Example The value of the equilibrium constant, K p , for the reaction given below is 1.9 x 10 3 atm -1 . Find the value of K at 25 ˚C. 2NO (g) + Cl 2 (g) 2NOCl (g)
What can K eq tell us? • Possible results of mixing similar concentrations of reactants together: • The reaction will go almost to completion • The reaction will barely go at all • The reaction will go about halfway • What would the value of K eq be (generally) in each case?
Homogeneous vs. Heterogeneous Equilibria • Homogeneous equilibria: all reactants and products are in one phase. • Heterogeneous equilibria: one or more reactants or products are in a different phase. • Consider: • experimentally, the equilibrium amount of CO 2 does not seem to depend on the amounts of CaO and CaCO 3 . Why? CaCO 3 ( s ) CaO( s ) + CO 2 ( g )
Which species belong in the equilibrium constant expression for the reaction below? 6CO 2 (g) + 6H 2 O(l) ↔ C 6 H 12 O 6 (s) + 6O 2 (g) A. O 2 only B. CO 2 and O 2 only C. C 6 H 12 O 6 , CO 2 and O 2 D. H 2 O, C 6 H 12 O 6 , CO 2 and O 2 A question
What is the equilibrium constant expression for the reaction below? 6CO 2 (g) + 6H 2 O(l) ↔ C 6 H 12 O 6 (s) + 6O 2 (g) [O 2 ] 6 [CO 2 ] 6 K c = A question
What else can we do with K eq ? • We’ve already seen how to write K eq and use it to predict the extent of a reaction • We can also use it to: • predict the direction of a reaction • calculate equilibrium concentrations
Predicting the direction of a reaction • Q = K eq only at equilibrium • If Q < K eq , forward reaction proceeds • Products are formed, reactants are consumed • If Q > K eq , reverse reaction proceeds • Products are consumed, reactants are formed.
Solving for equilibrium concentrations • If we know K eq and the initial reactant and product concentrations, we can calculate the equilibrium concentrations of all species • We need to know the appropriate K eq equation and use an ICE table Reactants Products I nitial Concentration C hange in Concentration E quilibrium Concentration
Using the ICE Table • For the reaction H 2 (g) + I 2 (g) 2 HI (g) , starting out with 0.05 M of each reactant and no product, the table would look like this: H 2 I 2 2 HI Initial Concentration 0.05 M 0.05 M ---- Change in Concentration -x -x +2x Equilibrium Concentration 0.05 M – x 0.05 M – x 2x +
• The equilibrium constant for the following reaction is K C = 5.88 x 10 -3 at 25°C: N 2 O 4 (g) 2 NO 2 (g) 15.6 g of N 2 O 4 is placed in a 5.00 L flask. Calculate the amount (number of moles) of each species at equilibrium. How many moles of N 2 O 4 are in the flask at equilibrium? A. 0.045 moles C. 0.023 moles B. 0.058 moles D. 2 moles Problem
problem For the following reaction, K C = 2.19 x 10 -10 at 100°C. COCl 2 (g) CO(g) + Cl 2 (g) If the initial concentration of COCl 2 is 4.50 x 10 -2 M, what is the equilibrium concentration of the products?
Le Châtelier’s Principle Le Châtelier’s Principle: if a system at equilibrium is disturbed, the system will move in such a way as to counteract the disturbance.
- Fall '16
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