Question

State and explain Le Chatelier's principle. Apply the principle to the following equilibria. (i) \(2 \mathrm{SO}_{2}+\mathrm{O}_{2} \rightleftarrows 2 \mathrm{SO}_{3}+\) Heat (ii) \(2 \mathrm{NH}_{3} \rightleftarrows \mathrm{N}_{2}+3 \mathrm{H}_{2}-\) Heat

Short answer

\(2 \mathrm{SO}_{2}+\mathrm{O}_{2} \rightleftarrows 2 \mathrm{SO}_{3}+ \text{Heat}\) Answer: When the temperature is increased in the given equilibrium, the concentration of SO3 decreases. This is because the reaction is exothermic and the system shifts to the left, favoring the reverse (endothermic) reaction to counteract the increased temperature.

Step-by-step solution

State Le Chatelier's principle

Le Chatelier's principle states that when a chemical system at equilibrium is disturbed by a change in temperature, pressure, or concentration of reactants or products, the system will shift its equilibrium position to counteract the change. In other words, the system will respond in such a way as to minimize the external stress imposed on it.

Explain Le Chatelier's principle

When a chemical reaction reaches equilibrium, the rates of the forward and reverse reactions are equal. If an external stress is applied to the system (e.g., by changing the temperature, pressure, or concentration of a reactant or product), the system will adjust its equilibrium position in such a way that it counteracts this change, by either favoring the forward or reverse reaction until a new equilibrium is established.

Apply Le Chatelier's principle to the first equilibrium

For the given equilibrium: \(2 \mathrm{SO}_{2}+\mathrm{O}_{2} \rightleftarrows 2 \mathrm{SO}_{3}+ \text{Heat}\) This reaction is exothermic since heat is released as the products are formed. If we increase the temperature (i.e., add heat), the equilibrium will shift to the left, favoring the reverse (endothermic) reaction, to reduce the temperature, resulting in a decrease in the concentration of \(\mathrm{SO}_{3}\). Conversely, if we decrease the temperature (i.e., remove heat), the equilibrium will shift to the right, favoring the forward (exothermic) reaction, to increase the temperature. This will result in an increase in the concentration of \(\mathrm{SO}_{3}\).

Apply Le Chatelier's principle to the second equilibrium

For the given equilibrium: \(2 \mathrm{NH}_{3} \rightleftarrows \mathrm{N}_{2}+3 \mathrm{H}_{2}- \text{Heat}\) This reaction is endothermic since heat is absorbed as the products are formed. If we increase the temperature (i.e., add heat), the equilibrium will shift to the right, favoring the forward (endothermic) reaction, to reduce the temperature, resulting in an increase in the concentration of \(\mathrm{N}_2\) and \(\mathrm{H}_2\). Conversely, if we decrease the temperature (i.e., remove heat), the equilibrium will shift to the left, favoring the reverse (exothermic) reaction, to increase the temperature. This will result in a decrease in the concentration of \(\mathrm{N}_2\) and \(\mathrm{H}_2\).

Key concepts

Chemical Equilibrium

Chemical equilibrium is a state in a chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction. This balance means that the concentrations of reactants and products remain constant over time, not that they are equal to each other. As an analogy, think of it like two teams playing tug of war with equal strength; the rope doesn't move, but both teams are still putting in effort.
It's critical to understand that equilibrium does not mean that the reactions have stopped. Both forward and reverse reactions continue to occur, but because they are happening at the same rate, there is no net change in the concentrations of the substances involved.

Exothermic Reaction

Imagine you're warming your hands by a campfire. The heat you feel is similar to what's released during an exothermic reaction. These reactions give off energy, usually in the form of heat, to the surroundings. In a chemical equation, we often see heat as a product, like in the combustion of propane in a grill.

Practical Uses

• Energy production: From power plants to our body's cells releasing energy from food.
• Everyday products: Hand warmers, self-heating cans for food, and many other consumer goods rely on exothermic reactions.

Exothermic reactions are generally more stable than their reactants because energy is released when bonds are formed in the products, making the final substances lower in energy.

Endothermic Reaction

When you mix baking soda and vinegar to inflate a balloon, you're engaging in an endothermic reaction, where the system absorbs energy from its surroundings, often as heat, causing the reaction mixture to feel cold. In these reactions, energy is a reactant, and the overall process requires an input of energy to proceed.

Everyday Examples

• Photosynthesis: Plants absorb sunlight to convert carbon dioxide and water into sugars.
• Cooking: Baking bread, where yeast causes dough to rise by absorbing energy.

The endothermic processes often involve breaking strong bonds, which requires more energy than is released when new bonds form in the products.

Equilibrium Shift

When you adjust the conditions of a chemical reaction at equilibrium, it's like pushing a balanced seesaw to one side. The reaction will respond to restore balance or reduce the effect of the change, a process known as the equilibrium shift. Changes in concentration, temperature, or pressure can all cause this shift to occur.
For example, in the synthesis of ammonia (\(2 \text{NH}_3\rightleftarrows \text{N}_2 + 3 \text{H}_2\text{ -Heat}\)), increasing the pressure would favor the shift towards the side with fewer gas molecules. As the principle of Le Chatelier suggests, the system will try to lessen the pressure increase by forming more ammonia, which has fewer gas molecules than the nitrogen and hydrogen reactants. Understanding this concept is crucial in industrial processes and synthesizing chemicals because it allows us to control the yield of our desired products.