Question

Consider the following exothermic reaction at equilibrium: $$ \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g) $$ Predict how the following changes affect the number of moles of each component of the system after equilibrium is reestablished by completing the table below. Complete the table with the terms increase, decrease, or no change.

Short answer

After analyzing the impact of each change on the given exothermic reaction at equilibrium using Le Chatelier's Principle, the completed table can be summarized as follows: Change: | N2 | H2 | NH3 --------------------------------------------------------- Increase in N2 concentration| No change| Decrease | Increase Increase in H2 concentration| No change| Decrease | Increase Increase in NH3 concentration| Increase | Increase | Decrease Increase in temperature | Increase | Increase | Decrease

Step-by-step solution

Change 1: Increase in N2 concentration

When N2 concentration is increased, the system tries to consume the excess N2 to maintain the equilibrium. Consequently, the reaction will proceed to the right (towards the products). This shift would lead to an increase in the moles of NH3, a decrease in the moles of H2, and no change in the moles of N2.

Change 2: Increase in H2 concentration

An increase in H2 concentration causes the system to adjust by consuming the excess H2 to regain equilibrium. The reaction proceeds to the right (towards the products), which leads to an increase in moles of NH3, no change in moles of N2, and a decrease in moles of H2.

Change 3: Increase in NH3 concentration

When NH3 concentration is increased, the system tries to reduce the excess NH3 to restore the equilibrium. Thus, the reaction will shift to the left (towards the reactants). This leads to a decrease in moles of NH3, an increase in moles of N2, and an increase in moles of H2.

Change 4: Increase in temperature

An increase in temperature for an exothermic reaction leads to the system trying to absorb the excess heat by shifting the reaction in the endothermic direction. In this case, the reaction will shift to the left (towards the reactants). The increase in temperature will lead to a decrease in the moles of NH3, an increase in moles of N2, and an increase in moles of H2. Now, we have analyzed the impact of each change on the equilibrium, and we can complete the table as follows: Change: | N2 | H2 | NH3 --------------------------------------------------------- Increase in N2 concentration| No change| Decrease | Increase Increase in H2 concentration| No change| Decrease | Increase Increase in NH3 concentration| Increase | Increase | Decrease Increase in temperature | Increase | Increase | Decrease

Key concepts

Le Chatelier's Principle

Understanding how a system at equilibrium responds to changes is crucial for grasping chemical reactions. This is where Le Chatelier's principle comes in.Le Chatelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. It’s like a seesaw of chemical reactions: if you add more weight to one side, the seesaw will tip that way until it finds a new balance. In a chemical context, these 'weights' can be changes in concentration of reactants or products, pressure, volume, or temperature.For example, if we increase the concentration of a reactant, the system will adjust to consume some of that excess reactant and produce more products. Conversely, if we remove some product, the system will try to replace it by converting more reactants into products. Le Chatelier's principle is a valuable tool in predicting how a change will affect the concentrations of reactants and products in a chemical reaction.

Exothermic Reaction

An exothermic reaction is a type of chemical reaction that releases energy, usually in the form of heat, to its surroundings, resulting in a net release of energy. This is opposed to endothermic reactions, which absorb energy from their surroundings.

Understanding Heat as a Product

In exothermic reactions, you can think of heat as one of the 'products' of the reaction. So, when we consider Le Chatelier’s principle, an increase in temperature is akin to increasing a product of the reaction. The system tries to consume some of this 'excess' product (heat), by shifting the equilibrium to favor the reactants, which absorb heat. Conversely, decreasing temperature would shift the balance towards the products, as the system tries to produce heat to compensate for the loss.Most day-to-day combustion processes, like a candle burning or a campfire, are examples of exothermic reactions. These reactions are also at play in hand warmers and self-heating food packaging, which rely on exothermic processes to provide heat.

Equilibrium Shifts

Equilibrium shifts are responses to the changes in a system that disturb its state of balance. They occur in accordance with Le Chatelier's principle to counteract the disturbance and re-establish equilibrium.

How Equilibrium Shifts Relate to Reactants and Products

The effect of increasing the concentration of a reactant or product is essentially pushing the equilibrium position either to the right (toward more products) or to the left (toward more reactants). For instance, in the reaction between nitrogen and hydrogen to form ammonia, increasing the concentration of nitrogen causes the system to shift right, making more ammonia to reduce the excess nitrogen.

• If a reactant is increased, the system shifts to produce more of the product(s).
• If a product is increased, the system shifts to consume the product(s) and produce more of the reactant(s).
• If the temperature is increased in an exothermic reaction, the system shifts to favor the reactants, decreasing product concentration.
• If pressure is increased (by decreasing volume), the system shifts toward the side with fewer gas molecules, if gases are involved.

Breaking down shifts into these bite-sized scenarios makes predicting the outcomes of changes to an equilibrium much more straightforward. This knowledge is not only foundational for chemistry students but also immensely applicable in industrial processes where managing reaction conditions is essential for optimal productivity.