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6.8.2. Effect of Pressure Change
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Today, we're diving into how pressure changes influence chemical equilibria, particularly for gaseous reactions. Can anyone tell me why pressure might affect a reaction?
I think it has to do with how many gas molecules there are on each side of the reaction.
Exactly! That’s a crucial point. When we alter the pressure, we can cause the reaction to favor the side with fewer moles of gas. This is summarized by Le Chatelier's principle which states that a system at equilibrium will adjust to minimize any applied change. Let's explore this in more depth.
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When the volume is decreased in a reaction system, what happens to the pressure?
The pressure increases since pressure and volume are inversely related.
Correct! An increased pressure will drive the equilibrium towards the side with fewer gas moles, thereby reducing the total pressure. So, if we apply this to our previous example of CO and H2 reaction, what direction will the reaction shift if we increase the pressure?
It will shift towards the products, since they have fewer moles of gas.
Great job!
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Create a free accountNow, let’s look deeper into our previous example. If we compress the gas mixture of CO and H2 in a cylinder, how do we expect the reaction to behave in terms of Qc?
The concentrations will increase, therefore Qc will also change.
Exactly. If Qc moves below Kc due to the concentration increase from pressure changes, the reaction will shift to the right to restore equilibrium. How does this apply in industry, say while producing ammonia in the Haber process?
In that case, we want to keep pressure high, since it helps form ammonia which has fewer moles of gas.
Precisely! Maximizing pressure can significantly enhance yield.
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Create a free accountBefore we wrap up, can anyone summarize why pressure changes affect equilibria?
Increasing pressure shifts equilibrium towards the side with fewer gas molecules.
Right! And decreasing pressure shifts it towards the side with more gas molecules.
It's all about balancing the reaction to minimize the effect of external changes, according to Le Chatelier’s principle.
Well done! Remember, these principles are fundamental for practical applications in fields like chemistry and engineering.
Overview
Short Summary
This section discusses how changes in pressure affect the position of equilibrium in gaseous reactions.
Medium Summary
The effect of pressure changes on equilibrium is significant only in reactions involving gases where the number of moles of reactants differs from products. When the pressure increases, the equilibrium shifts towards the side with fewer moles of gas, following Le Chatelier's principle.
Detailed Summary
Effect of Pressure Change
In chemical equilibria, pressure changes can profoundly affect gaseous reactions, especially when the number of moles of gaseous reactants differs from that of products. According to Le Chatelier's principle, if an external change is applied to a system at equilibrium, the system will adjust itself to counteract that change and restore a new equilibrium state.
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Pressure and Volume Relationship: When the volume of a gaseous reaction mixture is decreased, the pressure increases, resulting in an equilibrium shift toward the side with fewer gaseous moles. Conversely, increasing the volume lowers the pressure, leading the equilibrium to shift toward the side with more gaseous moles.
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Example Reaction: Considering the reaction:
CO(g) + 3H2(g) ⇌ CH4(g) + H2O(g)
This reaction starts with 4 moles of gaseous reactants (1 CO and 3 H2) and results in 2 moles of gaseous products (1 CH4 and 1 H2O). If the pressure is increased by compressing the volume, the equilibrium will shift to the right (producing more CH4 and H2O).
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Application of the Principle: The principle can also be illustrated using Qc, the reaction quotient. An increase in pressure causes the concentrations of gases to increase, shifting the equilibrium to the side that reduces total moles.
In summary, understanding the influence of pressure on equilibrium is crucial, especially in industrial and laboratory settings, where optimizing conditions can maximize product yield.
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Create a free accountA pressure change obtained by changing the volume can affect the yield of products in case of a gaseous reaction where the total number of moles of gaseous reactants and total number of moles of gaseous products are different.
Detailed Explanation
When we change the pressure of a system involving gases, it can shift the equilibrium position if there are different numbers of moles of gaseous reactants and products. According to Le Chatelier’s principle, when stress is applied to a system at equilibrium, the system will adjust to counteract that stress, aiming to re-establish equilibrium.
Examples & Analogies
Think of a crowded subway train: When the doors open at a station (analogous to an increase in volume), passengers may get off (create more space), leading to a change in the density of people inside the train. Similarly, if the station were to close the doors and increase pressure (decrease volume), people would be forced to press against one another, but some might find an opportunity to leave the crowded space.
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Create a free accountConsider the reaction, CO(g) + 3H2(g) CH4(g) + H2O(g). Here, 4 mol of gaseous reactants (CO + 3H2) become 2 mol of gaseous products (CH4 + H2O).
Detailed Explanation
This particular reaction reduces the number of gaseous moles from four to two when producing methane and water. If we compress the equilibrium mixture (decrease the volume), the pressure increases, causing the equilibrium to shift towards the products to reduce pressure. This is because the forward reaction results in fewer moles of gas, effectively lowering the pressure.
Examples & Analogies
You can visualize this with a balloon. When you squeeze a balloon (decrease its volume), the air inside becomes denser. The air pressure increases until some air escapes. In chemical reactions, if the pressure rises, systems will favor pathways that reduce the number of gas particles (like letting some air out).
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Create a free accountSuppose the equilibrium mixture kept in a cylinder fitted with a piston at constant temperature is compressed to one half of its original volume. Then, total pressure will be doubled. The direction in which the reaction goes to re-establish equilibrium can be predicted by applying Le Chatelier’s principle.
Detailed Explanation
When you compress the cylinder, the total pressure increases, and Qc (the reaction quotient) will shift. Since the system is no longer at equilibrium, the reaction will favor the direction that produces fewer gaseous moles, ultimately creating products. In this example, the forward reaction is favored because it reduces the mole count, thereby relieving the added pressure.
Examples & Analogies
Imagine a soda can. When carbonated drink in the can is shaken, pressure builds up due to carbon dioxide gas being trapped. Opening the can releases that pressure. In a similar way, compressing a reaction system forces equilibrium to adjust to reduce the 'pressure' of the system, ultimately favoring the production of fewer moles of gas.
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Create a free accountIn reaction C(s) + CO2(g) 2CO(g), when pressure is increased, the reaction goes in the reverse direction because the number of moles of gas increases in the forward direction.
Detailed Explanation
In this example, an increase in pressure shifts the equilibrium towards the side with fewer moles of gas, which is a common response in gaseous equilibria. In general, systems at equilibrium will shift to mitigate increases in pressure, prioritizing the production of fewer gas particles.
Examples & Analogies
Consider a big party with too many people (representing gas moles) in a small room. If the door opens (pressure decreases), people can leave freely. Similarly, when the pressure increases by a system change, it forces equilibria to adjust, just like people wanting to leave a crowded room.
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Key Concepts
Core takeaways and short definitions to help you quickly recall the key ideas from this section.
Le Chatelier's principle: Predicts the shift of equilibrium in response to changes in pressure, concentration, or temperature.
Reaction quotient (Qc): The ratio of product concentrations to reactant concentrations at any point in the reaction.
Pressure effects: Increased pressure favors the side of the reaction with fewer gaseous moles.
Examples
Step-by-step examples to apply the section's ideas and test your understanding.
In the reaction CO(g) + 3H2(g) ⇌ CH4(g) + H2O(g), increasing pressure shifts the equilibrium toward the right, minimizing the number of gas moles.
For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g), a decrease in volume (which increases pressure) drives the equilibrium to form more ammonia.
Memory Aids
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Flash Cards
Glossary
Equilibrium
A state in which the forward and reverse reactions occur at the same rate in a closed system.
Le Chatelier's principle
States that if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system adjusts to reduce the change.
Molarity (M)
A measure of the concentration of a solute in a solution, expressed in moles of solute per liter of solution.
Q_c
The reaction quotient, which can provide insight on the direction of the reaction by comparing it to K_c.
K_c
The equilibrium constant for a reaction at a specific temperature, representing the ratio of product concentrations to reactant concentrations.