1. Chapter 17
CHEMICAL EQUILIBRIUM

2. REVERSIBLE REACTIONS Reactions can proceed in both directions.
This is usually indicated by: N2O4 (g)  2 NO2 (g)
A chemical reaction in which the products regenerate the original reactants is called a REVERSIBLE REACTION.
In theory, all reactions are reversible.
Some are reversible on their own and others are reversible only under restricted conditions.
Do not assume that all reactions proceed until one of the reactants is entirely consumed. GOOD TEST QUESTION!
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REVERSIBLE REACTIONS

3. Heat is absorbed when N2O4 dissociates into 2 NO2 molecules
N2O4 (g)  2 NO2 (g)
colorless  red-brown
DEMO

4. Remember that the rate of a reaction depends upon:
the NATURE OF THE REACTANTS,
the TEMPERATURE and the CONCENTRATION,
for gases, the PRESSURE at which the reaction takes place.
In equilibrium, it is important to remember that reaction rate is proportional to concentration. Reaction rate is the speed at which a reaction occurs. When the reaction goes faster, the reaction rate increases, but the actual time decreases.
CHEMICAL EQUILIBRIUM

5. CHEMICAL EQUILIBRIUM How many times have we talked about equilibrium?
What are some examples?
CHEMICAL EQUILIBRIUM

6. Chemical equilibrium is the state in which the concentrations of reactants and products remain constant with time
The rate at which they are formed in each reaction equals the rate at which they are consumed in the opposite direction.
The concentration remains constant.
Chemical equilibrium is a dynamic, or changing, process. But the net change in concentration is zero.
REMEMBER DYNAMIC EQUILIBRIUM? Here it is again!
CHEMICAL EQUILIBRIUM

7. Rateforward = Ratereverse
CHEMICAL EQUILIBRIUM
Rateforward = Ratereverse

8. CHEMICAL EQUILIBRIUM Rateforward = Ratereverse
H2(g) + I2(g)  2HI(g)
Rateforward = Ratereverse 
Rateforward = Kf[H2][I2] Ratereverse = Kr[HI]2
 the coefficient becomes the exponent;
brackets mean concentration.
CHEMICAL EQUILIBRIUM

9. If the two rates are equal, see if your group can manipulate the two equations so the K constants are n the left and the concentrations are on the right.
Kf[H2][I2] = Kr[HI]2
CHEMICAL EQUILIBRIUM

10. CHEMICAL EQUILIBRIUM Kf[H2][I2] = Kr[HI]2 becomes Kf = [HI]2
Keq = [HI] product
Equilibrium Constant [H2][I2] reactant
CHEMICAL EQUILIBRIUM

11. Remember: products divided by reactants
TO SUM IT ALL UP
The mass action law states that if the system is at equilibrium at a given temperature, then the following ratio is a constant.
aA + bB  cC + dD
becomes
Keq = [C]c[D]d
[A]a[B]b
Remember: products divided by reactants
CHEMICAL EQUILIBRIUM

12. The Law of Chemical Equilibrium states that at a given temperature, a chemical system may reach a state in which a particular ratio of reactant and product concentrations has a constant value, Keq.
Keq is the numerical value of the ratio of product concentration to reactant concentration with each concentration raised to the power corresponding to its coefficient.
It is constant only at a specific temperature. If temperature changes, Keq changes.
CHEMICAL EQUILIBRIUM

13. Write the rate expressions for the following reactions. All are gases
Write the rate expressions for the following reactions. All are gases. Start out by writing “Keq = “ 1. 3H2(g) + N2(g)  2NH3(g) 2. SO2(g) + NO2(g)  NO(g) + SO3(g) 3. 2H2O(g)  2H2(g) + O2(g) 4. FeCl3(aq) + 3KSCN(aq)  3KCl(aq) + Fe(SCN)3(s)
CHEMICAL EQUILIBRIUM

14. CHEMICAL EQUILIBRIUM SOMETHING TO REMEMBER:
Use GASES and AQUEOUS SOLUTIONS in the equilibrium expressions. Leave everything else out. Explanation to follow.
On the Practice Problems handout provided, do #1-18 only. Due tomorrow.
CHEMICAL EQUILIBRIUM

15. DETERMINING A NUMERICAL VALUE FOR Keq
Experiments must be performed to determine the concentrations of reactants and products at equilibrium.
To determine the value of Keq, we will use the equilibrium expression we just learned.
STEPS:
Balance the chemical equation.
Write down the concentrations for all reactants and products. Use the equilibrium concentrations.
Use the relationship to solve the problem.
DETERMINING A NUMERICAL VALUE FOR Keq

16. DETERMINING A NUMERICAL VALUE FOR Keq
Table 1
Initial Initial EQ EQ Keq
Trial [R]M [P]M [R]M [P]M ?
Calculate Keq for the third trial. Use the equation: R  P
Keq = [P] = ( ) = 4.70
[R] (0.0310)2
Keq has no units
DETERMINING A NUMERICAL VALUE FOR Keq

17. NUMERICAL VALUE FOR Keq
EXAMPLE:
What is the equilibrium constant for the following reaction if the final concentrations are CH3COOH = 0.302M, CH3CH2OH = 0.428M, CH3CH2OOCH3 = 0.655M and H2O = 0.654M. Assume they are all gases if they do not tell you.
CH3COOH CH3CH2OH  CH3CH2OOCH H2O
2. and Keq = [CH3CH2OOCH3][H2O]
[CH3COOH] [CH3CH2OH]
= [0.655][0.654] = 3.31
[0.302][0.428]
NUMERICAL VALUE FOR Keq

18. TO REMEMBER SOME THINGS TO REMEMBER:
If Keq is small (less than one), then the reactants are favored and equilibrium is established before the product is formed.
If Keq is large (greater than one), then products are favored and equilibrium is established after lots of product is formed.
Each reaction has a unique Keq for every temperature. If the temperature changes, the Keq changes.
Keq does not indicate the time it takes to reach equilibrium. It only provides information about the mixture of reactants and products at equilibrium. GREAT TEST QUESTION!!
TO REMEMBER

19. You can now do the following: 1
You can now do the following: 1. Write a rate expression as the mass action law. 2. Solve for Keq if given the EQ concentrations 3. Solve for a missing species if given Keq and all but one EQ concentration.
WHAT YOU CAN DO…

20. It is not always obvious if a mixture has reached equilibrium or not.
If it is not at equilibrium, then it is helpful to know which direction the reaction must proceed in order to reach equilibrium.
Example:
N H2 ↔ 2 NH3
Keq = 0.105
REACTION QUOTIENT, Q

21. Suppose you are running the reaction and want to know if the reaction has reached equilibrium. You measure the concentrations to be: [NH3] = 0.15 M [H2] = 0.10 M [N2] = M The reaction quotient can be calculated by substituting these concentrations into the equilibrium expression.
REACTION QUOTIENT, Q

22. REACTION QUOTIENT, Q SOLVE FOR “Q”:
Q = [NH3]2 = (0.15) = =
[N2][H2]3 (0.0020)(0.10)3
This reaction is NOT at equilibrium. “Q” ≠ Keq.
The reaction proceeds towards the left, towards reactants.
REACTION QUOTIENT, Q

23. When trying to determine the direction for Q, compare Q to Keq
Then think of a number line.
Put Keq on the number line. Then put Q on it. Which direction do you need to travel for Q to become Keq?
Keq = 0.105
Q = 11000
REACTION QUOTIENT, Q

24. REACTION QUOTIENT, Q Deciding Direction – always compare Q to Keq:
If Q < Keq, this means that at the time of measurement, there were too much of the reactants and too little of the products. The reaction will proceed to the right, in the direction of the products.
If Q > Keq, this means that at the time of measurement, there were too little of the reactants and too much of the products. The reaction will proceed to the left, in the direction of the reactants.
If Q = Keq, the system is at equilibrium and there will be no shift in direction at this temperature.
REACTION QUOTIENT, Q

25. QUIZ FORMAT Quiz will have the following:
Write the equilibrium constant expression for any system at equilibrium.
Calculate Keq given the concentrations of products and reactants at equilibrium.
Given Keq, calculate the equilibrium concentrations of products and reactants at equilibrium (missing species).
Calculate the reaction quotient, Q, and relate it to Keq. Determine in which direction the reaction will proceed (shift left, shift right, at equilibrium).
QUIZ FORMAT

26. LECHATELIER’S PRINCIPLE
If a change in conditions is imposed on a system at equilibrium, the equilibrium position will shift in the direction that tends to reduce that change in conditions.
A STRESS is any kind of change in a system at equilibrium that upsets the equilibrium.
A reaction system will shift in the forward or reverse direction to “undo” the altering factor.
LECHATELIER’S PRINCIPLE

27. LECHATELIER’S PRINCIPLE
STRESSES
Concentration Pressure
Temperature
ANIMATION
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LECHATELIER’S PRINCIPLE

28. If you add more of a reactant, the reaction will proceed to the right
If you add more of a reactant, the reaction will proceed to the right. If you add more of a product, the reaction will proceed to the left.
If a substance is removed, its concentration decreases. The reaction will return to equilibrium by producing more of the substance that was removed.
**The position of equilibrium shifts, the equilibrium constant does not.
CONCENTRATION

29. If Q > Keq, the reaction will shift to the left in order to reach equilibrium.
If Q < Keq, the reaction will shift to the right to return to equilibrium.
Adding a substance to a system at equilibrium drives the system to consume that substance. Removing a substance from a system at equilibrium drives the system to produce more of that substance.
CONCENTRATION

30. NH4Cl (s)  NH3 (g) + HCl (g)
In gaseous systems, If you increase the pressure of the system, the system will shift to reduce that pressure by proceeding in the direction that produces the fewer molecules.
NH4Cl (s)  NH3 (g) + HCl (g)
Increased Pressure: O moles of gas on the left produce 2 moles of gas on the right. The reaction will run to the left in order to reduce the number of moles of gas produced.
PRESSURE (AND VOLUME)

31. If there is the same number of moles on both sides of a reaction in equilibrium, there will be no shift in equilibrium if pressure is added.
Pressure only affects gases, not solids and liquids.
Remember, if you change volume, you change pressure.
**The position of equilibrium shifts, the equilibrium constant does not.
PRESSURE

32. H2 (g) + I2 (g)  2 HI (g) + heat
Remember that the value of the equilibrium constant for a particular reaction depends on temperature.
Increasing the temperature causes some chemical reactions to proceed more completely to products, increasing the value of the equilibrium constant.
Sometimes increasing the temperature causes some chemical reactions to proceed less completely, lowering the value of the equilibrium constant.
H2 (g) + I2 (g)  2 HI (g) + heat
Keq (400oC) = 54.5 Keq (490oC) = 45.9
TEMPERATURE

33. H2 (g) + I2 (g)  2 HI (g) + heat
Keq (400oC) = 54.5 Keq (490oC) = 45.9
Keq is a measure of the extent to which the reaction proceeds.
In this example, raising the temperature causes the reaction to proceed less completely to products. Lowering the temperature would produce a higher yield.
TEMPERATURE

34. To understand the effect of temperature on equilibrium, you must know whether the reaction gives off heat or absorbs heat.
If heat is added to a system that is exothermic in the forward direction, the reaction tends to reestablish equilibrium by consuming the additional heat through the reverse (endothermic) reaction.
TEMPERATURE

35. Heat + NH4Cl (s)  NH3 (g) + HCl (g)
In this reaction, raising the temperature will drive the reaction in the forward direction and lowering the temperature will drive it in the reverse direction.
TEMPERATURE

36. PRACTICE Use the equation to solve the problems:
4 HCl + O2 + heat  Cl2 + 2 H2O
A. If the temperature increases, the reaction will favor formation of __________
B. If Cl2 is removed as it is formed, the reaction will favor formation of __________
C. If the pressure is increased, the reaction will favor formation of __________
PRACTICE

37. QUIZ FORMAT Quiz will have the following:
Write the equilibrium constant expression for any system at equilibrium.
Calculate Keq given the concentrations of products and reactants at equilibrium.
Given Keq, calculate the equilibrium concentrations of products and reactants at equilibrium (missing species).
Calculate the reaction quotient, Q, and relate it to Keq. Determine in which direction the reaction will proceed (shift left, shift right, at equilibrium).
QUIZ FORMAT

38. Fritz Haber studied the reaction of gaseous nitrogen and gaseous hydrogen to form ammonia.
3 H2 (g) + N2 (g) 
NH3 (g) + heat
THE HABER PROCESS

39. In reality, the reaction reached equilibrium before any sizable amount of ammonia could be formed.
Haber studied the reaction and determined the pressure and temperature of the reaction to maximize the yield of ammonia.
His work was important, but controversial, for the Germans used it in World War II to produce explosives.
THE HABER PROCESS

40. REVIEW Describe a reversible reaction.
Describe the state of dynamic chemical equilibrium and explain how it is achieved.
Write the equilibrium rate expression for any system at equilibrium.
Calculate Keq given the concentrations of products and reactants at equilibrium.
Given Keq, calculate the equilibrium concentrations of products and reactants at equilibrium.
Calculate the reaction quotient, Q, and relate it to Keq. Determine in which direction the reaction will proceed (shift left, shift right, at equilibrium).
REVIEW

41. 7. Describe Le Chatelier’s Principle and the stresses that can alter equilibrium (concentration, pressure, and temperature). 8. Use Le Chatelier’s Principle to predict shift in equilibrium in terms of relative concentrations of products and reactants. Understand what changes in temperature, pressure (volume), concentration, and moles (pressure) can do. 9. Describe the significance of the size of Keq in an equilibrium system. 10. Be able to determine the solubility of a compound if given the Ksp.
REVIEW