can i please get some help with this?
For the equilibrium reaction involving the formation of ammonia, which can be described by the reaction below, describe the effect of adding H2 gas to the reaction, increasing the pressure in the reaction, and decreasing the temperature of the reaction.
N2(g) + 3H2(g) <-> 2NH3(g) ΔH = –92.6 kJ
i don't know where to start nor do i get the point of the given delta h
N2(g) + 3H2(g) <-> 2NH3(g) ΔH = –92.6 kJ
Let me rewrite the equation. Since dH is negative that means the reaction is exothermic so we can write N2(g) + 3H2(g) <-> 2NH3(g) + heat
The POINT: This is a question about Le Chatelier's principle. That principle, in what I call "up tight" language is this. When a chemical system in equilibrium is subjected to a stress the reaction will shift in such a direction so as to relieve the stress. Simple. Straight forward. However, few students understand it. Here it is in simpler terms. A chemical system in equilibrium will try to undo what we do to it.
N2(g) + 3H2(g) <-> 2NH3(g) + heat
So we add H2 to the system. Now it has too much. It will try to get rid of the extra H2 that was added. How can it do that? By shifting to the right the reaction will use up the added H2. N2 will decrease. NH3 will increase, It will get hotter.
Add pressure: Remember that increasing pressure of gas makes it want to occupy a smaller volume. You have 4 mols gas on the left side and 2 mols gas on the right side. Which is the smaller volume. The right, of course. So the reaction will move (shift) to the right. N2 will get smaller, H2 will get smaller, NH3 will get larger, it will produce more heat. By now you have the heat thing figured out. If you increase temperature (increase heat) it will shift to the left because by doing that it will use up the heat that was added.
I hope this helps. In my opinion this is the simplest "law" in chemistry but it gives students the most trouble.I love talking about Le Chatelier's principle.
ah thank you so much for the detailed explanation! the delta h was really throwing me off. when explained well and in simpler terms, le chatelier's principle seems to be rather intuitive.
Certainly! I'd be happy to help you understand the effects of adding H2 gas, increasing the pressure, and decreasing the temperature on the equilibrium reaction involving the formation of ammonia.
Let's go through each scenario step-by-step:
1. Effect of Adding H2 Gas to the Reaction:
By adding H2 gas (reactant) to the reaction, you are increasing the concentration of a reactant. According to Le Chatelier's Principle, when the concentration of a reactant is increased, the equilibrium position will shift in the direction that reduces the concentration of the reactant. In this case, the equilibrium will shift to the right to produce more NH3 gas.
2. Effect of Increasing the Pressure in the Reaction:
Increasing the pressure of a system affects the equilibrium position. In the given reaction, the reaction includes four moles of gas on the left side (N2 and 3H2), and two moles of gas on the right side (2NH3). According to Le Chatelier's Principle, if the pressure is increased, the system will shift in the direction that reduces the total number of moles of gas. In this case, the equilibrium will shift to the right to produce more NH3 gas since it has fewer moles of gas.
3. Effect of Decreasing the Temperature of the Reaction:
Lowering the temperature of the reaction affects the equilibrium position. In an exothermic reaction like the one given (ΔH = -92.6 kJ), the forward reaction (formation of ammonia) is favored by higher temperatures. Therefore, if the temperature is decreased, the equilibrium will shift to the left to produce more reactants (N2 and H2) and less product (NH3).
Overall, the given ΔH represents the enthalpy change of the reaction. In this case, the negative sign indicates that the reaction is exothermic, releasing energy (heat). Understanding the enthalpy change can help determine the direction in which the equilibrium will shift when changes are made to the reaction conditions.
I hope this helps! Let me know if you have any further questions.
Of course! I'd be happy to help you with that. Let's start by understanding the given ΔH value in the equilibrium reaction.
The ΔH value, which stands for the change in enthalpy, indicates the amount of heat released or absorbed during a chemical reaction. In this case, since the ΔH value is negative (-92.6 kJ), it means that the forward reaction (formation of ammonia) is exothermic, meaning that it releases heat.
Now, let's discuss the effect of adding H2 gas, increasing the pressure, and decreasing the temperature on the equilibrium.
1. Effect of adding H2 gas:
When you add more H2 gas to the reaction, it increases the concentration of the reactant. According to Le Chatelier's Principle, in response to this increase, the equilibrium will shift to reduce the excess H2 gas. In this case, it will favor the forward reaction to consume the additional H2 gas and form more NH3. As a result, the amount of NH3 at equilibrium will increase.
2. Effect of increasing the pressure:
Increasing the pressure of a reaction system affects the equilibrium by favoring the side with fewer gas molecules. In the given reaction, there are four moles of gas on the left side (N2 + 3H2) and two moles of gas on the right side (2NH3). According to Le Chatelier's Principle, when pressure increases, the equilibrium will shift in the direction that reduces the total number of gas molecules. In this case, it will shift to the side with fewer gas molecules, which is the right side (2NH3). Thus, the equilibrium will favor the formation of more NH3.
3. Effect of decreasing the temperature:
Lowering the temperature of an exothermic reaction favors the production of products according to Le Chatelier's Principle. In this case, since the forward reaction (formation of NH3) is exothermic and releases heat, decreasing the temperature will shift the equilibrium to the right, resulting in the formation of more NH3.
To summarize:
- Adding H2 gas will increase the amount of NH3 at equilibrium.
- Increasing the pressure will favor the formation of NH3.
- Decreasing the temperature will also favor the formation of NH3.
These factors can be explained using Le Chatelier's Principle and the knowledge of how the equilibrium position can shift in response to changes in concentration, pressure, and temperature.