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Presentation on theme: "CHAPTER 14 CHEMICAL EQUILIBRIUM"— Presentation transcript:


2 I. Chemical Equilibrium
A. Definition / Explanation When a chemical reaction occurs, the conc. of reactants decrease and the conc. of products increase. Eventually chemical equilibrium is reached, in which the concentration of the products and reactants remain constant. “No net change in reaction mixture”. This does not mean [reactants] = [products] !!!

3 B. Achieving Equilibrium
1. Equilibrium is not reached instantaneously. How fast or slow determined by kinetics. 2. How much product and reactants are present at equilibrium is determined by thermodynamics (energetics of reaction). What is meant by “product-favored” reaction? What is meant by “reactant-favored” reaction?

4 C. Dynamic Equilibrium Chemical reaction reach dynamic equilibrium. State of a reaction in which its forward and reverse reactions occur at equal rates so that the conc. of the reactants and products do not change with time. ** Reactants and products formed at same rate. ** Does not mean [reactants] = [products].

5 Note following diagram
D. Direction of Approach Dynamic equilibrium is independent of the direction of approach. For a given temperature, the equilibrium state will be the same whether the reaction is approached from starting materials of pure reactants or pure products. Consider rxn: N2(g) + 3 H2(g)  2 NH3(g) Note following diagram Explain idea


7 E. Catalysts Do No Effect Equilibrium Conc.
They effect rate, but not equilibrium state !!!

8 How Can We Quantitatively Describe Equilibrium?
… the expression with products in numerator, reactants in denominator, where each concentration is raised to the power of its coefficient, appears to give a constant. Regardless of the starting concentrations; once equilibrium is reached …

9 II. Equilibrium Constant
A. Equilibrium Expression “Mass Action Expression” For rxn: aA + bB  cC + dD ** No units ** Exponents always = reaction coefficients

10 B. Species Not in Equilibrium Expression
1. Pure solids – (NaCl(s), Na(s) 2. Pure liquids - H2O(l) 3. Solvents - may be involved in reaction, but its concentration does not change. (dilute solutions) **General Rule – only include substances that are (g) or (aq)

11 Write equilibrium expressions for the following reactions:
1. CO(g) + 3 H2(g)  CH4(g) + H2O(g) 2. H2(g) + I2(g)  2 HI(g) 3. AgCl(s)  Ag+(aq) + Cl-(aq) 4. NH3(aq) + H2O(l)  NH4+(aq) + OH- (aq)

12 C. Kc for Related Reactions
1. Changing Stoichiometric Coefficients N2(g) H2(g)  2 NH3(g) Kc = 3.5 x 108 What if we alter the coefficients in equation? 2 N2(g) H2(g)  4 NH3(g) K’c = ? When the coefficients of an equation are multiplied by a common factor “n”, we raise the original Kc value to the power “n” to obtain the new equilibrium constant.

13 2. Changing Direction of Reaction
N2(g) H2(g)  2 NH3(g) Kc = 3.5 x 108 What if we consider the reverse reaction? 2 NH3(g)  N2(g) H2(g) K’c = ? When we reverse the equation for a chemical reaction for which the equilibrium constant is Kc, we invert the equilibrium constant. The reverse reaction has the equilibrium constant 1 / Kc.

14 D. Kc For Combining Reactions A + B  C Kc1 = 2.0 x 10-4
C + B  F Kc2 = 1.0 x 10-6 What is the Kc for the reaction? A B  F Kc = ? When we add the equations for individual reactions to obtain an overall equation, we multiply their equilibrium constants to obtain the equilibrium constant for the overall reaction.

15 E. Kc For Gas Phase Reactions
If all reactants and products are in gas phase, the concentration terms are replaced by partial pressures. (Derived from Ideal Gas Law, PV=nRT ) When Solving, Units Must Be in atm !!!!! What is the Kp expression for the rxn: 2 H2(g) + O2(g)  2 H2O(g)

16 What is the relationship between
Kp and Kc? Δn = (sum of moles of gas products) – (sum of moles of gas reactants) As determined from coefficients in chemical reaction.

17 F. Types of Equilibrium Expressions
Kc Type Reference Table Kc general rxns Ka acid Acid Table Kb base Base Table Kw water Ksp slight Solubility Product solubility Table

18 III. Meaning of Kc Kc Value Reaction very large product favored
Kc > 1 Larger Kc: more product favored very small reactant favored Kc < 1 Smaller Kc : more reactant favored Kc = 1 significant amounts of both reactants and products

19 1. If Kc = 1, does that mean that the concentration of products equal the concentrations of reactants? 2. Which reaction has the greatest tendency to go to completion (product favored)? a. 2 NO2(g)  N2O4(g) Kc = 1.7 x 102 b. N2(g) + O2(g)  2 NO(g) Kc = 1.7 x 10-3

20 IV. Equilibrium Problems
A. Solving For Equilibrium Constant 1. Guidelines a. Write equilibrium expression for reaction. b. Construct an “ICE” table with the Initial, Change, and Equilibrium concentrations of all reactants and products.

21 When determining concentrations for table:
1) Units of conc. = Molarity = mol / L (may use pressures if all gases and Kp) 2) You may need to use “x” to represent unknown changes in concentrations in table. 3) Remember changes in concentration of reactants and products occur in same ratio as equation coefficients. c. Always plug in equilibrium concentrations into the equilibrium expression to solve the problem.

22 2. Problems – Solving for Kc Value
1. At equilibrium, a 2.0 L flask is found to contain 0.10 mol of CO, 0.20 mol of CO2 , and 4.8 g of carbon. C(s) + CO2(g)  2 CO(g) What is Kc for this reaction?

23 2. A student placed 1. 00 mol of CO and 2. 00. mol of H2 in a 2
2. A student placed 1.00 mol of CO and 2.00 mol of H2 in a 2.00 L flask at 7800C. After reaching equilibrium, the CO concentration is found to be M. Determine the value of the equilibrium constant for this reaction. CO(g) H2(g)  CH3OH (g)

24 Predicting Direction of a Reaction 1. Reaction Quotient, Q
“Q” has the same form as Kc expression, but uses the actual concentrations in the reaction mixture. “Q” is not a constant. 2. Significance of Q By comparing the Reaction Quotient to the Kc value, one can determine the direction a reaction will occur as it proceeds towards equilibrium.

25 If Q > Kc [Product] greater than would be at equilibrium
Reaction would proceed to the left. If Q < Kc [Product] less than would be at equilibrium Reaction would proceed to the right. If Q = Kc Reaction is at equilibrium. Dynamic equilibrium. No net change occurs.

26 Example Problem A 50.0 L reaction vessel contains 1.00 mol N2, 3.00 mol H2, and mol NH3. Will more ammonia, NH3, be formed or will it dissociate when the mixture goes to equilibrium at 400oC? The reaction is: N2(g) H2(g)  2 NH3(g) Kc = 0.500

27 B. Solving for Equilibrium Concentrations
1. Guidelines a. Follow same guidelines as with Kc calculations. b. Additional Comments 1) Some equilibrium concentrations may contain a number (n) plus or minus “x”, (i.e., n+x, n-x). 2) (When Kc is very small or large, 10-5 > Kc > 105), “x” may be very small compared to “n” such that simplification can be made to problem. n ± x ≈n. 3) Verify assumption once “x” is determined to see that it is less than (± 5%) of “n”. 5 % RULE 4) If assumption is not true, you will need to solve problem using the quadratic equation (see later).

28 1. Problem Solved Since Perfect Square (or Quadratic Equation)
A 0.20 mol sample of H2 and a 0.20 mol sample of I2 were placed into a 1.0 L flask. What were the equilibrium concentrations of H2, I2, and HI? Reaction temperature was 500K. H2(g) + I2(g)  2 HI(g) Kc = 49.5

29 Another Example Problem:
Given the equilibrium reaction at 745 K: H2(g) + I2(g)  2 HI(g) Kc = 50.0. Suppose that 0.75 mol HI(g), mol H2(g), and mol I2(g) are placed into a 20.0 L flask and heated to 745 K. Calculate the equilibrium concentrations of all three substances.

30 2. Problems Solved Using Assumptions
(or Quadratic Equation) A 2.0 L flask contained initial concentrations of M N2 and M O2. What is the equilibrium concentration of NO in the flask? N2(g) + O2(g)  2 NO(g) Kc = 4.8 x 10-31

31 3. Problem Solved Using Quadratic Equation
(can only solve by using Quadratic Equation) Consider the following reaction at 4480C: I2(g) + Cl2(g)  2 ICl(g) Kp = 50.5 The reaction mixture initially contains pressure of I2 = 60.0 atm and pressure of Cl2 = atm. Determine the equilibrium pressure of ICl(g).

32 VI. Shifting Chemical Equilibriums
A. Le Chatelier’s Principle 1. Definition – When any change in concentration, temperature, pressure, or volume is imposed on a system at equilibrium, the system responds by attaning a new equilibrium condition that minimizes the impact of the imposed change. ** Reaction tries to maintain equilibrium state. ** Significance

33 2. Factors (Stress) Affecting Equilibrium
a. Changing Concentrations Reaction will try to maintain equilibrium or maintain a similar Kc value. A + B  C + D Adding reactants: shifts equilibrium to right Adding products: shifts equilibrium to left Removing reactants? Removing products? Will solids or pure liquids affect equilibrium? Significance

34 Le Chatelier’s Principle Related to “Q” and “Kc”

35 b. Changing Pressure in Gas Equilibria
Pressure increases (or volume decreases) for gaseous equilibrium cause shift in equilibrium. Think about it like a concentration factor. An increase in pressure (or decrease in volume) always drives the reaction in the direction of the fewer # of gas molecules. How will increase in pressure affect equilibrium? N2(g) H2(g)  2 NH3(g) H2(g) + I2(g)  2 HI(g)

36 N2O4(g)  2 NO2(g) Initial When pressure is increased …
… to give one molecule of N2O4, reducing the pressure increase. … two molecules of NO2 combine …

37 c. Changes in Temperature
When considering exothermic or endothermic reactions, consider heat as a product or reactant. (Use same principles as changes in concentration.) Add heat to exothermic: shifts equilibrium to left Add heat to endothermic: shifts right N2(g) H2(g)  2 NH3(g) ΔH = kJ N2(g) H2(g)  2 NH3(g) (+ heat on which side??)

38 How will the following factors affect the
B. Example Problem How will the following factors affect the final equilibrium amount of Cl2 ? PCl5(g)  PCl3(g) + Cl2(g) ΔH = 21 kcal/mol a. Add PCl3 b. Remove PCl5 c. Increase temperature d. Decrease volume e. Decrease pressure f. Adding inert gas to mixture


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