One of the core goals in chemical synthesis is to ensure the maximum conversion of reactants into products while using the least amount of energy. To achieve this efficiently, it is important to understand how different external factors influence the equilibrium of a reaction. Although the equilibrium constant (K) for a given reaction remains constant at a fixed temperature, the position of equilibrium can shift significantly when certain conditions—such as concentration, pressure, temperature, or catalysts—are altered. Each of these factors affects equilibrium in unique ways, governed by the well-known Le Chatelier’s Principle, which provides a predictive approach to understanding these changes.
Le Chatelier’s Principle
Le Chatelier’s Principle forms the foundation for understanding equilibrium shifts. It states that:
“When a system at equilibrium experiences a change in concentration, pressure, or temperature, the equilibrium will shift in a direction that tends to oppose that change and restore a new equilibrium.”
This means that whenever an equilibrium system is disturbed, it naturally adjusts itself to minimize the effect of the disturbance. Le Chatelier’s Principle applies universally to both chemical and physical equilibria and is a cornerstone concept for predicting the direction of chemical reactions in changing environments.
For example, in industrial chemistry—such as the Haber process for ammonia or the Contact process for sulphuric acid—this principle helps determine the best conditions to maximize yield and minimize costs.
Effect of Concentration Change
When the concentration of any reactant or product changes, the equilibrium shifts to counteract the disturbance and re-establish balance.
- Adding reactants: The equilibrium shifts toward products, i.e., the forward direction, to consume the added reactants.
- Removing reactants: The equilibrium shifts toward reactants, i.e., the reverse direction, to replenish what was removed.
- Adding products: The equilibrium shifts toward reactants, i.e., the backward direction, to decrease the product concentration.
- Removing products: The equilibrium shifts forward, forming more product to replace what was removed.
Example: Formation of Hydrogen Iodide
H₂(g) + I₂(g) ⇌ 2HI(g)
When hydrogen is added to the system, the equilibrium shifts to the right to form more hydrogen iodide (HI). If HI is removed as it forms, the reaction continues to produce even more HI. Conversely, if extra HI is introduced, the equilibrium shifts left to form more H₂ and I₂, demonstrating the system’s natural resistance to change.
Visual Example: Iron(III) Thiocyanate Reaction
Consider the reaction:
Fe³⁺(aq) + SCN⁻(aq) ⇌ [Fe(SCN)]²⁺(aq)
This equilibrium forms a deep red complex ion, [Fe(SCN)]²⁺, from colourless Fe³⁺ and SCN⁻ ions.
If Fe³⁺ or SCN⁻ ions are added, the solution becomes darker red, indicating a shift toward the formation of more product. If either ion is removed or diluted, the red colour fades, showing a shift in the reverse direction. This simple colorimetric experiment clearly demonstrates how concentration changes influence equilibrium shifts.
Effect of Pressure Change
Pressure affects equilibrium only in reactions involving gases, since gases occupy variable volumes. A change in pressure alters the total number of gas molecules, thereby shifting the equilibrium position.
- Increasing pressure: The equilibrium shifts to the side with fewer gas molecules to reduce pressure.
- Decreasing pressure: The equilibrium shifts to the side with more gas molecules to increase pressure.
Example: Methanation Reaction
CO(g) + 3H₂(g) ⇌ CH₄(g) + H₂O(g)
Here, four moles of gaseous reactants form only two moles of products. Increasing pressure favours the forward direction (formation of methane and water), while reducing pressure favours the reverse direction.
Example: Carbon Monoxide–Carbon Dioxide System
C(s) + CO₂(g) ⇌ 2CO(g)
In this case, an increase in pressure drives the equilibrium backward since the forward reaction increases the number of gas molecules (from one to two).
This relationship between pressure and equilibrium is especially important in designing chemical reactors where volume and gas flow need to be optimized for efficiency.
Effect of Inert Gas Addition
When an inert gas (like argon) is introduced at constant volume, it does not affect the equilibrium composition, because the partial pressures of reacting gases remain unchanged. However, if added at constant pressure, the total volume of the system increases, which reduces the partial pressures of all gases. This causes the equilibrium to shift toward the side with more moles of gas, attempting to restore balance.
This principle explains why industrial processes carefully control the introduction of non-reactive gases to avoid unintended equilibrium shifts.
Effect of Temperature Change
Temperature influences equilibrium in a dual manner—it affects both the position of equilibrium and the value of the equilibrium constant (K₍c₎). The direction of the equilibrium shift depends on whether the reaction is exothermic or endothermic.
- For exothermic reactions (ΔH < 0): Increasing temperature shifts the equilibrium backward, reducing product yield, while decreasing temperature favours the forward reaction.
- For endothermic reactions (ΔH > 0): Increasing temperature shifts the equilibrium forward, increasing product yield, while decreasing temperature favours the reverse direction.
Example: The Haber Process
N₂(g) + 3H₂(g) ⇌ 2NH₃(g); ΔH = -92.38 kJ mol⁻¹
In the Haber process, the reaction is exothermic. Raising the temperature reduces ammonia yield by driving the equilibrium backward. However, lowering temperature slows down the reaction rate. Hence, industrial conditions are carefully balanced at about 500°C and 200 atm to ensure a satisfactory yield at an acceptable rate.
Demonstration: Nitrogen Dioxide and Dinitrogen Tetraoxide
2NO₂(g) ⇌ N₂O₄(g); ΔH = -57.2 kJ mol⁻¹
At higher temperatures, the mixture appears brown (due to NO₂), while at lower temperatures it becomes colourless (due to N₂O₄ formation). This reversible colour change demonstrates how temperature influences equilibrium.
Similarly, an endothermic reaction like:
[Co(H₂O)₆]²⁺(aq) + 4Cl⁻(aq) ⇌ [CoCl₄]²⁻(aq) + 6H₂O(l)
shows a pink colour at room temperature and turns blue when heated, confirming that the forward reaction absorbs heat.
Effect of a Catalyst
A catalyst accelerates the attainment of equilibrium but does not change the equilibrium position or the value of K₍c₎ or K₍p₎. It simply provides an alternative pathway with a lower activation energy for both forward and reverse reactions. Therefore, equilibrium is achieved faster but with the same final composition.
Industrial Examples
- Haber Process: Iron acts as a catalyst for the synthesis of ammonia from nitrogen and hydrogen. It ensures faster equilibrium attainment even at moderate temperatures.
- Contact Process: In the manufacture of sulphuric acid, V₂O₅ or platinum serves as a catalyst in the oxidation of SO₂ to SO₃:
2SO₂(g) + O₂(g) ⇌ 2SO₃(g); K₍c₎ = 1.7 × 10²⁶
Although the equilibrium favours product formation, the reaction is slow without a catalyst. The presence of V₂O₅ drastically speeds it up.
Note: Catalysts cannot make a non-spontaneous reaction occur; if the equilibrium constant is extremely small, even a catalyst offers little effect because the reaction is thermodynamically unfavourable.
Practical Implications in Industry
Understanding equilibrium shifts is vital for designing efficient industrial processes:
- Ammonia Synthesis (Haber Process): Balance between pressure, temperature, and catalyst choice ensures maximum NH₃ yield.
- Sulphuric Acid Production (Contact Process): Temperature and catalyst optimization improves SO₃ output.
- Methanol Production: Adjusting pressure and temperature maximizes methanol yield from CO and H₂.
Industrial chemists rely on Le Chatelier’s Principle to fine-tune conditions for both cost efficiency and high productivity.
Key Takeaways for NEET & JEE Aspirants
- Le Chatelier’s Principle predicts equilibrium response to changes in external conditions.
- Pressure, temperature, and concentration are the main variables affecting equilibrium.
- Catalysts only influence the rate, not the position of equilibrium.
- Understanding equilibrium shifts is crucial for optimizing reaction conditions in real-world industrial chemistry.
- Practising equilibrium calculations using this principle strengthens conceptual understanding for competitive exams.
Conclusion
The factors affecting equilibrium define how chemical systems adapt to external changes. Le Chatelier’s Principle remains an essential guide for predicting these adjustments. Whether it’s maximizing industrial yields or solving complex equilibrium problems in exams, mastery of this topic empowers learners to connect theoretical chemistry with practical applications. For NEET and JEE aspirants, this understanding not only enhances problem-solving skills but also deepens comprehension of how thermodynamics governs chemical processes.
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