2. Chemical Equilibrium Section 18-1 Pp. 589 - 591

3. Equilibrium is… Equilibrium is not static Opposing processes occur at the same time and at the same rate Rate of forward reaction = Rate of reverse reaction For example 1: water at 0.0°C At 0.0, some liquid H2O is freezing and some ice (solid) H2O is melting. Their rates are equal, so equilibrium For example 2: adding sugar molecules into water At equilibrium, the rate of sugar molecules in solution crystallizing equals the rate of sugar crystals dissolving

4. Reversible Reactions Every reaction can proceed: Reactants  Products OR Reactants  Products Def. – chem rxn in which products can react to re-form reactants 2HCl  H2 (g) + Cl2 (g) Here most HCl is decomposing into hydrogen gas & chlorine gas However, some H2 and Cl2 are synthesizing into HCl At equilibrium, these rates are equal

5. Conclusion A reversible chemical reaction is in chemical equilibrium when rate of forward reaction = rate of reverse reaction The concentrations of the reactants and products are static [Reactants] DOES NOT = [Products] !!! Usually see double arrows to indicate reversibility:

6. Equilibrium lies to the right? Some reactions “favor” the formation of products At equilibrium: higher concentrations of Products than of Reactants [Products] > [Reactants] Sometimes the forward arrow will be longer than the reverse arrow, to indicate “product favored” Some reactions “favor” the formation of reactants At equilibrium: [Products] < [Reactants] Sometimes the reverse arrow will be longer than the forward arrow, to indicate “reactant favored”

7. Equilibrium Constant Given the general reaction: nA + mB  xC + yD Initially, there is only A & B but no C or D, so the forward reaction rate is at its maximum. Over time, C + D accumulate, so forward rate slows & reverse rxn rate increases Eventually the two rates become equal to each other Equilibrium !!!

9. Equilibrium Constant After equilibrium is reached, the individual concentrations of A, B, C, and D undergo no further change if conditions remain the same. A ratio of their concentrations should also remain constant. The equilibrium constant is designated by the letter K.

10. The constant K is independent of the initial concentrations. K is dependent on the temperature of the system. The Equilibrium Constant The numerical value of K for a particular equilibrium system is obtained experimentally.

11. If the value of K is small, the reactants are favored. A large value of K indicates that the products are favored. Only the concentrations of substances that can actually change are included in K. Pure solids and liquids are omitted because their concentrations cannot change.

12. Determining Keq for Reaction at Chemical Equilibrium

13. Shifting Equilibrium Section 18.2 Pp. 598 - 601

14. Le Chatelier’s Principle A system is happily at equilibrium, then a change in concentration, pressure, or temperature occurs What will happen to the equilibrium? Def. – if a system in equilibrium is stressed, then the equilibrium is shifted in a way to relieve that stress 3 main Stressors: pressure, concentration, and temperature

15. Pressure Change N2(g) + 3H2(g)  2 NH3(g) What does the (g) stand for? How many moles of gas on the left? On the right? So if there was a pressure INCREASE: Causes an increase in gas concentration Which side will experience the biggest change in gas pressure (concentration)? How would the system relieve the stress? Shifts toward side with fewer moles!! Pressure increase = shift to side with fewer moles

16. Pressure Change (Page 2) What is pressure was DECREASED? Shifted to side with more moles of gas INVERSE relationship between pressure change and side with number of moles INCREASE pressure = shift to side with FEWER moles DECREASE pressure = shift to side with MORE moles

17. Pressure change & K eq Even though changes in pressure may shift the equilibrium position, they do not affect the value of the equilibrium constant. Increasing pressure by adding a gas that is not a reactant or a product cannot affect the equilibrium position of the reaction system.

18. Concentration Change If one side INCREASES concentration, then the system will shift the reaction to the opposite side If one side DECREASES concentration, then the system will shift the reaction to its side If we increase the H2 concentration, what will happen? Will N2 increase or decrease?What about NH3? If we decrease the N2 concentration, what will happen? Will H2 increase or decrease?What about NH3?

19. Concentration & K eq Changes in concentration have no effect on the value of the equilibrium constant. Such changes have an equal effect on the numerator and the denominator of the chemical equilibrium expression.

20. Temperature Change Reactions are either exothermic or endothermic If endothermic, then energy is a REACTANT If exothermic, then energy is a PRODUCT So if a reaction is ENDOTHERMIC, then increasing temperature shifts equilibrium to products Increase TEMP  Shift to PRODUCTS Decrease TEMP  Shift to REACTANTS

21. Temperature Change If a reaction is exothermic, then heat is a PRODUCT In an EXOTHERMIC rxn, when temp is INCREASED, then the rxn shifts to REACTANTS In an EXOTHERMIC rxn, when temp is DECREASED, then the rxn shifts to PRODUCTS The value of the equilibrium constant (K eq ) for a given system is affected by the temperature.

23. Haber Process Artificial production of ammonia N2(g) + 3H2(g)  2 NH3(g)+92 kJ In this reaction, we want to produce as maximum amount of NH3 (Ammonia) There are how many moles of reactants? Products? So how do we create the maximum product concentration by changing the pressure?

24. Haber Process (Page 2) N2(g) + 3H2(g)  2 NH3(g)+92 kJ Temperature Is the reaction endothermic or exothermic? Is heat a reactant or product? So what should be done to maximize products? Concentration How would me maximize products by altering concentrations?

25. Last Example Bicarbonate Buffer in mammalian blood H2CO3  H+ + HCO3- What happens when pH decreases? What happens to the [H+]? So what happens to the equilibrium? What about the [H2CO3] & [HCO3-] What happens when pH increases? What happens to the [H+]? So what happens to the equilibrium? What about the [H2CO3] & [HCO3-]