What is the primary difference between an irreversible and a reversible reaction?

An irreversible reaction proceeds in one direction, typically until one of the reactants is completely consumed, and products cannot readily convert back into reactants under the given conditions. For example, burning wood. A reversible reaction, on the other hand, can proceed in both forward (reactants to products) and reverse (products to reactants) directions simultaneously. These reactions eventually reach a state of dynamic equilibrium where both processes occur at equal rates.

Does the equilibrium constant ($K$) change if we add more reactants or products to a system at equilibrium?

No, the equilibrium constant ($K$) is a constant value for a specific reaction at a given temperature. Adding more reactants or products will cause the system to shift its equilibrium position (according to Le Chatelier's Principle) to re-establish the same value of $K$. The concentrations of reactants and products will change, but their ratio, as defined by $K$, will remain the same, provided the temperature is constant.

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

Pure solids and pure liquids have constant molar concentrations (or densities) at a given temperature, regardless of the amount present. Their 'concentration' doesn't change during the reaction. Since they are constant, their values are effectively incorporated into the equilibrium constant itself. Including them explicitly would just make the constant appear different without adding any new information about the system's dynamic state.

What does a very large value of $K$ (e.g., $K > 10^3$) signify?

A very large value of $K$ indicates that at equilibrium, the concentration of products is significantly higher than the concentration of reactants. This means the reaction proceeds almost to completion in the forward direction. In essence, the equilibrium lies far to the right, favoring the formation of products.

How can we predict the direction of a reaction if it's not at equilibrium?

We use the reaction quotient ($Q$). The reaction quotient has the same mathematical form as the equilibrium constant but uses current (non-equilibrium) concentrations. If $Q K$, the reaction will proceed in the reverse direction. If $Q = K$, the system is already at equilibrium.

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Law of Chemical Equilibrium

Chemistry

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Version 1Updated 22 Mar 2026

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The Law of Chemical Equilibrium, also known as the Law of Mass Action, states that at a given temperature, the ratio of the product of the molar concentrations (or partial pressures) of the products, each raised to the power of its stoichiometric coefficient, to the product of the molar concentrations (or partial pressures) of the reactants, each raised to the power of its stoichiometric coefficie…

Quick Summary

The Law of Chemical Equilibrium, or Law of Mass Action, describes the quantitative relationship between reactants and products in a reversible reaction at equilibrium. Equilibrium is a dynamic state where the rates of forward and reverse reactions are equal, leading to constant macroscopic properties like concentrations.

The equilibrium constant ( $K_c$ for concentrations, $K_p$ for partial pressures) is the ratio of product concentrations (raised to stoichiometric powers) to reactant concentrations (raised to stoichiometric powers) at equilibrium.

This constant is temperature-dependent but independent of initial concentrations. A large $K$ indicates product-favored equilibrium, while a small $K$ indicates reactant-favored equilibrium. For gaseous reactions, $K_p = K_c(RT)^{Delta n_g}$ , where $Delta n_g$ is the change in moles of gaseous species.

The reaction quotient ( $Q$ ) is used to predict the direction a reaction will shift to reach equilibrium: if $Q < K$ , forward shift; if $Q > K$ , reverse shift; if $Q = K$ , at equilibrium. Pure solids and liquids are excluded from $K$ expressions as their concentrations are constant.

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Key Concepts

Equilibrium Constant (

K_c

)

The equilibrium constant $K_c$ is a fundamental value for any reversible reaction at a specific temperature.…

Relationship between

K_c

and

K_p

For reactions involving gases, the equilibrium constant can be expressed in terms of partial pressures…

Reaction Quotient (

Q

) and Reaction Direction

The reaction quotient ( $Q$ ) is a powerful tool to predict the direction a reaction will shift to reach…

Law of Mass Action: — For

aA + bB \rightleftharpoons cC + dD

, Rate

\propto [Reactants]^{coeff}

Equilibrium Constant ($K_c$): —

K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}

(at equilibrium, in molar concentrations).

Equilibrium Constant ($K_p$): —

K_p = \frac{P_C^c P_D^d}{P_A^a P_B^b}

(at equilibrium, in partial pressures).

Relation between $K_p$ and $K_c$: —

K_p = K_c (RT)^{Delta n_g}

, where

\Delta n_g = (c+d) - (a+b)

for gaseous species.

Reaction Quotient ($Q_c$): —

Q_c = \frac{[C]_t^c[D]_t^d}{[A]_t^a[B]_t^b}

(at any time

t

Predicting Reaction Direction: — If

Q < K

, forward; if

Q > K

, reverse; if

Q = K

, equilibrium.

Heterogeneous Equilibria: — Pure solids and liquids are excluded from

K

expressions.

King Queen Really Think Diamonds Nice Gems: $K_p = K_c (RT)^{Delta n_g}$

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