Question

Consider the exothermic reaction $$ \mathrm{CO}(g)+2 \mathrm{H}_{2}(g) \rightleftharpoons \mathrm{CH}_{3} \mathrm{OH}(l) $$ Predict three changes that could be made to the system that would decrease the yield of product over that produced by a system in which no change was made.

Short answer

To decrease the yield of CH3OH in the exothermic reaction \( CO(g) + 2 H_{2}(g) \rightleftharpoons CH_{3}OH(l) \), we can make the following changes based on Le Chatelier's principle: 1. Increase the temperature of the system, which will shift the reaction towards the reactants (left) and decrease the CH3OH yield. 2. Decrease the pressure of the system, which will also shift the equilibrium towards the reactants (left) and decrease CH3OH yield. 3. Remove one of the reactants (either CO(g) or H2(g)), causing the reaction to shift towards the reactants (left) and decreasing the yield of CH3OH.

Step-by-step solution

1. Increase in Temperature

Since the reaction is exothermic, it releases heat. According to Le Chatelier's principle, if we increase the temperature of the system, the reaction will try to counteract the change by shifting in the direction that consumes heat, which is the reverse (left) direction. So, if we increase the temperature of the system, the reaction will shift to the left, and the yield of CH3OH will decrease.

2. Decrease in Pressure

From the given reaction, we can notice that there are three moles of reactants (gas phase) and one mole of product (liquid phase). The decrease in pressure will cause the reaction to shift towards the direction with more gaseous moles, which is the reactant side. So, reducing the pressure on the system will shift the equilibrium to the left, and the yield of CH3OH will decrease.

3. Removal of a Reactant

If we remove one of the reactants, CO(g) or H2(g), the reaction will try to counteract the change by shifting in the direction that produces more of the removed reactant. In this case, the reaction will shift to the left (towards reactants), leading to a decrease in the yield of the product, CH3OH.

Key concepts

Le Chatelier's Principle

Le Chatelier's principle is a foundational concept in chemistry that describes how a chemical system at equilibrium responds to changes in concentration, temperature, or pressure. It essentially tells us that if an external change is applied to a system at equilibrium, the system will adjust itself to minimize that change.

Imagine a seesaw perfectly balanced with weights on both ends. If you add weight to one side, the seesaw will tip in that direction. Le Chatelier's principle suggests that the system (the seesaw in this case) will try to restore balance by reacting in such a way as to reduce the impact of the added weight.

In chemical terms, if a reactant or product is added to a system, the equilibrium will shift to oppose the addition. If a reactant is removed, the equilibrium will shift to replace it, and so on. The principle is crucial for understanding the behavior of chemical reactions under different conditions and for optimizing industrial chemical processes.

Exothermic Reactions

Exothermic reactions are chemical reactions that release heat to their surroundings. They can be thought of as cozy fireplaces, where the burning wood represents the reactants transforming into ashes, the products, and releasing energy in the form of heat. The heat is essentially a product of these reactions.

Since heat is released, the temperature of the surroundings increases, and if we consider the fire in a closed room (system), raising the room's temperature (by, for instance, turning up the thermostat) would be similar to adding more 'product' into the system. According to Le Chatelier's principle, the reaction will shift to absorb this extra heat, favoring the reactants. Thus, in exothermic reactions, increasing the temperature pushes the equilibrium towards the reactants and decreases the yield of the products.

Equilibrium Shift

An equilibrium shift occurs when the balance between the reactants and products in a chemical reaction is disturbed. This shift is the reaction's response to an external force, such as changes in temperature, pressure, or concentration, striving to establish a new balance or equilibrium.

The idea of a shift can be visualized by thinking of a tug-of-war contest. When one team pulls harder (change in conditions), the rope (equilibrium) moves towards that side. If conditions favor the formation of more products, we say the equilibrium shifts 'to the right', increasing the yield of the products. Conversely, if the conditions favor the reactants, the equilibrium shifts 'to the left', reducing the yield of products.

To provide a practical illustration, if we decrease the pressure in a reaction involving gases, the system will favor the direction that produces more gas molecules to counter the pressure change. This tug-of-war between the competing influences of reactants and products under different conditions is an essential concept for predicting the outcomes of chemical reactions.