1. TOPIC 7 EQUILIBRIUM
7.1
EQUILIBRIUM

2. ESSENTIAL IDEA
Many reactions are reversible. These reactions will reach a state of equilibrium when the rates of the forward reaction and reverse reaction are equal. The position of equilibrium can be controlled by changing the conditions.
NATURE OF SCIENCE (1.8)
Obtaining evidence for scientific theories – isotopic labeling and its use in defining equilibrium.
NATURE OF SCIENCE (5.5)
Common language across different disciplines – the term dynamic equilibrium is used in other contexts, but not necessarily with the chemistry definition in mind.

3. INTERNATIONAL-MINDEDNESS
The Haber process has been described as the most important chemical reaction on Earth as it has revolutionized global food production. However, it also had a large impact on weaponry in both world wars.

4. UNDERSTANDING/KEY IDEA 7.1.A
A state of equilibrium is reached in a closed system when the rates of the forward and reverse reactions are equal.

5. The system is at equilibrium when the rate of the forward reaction equals the rate of the reverse reaction and the concentration of products and reactants remains unchanged.
A reversible reaction is a chemical reaction in which the products can react to re-form the reactants.
Arrows going both directions (  ) indicates equilibrium in a chemical equation.
2HgO(s)  2Hg(l) + O2(g)

6. APPLICATION/SKILLS
Be familiar with the characteristics of chemical and physical systems in a state of equilibrium.

7. EQUILIBRIUM IN PHYSICAL SYSTEMS
A physical system is one in which a chemical reaction is not occurring.
Equilibrium will only occur in a “closed” system.
A “closed” system is one in which no matter can be exchanged with the surroundings.
A typical example is the evaporation-condensation cycle of a volatile liquid in a closed container.
When the rate of evaporation equals the rate of condensation, the system is at equilibrium.

8. EQUILIBRIUM IN CHEMICAL SYSTEMS
A chemical system is one in which a chemical reaction is occurring.
For the system to reach equilibrium, the forward reaction must equal that of the reverse reaction.
See the graphic on the next slide.

9. 2NO2(g)  2NO(g) + O2(g)

10. CHARACTERISTICS OF THE EQUILIBRIUM STATE
1. Equilibrium is dynamic – the reaction has not stopped but both reactions or processes are still occurring on the microscopic level even though it looks like the reaction or process has stopped.
2. Equilibrium is achieved in a closed system so that no matter can escape to the surroundings.
3. The concentrations of reactants and products remain constant at equilibrium – they are being produced or consumed at an equal rate.

11. 4. There is no macroscopic change in properties – meaning the observable properties do not change at equilibrium because they depend on the concentration of the mixture components.
5. Equilibrium can be reached from either direction – meaning that at the same conditions equilibrium will be reached whether the reaction started with all reactants, all products or a mixture of both.
6. It is important to note that even though the concentrations of reactants and products are constant at equilibrium, this does not imply they are equal.

12. Equilibrium Position
The equilibrium position is the proportion of the reactants and products in the equilibrium mixture.
A reaction with an equilibrium mixture with mostly products is said to “lie to the right”.
A reaction with an equilibrium mixture with mostly reactants is said to “lie to the left”.
We can quantitatively compare equilibrium mixtures with different conditions by using the equilibrium constant Kc.
There can be numerous equilibrium positions, but only one equilibrium constant, Kc, at a specified temperature.

13. UNDERSTANDING/KEY IDEA 7.1.B
The equilibrium law describes how the equilibrium constant (Kc) can be determined for a particular chemical reaction.

14. Equilibrium Constant Kc
For the reaction:
jA + kB  lC + mD
Where Kc is the equilibrium constant, and is unitless. The only thing that can change Kc for a reaction is a change in temperature.

15. APPLICATION/SKILLS
Be able to deduce the equilibrium constant (Kc) from an equation for a homogeneous reaction.

16. Writing an Equilibrium Expression
Write the equilibrium expression for the reaction:
2NO2(g)  2NO(g) + O2(g)
Remember that concentrations are written with square brackets [ ]. Kc is calculated using molar concentrations. Solids and liquids are not included in the K expression.
Kc = ???

17. UNDERSTANDING/KEY IDEA 7.1.C
The magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent.

18. Product Favored Equilibrium
Large values for K signify the reaction is “product favored”
When equilibrium is achieved, most of the reactants have been converted to products.

19. Reactant Favored Equilibrium
Small values for K signify the reaction is “reactant favored”
When equilibrium is achieved, very little reactant has been converted to product.

20. UNDERSTANDING/KEY IDEA 7.1.D
The reaction quotient (Q) measures the relative amount of products and reactants present during a reaction at a particular point in time. Q is the equilibrium expression with non-equilibrium concentrations. The position of the equilibrium changes with changes in concentration, pressure and temperature.

21. jA + kB  lC + mD The Reaction Quotient
For some time, t, when the system is not at equilibrium, the reaction quotient, Q takes the place of K, the equilibrium constant, in the law of mass action.
jA + kB  lC + mD

22. Significance of the Reaction Quotient
If Q = K, the system is at equilibrium
If Q > K, the system shifts to the left, consuming products and forming reactants until equilibrium is achieved
If Q < K, the system shifts to the right, consuming reactants and forming products until equilibrium is achieved

23. APPLICATION/SKILLS
Be able to determine the relationship between different equilibrium constants for the same reaction at the same temperature.

24. The equilibrium constant Kc has a fixed value at the same temperature.
You can derive Kc from different sets of data for the same reaction at constant temperature.
All Kc values should be equal no matter how you arrived at the value since only temperature affects Kc.

25. GUIDANCE
Be familiar with the relationship between Kc values for reactions that are multiples or inverses of each other.

26. Conclusions about Equilibrium Expressions
If an equation is written in reverse, the K expression is the reciprocal of the original K.
2NO2(g)  2NO(g) + O2(g)
2NO(g) + O2(g)  2NO2(g)

27. Conclusions about Equilibrium Expressions
When the balanced equation for a reaction is multiplied by a factor n, the equilibrium expression for the new reaction is the original expression, raised to the nth power.
2NO2(g)  2NO(g) + O2(g)
NO2(g)  NO(g) + ½O2(g)

28. If you add two reactions together, you will multiply the two K expressions to obtain the new K’.

29. UNDERSTANDING/KEY IDEA 7.1.E
A catalyst has no effect on the position of equilibrium or the equilibrium constant.

30. ADDITION OF A CATALYST
When a catalyst is added, both the rate of the forward and the rate of the reverse reactions are increased by the same amount so there is no effect on the equilibrium position.
It only speeds up how fast Kc is achieved.

31. APPLICATION/SKILLS
Be able to apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure, and concentration on the position of equilibrium and on the value of the equilibrium constant.

32. LeChatelier’s Principle
When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress.
Translated: The system undergoes a temporary shift in order to restore equilibrium.
1. When you add a substance or heat, the system shifts to the opposite side.
2. When you take out a substance or heat, the system shifts to the side of the take out.
When you increase pressure, the system shifts to the side with the least number of gaseous molecules.

33. Le Chatelier’s Principle – The Concentration Effect
N2 (g) +3 H2 (g)  2NH3 (g) DH = - 92 kJ
Kc =
[NH3]2
[N2] [H2]3
An increase in the concentration of NH3 results in a increase in N2 and an increase in H2 in such a way to keep the equilibrium constant the same.

34. Le Chatelier’s Principle – The Temperature Effect
N2 (g) +3 H2 (g)  2NH3 (g) DH = - 92 kJ
The reaction is exothermic in the forward direction.
An increase in the temperature would trigger a response in the heat consuming (endothermic) direction.
An increase in temperature therefore causes the reaction to shift in the reverse direction. Some NH3 decomposes to N2 and H2.

35. Le Chatelier’s Principle – The Pressure Effect
N2 (g) +3 H2 (g)  2NH3 (g) DH = - 92 kJ
All molecules in the equilibrium are gases
When the reaction proceeds in the forward direction the number of moles are reduced from 4 to 2.
Pressure is proportional to the number of moles of gas
An increase in pressure therefore causes the reaction to moved in the forward direction. Some N2 and H2 combine to form more NH3

36. Le Chatelier’s Principle – The Effect of Catalysts
N2 (g) +3 H2 (g)  2NH3 (g) DH = - 92 kJ
Catalysts lower the activation energy
A catalyst affects the forward and the reverse direction equally
There is no change in the equilibrium position from a catalyst
A catalyst decreases the time required for the system to achieve equilibrium

37. LeChatelier Example #1
A closed container of ice and water is at equilibrium. Then, the temperature is raised.
Ice + Energy  Water
The system temporarily shifts to the _______ to restore equilibrium.
right

38. LeChatelier Example #2
A closed container of N2O4 and NO2 is at equilibrium. NO2 is added to the container.
N2O4 (g) + Energy  2 NO2 (g)
The system temporarily shifts to the _______ to restore equilibrium.
left

39. LeChatelier Example #3
A closed container of water and its vapor is at equilibrium. Vapor is removed from the system.
water + Energy  vapor
The system temporarily shifts to the _______ to restore equilibrium.
right

40. LeChatelier Example #4
A closed container of N2O4 and NO2 is at equilibrium. The pressure is increased.
N2O4 (g) + Energy  2 NO2 (g)
The system temporarily shifts to the _______ to restore equilibrium, because there are fewer moles of gas on that side of the equation.
left

41. Industrial Applications
***Remember the only thing that can change the value of Kc is a change in temperature.
In reactions involving the manufacture of a chemical, the goal is to obtain the highest yield possible.
We use our knowledge of Le Chatelier’s principle and kinetics to maximize the yield.

42. HABER PROCESS: Production of NH3
N2(g) + 3H2(g)  2NH3(g) ∆H = -93 kJ/mol
First we need to get information from the balanced equation.
All products and reactants are in the gaseous state with 4 moles on the left and 2 on the right.
The reaction is exothermic so we want to re-write the equation as:
N2(g) + 3H2(g)  2NH3(g) + heat

43. N2(g) + 3H2(g)  2NH3(g) + heat
To maximize the amount of NH3 obtained we can do several things:
1. Remove NH3 as it is produced so the reaction will continue to the right producing more NH3 to achieve equilibrium.
2. Increasing the pressure pushes the equilibrium to the side of the least gaseous molecules which is to the right; therefore, producing more NH3.
3. We want a lower temperature since the reaction is exothermic, but not too low which would slow things down too far.
4. Adding a catalyst will speed up the reaction and compensate for the lower temperature.

44. CONTACT PROCESS: Production of sulfuric acid, H2SO4
Sulfuric acid has the highest production of any chemical in the world.
There are three steps in the production of H2SO4 with the second step being the rate determining step.
1. Combustion of sulfur to form sulfur dioxide.
2. Oxidation of sulfur dioxide to form sulfur trioxide.
3. Combination of sulfur trioxide produced in step 2 with water.

45. 2SO2(g) + O2(g)  2SO3(g) ∆H=-196 kJ/mol
Again, the reaction is exothermic and has fewer gaseous molecules on the right so we will basically follow the same industrial applications as with the Haber process.
1. We will not take out SO3 because it is what is needed for step 3 of the process.
2. Increase pressure to shift equilibrium to the right.
3. Lower temperature because the reaction is exothermic.
4. Add a catalyst to increase the rate of the reaction.

46. THEORY OF KNOWLEDGE
Scientists investigate the world at different scales; the macroscopic and microscopic. Which ways of knowing allow us to move from the macroscopic to the microscopic? Chemistry uses a specialized vocabulary; a closed system is one in which no matter is exchanged with the surroundings. Does our vocabulary simply communicate knowledge, or does it shape what we can know? The career of Fritz Haber coincided with the political upheavals of two world wars. He supervised the release of chlorine on the battlefield in World War I and worked on the production of explosives. How does the social context of scientific work affect the methods and findings of science? Should scientists be held morally responsible for the applications of their discoveries?

47. Citations
International Baccalaureate Organization. Chemistry Guide, First assessment Updated Brown, Catrin, and Mike Ford. Higher Level Chemistry. 2nd ed. N.p.: Pearson Baccalaureate, Print. Most of the information found in this power point comes directly from this textbook. The power point has been made to directly complement the Higher Level Chemistry textbook by Catrin and Brown and is used for direct instructional purposes only.