1. CHE1031 Lecture 11: Chemical equilibrium
Lecture 11 topics Brown chapter 1
1. Concept of equilibrium
Equilibrium reactions are reversible
2. The equilibrium constant
Law of mass action
Equilibrium constant expressions
3. Working with equilibrium expressions
What does Kc tell us?
Kc & direction of reaction
4. Le Chatelier’s Principle
Application to Haber reaction
Changes of concentration
Changes in volume & pressure
Changes in temperature
5. Catalysts & equilibrium

2. Equilibrium reactions are reversible.
The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws.
It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made.
This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory.
While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

3. What is equilibrium? Equilibrium: Static Equilibrium:
Dynamic Equilibrium:
Chemical Equilibrium:
to be in a state of balance
an object subject to equal & opposite forces is at rest (physics)
opposing processes occur simultaneously & at the same rate
So change is happening, but the situation seems static
reversible chemical reactions occur in the forward & backward directions simultaneously & at identical rates
In a closed system, the [reactants & products] appear static.
Actually, both are constantly being made & broken down, but their concentrations do
not change so the
Situation appears to
be static. p

4. Equilibrium reactions are reversible.
When a sealed test tube of dinitrogen tetroxide is placed into a beaker of warm water a reversible reaction starts. Because the tube is sealed (a closed system) it eventually reaches equilibrium.
N2O4(g) NO2(g)
colorless brown
We can write rate equations for each reaction:
Forward rxn:
Reverse rxn:
At equilibrium:
Rate f = kf[N2O4]
Rate r = kr[NO2]2
Rate f = Rate r kf[N2O4] = kr[NO2]2
Rearrange to: Kc = kf = [NO2]2
kr [N2O4]
Keq is a constant for
each & every reaction p

5. Graphical descriptions of equilibrium
N2O4(g) NO2(g)
colorless brown
Concentration notes:
Started w/ only N2O4 and slowly lost it.
Started w/ no NO2 and slowly made it.
At equilibrium both concentrations appeared to become static.
Rate notes:
Initial rate of product formation double the initial rate of loss of reactant.
At equilibrium the rates become equal & appear to become static. p

6. The Haber reaction
This reaction was developed by German chemists in It’s widely used to make chemical fertilizer and explosives. The reaction requires high temperature (500C) & pressure (200 atm). The US makes 40 billion pounds of NH3 annually!
N2(g) + 3H2(g) 2NH3(g)
Note that equilibrium can be reached from either direction:
From the left (starting with reactants)
From the right (starting from product) p