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Equilibrium Constant

The equilibrium constant expression is a vital concept in understanding the extent of a chemical reaction and predicting its outcomes. Denoted by K, this expression quantifies the ratio of products to reactants at equilibrium.

Derived from the law of mass action, it varies for different reactions and plays a crucial role in determining the direction and rate of a reaction. By analyzing the equilibrium constant expression, scientists can gain insights into the stability of chemical systems and make informed decisions about reaction conditions.

Factors Affecting the Equilibrium Constant


  • Temperature has a significant impact on the equilibrium constant (K) value.

  • Changes in temperature can alter the direction and magnitude of K.

  • An increase in temperature favors an endothermic reaction, shifting the equilibrium towards products and increasing K.

  • Conversely, a decrease in temperature favors an exothermic reaction, shifting the equilibrium towards reactants and decreasing K.

Pressure/Concentration Changes:

  • Changes in pressure or concentration can affect the position of equilibrium and consequently alter K.

  • Increasing pressure shifts the equilibrium towards the side with fewer moles of gas molecules, thereby increasing K.

  • Decreasing pressure has the opposite effect, shifting the equilibrium towards the side with more moles of gas molecules and decreasing K.

  • Similarly, changes in concentration can also shift the equilibrium position and change K based on Le Chatelier’s principle.


  • Catalysts do not directly affect the value of K but can increase reaction rates by providing alternative reaction pathways.

  • By lowering activation energy barriers, catalysts enable reactions to reach equilibrium faster without affecting K itself.


  • The stoichiometry of a balanced chemical equation determines how changes in reactant concentrations impact K.

  • Reactant concentrations are raised to their respective coefficients when calculating Q (reaction quotient) for comparison with K.

  • If Q > K, favoring reactants; if Q < K, favoring products; if Q = K, at equilibrium.

Solving for K: Substituting into the Equilibrium Expression

To determine the equilibrium constant (K) for a given system, we can substitute known concentrations or pressures into the equilibrium expression. By doing so, we can calculate the value of K and gain insights into the nature of the reaction.

When substituting values into the equation, it is important to include only species with non-zero coefficients. This means that if a reactant or product has a coefficient of zero in the balanced chemical equation, it should not be included in the equilibrium expression.

The form of the equilibrium expression may vary depending on whether the system is homogeneous or heterogeneous.

In a homogeneous system where all reactants and products are in the same phase (e.g., all gases or all aqueous solutions), we express K using concentrations.

On the other hand, in a heterogeneous system where reactants and products are present in different phases (e.g., gas and solid), we use partial pressures to represent K.

To ensure accurate calculations, it is crucial to maintain consistency with units when substituting values into the equation. For example, if concentrations are used in one part of an equation, all concentrations should be expressed using the same unit.

Determining Equilibrium Concentrations or Pressures

To determine equilibrium concentrations or pressures in a chemical system, we can use initial conditions and stoichiometry. By utilizing ICE tables (Initial, Change, Equilibrium), we can track the changes that occur during a reaction until it reaches equilibrium.

Equilibrium concentrations/pressures using initial conditions and stoichiometry

  • Start with the given initial concentrations or pressures of reactants and products.

  • Use stoichiometry to determine the change in concentration or pressure for each species involved in the reaction.

  • Incorporate these changes into the ICE table.

Applying principles of conservation and mole ratios

  • Apply principles such as conservation of mass and mole ratios to solve for unknowns in the ICE table.

  • Determine how much reactant is consumed or product is formed at equilibrium by analyzing the changes column in the ICE table.

Comparing Q (reaction quotient) with K (equilibrium constant)

  • Calculate Q by substituting concentration or pressure values into the equilibrium constant expression.

  • Compare Q with K to determine if a system is at equilibrium:

  • If Q = K, then the system is at equilibrium.

  • If Q < K, then more products need to be formed to reach equilibrium.

  • If Q > K, then more reactants need to be consumed to reach equilibrium.

By following these steps and principles, we can determine the equilibrium concentrations or pressures for a given system. It allows us to understand how a reaction progresses towards an equilibrium mixture. Remember that this process relies on factors such as dissociation in solution or dependence on solvent properties.

Thermodynamic and Kinetic Equilibrium Expressions:

This expressions provide different insights into chemical equilibria.

Thermodynamic expressions use standard state conditions (1 atm pressure, 298K temperature).

  • They are based on the concept of free energy change (∆G) and the equilibrium constant (K).

  • The equilibrium constant expression for a given system is determined by the stoichiometric coefficients of the balanced chemical equation.

  • It represents the relationship between the concentrations (or pressures) of reactants and products at equilibrium.

Kinetic expressions consider rate constants and activation energies instead of concentrations/pressures directly.

  • They focus on the rates at which reactions occur rather than their final concentrations.

  • The rate constant (k) relates to the forward reaction, while its reverse reaction has its own rate constant (k’).

  • These expressions involve terms related to reaction rates, activation energies, and temperature.

Thermodynamic expressions are more commonly used for equilibrium calculations because they provide a direct measure of how far a reaction proceeds towards completion. However, kinetic expressions are essential for understanding the factors that influence reaction rates.

Examples: Haber Process and Exercise

The Haber process is a well-known example of an equilibrium system. It involves the synthesis of ammonia from nitrogen and hydrogen gases. In this process, the equilibrium constant expression determines the yield of ammonia.

It can be helpful to work through an exercise problem. This practice activity reinforces our understanding of how to write the expression for a given system.

Let’s consider an exercise problem involving the formation of water from hydrogen and oxygen gases:

2H2(g) + O2(g) ⇌ 2H2O(g)

In this case, the equilibrium constant expression would be:

K = [H2O]^2 / [H2]^2 * [O2]

Solving numerical problems involving K enhances our proficiency in applying the concept. By working through different scenarios and calculations, we become more comfortable with using equilibrium constant expressions in various systems.

For instance, let’s say we have a reaction where 0.5 moles of H2O, 1 mole of H2, and 0.25 moles of O2 are present at equilibrium. We can use these values to calculate the value of K for this particular system.

By plugging the concentrations into the equation and performing the necessary calculations, we can determine the value of K for this specific case.

Understanding Equilibrium Constant Expression: CO(g) + H(g) → CHOH(g)

The equation we will be focusing on is the formation of methanol from carbon monoxide (CO) and hydrogen (H) gases, represented by the chemical equation

CO(g) + H(g) → CHOH(g).

To write the equilibrium constant expression for this system, we need to consider the concentrations or pressures of the reactants and products involved. In this case, it would involve the concentrations/pressures of CO, H, and CHOH only.

Here’s an overview of what we need to consider when writing the equilibrium constant expression for this system:

  • The concentration or pressure of carbon monoxide (CO)

  • The concentration or pressure of hydrogen (H)

  • The concentration or pressure of methanol (CHOH)

By incorporating these factors into our equilibrium constant expression, we can obtain a quantitative representation of how far the reaction proceeds towards forming methanol at a given set of conditions.


The equilibrium constant expression is a fundamental concept in chemistry that quantifies the extent of a chemical reaction at equilibrium. By understanding and utilizing this expression, scientists can gain valuable insights into the behavior of chemical systems.

The equilibrium constant (K) provides information about the relative concentrations or pressures of reactants and products in a system, allowing us to predict the direction and stability of a reaction.

By mastering the factors affecting the equilibrium constant, solving for K through substitution, determining equilibrium concentrations or pressures, and exploring thermodynamic and kinetic equilibrium expressions, you will have a solid foundation to analyze chemical reactions.

Studying real-world examples such as the Haber Process and exercises involving equilibria will enhance your understanding.


What is the purpose of an equilibrium constant expression?

The purpose of an equilibrium constant expression is to quantify the extent to which a chemical reaction reaches its state of balance or equilibrium. It provides valuable information about the relative concentrations or pressures of reactants and products in a system.

How do I determine the equilibrium concentrations or pressures?

To determine the equilibrium concentrations or pressures, you need to set up an ICE (Initial-Change-Equilibrium) table. By applying stoichiometry principles and considering any initial conditions given in the problem, you can calculate these values.

Can thermodynamic and kinetic factors affect an equilibrium constant?

Yes, both thermodynamic (temperature-dependent) and kinetic (rate-dependent) factors can influence an equilibrium constant. Changes in temperature can shift the position of an equilibrium while changes in reaction rates can impact how quickly it is reached.

What are some common examples where we use equilibrium constants?

Equilibrium constants are frequently used in various chemical processes, such as the Haber Process for ammonia synthesis, acid-base reactions, and solubility equilibria. They play a crucial role in understanding and predicting the behavior of these systems.

How can I apply equilibrium constant expressions to gas phase equilibria?

For gas phase equilibria, equilibrium constant expressions are often expressed in terms of partial pressures. By considering the stoichiometry of the reaction and applying the ideal gas law, you can relate partial pressures to concentrations and determine equilibrium constants.