1. Equilibrium

2. Non reversible reactions
Some chemical and physical reactions occur until one or all the reactants are used up
Example Evaporation of water from a beaker (a physical process)
Reaction is complete when all the H2O molecules are in the gas phase
gaseous molecules from evaporation
water in liquid phase
Evaporation continues
Fig 1: Physical process going to completion

3. Non reversible chemical reactions
Example 2 Magnesium in excess dilute acid solution
Eventually all the magnesium reacts and gas production stops
Hydrogen gas
Magnesium
Fig 2: Chemical process going to completion
In examples 1 and 2 the forward reaction continues until the reactant(s) are used up this is written as:
Reactants Products

4. Reversible Reactions
If we warm our beaker of water with glad wrap over the top we get a reversible reaction because the reaction is now in a closed system
Example 3
Evaporation is occurring but so is the opposing process of condensation
Gaseous water molecules from evaporation
Water in liquid phase
The forward reaction is from left to right
H2O (l) H2O (g)
This reaction has reached Equilibrium when evaporation and condensation occur at the same rate
The reverse reaction is
H2O (g) H2O (l)
The overall equilibrium reaction is written as
H2O (l) H2O (g)

5. Equilibrium
When reversible reactions reach a point where the rate of forward reaction equals the rate of the backward reaction the system is said to have reached equilibrium
The equilibrium is said to be dynamic which means the forward and backward reactions do not stop

6. Important Points so Far
Not all reactions are equilibrium reactions (ie some reactions are not reversible)
If an equilibrium reaction is to occur it must be in a closed system
When a reaction has reached dynamic equilibrium this means that the forward reaction is occurring at the same rate as the reverse reaction and this is written as:
reactants products

7. Chemical Equilibrium
Many chemical reactions behave as a dynamic equilibrium
An Example is the Iron(lll) / thiocyanate equilibrium
Fe3+ (aq) SCN-(aq)
yellow colourless ion
In the reaction
The red colour forms quickly then remains constant
The system is closed as the ions Fe3+, SCN- and FeSCN2+ are all contained in the same aqueous solution
FeSCN 2+(aq)
bright red complex

8. Iron(lll) / thiocyanate equilibrium continued
The reaction hasn’t gone to completion because it can be shown that their are unreacted Fe3+, SCN- ions in the red solution, this shows the reverse reaction is also occurring
FeSCN 2+(aq) Fe3+ (aq) SCN-(aq)
The equilibrium for the reaction is written:
Fe3+ (aq) SCN-(aq) FeSCN 2+(aq)
At equilibrium the rate at which the complex ions (FeSCN2+) form is equal to the rate at which the complex ions (FeSCN2+) break up

9. Part A
Fe3+ (Iron III) ions (yellow) & SCN- (thiocyanate) ions (colourless)
These two ions react to form a reddish complex ion:
Fe3+ + SCN-  FeSCN2+
Make some of this and place on 4 petrie dishes

10. Add more Fe3+
Control
Add more SCN-
Add F-

11. Why does a reaction stop?
Until now, you expect it to be because one or both of the reactants have run out
In this reaction, neither the Fe3+, nor the SCN- have run out, yet the reaction appears to stop
Why?

12. Two processes occur one to make FeSCN2+ the other to decompose FeSCN2+
The reaction stops when these two reactions balance out
Fe3+ + SCN- forming complex FeSCN2+
Fe3+ + SCN- decomposing FeSCN2+
This is written as:
Fe3+ + SCN- FeSCN2+

13. Question - Which one’s correct?
Equilibrium occurs only when:
The reactants are used up
The concentration of reactants is equal to the concentration of products
All chemicals stop reacting
The products react together to form the reactants at the same rate as reactants form products
Ans - D

14. Starter Questions
What two conditions do you need for an equilibrium reaction to occur?
The reaction must be a reversible reaction
The reaction must be in a closed system
What is meant by the term equilibrium?
The forward reaction is occurring at the same rate as the reverse reaction (rate forward = rate reverse)
What factors can affect an equilibrium reaction?
Concentration of reactants or products
Temperature
Surface area
Pressure

15. Equilibrium constant
Scientists wanted to put a number on the proportions of products and reactants when equilibrium is reached
They did heaps of experiments and eventually they came upon a method for finding the equilibrium constant (K) for each particular reaction

16. Equilibrium Constant
For a general reversible reaction:
aA bB cC dD
‘a’ moles of ‘b’ moles of ‘c’ moles of ‘d’ moles of
substance A substance B substance C substance D
The ratio of products to reactants is expressed as the equilibrium constant symbol K and is written as
remember as

17. Writing Equilibrium Equations
Where:
[B] is the concentration of B in molL-1 at equilibrium
[A] is the concentration of A in molL-1 at equilibrium
Writing the following Equilibrium expression for:
CO (g) + H2O (g) CO2 (g) + H2 (g)

18. Equilibrium Constants
Equilibrium constants (K) are specific for a given reaction and temperature.
An equilibrium constant is useful for determining how far a reaction has proceeded.
If the equilibrium constant (K) is large
eg > 105 then at equilibrium the reaction has proceeded almost to completion.

19. Equilibrium Constants (continued)
If the equilibrium constant (K) is small
eg < then at equilibrium only a small amount of reactant is used up and only a small amount of product has formed then the reaction has effectively not happened

20. Be very careful! Writing Equilibrium expressions:
Task: Write the equilibrium expression for
CH4 (g) H2O (g) CO (g) H2 (g)
Be very careful!

21. For equilibrium systems involving solid(s) and pure liquids
The solid(s) and liquids are not included in the equilibrium expression
Eg: write the equilibrium expression for
C (s) H2O (g) CO (g) H2 (g)
Be very careful!

22. Be very careful! Writing more Equilibrium expressions:
Eg: write the equilibrium expression for
S(s) O2 (g) SO2 (g)
Be very careful!

23. Starter : write the equilibrium expression for:
N2(g) H2(g) NH3(g)

24. write the equilibrium expression for:
2N2(g) H2(g) NH3(g)

25. Moving Equilibrium Systems
Le Chatelier’s Principle: When change is applied to a system at equilibrium, the system responds so that the effects of the change are minimised

26. Example: H2(g) + Cl2(g) 2HCl(g)
In the above reaction what would be the effect on the amount of HCl formed if Cl2 gas was added?
The amount of HCl formed would increase, because the system will try to minimise the change and remove Cl2 by shifting the equilibrium to the RHS producing more HCl.
What would be the effect on the amount of HCl formed if hydrogen gas was removed ?
The amount of HCl formed would decrease because the system would try to produce more H2 by shifting the equilibrium to the LHS.

27. Turn to page 147 in your lab book – Equilibrium Systems
Lets see how changing concentration and temperature affect the position of the equilibrium

28. Forward reaction is endothermic Backward reaction is exothermic
N2O4(l) 2NO2(g) ΔH = +54kJ
(colourless) (brown)
Or…
N2O4(l) + energy 2NO2(g)
Forward reaction is endothermic
Backward reaction is exothermic
Cooling favours exothermic
Heating favours endothermic

29. So heating this system will move the equilibrium to the right and the system will appear dark brown
Cool it and move the equilibrium system to the left and the system will appear pale brown

30. This would decrease the amount of NH3 produced.
Example
N2(g) + 3H2(g) NH3(g) H= -92 kJ
H= +92 kJ
The endothermic (+H) direction is the backward reaction, so an increase in temperature would move the equilibrium in the backward direction to take more energy in and therefore minimise the increase in temperature.
This would decrease the amount of NH3 produced.

31. Consider another form of the N2O4/NO2 system
N2O4(g) 2NO2(g)
(colourless) (brown)
Fewer molecules More molecules
Fewer molecules
More molecules

32. If the pressure is decreased, the system opposes this by increasing the number of molecules in the system
Forms more NO2(g)
colour will go from light brown to a darker brown as new equilibrium established
Fewer molecules
More molecules

33. Pressure N2(g) + 3H2(g) 2NH3(g)
Now what would happen to the amount of NH3 produced if the pressure was decreased?
For the same reaction at equilibrium, an increase in the volume of the container, (or decrease in pressure), will shift the equilibrium to the side with more gas particles.
This means the concentration of NH3 would decrease.

34. Reactant concentration
Increase – reaction moves forwards (to the right)
Decrease- reaction moves backwards (to the left)
Product concentration
Increase – reaction moves backwards
Decrease – reaction moves forwards
Catalyst No change in equilibrium position, but equilibrium reached more quickly.

35. Exercise
For the equilibrium
N2O4(g) NO2(g) rH= -92 kJ mol -1
what would happen to the number of NO2 molecules if the following changes are made to the system at equilibrium
a) an increase in pressure
b) the concentration of N2O4 is increased
c) the pressure of the system is decreased
d) the temperature is decreased
e) a catalyst is added

36. Starter Equilibrium Changes
PCl3 (g) + Cl2(g) PCl5 (g) H= +110 kJ mol -1
what would happen to the amount of PCl5 if the following changes are made to the system at equilibrium
a) an increase in pressure
b) the concentration of Cl2 is increased
c) the volume of the container is increased
d) the temperature is decreased

37. Equilibrium in
Industry

38. Strategies to lower costs include
Buy raw materials as cheaply as possible
Use energy efficiently
Recycle unused reactants and by-products, if possible

39. Heat exchangers may be used to remove energy from an exothermic process and deliver it to an endothermic process
Catalysts assist in the rapid establishment of an equilibrium

40. Fritz Haber developed the “Haber Process” used today to make ammonia
A severe shortage of nitrogen-based fertilisers at the start of the 20th century started hunt for a method of using nitrogen in the air
Nitrogen = 71% of air
Fritz Haber developed the “Haber Process” used today to make ammonia

41. Developed “Ecstasy”
Did fundamental research into the Bunsen flame
Developed a process for extracting gold from seawater
Noble prize 1918 for Haber process

42. Haber died in exile after standing up to Hitler’s anti-Semitic policies which he felt were morally wrong and were decimating German research institutes

43. The Equilibrium N2(g) + 3H2(g) 2NH3(g) -ΔH
To move the equilibrium to the right means a pressure drop and the production of heat
High temperatures favour the backward reaction but make the reaction rate faster
High pressures favour the forward reaction and increase the yield of ammonia

44. % Yield of Ammonia Using The Haber Process
Pressure (atm)
T(0C) 10 50
100
300
1000
200 51 74 82 90 98 15 39 52 71 93
400 4 25 47 80
500 1 6 11 26 57
600
0.5 2 5 14 13