1. Kinetics & Equilibrium
PACKET #11:
Kinetics & Equilibrium
Textbook: Chapter 17
Reference Table: I & F

2. Kinetics
concerned with rates of reaction (how fast, how many moles are consumed or produced in a unit of time) and mechanism (steps in a chemical reaction).
Recall that kinetic energy is the energy related to movement

3. Collision Theory
In order for a reaction to occur, the particles of the reactant must have enough energy, and must collide at the correct angles (proper orientation).
The collision theory explains the factors that affect the rate of a reaction.
The greater the rate of effective collisions, the greater the reaction rate is
The collision theory explains the factors that affect the rate of reaction

4. Factors that Affect the Rate of a Chemical Reaction
The following factors can increase the rate of a chemical reaction by increasing the number of effective collisions that a occur at a given time
Concentration
Temperature
Surface Area
Nature of Reactants
Catalyst

5. Concentration
When the concentration of one or more of the reactants increases, the rate of reaction increases (the greater the concentration, the more frequent the collisions)
When discussing gases, the greater the vapor pressure, the greater the concentration of the gas, and therefore an increase in reaction rate
Steel Wool glowing
in air (20% O2)
Steel wool burning
vigorously in pure O2

6. Temperature An increase in temperature increases the rate of reaction.
Increasing temperature causes particles to move faster (increase kinetic energy), and therefore more collisions (also more effective collisions)

7. Surface Area
Increasing surface area increases the rate of reaction. The larger the surface area, the more exposed the particles that can react (the more collisions that can occur).
Example: powdered Mg in HCl will react faster than a solid piece of Mg because powdered Mg has more surface area.

8. Nature of Reactants
Ionic substances react very fast because the ions just move close to each other and can collide at any angle. Recall that ionic bonds are when electrons are transferred.
Example: Pb+2 + S-2  PbS
Covalent substances react slower because bonds are broken and new ones are formed; the particles must collide at the correct angle. Recall that covalent bonds are when electrons are shared.
Example: H2 + O2  H2O

9. Catalyst
A catalyst speeds up chemical reactions by lowering the activation energy (amount of energy needed to start a reaction) and gives a faster rate of reaction.
Particles therefore require less energy to react (to have effective collisions).
Remember in Living Environment that an enzyme is defined as a biological catalyst.

10. Heat of Reaction (Enthalpy)
The amount of heat given off or absorbed in a chemical reaction. Heat of reaction is the difference in heat content of the products and reactants. Table I lists common reactions, and there heat of reactions.
∆H = Hproducts - Hreactants
∆H = heat of reaction
Hproducts = potential energy of the product(s)
Hreactants = potential energy of the reactant(s)
∆H is measured in kJ (kilojoules)

11. Heat of Reaction Exothermic Endothermic
Heat is released when bonds are formed
If the products have less potential energy or heat content than the reactants, then heat is given off and it is an Exothermic Reaction, and therefore ∆H is negative.
Endothermic
Heat is absorbed when bonds are broken
If the products have greater potential energy or heat content than the reactants, then heat is abdorbed and it is an Endothermic Reaction, and therefore ∆H is positive.

12. Table I
Table I is a chart that shows how much heat (∆H in kJ) is given off or absorbed in various chemical the reaction.
The reactions with a negative ∆H are exothermic reactions, and the reactions with a positive ∆H are endothermic.
Remember that heat travels from higher temperatures to lower temperatures. Look at the reaction on Table I that starts with NaOH (sodium hydroxide). The reaction is exothermic according to the table; therefore, when it is dissolved into water, heat travels from the NaOH to the water which will cause the water to get hotter.

13. Potential Energy Diagrams

14. Potential Energy Diagrams
If the potential energy of the reactants are greater than the potential energy of the products, then the reaction is EXOTHERMIC because the ∆H is a negative number
∆H = Hproducts - Hreactants

15. Potential Energy Diagrams ∆H = Hproducts - Hreactants
If the potential energy of the reactants are less than the potential energy of the products, then the reaction is ENDOTHERMIC because the ∆H is a POSITIVE number
∆H = Hproducts - Hreactants

16. Spontaneous Reactions
A reaction that takes place under a specific set of conditions spontaneously.
Occur in the direction of:
Less energy (lower enthalpy): favors exothermic reactions.
Greater entropy (disorder):
Solids have the least entropy (most order), liquids have more, and gases have the most entropy (disorder).
When a solid dissolves in water (salts or sugars), entropy increases.
At low temperature, energy is important; at high energy, entropy is important.

17. (double arrow represents an equilibrium reaction)
Every potential energy diagram depicts a reaction that is proceeding form left to right, that is in the forward direction. The reactants first collide to form the activated complex, and then they form products. Can the reverse ever happen? Can products collide to form the activated complex and become reactants? Not only can a reverse reaction occur, but both the forward and reverse reactions can occur at the same time, this is known as EQUILIBRIUM
A + B C + D
(double arrow represents an equilibrium reaction)

18. Equilibrium
When the forward and reverse reactions occur at the same rate.
The rates are equal, but the quantities (amounts) of reactants and products are not necessarily equal.
Types of Equilibrium:
Phase (Dynamic)
Solution
Chemical

19. Phase Equilibrium
Rate of particles returning to the original phase = the rate of particles escaping original phase.
Occurs at phase changes only!
Rate of melting equals the rate of freezing.
Rate of evaporating equals the rate of condensing.
Remember there may not be the same amounts of solid and liquid present, but the rate of melting will be equal to the rate of freezing

20. Example
Water in a closed system, at constant temperature, is in a dynamic equilibrium.
H2O(l) --> H2O(g) forward reaction
H2O(g) --> H2O(l) reverse reaction
At equilibrium H2O(l) <--> H2O(g)

21. Solution Eqilibrium Solids in Liquids:
At equilibrium, the rate at which sugar crystallizes (sugar on the bottom) equals the rate at which sugar dissolves. When these rates are equal it is known as a saturated solution (REMEMBER!! )
Gas in Liquid:
When the rate that gas dissolved in liquid goes up into the space above the liquid equals the rate that undissolved gas on top goes into the liquid. Can be disturbed by changes in temperature or pressure. Increase in temperature decreases solubility of gas
carbonated beverages stay carbonated best when tightly closed and chilled

22. Example
In a saturated solution of KCl(aq), at constant temperature, with excess KCl(s) the following reactions occur.
KCl(s) --> K+(aq) + Cl-(aq) forward reaction (dissolution)
K+(aq) + Cl-(aq) --> KCl(s) reverse reaction (crystallization)
At equilibrium KCl(s) <--> K+(aq) + Cl-(aq)   

23. Chemical Equilibrium
Rate of the forward reaction equals the rate of the reverse reaction.
Concentration of the products and reactants remains constant.
Ex: Haber Process
N2(g) + 3H2(g) ↔ 2NH3(g)

24. Example
What information on the graph indicates the system was initially at equilibrium?

25. Reactions Going to Completion
Some reactions go to completion; the reaction goes in only one direction, the reactants form products, products DO NOT form reactants.
Some indicators that a reaction has gone to completion when the following are produced:
(1) a gas
(2) an essentially unionized product (like water) (3) a precipitate is one of the products.
Remember that precipitates are insoluble - Table F

26. Important Facts about Equilibrium
Equilibria are dynamic.
Equilibria must occur in a closed system.
Equilibrium can only exist in reversible reactions. The reaction CANNOT go to completion.
Reaction rates are equal but not necessarily the amount of the reactants and/or products.

27. Le Chatelier’s Principle
Henry Louis Le Chatelier was a French chemist who devised a principle to predict the effect of change in conditions on a chemical equilibrium reaction
The principle states that if a system at equilibrium is subjected to stress, the equilibrium will shift in the direction that relieves the stress.
Types of stresses include: concentration, temperature, and pressure

28. Le Chatelier’s Principle
CHANGES IN CONCENTRATION:
A + B ‹—› C + D
An increase in concentration of anything on the left side (reactants) causes the reaction to go to the right (shifts the equilibrium to the right). An increase in concentration of anything on the right side (products) causes the reaction to go to the left (shifts equilibrium to the left). A decrease in concentration on the left will have the same affect as an increase on the right, and a decrease in concentration on the right will have the same affect as an increase on the left.

29. H2O(g) + CO(g) ↔ H2(g) + CO2(g)
Example
H2O(g) + CO(g) ↔ H2(g) + CO2(g)
State how the equilibrium would shift when each of the changes below are made? Each change is made independent of all others.
increase [CO]
decrease [water]
increase [hydrogen gas]
increase [carbon dioxide]
decrease [CO]

30. Example N2(g) + O2(g) ↔ 2NO(g) Given the reaction at equilibrium:
As the concentration of N2(g) increases, the concentration of O2(g) will
Decrease
Increase
remains the same

31. Le Chatelier’s Principle
CHANGES IN TEMPERATURE:
A + B ‹—› C + D + heat
Increasing the temperature on a system at equilibrium shifts the reaction so that the heat is absorbed. This favors the endothermic reaction.
When you are considering an equilibrium question with a change in temperature, it is very important that you note where the heat (energy) is. Is the reaction endothermic or exothermic.

32. CaCO3(s) + heat ↔ CaO(s) + CO2(g)
Example
CaCO3(s) + heat ↔ CaO(s) + CO2(g)
Is the forward reaction endothermic or exothermic? Why?
How would the equilibrium shift if the temperature was increased?
How would the equilibrium shift if the temperature was decreased?

33. 2HBr(g) + 17.4kcal ↔ H2(g) + Br2(g)
Example
Given the equilibrium reaction at constant pressure:
2HBr(g) kcal ↔ H2(g) + Br2(g)
When the temperature is increased, the equilibrium will shift to the
(1) right and the concentration of HBr(g) will decrease
(2) right and the concentration of HBr(g) will increase
(3) left and the concentration of HBr(g) will decrease
(4) left and the concentration of HBr(g) will increase

34. 2H2(g) + O2(g) ↔ 2H2O(g) + heat
Example
Given the reaction at equilibrium:
2H2(g) + O2(g) ↔ 2H2O(g) + heat
Which concentration changes occur when the temperature of the system is increased?
the [H2] decreases and the [O2] decreases
(2) the [H2] decreases and the [O2] increases
(3) the [H2] increases and the [O2] decreases
(4) the [H2] increases and the [O2] increases

35. Le Chatelier’s Principle
CHANGES IN PRESSURE:
2A (g) + B (g) ↔ C (g) + D (g)
Increasing the pressure in a gaseous equilibrium will shift the reaction to the side with less total moles.
Solids and liquids are NOT effected by pressure changes.

36. 2SO3(g) + heat ↔ 2SO2(g) + O2(g)
Example
2SO3(g) + heat ↔ 2SO2(g) + O2(g)
If the pressure on the system is increased, what will be the effect on the equilibrium?
If the pressure on the system is decreased, what will be the effect on the equilibrium?
If the pressure on the system is increased, how will the number of moles of oxygen change? How will the number of moles of SO3 change?

37. Example N2(g) + O 2(g) ↔ 2NO(g) Given the reaction at equilibrium:
If the temperature remains constant and the pressure increases, the number of moles of NO(g) will
Decrease
Increases
Remains the same

38. Le Chatelier’s Principle
ADDING A CATALYST:
Catalysts do NOT shift the equilibrium!
They only speed up the rate of the forward and reverse reactions equally.

39. Common-Ion Effect NaOH <-> Na+ + OH-
If you added NaCl to this equilibrium reaction what would happen to the reaction above?
In water, NaCl dissociates into Na+ and Cl-
Do any of the substances above increase with the addition of NaCl?
How can you explain in terms of Le Chatelier’s?

40. Equilibrium Constant (Keq)
When a system reaches equilibrium, it is useful to know the relative amounts of products and reactants present. Certain reactions produce hardly any products, while other reactions leave hardly any reactants remaining. The equilibrium constant (Keq) provides us with this information.
Keq = [C]c [D]d
[A]a [B]b
When writing an equilibrium constant expression, we omit the concentration expressions for solids and liquids!! Use only gas and aqueous.
aA + bB
cC + dD

41. Example 1. Write the Keq expressions for this reaction:
N2(g) + 3H2(g) <-> 2NH3(g).
(b) Evaluate Keq if:
[NH3] = 0.2 mole/liter
[N2] = 0.04 mole/liter
[H2] = 0.01 mole/liter
2. Write the equilibrium expression for this system:
S(s) + O2(g) <-> SO2(g)

42. Equilibrium Constant (Keq)
A large Keq (>1) means that the reaction goes almost to completion; at equilibrium there is a high concentration of product and we say that equilibrium “lies to the right”.
A small Keq (<1) means that the reaction is not very favorable; at equilibrium there are still a lot of reactants around and we say that the equilibrium “lies to the left”.

43. Reaction Mechanisms
Recall Hess’ Law states that if a reaction is carried out in a series of steps, ∆H for the reaction will be equal to the sum of the enthalpy changes for the individual steps.
A chemical reaction is actually a series of steps, and only the steps that are “slow” can account for any changes in the rate of reaction (rate-determining step). “A chain is only as strong as its weakest link” (worksheet).

44. Solubility Product Constant (Ksp)
The solubility product constant (Ksp) is a special type of equilibrium constant that measures the concentrations of ionic compounds in water.
A compound with a small Ksp value is only slightly soluble in water.
Example
Ksp for CaSO4 = 9.1 x 10-6
Ksp for BaSO4 = 9.1 x
Which of the above ionic compounds is less soluble in water?

45. Example Write the solubility equation and Ksp expression for BaSO4
Write the solubility equation and Ksp expression for PbI2
RECALL:
Ionic compounds are made of ions, break down the compounds into their appropriate aqueous ions

46. Solubility Product Constant (Ksp)
General Dissociation:
AmBn(s)  mAn+ (aq) + nBm- (aq)
Ksp = (mx)m(nx)n - x is the molar solubility
The solubility of KCl in aqueous solution was determined to be 1.6 x 10-3 mol/L. What is the value of the Ksp for KCl?
KCl(s)  K+(aq) + Cl-(aq)
Ksp = (1x)1 (1x)1  (x)(x)  x2

47. Example: The solubility of Fe(OH)3 in an aqueous solution was determined to be 4.5x10-10 mol/L. What is the value for Ksp for Fe(OH)3.
Fe(OH)3(s)  Fe+3(aq) + 3OH-

48. What are the concentrations of each of the ions in a saturated solution of PbBr2, given that the Ksp of PbBr2 is 2.1 x 10-6?
PbBr2(s) 

49. Review Questions 1) Which sample has the lowest entropy?
A) 1 mole of KNO3(s)
B) 1 mole of H2O(l)
C) 1 mole of KNO3(l)
D) 1 mole of H2O(g)
2) According to the Heats of Reaction at kPa and 298 K chemistry reference table, which potential energy diagram best represents the reaction that forms H2O(l) from its elements?

50. 3) Increasing the temperature increases the rate of a reaction by
A) lowering the frequency of effective collisions between reacting molecules
B) increasing the activation energy
C) increasing the frequency of effective collisions between reacting molecules
D) lowering the activation energy
4) Given the equilibrium reaction at STP:
N2O4(g) >2NO2(g)
Which statement correctly describes this system?
A) The concentrations of N2O4 and NO2 are equal.
B) The concentrations of N2O4 and NO2 are both increasing.
C) The forward and reverse reaction rates are both increasing.
D) The forward and reverse reaction rates are equal.

51. 5) Which statement correctly describes a chemical reaction at equilibrium?
A) The rate of the forward reaction is less than the rate of the reverse reaction.
B) The concentrations of the products and reactants are equal.
C) The concentrations of the products and reactants are constant.
D) The rate of the forward reaction is greater than the rate of the reverse reaction.
6) Given the reaction at equilibrium:
C2(g) + D2(g) > 2CD(g) + energy
Which change will cause the equilibrium to shift?
A) increase in volume
B) increase in pressure
C) addition of a catalyst
D) addition of heat

52. 7) The potential energy diagram below represents a reaction.
Which arrow represents the activation energy of the forward reaction?
A) A B)B C) C D) D
8) Given the reaction:
CH4(g) + 2O2(g)  2H2O(g) + CO2(g)
What is the overall result when CH4(g) burns according to this reaction?
A) Energy is absorbed and DH is positive.
B) Energy is released and DH is positive.
C) Energy is absorbed and DH is negative.
D) Energy is released and DH is negative.

53. 9) Given the reaction at 25°C:
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
The rate of this reaction can be increased by using 5.0 grams of powdered zinc instead of a 5.0-gram strip of zinc because the powdered zinc has
A) lower kinetic energy
B) more surface area
C) more zinc atoms
D) lower concentration
10) In most aqueous reactions as temperature increases, the effectiveness of collisions between reacting particles
A) increases B) remains the same C) decreases

54. 11) Which equation shows an increase in entropy?
A) CH3OH(l)  CH3OH(s)
B) CO2(g)  CO2(s)
C) CO2(l)  CO2(g)
D) CH3OH(g)  CH3OH(l)
12) Given the equilibrium reaction in a closed system:
H2(g) + I2(g) + heat <-> 2HI(g)
What will be the result of an increase in temperature?
A) The equilibrium will shift to the left and [H2] will decrease.
B) The equilibrium will shift to the right and [HI] will increase.
C) The equilibrium will shift to the right and [HI] will decrease.
D) The equilibrium will shift to the left and [H2] will increase.

55. Questions 13 and 14 refer to the following:
The diagram below represents the changes in potential energy that occur during the given reaction: A + B  C.
13) Does the diagram illustrate an exothermic or an endothermic reaction? [State one reason, in terms of energy, to support your answer.]
14) On the diagram above, draw a dashed line to indicate a potential energy curve for the reaction if a catalyst is added.