CHE1031 Lecture 11: Chemical equilibrium Lecture 11 topics Brown chapter 1 1. Concept of equilibrium 15.1 Equilibrium reactions are reversible 2. The equilibrium constant 15.2 Law of mass action Equilibrium constant expressions 3. Working with equilibrium expressions 15.3 What does Kc tell us? Kc & direction of reaction 4. Le Chatelier’s Principle 15.7 Application to Haber reaction Changes of concentration Changes in volume & pressure Changes in temperature 5. Catalysts & equilibrium

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.

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. 627-8

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) 2NO2(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. 628-9

Graphical descriptions of equilibrium N2O4(g) 2NO2(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. 628-9

The Haber reaction This reaction was developed by German chemists in 1912. 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. 630-1