An equilibrium constant is a number that describes the ratio of product concentrations to reactant concentrations when a reversible chemical reaction has reached a steady state. It tells you, in a single value, which side of a reaction is favored once the system settles into balance. A large equilibrium constant means the reaction produces mostly products; a small one means reactants dominate.
The Basic Formula
For any balanced chemical reaction where reactants A and B form products C and D:
aA + bB ⇌ cC + dD
the equilibrium constant expression is:
K = [C]c[D]d / [A]a[B]b
The square brackets represent concentrations (usually in moles per liter), and the lowercase letters are the coefficients from the balanced equation, used as exponents. Products always go in the numerator, reactants in the denominator. This relationship is known as the law of mass action: at a given temperature, this ratio of concentrations is always the same number, no matter how much of each substance you started with.
What the Size of K Tells You
The numerical value of K reveals which direction a reaction “prefers” at equilibrium:
- K greater than 1: Products are present in higher concentrations than reactants. The larger K is, the more completely the reaction proceeds toward products. As K approaches infinity, the reaction is essentially irreversible.
- K less than 1: Reactants dominate at equilibrium. A very small K (like 10-10) means the reaction barely produces any products at all.
- K approximately equal to 1: Significant amounts of both reactants and products coexist. Neither side is strongly favored.
This makes K a powerful shorthand. If someone tells you a reaction has a K of 4.0 × 108, you immediately know it goes nearly to completion. A K of 2.3 × 10-5 tells you the opposite: very little product forms.
Kc vs. Kp for Gas Reactions
When a reaction involves gases, you can express the equilibrium constant in two ways. Kc uses molar concentrations (mol/L), while Kp uses partial pressures (usually in atmospheres). The two are related by a simple equation:
Kp = Kc(RT)Δn
Here, R is the gas constant (0.0821 L·atm·mol-1·K-1), T is the temperature in Kelvin, and Δn is the total moles of gas on the product side minus the total moles of gas on the reactant side. When Δn equals zero (the same number of gas molecules on both sides), Kp and Kc are numerically equal.
Why Pure Solids and Liquids Are Left Out
In reactions that involve a mix of phases (called heterogeneous equilibria), you’ll notice that pure solids and pure liquids don’t appear in the equilibrium expression. The reason is straightforward: their concentrations don’t change during the reaction. A block of calcium carbonate, for example, has the same density whether you have a gram or a kilogram of it. Since its concentration is effectively constant, it’s already baked into the value of K and doesn’t need a separate term.
Formally, this is handled through the concept of “activity.” The activity of any pure solid or pure liquid is defined as 1, so including them in the expression would just mean multiplying by 1, which changes nothing.
Are Equilibrium Constants Truly Unitless?
In introductory chemistry courses, you’ll often see K written with apparent units (like mol/L raised to some power). Strictly speaking, though, thermodynamic equilibrium constants are dimensionless. Each concentration or pressure term in the expression is actually a ratio comparing the measured value to a standard reference value (1 mol/L for dissolved species, 1 atm for gases). Since you’re dividing a concentration by a concentration, the units cancel. In practice, most courses don’t worry about this distinction, but it explains why you sometimes see K reported without units.
The Reaction Quotient: Predicting Which Way a Reaction Shifts
The reaction quotient, Q, uses the exact same formula as K, but with concentrations measured at any point in time, not just at equilibrium. Comparing Q to K tells you what the system will do next:
- Q = K: The system is already at equilibrium. No net change occurs.
- Q < K: There are too many reactants relative to products. The reaction will shift forward, producing more products until Q rises to match K.
- Q > K: There are too many products. The reaction will run in reverse, converting products back into reactants until Q falls to match K.
This comparison is one of the most practical tools in chemistry. If you mix a set of chemicals together at known concentrations, calculating Q and comparing it to K immediately tells you whether the reaction will proceed forward, backward, or not at all.
Temperature Is the Only Thing That Changes K
A common source of confusion: changing concentrations, pressures, or volumes does not change the equilibrium constant. Those changes shift the position of equilibrium (the amounts of each substance present), but K itself stays the same. The system simply readjusts until the ratio of concentrations matches K again.
Adding a catalyst also has no effect on K. A catalyst speeds up both the forward and reverse reactions by the same amount, so the system reaches equilibrium faster, but the final balance point is identical.
Temperature, however, does change K. The relationship between K and temperature is captured by the van ‘t Hoff equation, which shows that K depends exponentially on temperature and on the heat absorbed or released by the reaction. For an exothermic reaction (one that releases heat), raising the temperature decreases K, shifting the balance toward reactants. For an endothermic reaction (one that absorbs heat), raising the temperature increases K, favoring products.
The Link to Free Energy
At a deeper level, the equilibrium constant is connected to the thermodynamic favorability of a reaction through the standard Gibbs free energy change (ΔG°). The relationship is:
ln K = −ΔG° / RT
where R is the gas constant (8.314 J·mol-1·K-1) and T is the temperature in Kelvin. A negative ΔG° (a spontaneous reaction under standard conditions) corresponds to K greater than 1. A positive ΔG° corresponds to K less than 1. This equation lets you calculate K from thermodynamic tables, or work backward from a measured K to determine the energy profile of a reaction.
This connection is what makes the equilibrium constant more than just a ratio of concentrations. It encodes the energy landscape of the reaction, reflecting how strongly nature drives the system toward one side or the other at a given temperature.