Le Chatelier's principle
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Intros
Examples
Lessons
- Apply Le Chatelier's principle to predict changes in equilibrium position.
The reaction equation below shows an EXOTHERMIC reaction at equilibrium:2A(g) + B(g) C(g) + D(g) = -89 kJ mol-1
Predict and explain the change in equilibrium position with the following changes of conditions happening separately:
Topic Notes
Introduction to Le Chatelier's Principle
Welcome to our exploration of Le Chatelier's Principle, a fundamental concept in chemistry that helps us understand how chemical systems respond to changes. This principle, named after French chemist Henry Louis Le Chatelier, is crucial for predicting the behavior of reactions at dynamic equilibrium. When we alter reaction conditions such as temperature, pressure, or concentration, Le Chatelier's Principle tells us how the system will adjust to counteract these changes. Our introductory video will guide you through this fascinating principle, making it easier to grasp its applications in real-world scenarios. You'll learn how chemists use this principle to optimize industrial processes and how it applies to natural phenomena. By understanding Le Chatelier's Principle, you'll gain valuable insights into the delicate balance of chemical reactions and how small changes can have significant effects. This knowledge is essential for anyone studying chemistry or related fields, providing a powerful tool for analyzing and controlling chemical systems.
Understanding Dynamic Equilibrium
Introduction to Dynamic Equilibrium
Dynamic equilibrium is a fundamental concept in chemistry that describes a state where two opposing processes occur at the same rate, resulting in no net change in the system. This balance is crucial in many chemical reactions and plays a significant role in various natural and industrial processes.
The Example Reaction: A + 2B C + D
To illustrate dynamic equilibrium, let's consider the reaction A + 2B C + D. In this reaction, reactants A and B combine to form products C and D, while simultaneously, C and D can react to reform A and B. The double arrow () indicates that the reaction can proceed in both directions.
Forward and Reverse Reactions
In the context of our example, the forward reaction is the process where A and B combine to form C and D. This can be represented as A + 2B C + D. The reverse reaction, on the other hand, is when C and D react to reform A and B, represented as C + D A + 2B. Both these reactions occur simultaneously in a closed system.
Reaction Rates and Equilibrium
The rate of a reaction refers to the speed at which reactants are converted into products or vice versa. In our example, the forward reaction rate is the speed at which A and B are consumed to form C and D. The reverse reaction rate is the speed at which C and D are consumed to reform A and B. As the reaction progresses, these rates change.
Achieving Dynamic Equilibrium
Initially, the forward reaction rate may be higher as there are more reactants (A and B) available. As the reaction proceeds, the concentration of reactants decreases while the concentration of products increases. This leads to a decrease in the forward reaction rate and an increase in the reverse reaction rate. Eventually, a point is reached where the forward and reverse reaction rates become equal. This state is known as dynamic equilibrium.
Characteristics of Dynamic Equilibrium
At dynamic equilibrium, several key features are observed:
- The concentrations of reactants and products remain constant over time.
- The forward and reverse reaction rates are equal.
- The system appears macroscopically static, but microscopic changes continue to occur.
- The equilibrium can be approached from either direction (reactants or products).
Importance of Equal Reaction Rates
The equality of forward and reverse reaction rates at equilibrium is crucial. It ensures that for every molecule of A and B that react to form C and D, an equal number of C and D molecules react to reform A and B. This balance maintains the constant concentrations observed at equilibrium.
Dynamic Nature of Equilibrium
Despite the constant concentrations, it's important to understand that equilibrium is dynamic. Molecules are continuously reacting and reforming. The term "dynamic" in dynamic equilibrium emphasizes this ongoing molecular activity, even though no net change is observed at the macroscopic level.
Factors Affecting Dynamic Equilibrium
Several factors can influence dynamic equilibrium:
- Changes in concentration of reactants or products
- Temperature variations
- Pressure changes (for gaseous reactions)
- Addition of catalysts
These factors can shift the equilibrium position, altering the relative concentrations of reactants and products while maintaining the dynamic nature of the equilibrium.
Conclusion
Understanding dynamic equilibrium is essential in chemistry and related fields. It explains how chemical systems achieve balance and helps predict how they will respond to changes. The concept of equal forward and reverse reaction rates at equilibrium provides insight into the molecular-level processes occurring in seemingly stable systems, highlighting the dynamic nature.
Le Chatelier's Principle: Temperature Changes
Le Chatelier's Principle is a fundamental concept in chemistry that helps predict how chemical equilibrium systems respond to external changes. One of the key factors that can influence equilibrium is temperature. Understanding how temperature changes affect equilibrium is crucial for various chemical processes in industry and nature.
Temperature and Equilibrium: An Overview
When a chemical system at equilibrium experiences a temperature change, it will shift to counteract that change, following Le Chatelier's Principle. The direction of this shift depends on whether the reaction is exothermic or endothermic.
Exothermic Reactions
An exothermic reaction is one that releases heat to its surroundings. In these reactions:
- Temperature Increase: When the temperature is raised, the equilibrium shifts to favor the reactants. This shift absorbs some of the added heat, counteracting the temperature increase.
- Temperature Decrease: Conversely, when the temperature is lowered, the equilibrium shifts towards the products. This shift releases heat, partially offsetting the temperature decrease.
Example: Consider the exothermic reaction of nitrogen and hydrogen to form ammonia:
N + 3H 2NH + heat
If the temperature is increased, the equilibrium will shift left, producing more N and H to absorb the excess heat.
Endothermic Reactions
An endothermic reaction absorbs heat from its surroundings. For these reactions:
- Temperature Increase: When the temperature rises, the equilibrium shifts to favor the products. This shift uses up some of the added heat, counteracting the temperature increase.
- Temperature Decrease: When the temperature falls, the equilibrium shifts towards the reactants. This shift absorbs heat from the surroundings, partially offsetting the temperature decrease.
Example: The endothermic decomposition of calcium carbonate:
CaCO + heat CaO + CO
If the temperature is increased, the equilibrium will shift right, producing more CaO and CO to absorb the excess heat.
Equilibrium Shift and Its Implications
The equilibrium shift in response to temperature changes has significant implications in various fields:
- Industrial Processes: Understanding these principles allows for optimization of industrial reactions. For instance, in the Haber process for ammonia production, lower temperatures favor the formation of ammonia (the desired product in an exothermic reaction).
- Biological Systems: Temperature changes can affect biochemical equilibria in living organisms, influencing metabolic processes and enzyme activities.
- Environmental Chemistry: Temperature fluctuations in the environment can shift chemical equilibria, affecting processes like the solubility of gases in water or the formation of smog.
Quantitative Aspects
The relationship between temperature and equilibrium constant (K) is described by the van 't Hoff equation:
d(ln K)/dT = ΔH°/(RT²)
Where ΔH° is the standard enthalpy change of the reaction, R is the gas constant, and T is the absolute temperature. This equation quantifies how the equilibrium constant changes with temperature.
Practical Applications
Understanding the effect of temperature on equilibrium is crucial in various applications:
- Chemical Manufacturing: Optimizing reaction conditions to maximize yield and efficiency.
- Environmental
Le Chatelier's Principle: Pressure Changes
Understanding Pressure Effects on Gas Reactions
Le Chatelier's Principle is a fundamental concept in chemistry that helps predict how chemical systems at equilibrium respond to changes in conditions. One of the key factors that can influence equilibrium in gas reactions is pressure. Understanding how pressure changes affect equilibrium is crucial for many industrial processes and natural phenomena.
The Relationship Between Gas Moles and Pressure
Before delving into the effects of pressure on equilibrium, it's important to understand the relationship between the number of gas moles and pressure. In a closed system, pressure is directly proportional to the number of gas moles present. This relationship is described by the ideal gas law: PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the gas constant, and T is temperature.
Pressure Increase and Equilibrium Shift
When a system at equilibrium experiences a pressure increase, Le Chatelier's Principle predicts that the system will shift to counteract this change. In gas reactions, this typically means the equilibrium will shift towards the side with fewer gas moles. This shift reduces the overall pressure by decreasing the total number of gas particles.
For example, consider the Haber process for ammonia synthesis:
N (g) + 3H (g) 2NH (g)
In this reaction, four moles of reactant gases (one mole of N and three moles of H) produce two moles of product gas (NH). If the pressure is increased, the equilibrium will shift towards the product side (right) to reduce the number of gas moles, thus favoring ammonia production.
Pressure Decrease and Equilibrium Response
Conversely, when a system experiences a pressure decrease, the equilibrium shifts to counteract this change by favoring the side with more gas moles. This shift increases the overall pressure by increasing the total number of gas particles.
Using the same Haber process example:
2NH (g) N (g) + 3H (g)
If the pressure is decreased, the equilibrium will shift towards the reactant side (left) to increase the number of gas moles, favoring the decomposition of ammonia into nitrogen and hydrogen.
Exceptions to Pressure Effects
It's important to note that pressure changes only affect equilibria involving gases, and only when there is a difference in the number of gas moles between reactants and products. For reactions where the number of gas moles is equal on both sides, pressure changes have no effect on the equilibrium position.
For instance, consider the reaction:
H (g) + I (g) 2HI (g)
In this case, there are two moles of gas on both sides of the equation. Pressure changes will not shift the equilibrium in either direction because the system cannot reduce or increase the number of gas moles by favoring either side.
Practical Applications of Pressure Effects
Understanding how pressure changes affect equilibrium is crucial in many industrial processes. Engineers and chemists can manipulate pressure to optimize reaction conditions and improve product yields. For example, in the Haber process, high pressures are used to increase ammonia production. Similarly, in the production of methanol from carbon monoxide and hydrogen, high pressures are employed to shift the equilibrium towards the product side.
Conclusion
Le Chatelier's Principle provides a powerful framework for predicting and understanding how chemical systems respond to pressure changes. By considering the number of gas moles on each side of a reaction, we can anticipate how equilibrium will shift in response to pressure increases or decreases. This knowledge is not only crucial for industrial applications but also
Predicting Equilibrium Shifts
Understanding how to predict equilibrium shifts using Le Chatelier's Principle is a crucial skill in chemistry. This principle helps us understand how chemical systems respond to changes in reaction conditions. Let's dive into a step-by-step guide on how to apply this principle effectively.
Step 1: Identify the equilibrium system
First, recognize the chemical reaction at equilibrium. For example, consider the formation of ammonia: N + 3H 2NHStep 2: Determine the change in reaction conditions
Identify what factor is being altered. This could be concentration, pressure, temperature, or the addition of a catalyst.Step 3: Predict the system's response
Apply Le Chatelier's Principle, which states that when a system at equilibrium is disturbed, it will shift to counteract the change.Step 4: Analyze the effect on the forward and reverse reactions
Consider how the change affects both the forward reaction and the reverse reaction. The system will favor the direction that opposes the disturbance.Let's explore some scenarios to better understand this process:
Scenario 1: Concentration Change
If we add more N to our ammonia reaction, the system will shift to reduce this excess. The forward reaction will be favored, producing more NH.Scenario 2: Pressure Change
Increasing pressure in a gaseous equilibrium system will favor the side with fewer gas molecules. In our ammonia example, the forward reaction would be favored as it reduces the number of gas molecules (4 reactant molecules to 2 product molecules).Scenario 3: Temperature Change
Temperature changes affect the equilibrium constant. For exothermic reactions, cooling favors the forward reaction, while for endothermic reactions, heating favors the forward reaction. The ammonia formation is exothermic, so lowering the temperature would shift the equilibrium to the right.Scenario 4: Catalyst Addition
Adding a catalyst doesn't shift the equilibrium position but speeds up both the forward and reverse reactions equally, helping the system reach equilibrium faster.Remember, when predicting equilibrium shifts, it's crucial to consider the specific reaction conditions and how they impact the system. Always think about how the system can counteract the imposed change.
Practice applying these steps to various chemical reactions to improve your understanding. Start with simple systems and gradually move to more complex ones. With time, you'll develop an intuition for predicting equilibrium shifts accurately.
Key takeaways:
1. Le Chatelier's Principle is a powerful tool for predicting equilibrium shifts.
2. Always consider how the system will oppose the change in reaction conditions.
3. Analyze the effect on both the forward reaction and reverse reaction.
4. Different factors (concentration, pressure, temperature) can influence the equilibrium in unique ways.
5. Catalysts affect reaction rates but not the equilibrium position.By mastering these concepts, you'll be well-equipped to tackle equilibrium problems in your chemistry studies and beyond. Keep practicing, and don't hesitate to apply these principles to real-world chemical processes to deepen your understanding.
The Effect of Catalysts on Equilibrium
Catalysts in chemical reactions play a crucial role in chemical reactions, significantly influencing the rate at which reactions occur without altering the final equilibrium position. These substances work by providing an alternative pathway for the reaction, typically with a lower activation energy, allowing reactants to transform into products more rapidly. It's essential to understand that while catalysts accelerate the reaction rate, they do not change the thermodynamics of the reaction or the equilibrium position.
In the context of chemical equilibrium, catalysts in chemical reactions affect both the forward and reverse reactions equally. This means that while a catalyst increases the rate at which equilibrium is reached, it does not shift the equilibrium position in favor of either the reactants or products. The equilibrium position is determined by thermodynamic factors, such as the Gibbs free energy change of the reaction, which remain unaffected by the presence of a catalyst.
To illustrate this concept, let's consider the Haber process for ammonia production. In this reaction, nitrogen and hydrogen gases combine to form ammonia: N + 3H 2NH. Iron catalyst in Haber process is used as a catalyst in this process. The catalyst increases the rate at which nitrogen and hydrogen molecules break apart and recombine to form ammonia. However, it also increases the rate of the reverse reaction at the same proportion. As a result, the equilibrium composition of the reaction mixture remains the same, but the time taken to reach this equilibrium is significantly reduced.
Another example is the decomposition of hydrogen peroxide into water and oxygen: 2HO 2HO + O. Manganese dioxide can be used as a catalyst for this reaction. While MnO dramatically increases the rate of hydrogen peroxide decomposition, it does not change the final amount of water and oxygen produced at equilibrium. The catalyst simply allows the reaction to reach its equilibrium state more quickly.
Understanding the role of catalysts in chemical reactions and their effect on equilibrium is crucial in various fields, from industrial processes to biological systems. Enzymes, for instance, are biological catalysts that enable vital reactions in living organisms to occur at rates compatible with life. In industrial applications, catalysts are employed to optimize production processes, reducing energy requirements and improving efficiency without compromising the yield dictated by the reaction's equilibrium position.
Practical Applications of Le Chatelier's Principle
Le Chatelier's Principle is a fundamental concept in chemistry that has far-reaching implications in both industrial applications and everyday life. This principle, which states that a system in equilibrium will adjust to counteract any change imposed upon it, plays a crucial role in optimizing chemical processes across various industries. Let's explore some fascinating real-world examples of how this principle is applied to enhance efficiency and productivity.
In the industrial realm, Le Chatelier's Principle is a game-changer for reaction optimization. Take the Haber process, for instance, which is used to produce ammonia on a massive scale. By manipulating pressure and temperature based on Le Chatelier's Principle, manufacturers can significantly increase ammonia yield. This process is vital for producing fertilizers that feed billions of people worldwide.
Another industrial application is found in the production of sulfuric acid through the Contact process. Here, Le Chatelier's Principle guides engineers in finding the optimal conditions to maximize the conversion of sulfur dioxide to sulfur trioxide. By carefully controlling temperature and using a catalyst, they can push the equilibrium towards the desired product, enhancing efficiency and reducing waste.
But it's not just heavy industry that benefits from this principle. In everyday life, we encounter Le Chatelier's Principle more often than we might realize. Consider carbonated beverages: when you open a soda can, you're witnessing the principle in action. The sudden decrease in pressure causes the dissolved CO2 to escape, creating those familiar bubbles. Understanding this helps beverage manufacturers maintain the perfect fizz in their products.
Even in our bodies, Le Chatelier's Principle is at work. The process of oxygen binding to hemoglobin in our blood is an equilibrium reaction influenced by factors like pH and CO2 concentration. This delicate balance ensures efficient oxygen delivery to our tissues, adapting to our changing needs during exercise or at high altitudes.
In the realm of environmental science, Le Chatelier's Principle helps us understand and address issues like ocean acidification. As more CO2 dissolves in the oceans, it shifts the carbonate equilibrium, affecting marine ecosystems. Scientists use this knowledge to develop strategies for mitigating these effects.
From industrial chemical processes to the air we breathe, Le Chatelier's Principle is a powerful tool for understanding and manipulating chemical equilibria. Its applications in reaction optimization continue to drive innovation across industries, making our world more efficient and sustainable. By embracing this principle, we unlock the potential to solve complex challenges and improve countless aspects of our daily lives.
Conclusion
Le Chatelier's Principle is a fundamental concept in chemistry that explains how chemical systems at equilibrium respond to changes in conditions. This principle states that when a system at equilibrium is disturbed, it will shift to counteract the change and establish a new equilibrium. Understanding Le Chatelier's Principle is crucial for predicting chemical reactions and controlling chemical reactions in various fields, from industrial processes to environmental science. The introduction video provides a clear and concise explanation of this principle, making it easier for students to grasp its applications. By applying Le Chatelier's Principle, chemists can manipulate reaction conditions to maximize product yield or minimize unwanted side reactions. Students are encouraged to explore this principle further and apply it in their chemistry studies, as it forms the basis for understanding complex chemical systems. To deepen your understanding of equilibrium and reaction control, we invite you to engage with additional resources and practice problems related to predicting chemical reactions and controlling chemical reactions.
Le Chatelier's Principle: What happens at equilibrium?
What happens at equilibrium? Recall dynamic equilibrium.
Step 1: Introduction to Le Chatelier's Principle
Le Chatelier's Principle is an extension of the concept of Dynamic Equilibrium. In the previous lesson, we discussed Dynamic Equilibrium, which is a state where the rate of the forward reaction equals the rate of the reverse reaction. This principle helps us understand how changes in reaction conditions can affect the equilibrium position of a chemical reaction.
Step 2: Understanding Dynamic Equilibrium
Dynamic Equilibrium is characterized by two main points: the dynamic nature and the equilibrium state. The term "dynamic" indicates that the system is in a state of constant change, while "equilibrium" means that the rates of the forward and reverse reactions are equal. This balance creates a stable system where the concentrations of reactants and products remain constant over time.
Step 3: Defining Equilibrium
Equilibrium is defined as the point where the rate of the forward reaction equals the rate of the reverse reaction. For example, consider a reaction where A and 2B react to form C and D. At equilibrium, the rate at which A and 2B form C and D is equal to the rate at which C and D revert to A and 2B. This balance ensures that the system remains stable, with no net change in the concentrations of reactants and products.
Step 4: The Role of Reaction Conditions
Changes in reaction conditions, such as temperature and pressure, can affect the equilibrium position. For instance, increasing the temperature or pressure can shift the equilibrium towards the products or reactants, depending on the nature of the reaction. Understanding how these changes influence the equilibrium position is crucial for controlling chemical reactions and optimizing product yields.
Step 5: Applying Le Chatelier's Principle
Le Chatelier's Principle states that if a system at equilibrium is subjected to a change in conditions, the system will adjust itself to counteract the change and restore a new equilibrium. For example, if the concentration of a reactant is increased, the system will shift the equilibrium towards the products to reduce the effect of the added reactant. Similarly, if the temperature is increased, the system will shift the equilibrium in the direction that absorbs heat, thereby counteracting the temperature change.
Step 6: Practical Applications
Le Chatelier's Principle is widely used in industrial processes to optimize the production of desired products. By manipulating reaction conditions, chemists can shift the equilibrium towards the products, increasing the yield and efficiency of chemical reactions. This principle is also essential in understanding natural processes and designing chemical experiments.
Step 7: Conclusion
In summary, Le Chatelier's Principle provides a framework for understanding how changes in reaction conditions affect the equilibrium position of a chemical reaction. By recalling the concept of Dynamic Equilibrium and applying this principle, we can predict and control the behavior of chemical systems, optimizing product yields and enhancing our understanding of chemical processes.
FAQs
Here are some frequently asked questions about Le Chatelier's Principle:
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What is a forward reaction in chemical equilibrium?
A forward reaction in chemical equilibrium is the process where reactants combine to form products. For example, in the reaction A + B C + D, the forward reaction is A + B C + D. At equilibrium, the forward reaction rate equals the reverse reaction rate.
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What happens to an exothermic reaction at equilibrium when heated?
When an exothermic reaction at equilibrium is heated, Le Chatelier's Principle predicts that the equilibrium will shift towards the reactants (left) to counteract the temperature increase. This shift absorbs some of the added heat, as the reverse reaction is endothermic.
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How does equilibrium shift when temperature is increased for an endothermic reaction?
For an endothermic reaction, increasing the temperature shifts the equilibrium towards the products (right). This shift favors the forward reaction, which absorbs heat, thus counteracting the temperature increase in accordance with Le Chatelier's Principle.
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What factors can favor a reverse reaction?
Factors that can favor a reverse reaction include: decreasing the concentration of products, increasing the concentration of reactants, changing temperature (cooling for exothermic reactions, heating for endothermic reactions), and changing pressure (for gas reactions with more moles on the reactant side).
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How can you determine if a reaction is exothermic or endothermic at equilibrium?
To determine if a reaction is exothermic or endothermic at equilibrium, observe how the equilibrium shifts with temperature changes. If increasing temperature favors the reverse reaction (shifts left), the forward reaction is exothermic. If increasing temperature favors the forward reaction (shifts right), the forward reaction is endothermic.
Prerequisite Topics for Understanding Le Chatelier's Principle
To fully grasp Le Chatelier's principle, it's crucial to have a solid foundation in several key chemistry concepts. One of the most fundamental prerequisites is understanding dynamic equilibrium. This concept forms the basis of Le Chatelier's principle, as it describes the state in which forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products.
Another essential prerequisite is knowledge of the factors affecting reaction rate. Le Chatelier's principle deals with how chemical systems respond to changes in conditions, and understanding these factors is crucial for predicting and explaining these responses. This ties in closely with the concept of activation energy, which plays a significant role in determining how easily a reaction can proceed and how it might be affected by changes in temperature or the addition of a catalyst.
A practical application of Le Chatelier's principle can be seen in the Haber process for ammonia production. This industrial process relies heavily on manipulating equilibrium conditions to maximize yield, making it an excellent real-world example of Le Chatelier's principle in action.
Having a solid introduction to chemical reactions is also vital. This includes understanding how catalysts work in chemical reactions, as catalysts can significantly impact the rate at which equilibrium is established without affecting the equilibrium position itself.
Lastly, knowledge of predicting chemical reactions using cell potential can provide valuable insights into the direction of spontaneous change in chemical systems, which is closely related to the principles outlined by Le Chatelier.
By mastering these prerequisite topics, students will be well-equipped to understand and apply Le Chatelier's principle. This principle is not just an isolated concept but a powerful tool that integrates various aspects of chemical equilibrium and kinetics. It allows chemists to predict how chemical systems will respond to changes in concentration, pressure, temperature, and other factors.
Understanding Le Chatelier's principle is crucial for many areas of chemistry, from industrial processes to environmental science. It helps explain phenomena such as why carbonated drinks fizz more when warm, how blood maintains its pH, and how to optimize industrial reactions for maximum efficiency. By building a strong foundation in these prerequisite topics, students will be better prepared to tackle more complex chemical systems and real-world applications of equilibrium principles.
In this lesson, we will learn:
- To recall Le Chatelier's principle when studying equilibria.
- How to predict changes to equilibria given changes in reaction conditions.
- How to explain changes in equilibrium position due to changes in reaction conditions.
Notes:
- Continuing from Dynamic Equilibrium, when more chemical reactions were studied, and more equilibria were found, chemists started changing the conditions (such as temperature and pressure) of the reaction to see what this did to the equilibrium they saw.
This led to them finding Le Chatelier's principle:- When a reaction at equilibrium is disturbed by a change in conditions, the system will respond in a way that COUNTERACTS the disturbance and then re-establish a new equilibrium.
- The system 'counteracts the disturbance' by favoring the forward or the backward reaction; it will do the opposite of whatever effect the disturbance had on the reaction! 'Favoring' a reaction means that reaction rate increases (causing a change in the ratio of reactants to products in the system), before a new equilibrium is reached as the two rates become equal again. Once this happens, the new ratio of reactant to product concentration won’t change unless equilibrium is disturbed again - in this case, the above happens again.
- Remember, 'at equilibrium' only tells us that the forward and reverse reactions are happening at the same rate. Equilibrium does not tell us how much product and reactant there is in the system!
- When a reaction at equilibrium is disturbed, it will work to re-establish equilibrium under these new conditions. If some facts about the reaction are known, the change to equilibrium position can be predicted from a change in conditions:
- If temperature is increased:
- In an exothermic reaction: The system will respond to favor the reverse reaction which leads to a greater concentration of reactants being observed in the reaction. An exothermic forward reaction, by definition, causes a net release of heat energy to the surroundings. The system therefore counteracts the disturbance of the temperature increase by temporarily favoring the reverse reaction, which will be the endothermic opposite of the exothermic forward reaction. This endothermic reverse reaction has a net absorbing of heat effect, absorbing the increase in temperature that initially disturbed the system. This is how Le Chatelier's principle works! When the reverse reaction is favored, we say the equilibrium has shifted to the left as the left-hand side of the reaction equation shows the reactants.
- In an endothermic reaction: The system will respond by favoring the forward reaction. This is for the same reasons as above; the system will favor the endothermic forward reaction, which absorbs heat and counteracts the increase in temperature that it was disturbed by. Favoring the forward reaction means the equilibrium shifts to the right. In this case, more products will be produced in the system.
- If temperature is decreased:
- In an exothermic reaction: The system will respond to favor the forward reaction. This is because the exothermic forward reaction has a net releasing of heat effect, which will raise the temperature and counteract the decrease in temperature that initially disturbed it. In this case we say the equilibrium has shifted to the right and more products will be produced.
- In an endothermic reaction: The system will respond to favor the reverse reaction. This is because the reverse reaction will be exothermic, causing a net release of heat energy to the system, raising the temperature that initially disturbed it. The equilibrium will shift to the left and more reactants will be present in the system.
- If pressure is increased, count the number of moles of gas on both sides of the reaction equation:
- When there are more moles of gas in the products than the reactants, the system will respond to favor the reverse reaction and the equilibrium will shift left. This has the effect of 'cutting back' the number of gas particles in the system as there are more of them in the products than the reactants. This leads to less gas particle collisions which decreases pressure, which counteracts the initial disturbance of a rise in pressure!
- When there are less moles of gas in the products than the reactants, the system will respond to favor the forward reaction and the equilibrium will shift right. This counteracts the disturbance in the same way as explained above.
- If pressure is decreased, count the number of moles of gas on both sides of the reaction equation:
- When there are more moles of gas in the products than the reactants, the system will respond to favor the forward reaction and the equilibrium will shift right This has the effect of increasing the number of gas particles in the system as there are more gas moles in the products than the reactants. This leads to more gas particle collisions which increases pressure, counteracting the initial disturbance of a pressure decrease!
- When there are less moles of gas in the products than the reactants, the system will favor the reverse reaction and the equilibrium will shift left. This leads to an increase in the number of gas particles in the system because there are more gas moles in the reactants than the products. This counteracts the initial disturbance of the pressure decrease.
- If a catalyst is added to the reaction mixture, there is no change to the position of equilibrium. A catalyst does not change equilibrium; it simply allows equilibrium to be reached quicker.
- The table below summarizes the changes to equilibrium position caused by a change in conditions:
- Even though chemists know how to change a reaction at equilibrium to make the largest amount of product, there are practical issues with shifting the equilibrium.
- Increasing the pressure on a closed system can be very expensive and there are safety considerations at very high pressure.
- Many reactions, such as the Haber process, have an exothermic forward reaction. This strangely means that to shift the equilibrium to make more product, you need to cool the reaction down which will reduce reaction rate.
