΔGº and KP ΔGº = –RT lnKP These two plots show this relation linearly (left) and logarithmically (right). Notice that an equilibrium constant of unity implies a standard free energy change of zero, and that positive values of ΔG° lead to values of K less than unity.

6. Free Energy Changesin the Dissociation of N2O4 The gas-phase reaction: N2O4 2NO2 furnishes a simple example of the free energy relationships in a homogeneous reaction.

7. Relationship between ΔGº, ΔG and KP

8. Relationship between ΔGº and KP

9. Relationship between ΔGº and KP

10. Heterogeneous Equilibrium

11. Test for Chemical Equilibrium • For a chemical reaction of the form: • aA + bB + … … + yY + zZ • The equilibrium constant is given by: • One of the most difficult task, especially for extremely slow reaction, is to tell whether a reaction is at equilibrium or not. • Example : 2H2 + O2 2H2O • Two practical ways to test for the equilibrium condition: • Adding catalysts: Catalysts don’t change the position of the equilibrium, but they increase the reaction rate and accelerate its approach to equilibrium. • Adding a small amount of a reactant or a product and observe the change in the reaction course.

12. Le ChâtelierPrinciple • It can be used to predict the effect of a change in conditions on a chemical equilibrium. • If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or total pressure, then the equilibrium shifts to counteract, as far as possible, the effect of that change. • Example 1: • CO + 2 H2 ⇌ CH3OH • If we increase the concentration of CO in the system, using Le Châtelier's principle we can predict that the amount of methanol will increase, decreasing the total change in CO. • Example 2: • N2 + 3 H2 ⇌ 2 NH3 ΔH = −92kJ • This is an exothermic reaction when producing ammonia. If we were to lower the temperature, the equilibrium would shift in such a way as to produce heat. Since this reaction is exothermic to the right, it would favor the production of more ammonia. In practice, in the Haper process the temperature is instead increased to speed the reaction rate at the expense of producing less ammonia.

13. Le ChâtelierPrinciple

14. Coupling of Reactions

15. Van’t Hoff Equation Dependence of Equilibrium Constant on Temperature ΔSº/R

16. Dependence of Equilibrium Constant on Pressure

17. Dependence of Equilibrium Constant on Pressure