A dynamic equilibrium is a state of a reversible reaction where the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of reactants and products remain constant. This means that the reaction is constantly occurring in both directions, but the net change in the concentrations of the reactants and products is zero.
Understanding the Concept of Dynamic Equilibrium
In a dynamic equilibrium, the concentrations of the reactants and products are not necessarily equal, but they are constant. This can be represented by the equation:
$\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$
where the concentrations of A, B, C, and D are not changing, but they may not be equal.
The concept of dynamic equilibrium can be further illustrated using the example of a beaker of water with a small amount of food coloring added. As the food coloring spreads throughout the water, the system reaches a state of dynamic equilibrium, where the rate of diffusion of the food coloring in one direction is equal to the rate of diffusion in the opposite direction.
It’s important to note that dynamic equilibria only occur in closed systems, where no matter or energy can enter or leave the system. In an open system, the concentrations of reactants and products can change over time due to the input or output of matter or energy.
Characteristics of Dynamic Equilibrium

Constant Concentrations: In a dynamic equilibrium, the concentrations of the reactants and products remain constant over time, even though the forward and backward reactions are continuously occurring.

Equality of Reaction Rates: The rate of the forward reaction is equal to the rate of the backward reaction, resulting in a net change of zero in the concentrations of the reactants and products.

Reversibility: The reaction is reversible, meaning that the reactants can form the products, and the products can reform the reactants.

Closed System: Dynamic equilibria only occur in closed systems, where no matter or energy can enter or leave the system.

Equilibrium Constant: The equilibrium constant, denoted as $K_{eq}$, is a measure of the relative concentrations of the reactants and products at equilibrium. It is defined as the ratio of the product concentrations raised to their stoichiometric coefficients to the reactant concentrations raised to their stoichiometric coefficients.
For the reaction $\text{A} + \text{B} \rightleftharpoons \text{C} + \text{D}$, the equilibrium constant is given by:
$K_{eq} = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
where $a$, $b$, $c$, and $d$ are the stoichiometric coefficients of the reactants and products, respectively.
Examples of Dynamic Equilibria
 Haber Process: The Haber process is an industrial process used to produce ammonia (NH3) from nitrogen (N2) and hydrogen (H2) in a reversible reaction:
$\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$
The equilibrium constant for this reaction is given by:
$K_{eq} = \frac{[NH_3]^2}{[N_2][H_2]^3}$
 Dissociation of Acetic Acid: The dissociation of acetic acid (CH3COOH) in water is a reversible reaction:
$\text{CH}_3\text{COOH} + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COO}^ + \text{H}^+$
The equilibrium constant for this reaction is given by:
$K_{eq} = \frac{[CH_3COO^][H^+]}{[CH_3COOH]}$
 Evaporation and Condensation of Water: The evaporation and condensation of water in a closed container is a dynamic equilibrium:
$\text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{O}(g)$
The equilibrium constant for this reaction is given by the ratio of the partial pressure of water vapor to the vapor pressure of pure water at the same temperature.
Factors Affecting Dynamic Equilibrium
The position of a dynamic equilibrium can be affected by various factors, such as:

Temperature: Changes in temperature can shift the equilibrium position according to the Le Chatelier’s principle. For example, an increase in temperature will favor the endothermic (backward) reaction, while a decrease in temperature will favor the exothermic (forward) reaction.

Pressure: Changes in pressure can shift the equilibrium position for reactions involving gases, as described by Le Chatelier’s principle. Increasing the pressure will favor the reaction that produces fewer moles of gas.

Concentration: Adding or removing reactants or products can shift the equilibrium position, as described by Le Chatelier’s principle. Increasing the concentration of a reactant will favor the forward reaction, while increasing the concentration of a product will favor the backward reaction.

Catalysts: The addition of a catalyst can increase the rates of both the forward and backward reactions, but it does not affect the equilibrium position.
Numerical Problems and Calculations
 Problem: Consider the reversible reaction: $2\text{SO}2 + \text{O}_2 \rightleftharpoons 2\text{SO}_3$. If the equilibrium concentrations are $[SO_2] = 0.2 \text{M}$, $[O_2] = 0.1 \text{M}$, and $[SO_3] = 0.4 \text{M}$, calculate the equilibrium constant $K{eq}$.
Solution:
The equilibrium constant is given by:
$K_{eq} = \frac{[SO_3]^2}{[SO_2]^2[O_2]}$
Substituting the given values:
$K_{eq} = \frac{(0.4)^2}{(0.2)^2(0.1)} = 8$
 Problem: The equilibrium constant for the reaction $\text{N}2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ at a certain temperature is $K{eq} = 1.0 \times 10^4$. If the initial concentrations are $[N_2] = 0.10 \text{M}$ and $[H_2] = 0.30 \text{M}$, calculate the equilibrium concentrations of $\text{N}_2$, $\text{H}_2$, and $\text{NH}_3$.
Solution:
Let the change in concentration of $\text{N}2$ and $\text{H}_2$ be $x$, and the change in concentration of $\text{NH}_3$ be $2x$.
At equilibrium:
$[N_2] = 0.10 – x$
$[H_2] = 0.30 – 3x$
$[NH_3] = 2x$
Substituting these values into the equilibrium constant expression:
$K{eq} = \frac{[NH_3]^2}{[N_2][H_2]^3} = \frac{(2x)^2}{(0.10 – x)(0.30 – 3x)^3} = 1.0 \times 10^4$
Solving this equation numerically, we get $x = 0.0447 \text{M}$.
Therefore, the equilibrium concentrations are:
$[N_2] = 0.0553 \text{M}$
$[H_2] = 0.1359 \text{M}$
$[NH_3] = 0.0894 \text{M}$
Conclusion
In summary, a dynamic equilibrium is a state of a reversible reaction where the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of reactants and products remain constant. This concept is crucial for understanding the behavior of chemical systems and predicting the outcomes of chemical reactions. By understanding the characteristics, factors, and calculations involved in dynamic equilibria, you can gain a deeper insight into the fundamental principles of chemistry.
References
 Chemical Equilibria Flashcards – Quizlet. Retrieved from https://quizlet.com/569483296/chemicalequilibriaflashcards/
 Dynamic Equilibrium – Class 11 Chemistry MCQ – Sanfoundry. Retrieved from https://www.sanfoundry.com/chemistryquestionsanswersequilibriumchemicalprocessesdynamicequilibrium/
 Dynamic equilibrium (video) – Khan Academy. Retrieved from https://www.khanacademy.org/science/apchemistrybeta/x2eef969c74e0d802:equilibrium/x2eef969c74e0d802:introductiontoequilibrium/v/dynamicequilibrium
 Chem.libretexts.org. (2022). 15.3: The Idea of Dynamic Chemical Equilibrium. Retrieved from https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Introductory_Chemistry/15:_Chemical_Equilibrium/15.03:_The_Idea_of_Dynamic_Chemical_Equilibrium
 StudySmarter.co.uk. (n.d.). The Dance of Balance: Understanding Dynamic Equilibrium. Retrieved from https://www.studysmarter.co.uk/explanations/chemistry/physicalchemistry/dynamicequilibrium/.
I am Keerthi K Murthy, I have completed post graduation in Physics, with the specialization in the field of solid state physics. I have always consider physics as a fundamental subject which is connected to our daily life. Being a science student I enjoy exploring new things in physics. As a writer my goal is to reach the readers with the simplified manner through my articles.