Predicting the Direction of the Reaction

The figure is a simple but powerful visual summary of how the reaction quotient $Q_c$ compares to the equilibrium constant $K_c$ to decide which way a reaction will shift. It shows three separate bar charts, each with two bars side by side: a green bar for $Q_c$ and a blue bar for $K_c$. The height of each bar represents the numerical value of that quantity.

In the first panel, the green bar ($Q_c$) is shorter than the blue bar ($K_c$). An arrow points from left to right, indicating that the net reaction moves forward — from reactants to products — until $Q_c$ rises to meet $K_c$. In the second panel, the two bars are exactly the same height, and there is no arrow: the system is at equilibrium, with no net change. In the third panel, the green bar ($Q_c$) is taller than the blue bar ($K_c$). The arrow now points from right to left, showing that the net reaction goes backward — from products to reactants — until $Q_c$ falls back down to $K_c$.

The physical idea is that $K_c$ acts like a target or a set point. No matter where you start, the reaction mixture will always adjust its composition until $Q_c$ equals $K_c$. The figure makes this directional rule immediate: if $Q_c < K_c$, the reaction proceeds forward; if $Q_c > K_c$, it proceeds in reverse; if $Q_c = K_c$, the system is already at equilibrium.

The key formula that the textbook develops alongside this figure is the definition of the reaction quotient. For a general reaction

the quotient is written exactly like the equilibrium constant, but using the concentrations present at any moment — not necessarily at equilibrium:

Here, $[\text{A}]$, $[\text{B}]$, $[\text{C}]$, and $[\text{D}]$ are the molar concentrations at the instant you measure them. The exponents $a$, $b$, $c$, $d$ are the stoichiometric coefficients from the balanced equation. The same idea applies to gases using partial pressures, giving $Q_p$.