What is the primary difference between $K_c$ and $K_p$?

$K_c$ is the equilibrium constant expressed in terms of molar concentrations of reactants and products, typically used for reactions in solution or when all species are gases and concentrations are preferred. $K_p$ is the equilibrium constant expressed in terms of partial pressures of gaseous reactants and products, exclusively used for reactions involving gases. The relationship between them is $K_p = K_c(RT)^{Delta n_g}$, where $Delta n_g$ is the change in the number of moles of gaseous species.

Does the equilibrium constant change if I add more reactants or products to a system at equilibrium?

No, the value of the equilibrium constant ($K$) itself does not change if you add more reactants or products. According to Le Chatelier's principle, the system will shift its equilibrium position to counteract the disturbance (i.e., consume the added reactant/product or produce more of it) until a new equilibrium is established. At this new equilibrium, the ratio of product to reactant concentrations (or partial pressures) will once again equal the original $K$ value, provided the temperature remains constant.

How does temperature affect the equilibrium constant?

Temperature is the only factor that can change the numerical value of the equilibrium constant ($K$) for a specific reaction. For an endothermic reaction ($Delta H > 0$), increasing the temperature increases the value of $K$, favoring product formation. For an exothermic reaction ($Delta H < 0$), increasing the temperature decreases the value of $K$, favoring reactant formation. This relationship is quantitatively described by the Van't Hoff equation.

What does a very large or very small value of K signify?

A very large value of $K$ (e.g., $K > 10^3$) indicates that at equilibrium, the reaction mixture consists predominantly of products. The reaction proceeds almost to completion in the forward direction. Conversely, a very small value of $K$ (e.g., $K < 10^{-3}$) signifies that at equilibrium, the reaction mixture consists predominantly of reactants. The reaction barely proceeds in the forward direction, and very little product is formed. A $K$ value close to 1 suggests significant amounts of both reactants and products are present at equilibrium.

Why are solids and pure liquids excluded from the equilibrium constant expression?

Solids and pure liquids have constant concentrations (or activities) at a given temperature, regardless of the amount present. Their molar concentration is essentially their density divided by their molar mass, which is a fixed value. Since these values are constant, they are implicitly incorporated into the value of the equilibrium constant itself. Including them explicitly would only complicate the expression without adding new information about the relative amounts of species that *can* change.

What is the Reaction Quotient (Q) and how is it different from K?

The Reaction Quotient (Q) has the same mathematical form as the equilibrium constant (K), but it is calculated using concentrations or partial pressures of reactants and products at *any* point during the reaction, not necessarily at equilibrium. K, on the other hand, is specifically calculated when the system *is* at equilibrium. By comparing Q with K, we can predict the direction a reaction will shift to reach equilibrium: if Q K, it proceeds backward; if Q = K, the system is at equilibrium.

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Q: What does a very large or very small value of K signify?

A very large value of $K$ (e.g., $K > 10^3$) indicates that at equilibrium, the reaction mixture consists predominantly of products. The reaction proceeds almost to completion in the forward direction. Conversely, a very small value of $K$ (e.g., $K < 10^{-3}$) signifies that at equilibrium, the reaction mixture consists predominantly of reactants. The reaction barely proceeds in the forward direction, and very little product is formed. A $K$ value close to 1 suggests significant amounts of both reactants and products are present at equilibrium.

Q: Why are solids and pure liquids excluded from the equilibrium constant expression?

Solids and pure liquids have constant concentrations (or activities) at a given temperature, regardless of the amount present. Their molar concentration is essentially their density divided by their molar mass, which is a fixed value. Since these values are constant, they are implicitly incorporated into the value of the equilibrium constant itself. Including them explicitly would only complicate the expression without adding new information about the relative amounts of species that *can* change.

Q: What is the Reaction Quotient (Q) and how is it different from K?

The Reaction Quotient (Q) has the same mathematical form as the equilibrium constant (K), but it is calculated using concentrations or partial pressures of reactants and products at *any* point during the reaction, not necessarily at equilibrium. K, on the other hand, is specifically calculated when the system *is* at equilibrium. By comparing Q with K, we can predict the direction a reaction will shift to reach equilibrium: if Q K, it proceeds backward; if Q = K, the system is at equilibrium.

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Version 1Updated 22 Mar 2026

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The equilibrium constant, denoted by $K$ , is a fundamental quantitative measure that expresses the ratio of product concentrations (or partial pressures) to reactant concentrations (or partial pressures) at chemical equilibrium, with each concentration or partial pressure raised to the power of its stoichiometric coefficient in the balanced chemical equation. It provides crucial information about …

Quick Summary

The equilibrium constant ( $K$ ) is a quantitative measure of the extent of a reversible chemical reaction at equilibrium. For a general reaction $aA + bB \rightleftharpoons cC + dD$ , $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$ (using molar concentrations) and $K_p = \frac{(P_C)^c(P_D)^d}{(P_A)^a(P_B)^b}$ (using partial pressures for gases).

The value of $K$ is constant for a given reaction at a specific temperature and indicates the relative amounts of products and reactants at equilibrium. A large $K$ means products are favored, while a small $K$ means reactants are favored.

Only temperature changes the value of $K$ . Solids and pure liquids are excluded from the $K$ expression because their concentrations are constant. The relationship $K_p = K_c(RT)^{Delta n_g}$ connects the two constants, where $Delta n_g$ is the change in moles of gaseous species.

The reaction quotient ( $Q$ ) is calculated similarly to $K$ but for non-equilibrium conditions, allowing prediction of reaction direction.

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Key Concepts

Writing

K_c

Expressions

To write the $K_c$ expression, identify the balanced chemical equation. Products go in the numerator,…

Calculating

K_p

from

K_c

The relationship $K_p = K_c(RT)^{Delta n_g}$ is crucial for interconverting between the two constants. First,…

Using Reaction Quotient (Q) to Predict Direction

The reaction quotient $Q$ is calculated using the same formula as $K$ , but with current (non-equilibrium)…

$K_c$ expression: —

K_c = \frac{[Products]^{coefficients}}{[Reactants]^{coefficients}}

(for aqueous/gaseous species)

$K_p$ expression: —

K_p = \frac{(P_{Products})^{coefficients}}{(P_{Reactants})^{coefficients}}

(for gaseous species)

Relationship: —

K_p = K_c(RT)^{Delta n_g}

- $Delta n_g = (\text{moles gaseous products}) - (\text{moles gaseous reactants})$ - $R = 0.0821 \text{ L atm mol}^{-1}\text{ K}^{-1}$ (for $P$ in atm)