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For reactions involving gases, however, it is usually more convenient to express the equilibrium constant in terms of
- a)temperature
- b)molar concentration of the reactants
- c)molar concentration of the products
- d)partial pressure
Correct answer is option 'D'. Can you explain this answer?
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For reactions involving gases, however, it is usually more convenient ...
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For reactions involving gases, however, it is usually more convenient ...
Understanding Equilibrium Constant for Gaseous Reactions
In chemical reactions involving gases, expressing the equilibrium constant in terms of partial pressures is often more convenient and accurate. Here’s why:
1. Definition of Equilibrium Constant (K)
- The equilibrium constant (K) quantifies the ratio of the concentrations (or pressures) of products to reactants at equilibrium.
- For gaseous reactions, it is represented as Kp when using partial pressures.
2. Importance of Partial Pressure
- Direct Relationship: The behavior of gases is closely related to their partial pressures due to the ideal gas law (PV=nRT).
- Simplified Calculations: Using partial pressures simplifies calculations, particularly when dealing with reactions under different volumes or temperatures.
3. Real-World Applications
- Dynamic Equilibrium: In dynamic systems, as gases react, their partial pressures change. Using Kp allows for a straightforward understanding of how these changes affect the reaction.
- Gas Mixtures: In reactions involving multiple gases, partial pressures provide a clear picture of how each component influences the equilibrium state.
4. Comparison with Other Methods
- Molar Concentration: While molar concentration can be used (Kc), it is less practical for gases due to varying volumes and densities.
- Temperature: Temperature is a factor that affects equilibrium but does not directly express the relationship between reactants and products.
Conclusion
In summary, using partial pressure to express the equilibrium constant for gaseous reactions offers a more precise and practical approach, facilitating easier calculations and understanding of gas behavior in dynamic systems.
In chemical reactions involving gases, expressing the equilibrium constant in terms of partial pressures is often more convenient and accurate. Here’s why:
1. Definition of Equilibrium Constant (K)
- The equilibrium constant (K) quantifies the ratio of the concentrations (or pressures) of products to reactants at equilibrium.
- For gaseous reactions, it is represented as Kp when using partial pressures.
2. Importance of Partial Pressure
- Direct Relationship: The behavior of gases is closely related to their partial pressures due to the ideal gas law (PV=nRT).
- Simplified Calculations: Using partial pressures simplifies calculations, particularly when dealing with reactions under different volumes or temperatures.
3. Real-World Applications
- Dynamic Equilibrium: In dynamic systems, as gases react, their partial pressures change. Using Kp allows for a straightforward understanding of how these changes affect the reaction.
- Gas Mixtures: In reactions involving multiple gases, partial pressures provide a clear picture of how each component influences the equilibrium state.
4. Comparison with Other Methods
- Molar Concentration: While molar concentration can be used (Kc), it is less practical for gases due to varying volumes and densities.
- Temperature: Temperature is a factor that affects equilibrium but does not directly express the relationship between reactants and products.
Conclusion
In summary, using partial pressure to express the equilibrium constant for gaseous reactions offers a more precise and practical approach, facilitating easier calculations and understanding of gas behavior in dynamic systems.
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For reactions involving gases, however, it is usually more convenient to express the equilibrium constant in terms ofa)temperatureb)molar concentration of the reactantsc)molar concentration of the productsd)partial pressureCorrect answer is option 'D'. Can you explain this answer?
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