State Le Chateliers Principle, Examples, And Its Applications
Le Chatelier's Principle is a fundamental concept in chemistry that is essential for JEE Main 2025 preparation. It explains how a system at equilibrium responds to external changes like concentration, temperature, or pressure. This principle helps predict the direction of a reaction when these factors are altered, making it a key topic for understanding chemical reactions. For JEE Main 2025, mastering this principle is crucial, as it frequently appears in theoretical and numerical questions.
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Le Chatelier’s Principle
When a system's equilibrium state is disturbed, a net reaction occurs in some direction causing the system to return to its original equilibrium state. Le Chatelier’s principle gives insight into how the system responds when equilibrium is altered. It states that if there is a change in any of the factors that govern the equilibrium conditions of a system, then the system will undergo a change to reduce the effect of the change. Le Chatelier’s law helps to give a qualitative prediction about the effect of a change in equilibrium conditions. It can be applied to both physical and chemical systems.
Law of Chemical Equilibrium
The Law of chemical equilibrium states that at a particular temperature, the equilibrium constant Kc is equal to the product of concentrations of products raised to their respective stoichiometric coefficients divided by the product of concentrations of reactants raised to their respective stoichiometric coefficients in the balanced chemical equation.
For a general reaction:
\[aA\text{ + bB }\rightleftharpoons \text{ cC + dD}\]
\[{{\text{K}}_{c}}\text{ = }{{\dfrac{{{\left[ C \right]}^{c}}\left[ D \right]}{{{\left[ A \right]}^{a}}{{\left[ B \right]}^{b}}}}^{d}}\]
Here, $[A]$ and $[B]$ are concentrations of reactants at equilibrium, while $[C]$ and $[D]$ are the equilibrium concentration of products. The equilibrium constant $K_c$ can be used to predict the direction of the reaction. It is only applicable at the equilibrium concentrations. Another constant called the reaction quotient Qc can also be defined similarly to the equilibrium constant except that Qc can be obtained anytime in the reaction.
If Qc > Kc, when decreasing the value of Qc, the reaction proceeds in the reverse direction.
If Qc < Kc, when increasing the value of Qc, the reaction proceeds in forward direction.
If Qc = Kc, the reaction will be in equilibrium.
Discussion on Le Chatelier’s Principle
Effect of Concentration Change
If the equilibrium is disturbed by the addition/removal of any reactant/ products, according to Le Chatelier’s principle, the reaction will proceed to minimise the effect of concentration changes by the following methods:
The concentration stress of an added reactant/product is nullified by changing the course of net reaction in the direction that consumes the added substance.
The concentration stress of a reactant/product which is removed is nullified by changing the direction of net reaction in such a way that it replenishes the removed substance.
Consider the reaction:
\[{{N}_{2}}\left( g \right)+3\text{ }{{H}_{2}}\left( g \right)\rightleftharpoons 2\text{ }N{{H}_{3}}\left( g \right)\]
Addition of a Reactant: When we add more nitrogen to the system, the concentration stress is relieved by shifting the reaction towards more nitrogen being consumed, i.e., the reaction proceeds in a forward direction. Equilibrium shifts toward a forward direction. The addition of a reactant at equilibrium results in a lower Qc compared to Kc. This favours a forward reaction.
Removal of Product: To reduce the effect, the removal of product equilibrium shifts to the forward direction.
Removal of Reactant/ Addition of Product: When a reactant is removed or the product formed is added in the reaction mixture, the reaction shifts to a backward direction, which favours the replenishment of reactant or consumption of the excess product.
For heterogeneous reactions like dissociation of calcium carbonate:
$\mathrm{CaCO}_{3}(\mathrm{~s}) \rightleftharpoons \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g})$
The equilibrium doesn’t depend on the amount of solids as the concentration of pure solids doesn't change.
Effect of Pressure Change
A pressure change can affect the yield of reactants and products in the gaseous state when different moles of reactants and products are present in the gaseous state.
Consider the reaction:
\[{{N}_{2}}\left( g \right)+3\text{ }{{H}_{2}}\left( g \right)\rightleftharpoons 2\text{ }N{{H}_{3}}\left( g \right)\]
Here, 4 moles reactants in the gaseous state yield 2 moles of product in the gaseous state.
Increase in Pressure: At constant temperature and volume, the amount of a gas is directly proportional to the pressure of the gas. If we increase the system pressure in the above written gas phase reaction, the equilibrium will shift to where there is less pressure. Here, the reaction shifts towards a forward direction as it produces less moles of gas molecules and hence decreases the pressure.
Decrease in Pressure: It will have a reverse effect and the system will shift towards increased pressure /a greater number of moles of gaseous reactant produced, which means backward reaction will increase.
Effect of Inert Gas
If inert gas is added at a constant volume, then there will be no effect on the equilibrium state. This is because, at constant volume, if inert gas is added, then total pressure increases but the number of moles of reactants and products per unit volume won’t change.
If inert gas is added at constant pressure, then the number of moles of reactants and products per unit volume will decrease, which decreases the individual partial pressure of each gas, and reaction will shift in direction where the number of moles increases.
For example: in reaction $\mathrm{N}_{2}+3 \mathrm{H}_{2} \rightleftharpoons 2 \mathrm{NH}_{3}$, the addition of inert gas at constant pressure will shift the reaction to the left (More number of moles) and will result in a decrease in product formation.
Effect of Temperature Change
When a temperature change occurs during a reaction, the value of the equilibrium constant Kc is changed. The change depends on the sign of enthalpy change $\Delta H$ for the reaction.
Exothermic Reaction (negative $\Delta H$): For this reaction, the equilibrium constant decreases as the temperature increases.
Endothermic Reaction (positive $\Delta H$): Here, the equilibrium constant increases as the temperature increases.
Ammonia synthesis is an exothermic reaction. According to Le Chatelier’s principle, an increase in temperature shifts the equilibrium to a backward reaction and decreases the equilibrium concentration of ammonia in ammonia synthesis. So, we can say that low temperature gives a high yield of ammonia. Practically, very low temperatures slow down the reaction. Thus, a catalyst is used to get a high yield.
Effect of a Catalyst
A catalyst increases the rate of the chemical reaction without taking part in the reaction. It is possible because the catalyst shifts the reaction to a low energy pathway. It does not affect equilibrium. The catalyst decreases the activation energy for the forward and reverse reactions by exactly the same amount.
Effect of Catalyst
Le Chatelier’s Principle Examples
These are some examples of the application of Le Chatelier’s principle:
When we put clothes for drying, an equilibrium is established between the gaseous molecules of water evaporated and water in the liquid form present in clothes. When air breezes, this equilibrium is disturbed as gaseous molecules decrease. Hence, in order to increase gas molecules, the reaction shifts in the direction where gaseous molecules increase. Thus, clothes dry quickly.
When blood carrying oxyhemoglobin reaches a tissue where oxygen partial pressure is less, oxyhemoglobin breaks down and gives oxygen to the tissue.
In Haber's process for the manufacture of ammonia in industries, Le Chatelier's principle is often used in order to increase the yield of ammonia by removing it. This leads to an increase in the yield of ammonia as the reaction moves forward.
$\mathrm{N}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})$
Significance of Le Chatelier’s Principle for JEE Main 2025
Le Chatelier’s Principle is fundamental to understanding chemical equilibrium and reaction dynamics, making it a frequently tested concept in JEE Main.
This principle forms the basis of 2-3 questions in JEE Main every year, appearing in both numerical and theoretical formats.
It is crucial for predicting the effects of changes in pressure, temperature, and concentration in various reactions, including industrial processes like the Haber process.
Le Chatelier’s Principle is a stepping stone to mastering topics such as buffer solutions, solubility equilibria, and phase equilibrium, which are often part of JEE Main questions.
The principle's real-world applications, like the production of ammonia and carbonated drinks, make it relatable and easy to understand, helping students grasp the concept better.
Questions related to this principle are usually straightforward, providing students with an excellent opportunity to score marks if prepared well.
Tips for Preparing Le Chatelier’s Principle for JEE Main 2025
Understand Key Concepts: Focus on how equilibrium shifts with changes in pressure, temperature, and concentration.
Practice Numerical Problems: Solve questions on equilibrium constants (Kc, Kp) and reaction quotients (Qc).
Learn Common Reactions: Study examples like the Haber process and solubility equilibria for real-world applications.
Memorise Effects: Know how endothermic and exothermic reactions respond to temperature changes.
Use Diagrams: Visualise equilibrium shifts with graphs and reaction setups.
Mock Tests & PYQs: Regularly solve previous year questions and mock tests to build confidence.
Revise Often: Use summary notes and flashcards for quick revision of formulas and principles.
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Conclusion
The equilibrium constant has a constant value at a fixed temperature and helps to predict the direction of a reaction. Le Chatelier’s principle states that the change in any condition of a reaction system such as temperature, pressure, concentration, etc. will cause the equilibrium to change the direction of reaction so that it will reduce the effect of the change. It is very helpful to determine the direction of equilibrium and to control the yield of products by controlling these factors.
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FAQs on Chemistry Le Chatelier's Principle: JEE Main 2025
1. State the law governing gas-liquid equilibrium.
Some gases which do not react with liquids may get dissolved in liquid in direct proportion to the liquid pressure. When the liquid and gas under pressure are kept in a closed container, an equilibrium is reached between the gas in the container and the gas dissolved in the liquid. For example, the carbon dioxide gas will be in equilibrium with that dissolved in soft drinks. This equilibrium is governed by Henry’s law, which states that the mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the solvent. This amount decreases with the increase in temperature.
2. Give examples of equilibrium in everyday life.
Since the amount of heat lost by the coffee to the environment balances the amount of heat acquired by the environment when hot coffee cools down to room temperature, it is considered to have reached thermal equilibrium with the space. When the rate of the forward reaction equals the rate of the backward reaction, reversible processes reach chemical equilibrium (converting products to reactants). Reproduction of a species at a rate that is equal to or greater than its mortality rate is an illustration of biological equilibrium.
3. What is Le Chatelier’s Principle?
Le Chatelier’s Principle states that a system at equilibrium will adjust to counteract changes in concentration, temperature, or pressure to maintain balance.
4. How to state Le Chatelier’s Principle?
To state Le Chatelier’s Principle, say that any disturbance in a system at equilibrium will cause the system to shift in a direction that opposes the change.
5. What is Le Chatelier’s Principle pressure?
Le Chatelier’s Principle pressure explains that increasing pressure favours the reaction side with fewer gas molecules, while decreasing pressure favours the side with more gas molecules.
6. What are Le Chatelier’s Principle examples?
Examples of Le Chatelier’s Principle include the Haber process for ammonia synthesis and changes in solubility equilibrium with concentration shifts.
7. What are real world examples of Le Chatelier’s Principle?
Real world examples include carbonation in soft drinks, industrial ammonia production, and the balance of oxygen and haemoglobin in the blood.
8. How does Le Chatelier’s Principle temperature work?
Le Chatelier’s Principle temperature states that increasing temperature favours endothermic reactions, while decreasing temperature favours exothermic reactions.
9. Why is Le Chatelier’s Principle important in chemistry?
It helps predict the behaviour of chemical reactions under changing conditions, crucial for equilibrium and reaction dynamics.
10. What are the limitations of Le Chatelier’s Principle?
Le Chatelier’s Principle does not consider the reaction rate or kinetic factors and assumes the system reaches a new equilibrium.
11. How to apply Le Chatelier’s Principle in numerical problems?
To apply, identify the change (pressure, temperature, or concentration), predict the direction of shift, and solve for equilibrium constants or concentrations.
12. What are common misconceptions about Le Chatelier’s Principle?
A common misconception is that the principle applies only to gaseous systems; it also applies to solutions and other equilibria.
