# D23.1 Reaction Quotient

Starting a reaction with only reactants present means that the reaction must initially be spontaneous in the forward direction, even if only very small concentrations of products are present when equilibrium is reached. Similarly, starting a reaction with only products present requires that the reaction proceed in the reserve direction.

However, which direction does a reaction go when both reactants and products are present but the reaction is not at equilibrium? The reaction quotient can answer this question. For a generic reaction:

mA + nB ⇌ xC + yD

the reaction quotient (Q) is defined as: The subscript “c” in Qc indicates the reaction quotient is in terms of concentrations; for a gas-phase reaction we could write Qp similarly in terms of partial pressures. The concentrations are represented by “[…]t” where the subscript “t” emphasizes that Qc for a reaction depends on the concentrations present at the time when Qc is determined. This is usually not at equilibrium.

When only reactants are present, Qc = 0. As the reaction proceeds, Qc increases because product concentrations increase and reactant concentrations decrease (see Figure below, panel (a) and (b)). Figure: Reaction Quotient. The changes in the concentrations of reactants and products are depicted for the 2 SO2(g) + O2(g) ⇌ 2 SO3(g) reaction. (Left) The reaction starts with only reactants present. Graph (a) shows changes in concentration and graph (b) shows the change in Qc as the reaction approaches equilibrium over time. (Right) The reaction starts with only products present. Graph (c) shows changes in concentration and graph (d) shows the change in Qc as the reaction approaches equilibrium over time.

When the reaction reaches equilibrium, Qc no longer changes over time because the concentrations no longer change, and, at equilibrium, Qc = Kc: A system that is not at equilibrium proceeds spontaneously in the direction that establishes equilibrium (Qc changes until it equals Kc). Hence, we can predict directional shifts of a reaction by comparing Qc to Kc: when Qc < Kc, the reaction proceeds spontaneously in the forward direction (from left to right; to the product side); when Qc > Kc, the reaction proceeds spontaneously in the reverse direction (from right to left; to the reactant side).

For example, for the water-gas shift reaction,

CO(g) + H2O(g) ⇌ CO2(g) + H2(g)          Kc (800 °C) = 0.64

different starting mixtures of CO, H2O, CO2, and H2 react (and the concentrations of reactants and products change) until the compositions reach the same value of Qc; that is, until Qc = Kc.

Figure: water-gas shift. Concentrations of four different mixtures are shown before and after reaching equilibrium at 800 °C for the water-gas shift reaction: CO(g) + H2O(g) ⇌ CO2(g) + H2(g).

It is important to recognize that Qc reaches the same equilibrium value (Kc) whether the reaction starts from all reactants, from all products, or from a mixture of both. In fact, one technique to determine whether a reaction is truly at equilibrium is to start with only reactants in one experiment and start with only products in another. If the same value of the reaction quotient is observed when the concentrations have stopped changing in both experiments, then it is highly likely that the system has reached equilibrium.

Exercise: Determining Which Direction a Reaction Goes 