1. Chemical Kinetics (Reaction Rates) and Equilibrium

2. Standards: Reaction Rates
8a. rate of reaction is the decrease in concentration of reactants or the increase in concentration of products with time.
8b. reaction rates depend on such factors as concentration, temperature, and pressure.
8c. the role a catalyst plays in increasing the reaction rate.

3. Chemical Reactions and Kinetics
* What is your favorite dish? What are the main ingredients and how do you prepare this dish?
Lets connect to Chemistry
*What is a chemical reaction?
*What does the word “kinetics” remind you of?
*What are the parts of a chemical reaction?

4. REVIEW: Parts of a chemical reaction
Reaction #1: A + B  AB
Reactants
Product
Reaction #2: 2A + B  A2B
Coefficient
Subscript

5. REVIEW: Types of a chemical reaction
Direction: Match the chemical reaction with the name of the type of reaction
1. Double Replacement
2. Single Replacement
3. Combustion
4. Synthesis Reaction
D. Na3PO4 + 3 KOH  3 NaOH + K3PO4
5. Acid-Base
6. Decomposition
A. P O2  2 P2O3
B. 2 NO2  2 O2 + N2
C. Pb + FeSO4  PbSO4 + Fe
E. C6H O2  6 CO2 + 6 H2O +Energy
F. HNO3 + NaOH  H2O + NaNO3

6. NaCl (s)  Na+ (aq) + Cl- (aq)
Sketch of what is happening on a molecular model of a chemical reaction (specific vs. general)
NaCl (s)  Na+ (aq) + Cl- (aq)  + Cl Na Cl
Have some colored salts and solutions to mix Na  + B A B A

7. Direction: Using the balanced equation below, create a molecular model of this chemical reaction (use 3 different colored pencils)
Mg + 2HCl  MgCl2 + H2 +

8. Collision Theory: reactants forming products using energy and proper orientation

9. Rate of reaction-- decrease in [reactants] or the increase [products] over time.
REAGENTS  PRODUCTS

10. What are the factors that determines the rate of reaction?

11. Explain how reaction rates depend on such factors as concentration, temperature, and pressure (use the pictures below to support your argument)
Pressure affects the rate of reaction, especially when you look at gases. When you increase the pressure, the molecules have less space in which they can move. That greater density of molecules increases the number of collisions
When you raise the temperature of a system, the molecules bounce around a lot more because they have more energy. When they bounce around more, they are more likely to collide.
If there is more of a substance in a system, there is a greater chance that molecules will collide and speed up the rate of the reaction. If there is less of something, there will be fewer collisions and the reaction will probably happen at a slower speed.
Low Pressure
High
Pressure

13. -- the frequency of collisions between reactants
Catalysts
the frequency of collisions between reactants
-- rearrange the orientation of reactants
-- speeds up reaction, not consumed
Ex. biological enzymes, catalase etc.
--reducing intramolecular bonding within reactant molecules, or donating electron density to the reactants.

14. Energy Diagram
*Reaction needs to surpass activation energy (Ea)

15. AP Chem Kinetics Activity: Rate Order
Activity 1: The Burning Candle Activity 2: The Empty Bucket Activity 3: Penny Decomposition Goals for each activity: *Graph data and determine which one corresponds to zero/1st/2nd rate order *What are the formulas for each rate order

16. AP Chem Kinetics Activity: Decomposition of Crystal Violet with NaOH (Colorimetry)
Goal: Determine the rate order of the decomposition of crystal violet by colorimetry (if possible spectro-vis) Direction: *Groups will collect at least 3 data time points *Record time and height/ calculate concentration *Materials: --Crystal Violet 25µM and NaOH 0.2M
CV + NaOH Decomposing
-- add 10 mL from class set-up
Standard
--CV 6mL + 4mL of water
(concentration =1.5µM)

17. Reaction Mechanism
*Rate determining step-- gives us the rate law (slow step) hint: look at coefficient (ex. A2 because A + A) *Intermediate-- it is produced then consumed *Catalyst – it is consumed first then produced reaction favors simpler (bi-molecular) 1st rxn: A + A  E If A + B  C 2nd rxn: E + B  G 3rd rxn: G  C + A *To check mechanism, if 1st rxn is slow, what is the rate law, is it similar to the predicted rate law expression? If so, then the mechanism is correct. *Draw molecular model for each rxn step.

18. Reaction Mechanism
Rate = k [H2O2] m

19. Kinetics Practice Problem
The reaction A + B  (products) is known to follow the mechanism:
A + B  C + E
A + E  G
G + A  D + F
If the reaction rate is observed to be the first order in A and first order in B, write the rate law for the reaction.  
Which step must be the slow step?
If step III were the slow step, what would the rate law become?  
Which letters from the reaction mechanism are used to represent the products of this reaction?
Which letters from the reaction mechanism are used to represent the unstable intermediate products?

20. Kinetics Practice Problem
For the reaction A + B  C the following rate data were obtained at constant temperature:
Exp. Rate (moles/sec) [A]0 [B]0
I x
II x
III x  
Determine the rate order and write the rate law for this reaction.
b) Calculate the rate constant, k. Include units.
c) If each reactant has a molarity of 0.3M, determine the rate
d) Propose a reasonable TWO step mechanism for this reaction and designate the slow step in the reaction.
_________________________
_______________________

21. Chemical Equilibrium Chp. 14 Concept Map
*Equilibrium Constant
*Heterogeneous Equilibrium
*Homogeneous Equilibrium
*Law of Mass Action
*Le Chatelier’s Principle
*Physical Equilibrium
*Reaction Quotient
*Reverse Reaction
*Forward Reaction
*K<1 vs. K>1
*Kp vs. Kc
*Multiple Equilibria
*Predicting Direction of Reaction
*How to calculate Eq. Concentration
*Factors that affect Chem Eq.

22. Standards: Equilibrium
9a. Students know how to use Le Chatelier's principle to predict the effect of changes in concentration, temperature, and pressure.
9b. Students know equilibrium is established when forward and reverse reaction rates are equal.
9c. Students know how to write and calculate an equilibrium constant expression for a reaction.

23. What can we conclude from each chart?
How are the two charts defining what it means for a reaction to be in equilibrium?

24. Which reaction is in equilibrium? Explain why.
H2 + F HF
HCl + NaOH H2O + NaCl
Equilibrium --established when forward and reverse reaction rates are equal and the concentration of reactants and products remain constant

25. Equilibrium Activity Q: endothermic or exothermic?
*Each student picks an index card (attach with paper clip on their shirt or ID badge). Card is either reactant A or reactant B
A + B+ Heat A-B
Q: endothermic or exothermic?

26. Factor that affects equilibrium
For each of the following situations, determine which side of the reaction is favored. Explain your choice.
Factor that affects equilibrium
Which side of equation is favored (Product or Reactant) AND which way is favored (Forward or Reverse Reaction)
Effect on K value
(Explain)
1. Add more products (change concentration)
2. Add more heat for an endothermic reaction (hint: which side of the equation will heat be part of)
3. Increase the pressure in the container that contains this reaction
N2(g)+3H2(g) 2NH3(g)

27. Disturbance # 1. Take away products
A + B+ Heat A-B
Disturbance # 1. Take away products

28. Disturbance # 2. Cool down this endothermic reaction
A + B+ Heat A-B
Disturbance # 2. Cool down this endothermic reaction

29. A (g) + B (g) + Heat A-B (g)
Disturbance # 3.
Decrease the pressure in the container that contains this reaction

30. Le Chatelier’s Principle
In a reversible chemical reaction, when stress is applied to a system in an equilibrium, the reaction will shift in a direction that relieves the stress and a new equilibrium will be established. Stress: Changing Concentration Pressure Temperature

31. Graphing Equilibrium:
Ex. The Haber Process is used to produce ammonia.
Direction: Create 2 graphs, X-axis is time, and Y-axis is concentration. For Graph #1: plot the forward reaction and Graph #2: plot the reverse reaction (label your lines with specific compounds)

32. How is Kinetics Related to Equilibrium?
Consider the rxn:
Write the rate law for the forward and reverse reaction (1st order)
Rate forward =
Rate reverse =
At equilibrium:
The EQ constant of the reverse reaction is simply the inverse/reciprocal of the EQ constant of the forward reaction.

33. Equilibrium Expression and Calculation Favors this side of reaction
Using the reaction: bR#1 + aR#2  cP
Keq= [products]c
[reactant #1 ]b [reactant #2]a
K eq
Favors this side of reaction
Keq>>>1
Keq<<<1
Products
Reactants

34. Additional Info:
*In heterogeneous equilibrium, we DON’T include solids and liquids in calculating Keq. Why? *If temperature is constant, then partial pressure of a gas directly related to the concentration (mol/L) *Kp = Kc (RT)Δn where R = Univ. gas constant ( L.atm/K.mol) Δn = moles of gas products – moles of gas reactants

35. Practice Problem
1. Which statement correctly describes a chemical reaction at equilibrium? (A) The concentrations of the products and reactants are equal. (B) The concentrations of the products and reactants are constant. (C) The rate of the forward reaction is less than the rate of the reverse reaction. (D) The rate of the forward reaction is greater than the rate of the reverse reaction.

36. AP Chem FR Practice #1
REMINDERS: Solids and liquids are not counted/ Partial pressure is proportional to # moles
NH4Cl(s)  NH3(g) + HCl(g) H = kilocalories
Suppose the substances in the reaction above are at equilibrium at 600K in volume V and at pressure P. State whether the partial pressure of NH3(g) will have increased, decreased, or remained the same when equilibrium is reestablished after each of the following disturbances of the original system. Some solid NH4Cl remains in the flask at all times. Justify each answer with a one-or-two sentence explanation.
(a) A small quantity of NH4Cl is added.
(b) The temperature of the system is increased.
(c) The volume of the system is increased.
(d) A quantity of gaseous HCl is added.
(e) A quantity of gaseous NH3 is added.

37. AP Chem FR Practice #1: ANSWER
(a) PNH3 does not change. Since NH4Cl(s) has constant concentration (a = 1), equilibrium does not shift.
(b) PNH3 increases. Since the reaction is endothermic, increasing the temperature shifts the equilibrium to the right and more NH3 is present.
(c) PNH3 does not change. As V increases, some solid NH4Cl decomposes to produce more NH3. But as the volume increases, PNH3 remains constant due to the additional decomposition.
(d) PNH3 decreases. Some NH3 reacts with the added HCl to relieve the stress from the HCl addition.
(e) PNH3 increases. Some of the added NH3 reacts with HCl to relieve the stress, but only a part of the added NH3 reacts, so PNH3 increases.

38. Reaction Quotient
The reaction quotient, Q is similar to the equilibrium constant, but describes where the reaction is at a point in time.
When Q > Keq:
not at equilibrium, too many products
reaction will move to the left, products will form reactants
When Q < Keq:
not at equilibrium, too many reactants
reaction will move to the right, reactants will form products
When Q = Keq: system at equilibrium

39. Calculating concentrations at equilibrium  ICE Method (see pg
Calculating concentrations at equilibrium  ICE Method (see pg for steps and example)
Ex. Calculate concentrations at equilibrium if you have 1 M of each reactant in this reaction: Step 1: A + B 2AB
ICE A B 2C
INITIAL (M)
CHANGE (M)
EQUILIBRIUM (M) 1 1 -x -x
+2x
1-x
1-x
0+2x

40. Calculating concentrations at equilibrium  ICE Method (see pg
Calculating concentrations at equilibrium  ICE Method (see pg for steps and example)
Step 2: Based on your ICE chart, set up your Keq Step 3: Solve for X  you might have to use the quadratic equation Step 4: Evaluate x-values (you will have 2 one will follow the parameters set in the problem) Step 5: Plug in x-value to solve for concentrations at equilibrium

41. AP Chem: Practice FR #2
CO2(g) + H2(g) H2O(g) + CO(g) When H2(g) is mixed with CO2(g) at 2,000 K, equilibrium is achieved according to the equation above. In one experiment, the following equilibrium concentrations were measured. [H2] = 0.20 mol/L [CO2] = 0.30 mol/L [H2O] = [CO] = 0.55 mol/L (a) What is the mole fraction of CO(g) in the equilibrium mixture? (b) Using the equilibrium concentrations given above, calculate the value of Kc, the equilibrium constant for the reaction. (c) Determine Kp in terms of Kc for this system. (d) When the system is cooled from 2,000 K to a lower temperature, 30.0 percent of the CO(g) is converted back to CO2(g). Calculate the value of Kc at this lower temperature. (e) In a different experiment, 0.50 mole of H2(g) is mixed with 0.50 mole of CO2(g) in a 3.0-liter reaction vessel at 2,000 K. Calculate the equilibrium concentration, in moles per liter, of CO(g) at this temperature.

42. AP Chem: Practice FR #1 Solution
(a) CO = = 0.34 (b) Kc = = 5.04 (c) Since Δn = 0, Kc = Kp (d) [CO] = % = = M [H2O] = = M [H2] = = M [CO2] = = M K = (0.385)2/(0.365 X 0.465) = 0.87 (e) let X = Δ[H2] to reach equilibrium [H2] = 0.50 mol/3.0L - X = X [CO2] = 0.50 mol/3.0L - X = X [CO] = +X ; [H2O] = +X K = X2/( X)2 = 5.04 ; X = [CO] = 0.12 M

43. MgF2(s)  Mg2+(aq) + 2 F-(aq)
AP Chem: Practice FR #3
MgF2(s)  Mg2+(aq) + 2 F-(aq)
#2.) In a saturated solution of MgF2 at 18°C, the concentration of Mg2+ is 1.21X10-3 molar. And F- is 2.42X10-3. The equilibrium is represented by the equation above.
(a) Write the expression for the solubility-product constant, Ksp, and calculate its value at 18°C.
(b) Calculate the equilibrium concentration of Mg2+ in liter of saturated MgF2 solution at 18°C to which mole of solid KF has been added. The KF dissolves completely. Assume the volume change is negligible.

44. AP Chem: Practice FR #2 Solution
(a) Ksp = [Mg2+][F-]2 = (1.21X10-3)(2.42X10-3)2= 7.09X10-9
(b) X = concentration loss by Mg2+ ion
2X = concentration loss by F- ion
[Mg2+] = (1.21X X) M
[F-] = ( X X) M
since X is a small number then ( X X)  0.100
Ksp = 7.09X10-9 = (1.21X X)(0.100)2
X = X10-3
[Mg2+]= 1.21X X´10-3 = 7.09X10-7M

45. AP Chem: Practice FR #4
C(s) + H2O(g)  CO(g) + H2(g) Hº = +131kJ #2. A rigid container holds a mixture of graphite pellets (C(s)), H2O(g), CO(g), and H2(g) at equilibrium. State whether the number of moles of CO(g) in the container will increase, decrease, or remain the same after each of the following disturbances is applied to the original mixture. For each case, assume that all other variables remain constant except for the given disturbance. Explain each answer with a short statement. (a)Additional H2(g) is added to the equilibrium mixture at constant volume. (b)The temperature of the equilibrium mixture is increased at constant volume. (c)The volume of the container is decreased at constant temperature. (d) The graphite pellets are pulverized.

46. AP Chem: Practice FR #4 Solution
(a) CO will decrease. An increase of hydrogen gas molecule will increase the rate of the reverse reaction which consumes CO. A LeChatelier Principle shift to the left.
(b) CO will increase. Since the forward reaction is endothermic (a ΔH > 0) an increase in temperature will cause the forward reaction to increase its rate and produce more CO. A Le Chatelier Principle shift to the right.
(c) CO will decrease. A decrease in volume will result in an increase in pressure, the equilibrium will shift to the side with fewer gas molecules to decrease the pressure, , a shift to the left.
(d) CO will remain the same. Once at equilibrium, the size of the solid will affect neither the reaction rates nor the equilibrium nor the concentrations of reactants or products.