Reaction Quotient 

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The equilibrium constant of a reaction is the ratio of products to reactants at an equilibrium state. Similarly, the reaction quotient is a ratio of the concentration of products to the concentration of reactants. The written expression for Q is the same expression as Keq, but the difference is that Q is not a constant ratio of products to reactants. Instead, the reaction quotient is calculated by examining a reaction system, determining the current concentrations of products and reactants, and finally plugging those values into the expression for reaction quotient. We will examine the reaction quotient of the chemical reaction below. Note that solids and liquids are not participatory in determining the equilibrium ratio, and that the reaction order of reactant A is determined by it’s coefficient. Calculating a value for Q would require us to be provided with concentrations of A, B, and D. For instance, if the concentrations were measured to be [D] = 2 M, [A] = 3 M, and [B] = 1 M, substituting those values into the equilibrium expression would yield a value of Q equal to 2/9.
Relating Q and Keq
Comparing the transient value Q to the constant Keq tells you how far away your current reaction system is from equilibrium. There are three cases to consider, and in each case there is a different implication about how the reaction will be affected.
When Q greater than Keq, the ratio of products to reactants is too large. The reaction will therefore proceed in the reverse direction to form more reactant.
When Q = Keq, the ratio of products to reactants in the reaction system is at the equilibrium ratio. The forward and reverse reaction rates will equal one another.
When Q less than Keq, the ratio of products to reactants in the reaction system is too small. The forward reaction will be favored, in order to form more products.
ΔG, Q, and Keq
The values of ΔG, Q, and Keq are associated through the equation below [1]. Recall also that Gibbs free energy under standard conditions is equal to -RTLn(Keq) [2]. If a reaction is carried out under standard conditions, ΔG = ΔG° . Under non-standard conditions, the equations below need to be used in order to determine the spontaneity of a reaction.
Considering a reaction carried out not at standard conditions, there are three cases to consider, for which the reaction will be exergonic, endergonic, or at an equilibrium.
1. If Q greater than Keq, then RTLn(Q) greater than -RTLn(Keq). In this instance, ΔG will be greater than 0, so the reaction is necessarily endergonic. Being an endergonic reaction, the reverse reaction will be favored over the forward reaction. Products will largely breakdown into reactants.
2. If Q = Keq, then RTLn(Q) = -RTLn(Keq). In this instance, ΔG will be equal to 0, so the reaction is at an equilibrium, and the forward and reverse reactions will proceed at equal rates.
3. If Q less than Keq, then RTLn(Q) less than -RTLn(Keq). In this instance, ΔG will be less than 0, so the reaction is necessarily exergonic. Being an exergonic reaction, the forward reaction will be favored over the reverse reaction. Reactants will largely be converted into products.
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