1. Kinetics, Potential Energy Diagrams, and Equilibrium

2. Remember Heat Calculations
We used these three formulas to determine how much heat a reaction absorbed or released

3. A Few Reminders q = mCΔT q = mHf q = mHv
Used when temp. changes
Temp change means NO phase change
q = mCΔT
q = mHf
q = mHv
The sign (+/-) in front of “q” tells us whether the system gained heat or lost heat.
When we feel the beaker as hot/cold we are feeling the surroundings.
Used for phase change
NO temp. change

4. Reactions that Release Energy
These reactions are Exothermic
- q [negative value]
the system loses heat
The surroundings heat up

5. Reactions that Absorb Energy
These reactions are Endothermic
+ q [positive value]
system gains heat
the surroundings cool down

6. Graphing Heat in Reactions
We’ve already looked at one graphical representation…
q = mHv
q = mCΔT
Temperature
q = mHf
Heat (Joules) or Time

7. What Heating Curves Tell Us
Heating Curves tell us how temperature changes when energy is absorbed or released.
For Example:
Energy Added  Kinetic Energy  ΔTemp
Energy Added  Potential Energy  Δ Phase
Heating Curves do NOT tell us how much energy was needed to get the rxn started.
Heating Curves do NOT tell us how the energy used to start the reaction compares to the energy produced by the reaction.

8. However, only a small fraction of collisions produces a reaction. Why?
We know that molecules must collide to react.
However, only a small fraction of collisions produces a reaction. Why?
Particles must collide with the correct orientation and the right amount of energy
Particles lacking the correct site or necessary kinetic energy to react bounce apart unchanged when they collide.

9. How Reactions Happen
Activation energy is the amount of energy needed to cause a reaction.
Collisions must equal or exceed the activation energy
Reactions can occur:
Very fast – such as a firecracker
Very slow – such as the time it took for dead plants to make coal
In chemistry, reaction rate is measured in terms of moles of reactant consumed (or product formed) per liter per second

10. Potential Energy Diagrams
A complete picture of the system’s
net change in energy (HEAT of REACTION) is best illustrated by a POTENTAL ENERGY DIAGRAM
In a Potential Energy Diagram potential energy of the reactants and products is drawn with reference to a time or rate sequence.
ΔH = q, (HEAT of Reaction (AKA enthalpy)
ΔH = Energy Contained in of Products – Energy contained in of Reactants

11. Ways to Express ΔH – Heats of Reaction
Endothermic Reactions
ΔH is a positive value
Example – N2 + O2 2NO ΔH = kJ
ΔH can be written as a reactant in an equation
Example - N2 + O kJ 2NO
Exothermic Reactions
1. ΔH is negative value
Example – 2CO + O2 2CO2 ΔH = kJ
2. ΔH can be written as a product in an equation
Example – 2CO + O2 2CO kJ

12. Practice Problem

13. http://player. discoveryeducation. com/index. cfm
Standard Deviants Chapter 17 – Chemistry Videos - Reaction Energy Diagrams (2:27min)
Chapter 17 – Chemistry Videos – Exothermic & Endothermic (2:39min),
Chapter 17 – Chemistry Videos - Transition Time & Energy (1:23 min)

14. Table I: Heats of Reaction
kJ = 1000 Joules
Exo-
thermic
Endo-
thermic DH
= (
PE of Products) -
(PE of Reactants)

15. 2 2

18. A POTENTIAL ENERGY DIAGRAM: C + O2 → CO2 + 395 KJ
Products have lower energy than reactants
Energy released
Exothermic
ΔH = -395kJ
(Heat of Reaction)
C + O2
Energy
395kJ
CO2
Reactants
Products

19. A POTENTIAL ENERGY DIAGRAM: CaCO3 + 176kJ → CaO + CO2
Products have higher energy than reactants
Energy absorbed
Endothermic
ΔH = +176kJ
(Heat of Reaction)
CaO + CO2
Energy
176kJ
CaCO3
Reactants
Products

20. Potential Energy Diagrams
One thing that is not represented in these simple examples…
Activation energy
Minimum amount of energy required to start a reaction

21. Endothermic with Activation Energy
However, adding 150kJ won’t make the reaction happen
100 kJ
250 kJ
Products have 350kJ of PE
Activation Energy
ΔH = Heat of Reaction
150 kJ
Reactants have 200kJ of PE

22. Endothermic with Activation Energy
Activated Complex – temporary, unstable, intermediate union of reactants
100 kJ
250 kJ
150 kJ
450 kJ
ΔH = 150kJ
Activation Energy

23. Exothermic with Activation Energy
Activation energy = 10kJ
Activation energy
Reactants
40kJ
ΔH = 25kJ ΔH
Products15kJ

24. Catalyst: something that reduces the amount of activation energy needed to start a reaction – speeds up a reaction
Activation energy with a catalyst

25. Endothermic Reaction with a Catalyst

26. Exothermic Reaction with a Catalyst

27. - Page 543

28. “Reaction Rates and Equilibrium”

29. Five Factors Affecting Reaction Rate
Nature of the Reactants: what the reactants are (solids, liquids or gases)
Temperature:
Increasing temperature always increases the rate of a reaction.
Surface Area:
Increasing surface area increases the rate of a reaction

30. Five Factors Affecting Reaction Rate
Concentration:
Increasing concentration USUALLY increases the rate of a reaction
Presence of Catalysts
Catalyst: A substance that speeds up a reaction, without being consumed itself in the reaction

31. Reversible Reactions SOME reactions do not go to completion
These reactions may be reversible
a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously
Forward: 2SO2(g) + O2(g) → 2SO3(g)
Reverse: 2SO3(g)  2SO2(g) + O2(g)

32. Reversible Reactions 2SO2(g) + O2(g) ↔ 2SO3(g)
The two equations can be combined into one, by using a double arrow, which tells us that it is a reversible reaction:
2SO2(g) + O2(g) ↔ 2SO3(g)

33. Reversible Reactions
For a reversible reaction, the rates of the forward and reverse reactions are equal
Thus a chemical equilibrium occurs
A dynamic condition in which the forward and reverse reactions are proceeding at equal rates, producing an apparent constant condition or no net change in the actual amounts of the components of the system
Equilibrium is recognized when observable changes such as color, temperature or pressure no longer occur

34. 2SO2(g) + O2(g) ↔ 2SO3(g)
What actually happens when sulfur dioxide and oxygen gas are mixed in a sealed container?
Initially the two reactants begin to react to form the product but no sulfur trioxide exists at the beginning of the reaction
The initial rate of the reverse reaction is zero
As product is formed the decomposition of sulfur trioxide begins
The reverse reaction proceeds slowly at first but it’s rate increases as the concentration of sulfur trioxide increases
Simultaneously the rate of the forward reaction decreases because sulfur dioxide and oxygen are being used up
Eventually, the synthesis and decomposition occur at the same rate
Equilibrium is reached
Both forward and reverse reactions continue at the same rate and the concentrations of the reactants and products remains the same or CONSTANT

35. Reversible Reactions Equilibrium is reached
Both forward and reverse reactions continue at the same rate and the concentrations of the reactants and products remains the same
This does not necessarily mean the concentrations of the components on both sides of the chemical equation are equal
In fact the concentrations can be dramatically different
It means the concentrations of the components remain constant

36. Factors Affecting Equilibrium
Chemical equilibrium is a delicate balance
Changes of almost any kind disrupt the balance
When the equilibrium of a system is disrupted the system makes adjustments to restore the equilibrium
However the equilibrium position of the restored equilibrium is different than the original equilibrium position
The amounts of reactants or products may have increased or decreased
The French chemist Henri Le Chatelier ( ) studied how the equilibrium position shifts as a result of changing conditions

37. Le Chatelier’s Principle
If stress is applied to a system in equilibrium, the system changes in a way that relieves the stress
Stresses on the equilibrium are considered to be:
Changes in concentration of reactants or products
Changes in temperature
Changes in pressure (only if a gas is present)

38. Le Chatelier’s Principle
Concentration
Increasing the concentration of one substance in a reaction at equilibrium will cause the reaction to proceed in the direction that will use up the increase
Adding more reactant produces more product
Removing the product as it forms will produce more product

39. Le Chatelier’s Principle
Example – Concentration Change
Decomposition of carbonic acid in aqueous solution to form the products carbon dioxide and water
Add CO2
Direction of shift
H2CO3(aq)
CO2(aq) + H2O(l)
Direction of shift
Remove CO2

40. Le Chatelier’s Principle
Temperature
Increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat – endothermic
Decreasing the temperature causes the equilibrium to shift in the direction that releases heat – exothermic
If heat is one of the products – exothermic reaction
Cooling an exothermic reaction will produce more product
Heating an exothermic reaction will shift the reaction to the reactant side of the equilibrium

41. Le Chatelier’s Principle
Example – Temperature Change
Sulfur trioxide is produced in an exothermic reaction of sulfur dioxide and oxygen
Add heat
Direction of shift
2SO2(g)
+ O2(g)
2SO3 + heat
Direction of shift
Remove heat (cool)

42. Le Chatelier’s Principle
Pressure
Changes in pressure will only effect gaseous equilibria
Increasing the pressure will usually favor the direction that results in a smaller volume caused by fewer molecules

43. Le Chatelier’s Principle
Example – Pressure Change
N2(g) + 3H2(g) NH3(g)
Increase pressure
Direction of shift
Direction of shift
Decrease pressure
For every two molecules of ammonia made, four molecules of reactant are used up – the equilibrium shifts to the right with an increase in pressure and shifts to the left with a decrease in pressure

44. N H Equilibria in the Real World
–The Haber Process – Production of Ammonia N H
Major Uses
Fertilizers (for example ammonium nitrate (NH4NO3)
Synthesis of Nitric Acid (HNO3)
Dyes
Drugs
Structure is 1 Nitrogen atom to 3 Hydrogens.
Oxidation (combustion, burning of ammonia makes nitric acid)
Ammonium nitrate is made by neutralisation of nitric acid with ammonia.

45. Haber Process - Gas Phase Equilibrium
catalyst
N2(g) + 3H2(g) <=====> 2NH3(g) + heat
high pressure and temperature
ΔH= - 92 kJ
See Reference Table I

46. The Principle of Le Chatelier
Changes in Concentration or Pressure
N2(g) H2(g) NH3(g)
an increase in N2 and/or H2 concentration or pressure, will cause the equilibrium to shift towards the production of NH3

47. The Principle of Le Chatelier
Changes in Concentration or Pressure
N2(g) H2(g) NH3(g)
likewise, a decrease in NH3 concentration will cause more NH3 to be produced
A decrease in pressure will cause more N2 and H2 to be produced

48. The Principle of Le Chatelier
Changes in Temperature
N2(g) H2(g) 2 NH3(g) + heat
for an exothermic reaction, an increase in temperature will cause the reaction to shift back towards reactants
Increase in temperature favors the endothermic reaction

49. Practice Problem
The cobalt complexes participating in the equilibrium below comprise a humidity sensor. From Le Châtelier's principle, when the sensor is moist (excess H2O), what color is the cobalt complex?
pink, blue

50. Hb(O2)4 + 4 CO (g)   Hb(CO)4 + 4O2 (g)
Practice Problem
A competition experiment involves O2 and CO vying for hemoglobin (Hb) sites, defined by the equilibrium
Hb(O2)4 + 4 CO (g)   Hb(CO)4 + 4O2 (g)
From Le Châtelier's principle, how is CO poisoning reversed?
decrease O2 pressure, increase O2 pressure, remove Hb(CO)4

51. Energy in Reactions Things in Nature move toward:
greater stability thus lower energy (exothermic energy changes)
less organization MORE CHAOS (increased randomness)

52. 1. Energy Changes
Systems in condition of great energy tend to change to conditions of low energy.
Thus exothermic reactions are favored in nature

53. 2. Randomness or Entropy (S)
Entropy is a measure of disorder or randomness,
Represented by the symbol S
Systems tend to change from conditions from of great order to conditions of low order. Thus entropy changes from low to high.
Your room NEVER cleans itself (disorder to order?)

54. Entropy When is there more disorder in a system?
1. Entropy of gases is greater than the entropy of liquids and solids
entropy increases in reactions in which solid reactants form liquids or gaseous products
S l g ΔS

55. Entropy 2. Entropy increases when a substance is divided into parts
Entropy increases when ionic compounds dissolve in water
Decomposition reactions
3. Entropy increases in chemical reactions in which total number of product molecules is greater than the total number of reactant molecules

56. Entropy 4. Entropy tends to increase when temperature increases
As temperature increases molecules move faster and faster which increases disorder

57. LeChatlier Principle Animation: http://www. mhhe