Name: chemical equilibrium and kc rules
Uploaded: 2025-08-11T09:37:07.000Z
Duration: 3 min 10 s
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K12 Standards

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NC K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.5.2

The Equilibrium Constant

The law of mass action introduces the equilibrium constant to relate the equilibrium concentrations of the reactants and products.
For a balanced chemical equation where A and B are reactants, C and D are products, and lowercase a, b, c, and d, are their respective stoichiometric coefficients, the equilibrium constant expression is given by the molar concentrations of the products divided by the molar concentrations of the reactants, all raised to their stoichiometric coefficients.
The equilibrium constant is often expressed with the symbol K or Kc, where the subscript c indicates the equilibrium constant with respect to molar concentration. Being a ratio of molar concentrations, the equilibrium constant is unitless.
A large equilibrium constant much greater than 1 means that the numerator, that is the concentration of the products, is greater than the denominator, the concentration of the reactants. This indicates that equilibrium lies towards the products and favors the forward reaction.
Conversely, a small equilibrium constant much less than 1 means that the denominator is greater than the numerator. This signifies that the equilibrium lies towards the reactants and favors the reverse reaction.
For an equilibrium constant of magnitude equal to or close to 1, the relative concentrations of the reactants and products are nearly the same. This indicates that neither reaction, forward nor reverse, is favored.
Modifications to the chemical equation also change the equilibrium constant.
If a chemical reaction with an equilibrium constant x is reversed, the equilibrium constant expression for the reverse reaction is the reciprocal of the forward reaction. So, the new equilibrium constant would be 1 over x.
If the coefficients are multiplied by another factor, n, the equilibrium constant is raised to the same factor. The new equilibrium constant would now be xn.
When two or more individual reactions are added together, the equilibrium constant for the overall reaction is the product of the individual equilibrium constants. So, the new equilibrium constant would be y &#215; z.
For a heterogeneous equilibrium, the equilibrium constant expression excludes pure solid or pure liquid entities.
For example, the equilibrium constant expression for the decomposition of solid ammonium nitrate to gaseous nitrous oxide and water vapor, excludes ammonium nitrate&#8212;a pure solid.
This is because the relative molar concentration of a pure solid or liquid remains constant during the reaction.

the equilibrium constant

Consider the oxidation of sulfur dioxide:
<img alt="Synthetic division process; polynomial long division diagram; mathematics educational illustration." src="/files/ftp_upload/11386/11386_Equation_1.png" style="text-align: center; height: 50px;" data-altupdatedat="1754904545000" data-alt="Eq1" />
For a reaction that begins with a mixture of reactants only, the product concentration is initially equal to zero. As the reaction proceeds toward equilibrium in the forward direction, the reactant concentrations decrease and the product concentration increases. When equilibrium is achieved, the concentrations of reactants and products remain constant.
If the reaction begins with only the products present, the reaction proceeds toward equilibrium in the reverse direction. The product concentration decreases with time and the reactant concentrations increase until the concentrations become constant at equilibrium.
The law of mass action states that the ratio of the concentration of products to the concentration of reactants at equilibrium, raised to their respective stoichiometric coefficients, is equal to a constant, called the equilibrium constant, K or Kc.
Thus, the equilibrium constant expression for the above reaction is written as:
<img alt="Static equilibrium diagram, &#931;Fx=0, &#931;Fy=0 equations, illustrating force balance concepts." src="/files/ftp_upload/11386/11386_Equation_2.png" style="text-align: center; height: 100px;" data-altupdatedat="1754904545000" data-alt="Eq2" />
where, the subscript &#8216;c&#8217; indicates that the equilibrium constant considers the molar concentration of reactants and products.
The magnitude of equilibrium constant explicitly reflects the composition of a reaction mixture at equilibrium. A reaction exhibiting a large K will reach equilibrium when most of the reactant has been converted to product, whereas a small K indicates the reaction achieves equilibrium after very little reactant has been converted. It&#8217;s important to keep in mind that the magnitude of K does not indicate how rapidly or slowly equilibrium will be reached. Some equilibria are established so quickly as to be nearly instantaneous, and others so slowly that no perceptible change is observed over the course of days, years, or longer. The equilibrium constant for a reaction can be used to predict the behavior of mixtures containing its reactants and/or products. As demonstrated by the sulfur dioxide oxidation process described above, a chemical reaction will proceed in whatever direction is necessary to achieve equilibrium.
Coupled Equilibria
Many equilibrium systems involve two or more coupled equilibrium reactions, those which have in common one or more reactant or product species. The K value for a system involving coupled equilibria can be related to the K values of the individual reactions. Three basic manipulations are involved in this approach, as described below: 
 
&#8226; Changing the direction of a chemical equation essentially swaps the identities of &#8220;reactants&#8221; and &#8220;products,&#8221; and so the equilibrium constant for the reversed equation is simply the reciprocal of that for the forward equation. 
 
&#8226; Changing the stoichiometric coefficients in an equation by some factor x results in an exponential change in the equilibrium constant by that same factor. 
 
&#8226; Adding two or more equilibrium equations together yields an overall equation whose equilibrium constant is the mathematical product of the individual reaction&#8217;s K values.
Equilibrium Constant expression for heterogeneous equilibria
For heterogeneous equilibria, involving reactants and products in two or more different phases, the concentrations of pure solids or pure liquids are not included in the equilibrium constant expression, as illustrated by the following example:
<img alt="Static equilibrium diagram with forces &#931;Fx=0, &#931;Fy=0 equations for mechanical analysis." src="/files/ftp_upload/11386/11386_Equation_3.png" style="text-align: center; height: 150px;" data-altupdatedat="1754904545000" data-alt="Eq3" />
This is because the relative concentrations for pure liquids and pure solids remain constant during the reaction.
This text has been adapted from <a href="https://openstax.org/books/chemistry-2e/pages/13-2-equilibrium-constants">Openstax, Chemistry 2e, Section 13.2 Equilibrium Constants</a>.

Consider the oxidation of sulfur dioxide:

For a reaction that begins with a mixture of reactants only, the product concentration is initially equal to ...

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JoVE Core

Chemical Equilibrium

KEY TERMS AND CONCEPTS

chemical equilibrium and kc rules

Chemical equilibrium and the equilibrium constant Kc describe how reactants and products are balanced in a reaction. In the oxidation of sulfur dioxide, a mixture that starts with reactants only has no product at first. As the reaction moves forward toward equilibrium, reactant concentrations decrease and product concentration increases. When equilibrium is reached, the concentrations of both stay constant.

A reaction can also begin with only products present. In that case, the reaction moves in reverse toward equilibrium. The product concentration falls with time, and the reactant concentrations rise until both become constant. This shows that a chemical reaction will move in whatever direction is needed to reach equilibrium.

The law of mass action links equilibrium to a constant value. It states that the ratio of product concentrations to reactant concentrations, each raised to their stoichiometric coefficients, equals the equilibrium constant, K or Kc. The subscript c means the expression uses molar concentration. The size of K shows the mixture makeup at equilibrium. A large K means most reactant has changed to product before equilibrium is reached. A small K means only a little reactant changes to product. The value of K does not tell how fast equilibrium is reached.

Some equilibria form very quickly. Others take so long that no visible change may be seen for days, years, or longer. The equilibrium constant can still be used to predict how a mixture with reactants and/or products will behave. For coupled equilibria, two or more equilibrium reactions share one or more species. The K value for the overall system can be found from the K values of the separate reactions. Reversing a chemical equation changes K to its reciprocal. Multiplying the stoichiometric coefficients by a factor x changes K by that same factor as an exponent. Adding equilibrium equations gives an overall K equal to the product of the individual K values.

For heterogeneous equilibria, the reactants and products are in different phases. Pure solids and pure liquids are not included in the equilibrium constant expression. Their concentrations stay constant during the reaction, so they do not change the value of K. This rule is an important part of writing equilibrium expressions correctly.

Chemical Equilibrium and Kc Rules

Chemical equilibrium and the equilibrium constant Kc describe how reactants and products are balanced in a reaction. In the oxidation of sulfur dioxide, a mixture that starts with reactants only has no product at first. As the reaction moves forward toward equilibrium, reactant concentrations decrease and product concentration increases. When equilibrium is reached, the concentrations of both stay constant. A reaction can also begin with only products present. In that case, the reaction moves...

le chatelier's principle: changing concentration

le-chatelier-s-principle-changing-concentration

Shifting Equilibrium by Changing Concentration

Chemical equilibrium is a dynamic balance. At equilibrium, the forward and reverse reactions happen at equal rates. If the system is stressed in a way that changes one rate more than the other, it is no longer at equilibrium and will shift until balance is restored. This behavior is described by Le Châtelier’s principle. Changes in concentration are a major stress on an equilibrium system. Reaction rates depend mainly on concentration and temperature. For that reason, changing the amount of a...

homogeneous equilibria for gaseous reactions

homogeneous-equilibria-for-gaseous-reactions

Relating Kc and Kp in Gas Equilibria

Gas-phase equilibrium can be written using either concentration or pressure. The equilibrium constant may be shown as Kc, which uses molar concentration, or as Kp, which uses partial pressure of the reactants and products. The link between Kc and Kp comes from the ideal gas equation and the definition of molarity. Molarity is the number of moles divided by volume. In the ideal gas equation, P stands for partial pressure, V for volume, n for number of moles, R for the gas constant, and T for...

calculating the equilibrium constant

Finding Kc from Equilibrium Data

The equilibrium constant, Kc, can be found from the equilibrium concentrations of reactants and products. When those concentrations are known, the calculation is straightforward. The values are placed into the Kc expression for the reaction. One example uses nitrogen dioxide and dinitrogen tetroxide. When 0.10 mol NO2 is added to a 1.0-L flask at 25 °C, the concentrations change until equilibrium is reached. At equilibrium, [NO2] is 0.016 M and [N2O4] is 0.042 M, and these values are used to...

reaction quotient

Using Q to Track Equilibrium Changes

Reaction quotient Q helps track how far a reversible reaction has moved toward equilibrium. For a balanced reaction such as m A + n B ⇌ x C + y D, Q comes directly from the stoichiometric coefficients in the equation. When concentrations are used, the expression is written as Qc. For gases, a similar form uses partial pressures. These expressions use relative concentration and pressure values, not absolute values. That means reaction quotients are dimensionless and have no units. As a reaction...

calculating equilibrium concentrations

Equilibrium Concentrations from an ICE Table

Equilibrium concentrations can be found from initial concentrations and an equilibrium constant. This skill matters in many areas of science and technology, including pharmaceutical product formulation and dosing. After a drug is ingested or injected, it may take part in several chemical equilibria that change its final concentration in the body system of interest. Scientists use these quantitative equilibrium relationships to help calculate a dosage that will produce the desired therapeutic...

dynamic equilibrium

Chemical Equilibrium in Reversible Reactions

Chemical equilibrium in reversible reactions happens when a reaction moves in both directions at the same time. A reversible reaction uses a double arrow to show that products can turn back into reactants. At equilibrium, the concentrations of reactants and products stay constant over time because the forward and reverse reaction rates are equal. The video uses the reversible decomposition of colorless dinitrogen tetroxide to brown nitrogen dioxide as an example. At time zero, there is finite...

le chatelier's principle: changing volume (pressure)

le-chatelier-s-principle-changing-volume-pressure

Gas Equilibrium Under Volume Changes

Gas-phase equilibrium can shift when the volume of the system changes. In these reactions, a change in pressure changes concentration because the partial pressure of an ideal gas is proportional to its molar concentration. That means pressure changes affect equilibrium in the same way as concentration changes. When the volume of a gas mixture changes, the concentrations of reactants and products change together. Since every gas in the equilibrium occupies the same space, the same volume change...

le chatelier's principle: changing temperature

le-chatelier-s-principle-changing-temperature

Le Chatelier's Principle in Heating and Cooling

Le Chatelier's principle helps explain how chemical equilibrium responds to heating and cooling. A system at equilibrium will shift to restore balance after a temperature change, but the equilibrium constant, K, will not stay the same. Instead, the mixture reaches a new equilibrium with a different composition. This temperature effect is different from a concentration change, which shifts equilibrium without changing K. To understand why temperature matters, the reaction can be written with...

the small x assumption

Estimating Equilibrium with Small Changes

A small equilibrium constant can make equilibrium calculations easier. In these reactions, the equilibrium position favors the reactants. When the initial reactant concentration is high and the Kc value is small, the change in concentration may be negligible. In that case, the equilibrium concentration is close to the initial concentration. This idea helps simplify weak acid and weak base problems. Chemists can assume that some equilibrium concentrations are equal to the initial concentrations.

key terms and concepts

Explore chemical equilibrium, Kc, and the law of mass action. Learn how concentration changes, reaction direction, and mixed-phase equilibria shape equilibrium expressions.

Chemical Equilibrium and Kc Rules | JoVE K12

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Pennsylvania State Academic Standards

Physical Science

Chemical Reactions

3.2.9-12.F

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Ohio’s Learning Standards and Model Curriculum

Interactions Of Matter

Chemical reactions

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Matter and Its Interactions

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Chemical Equilibrium and Kc Rules

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