Equilibrium is said to occur when a reaction and its reverse reaction proceed at the same rate. In a reaction that is at equilibrium, the amount of reactants and products remains constant. The rate of reaction is equivalent to the rate constant (K) multiplied by the concentration of the species.
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| Equilibrium is reached when the amount of product and reactant becomes constant. |
The equilibrium constant (Keq) is central to the concept of equilibrium. We simulated Keq) in the Phet simulation by taking the number of B molecules and dividing it by the number of A molecules. In reality, K is calculated in a similar manner. Consider the following hypothetical reaction:
Aa + bB <–> cC + dD
(where lowercase letters are stoichiometric coefficients and uppercase letters are hypothetical elements or compounds)
In this reaction, the equilibrium constant (in this case the concentration constant, Kc) is equivalent to the concentrations of the products raised to the power of their respective stoichiometric coefficients divided by the concentrations of the reactants raised to the power of their respective stoichiometric coefficients. This can be modeled by the following equation:
Kc = [C]^c[D]^d / [A]^a[B]^b
For gases, pressure is proportional to concentration in a closed system. Thus, there is also a pressure constant that is calculated in a similar manner:
Kp = (PC^c)(PD^d) / (PA^a)(PB^b)
If the K value for a reaction is much greater than one, the reaction is product favored. There is more product than reactant at equilibrium. If the K value is much less than one, the reaction is reactant favored. There is more reactant than product at equilibrium. The equilibrium constant is not affected by changes in number of moles, volume, or pressure. Keq is only affected by temperature. I found this concept to be fairly straightforward. The constant equals products over reactants. To cover this subject I completed the Phet simulation activity, the Equilibrium I lecture, as well as the Equilibrium I worksheet.
The equilibrium constant is similar to another aspect of equilibrium, the reaction quotient (Q). The reaction quotient gives the same ratio that the equilibrium constant gives (products over reactants) but for a system that is not at equilibrium. Q is calculated by substituting the initial concentrations for reactions and products into the equilibrium expression:
Qc = [C]^c[D]^d / [A]^a[B]^b
If Q > K, there is less reactant and more product in the initial conditions than at equilibrium. If Q < K, there is more reactant and less product in the initial conditions than at equilibrium. When Q = K, the reaction is at equilibrium. As it is very similar to calculating the normal equilibrium constant, I also understood the reaction quotient well. The reaction quotient was covered in the Equilibrium II worksheet as well as the Equilibrium Calculations I lecture.
Another central part of the concept of equilibrium is Le Châtelier's principle. Le Châtalier's principle states that if a system at at equilibrium is disturbed by a change in temperature, pressure, or the concentrations of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance. Once a reaction is at equilibrium is at equilibrium, it is possible to change the concentrations of the products and reactants by changing the external conditions in three ways: adding/removing reactants or products, expanding/contracting a reaction system, and changing the temperature. When equilibrium is disturbed, it is reestablished when the reaction proceeds in the direction where the number of moles or pressure has dropped. Le Châtelier's principle was covered in the Equilibrium Part II lecture as well as in the Equilibrium I and II worksheet. I found Le Châtelier's principle to be quite intuitive, especially after completing the Phet simulation. The Phet simulation really helped me to visualize the effects of a change in temperature or number of moles.
Additionally, this week I was introduced to the concept of RICE charts. RICE charts (standing for Reaction, Initial, Change, Equilibrium). It is a simple way of figuring out the number of moles of reactants and/or products at equilibrium, as well as the equilibrium constant. RICE charts were a central part of the latter half of the Equilibrium I worksheet.
The final aspect of the equilibrium topics covered this week was the relationship between thermodynamics and equilibrium. I studied how Gibbs Free Energy (∆G) relates to equilibrium. At standard conditions, Qp = 1 and ∆G˚ = ∆G. However, conditions are not always standard for reactions. When conditions are not standard, change in Gibbs Free Energy can be calculated using this equation:
∆G = ∆G˚ + RTlnQ
(where R = 8.314 J/mol-K, T = temperature in kelvin, and Q = the reaction quotient at the moment)
At equilibrium, ∆G = 0. The standard state free energy of a reaction (∆G˚) is a measure of how far a reaction is from equilibrium. The smaller the ∆G˚ value, the closer the standard state is to equilibrium. The larger the ∆G˚ value, the closer the standard state is to equilibrium. When ∆G > 0 and Keq < 1, there are mostly reactants at equilibrium. When ∆G is much greater than 1 and Keq is much less than 1, there are almost solely reactants at equilibrium. When ∆G < 0 and Keq > 1, there are mostly products at equilibrium. When ∆G is much less than 1 and Keq is much greater than 1, there are almost all products at equilibrium. The relationship between thermodynamics and equilibrium was covered in the lecture of the same name as well as in the Equilibrium III worksheet. Of all of the concepts covered, this is the concept that I understood the least. I am able to solve the problems by blindly plugging in values to the given equations, but I do not really understand the concept. This website helped a bit, but I still do not have a very developed understanding of the topic.
