CHAPTER 16: PRINCIPLES OF REACTIVITY: Chemical Equilibria

16.0 OBJECTIVES State the characteristics of a system in dynamic equilibrium as applied to reversible reactions. Calculate the value of and interpret the meaning of the reaction quotient based on either partial pressures or concentrations. Understand the factors that will change the values of the reaction quotient. Use LeChatelier’s Principle to describe changes in reactions based on stresses to the equilibrium system.

Homework #1 - 9, 11, 13, 15, 37, 39, 45 Basic Equilibrium Calculations #2 - 17, 19, 21, 23, 41 ICE Problems #3 - 27, 29, 31, 43, 51, 53, 55, 57 More ICE and Conceptual Equilibrium #4 - 67, 33, 35, 47, 49 Le’Chatelier’s Principle

16.1 NATURE OF THE EQUILIBRIUM STATE A. Requirements Reaction that forms products to a certain point (equilibrium!), then reforms reactants NOT ALL RXNS will go to completion and only form products Equ. will not occur as soon as the reversible reaction begins  Like all good things, it takes time!

16.1 NATURE OF THE EQUILIBRIUM STATE B. Characteristics of a Specific System Double Headed Arrow Beginning of Rxn  Form lots of products (rate laws) Moments Later  Forming both products and reactants (equilibrium)  **Not Necessarily at Equal Rates! Dynamic Equilibrium- “two opposing processes occur at exactly the same rate” Reactants  Products = Products  Reactants

16.1 NATURE OF THE EQUILIBRIUM STATE C. Graphs

16.1 NATURE OF THE EQUILIBRIUM STATE N 2 O 4  2NO 2 Equilibrium Demo Decomposition Reaction NO 2 = Smog (dark brown color) N 2 O 4 = Propellant (light brown/clear) **Observe and Write down some observations!

16.2 THE REACTION QUOTIENT AND EQUILIBRIUM CONSTANT A. For aA(g) + bB(g)  cC(g) + dD(g) K c = [C] c [D] d …_ [A] a [B] b...

16.2 THE REACTION QUOTIENT AND EQUILIBRIUM CONSTANT B. Example 16.1 Write the K c for a. H 2 (g) + I 2 (g)  2 HI (g) b. 2 SO 2 (g) + O 2 (g)  2SO 3 (g)

16.2 THE REACTION QUOTIENT AND EQUILIBRIUM CONSTANT C. K c independent of Concentration dependent on Temperature Reaction/Stoichiometry

16.2 THE REACTION QUOTIENT AND EQUILIBRIUM CONSTANT D. No term for pure liquids, solvents, or solids Concentration does not change during the reaction

16.2 THE REACTION QUOTIENT AND EQUILIBRIUM CONSTANT

16.3 DETERMINING AN EQUILIBRIUM CONSTANT A. Determine the value of K c when all equilibrium concentrations are known: Products over Reactants raised to their coefficients

16.3 DETERMINING AN EQUILIBRIUM CONSTANT 1. Example 16.3 For NH 4 Cl(s)  NH 3 (g) + HCl(g) at 500 o C, at equilibrium there are 2.0moles of ammonia, 2.0 moles of hydrochloric acid and 1.0 mole of ammonium chloride present in a 5.0L container. Calculate the equilibrium constant of the system at this temperature.

16.3 DETERMINING AN EQUILIBRIUM CONSTANT 2. Example 16.4 For the system 2SO 3 (g)  2SO 2 (g) + O 2 (g) at a given temperature the equilibrium concentrations are [SO 2 ] = [O 2 ] = 0.10M and [SO 3 ] = 0.20M. Calculate the K c at this temperature.

16.3 DETERMINING AN EQUILIBRIUM CONSTANT

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS A. Predicting the direction of shift of a reaction to reach equilibrium 1. For aA(g) + bB(g)  cC(g) + dD(g) K c = [C] c [D] d …_ [A] a [B] b... Q = [C] c [D] d …_ [A] a [B] b...

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 2. Relationship between K c and Q Q is the ratio of Products/Reactants at ANY given point in time, not necessarily at equilibrium Q values can change as reaction proceeds K c is the equilibrium constant which is established for EQUILIBRIUM CONDITIONS ONLY!

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 3. Example 16.7 N 2 (g) + 3H 2 (g)  2NH 3 (g) K c = 5 x 10 8 at 25 o C If the original concentrations are [N 2 ] = [H 2 ] = 2.0M, which way will the reaction run?

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 4. Example 16.8 For the system N 2 O 4 (g)  2NO 2 (g) K c = 0.36 at 100 o C. Predict the direction the reaction will shift to reach equilibrium if the original concentrations are: a moles of N 2 O 4 in a 4.0L container b moles of N 2 O 4 and 0.20 moles of NO 2 in a 4.0L container c. 1.00M NO 2 and N 2 O 4

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS B. Determine the equilibrium concentration of one reactant or product: Example 16.9 For N 2 (g) + O 2 (g)  2NO(g) K c = 1.0 x at 25 o C Calculate the equilibrium concentration of NO(g) if at equilibrium the concentration of N 2 (g) is 0.04M and that of O 2 is 0.01M.

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS C. Determine the equilibrium concentrations of all reactants and products from original concentrations and K c. Example For CO 2 (g) + H 2 (g)  CO(g) + H 2 O(g) K c = 0.64 at 900K. Starting with both reactants at a concentration of 0.100M, what are the equilibrium concentrations of all species?

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS Example For the same system as above, calculate the equilibrium concentrations of all species if the reaction system is started with [CO 2 ] = 0.100M and [H 2 ] = 0.200M.

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS Example For N 2 O 4 (g)  2NO 2 (g) K c = 0.36 at 100 o C. Starting with a concentration of 0.100moles/L for N 2 O 4, what are the equilibrium concentrations of both species?

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS A. K P Ratio of Partial Pressures of reactants/products at Equilibrium K P = P C c P D d …_ P A a P B b...

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS B. Equation K p = K c (RT)  n gas

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS C. Example For N 2 (g) + 3H 2 (g)  2NH 3 (g) K c = 9.5 at 300 o C. Calculate the K P at this same temperature.

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS D. Equilibrium constants and different forms of the equation For N 2 O 4 (g)  2NO 2 (g) K c = 0.36 at 100 o C 1. K c of 2NO 2 (g)  N 2 O 4 (g) Reverse Equation: 1/K

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS 2. K c of 4NO 2 (g)  2N 2 O 4 (g) Multiply by coefficent: K coefficient 3. K c of 1/2N 2 O 4 (g)  NO 2 (g) See above 4. Summation of Equations Rule Multiply K values for all equations

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS 5. Example Given: 3/2O 2 (g)  O 3 (g) K = 2.5 x State the K for a. 3O 2 (g)  2O 3 (g) and b. 2O 3 (g)  3O 2 (g)

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS 6. Example Given at 500K: H 2 (g) + Br 2 (g)  2HBr(g)K p = 7.9 x H 2 (g)  2H(g)K p = 4.8 x Br 2 (g)  2Br(g)K p = 2.2 x Calculate the K p for H(g) + Br(g)  HBr(g) at the same temperature

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle A. Statement “When a system is stressed (changing conc., temp, pressure, or volume), the system will respond by attaining new equilibrium conditions that counteract the change”

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle B. Adding of removing a reactant or product Equilibrium shifts in the direction opposite of the stress (ex. add products  shifts to reactants)  Add Reactant  Shifts toward products  Remove Reactant  Shifts toward reactants  Add Product  Shifts toward reactants  Remove Product  Shifts toward products

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle C. Example For 2HI(g)  H 2 (g) + I 2 (g) K c = at 520 o C Starting at the following concentrations at equilibrium [H 2 ] = [I 2 ] = 0.010M and [HI] = 0.080M.. After reaching the above equilibrium, enough HI is added to raise its concentration of 0.096M. What are the equilibrium concentrations of all species?

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle D. Change in amounts of solid, pure liquids, or solvents Has NO effect on equilibrium

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle E. Change in volume or pressure (for gases) Pressure increased (Volume dec.)  Equ. shifts in the direction producing the smaller number of moles of gas Pressure is decreased (Volume inc.)  Equ. Shifts in the direction producing the larger number of moles of gas NO CHANGE in moles  no effect on an equilibrium

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle F. Change in temperature Increase Temp  Equ. shifts in the direction of the endothermic rxn Decrease Temp  Equ. shifts in the direction of the exothermic rxn

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle G. Adding a catalyst Has NO effect on equilibrium

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle

END OF CHAPTER 16