Question

What effect do the listed changes have on the position of the equilibrium in the reaction of carbon with hydrogen? $$ \mathrm{C}(s)+2 \mathrm{H}_{2}(g) \rightleftarrows \mathrm{CH}_{4}(g) \quad \Delta H=-75 \mathrm{~kJ} / \mathrm{mol} $$ (a) Increasing temperature (b) Increasing pressure by decreasing volume (c) Allowing \(\mathrm{CH}_{4}\) to escape continuously from the reaction vessel

Short answer

(a) Shifts left, (b) Shifts right, (c) Shifts right.

Step-by-step solution

Identify the Reaction Type

The reaction given is \( \mathrm{C}(s) + 2 \mathrm{H}_2(g) \rightleftarrows \mathrm{CH}_4(g) \) with a \( \Delta H = -75 \mathrm{~kJ/mol} \). Since \( \Delta H \) is negative, the reaction is exothermic, releasing heat as it proceeds forward.

Analyze Temperature Increase

For exothermic reactions, increasing the temperature adds heat to the system. According to Le Chatelier's principle, the system will try to absorb this added heat by favoring the endothermic process (reverse reaction). Therefore, increasing the temperature shifts the equilibrium to the left, towards the reactants.

Analyze Pressure Increase by Decreasing Volume

For reactions involving gases, increasing pressure by reducing volume shifts the equilibrium toward the side with fewer gas molecules. The reaction has 2 moles of \( \mathrm{H}_2(g) \) on the reactant side and 1 mole of \( \mathrm{CH}_4(g) \) on the product side. Therefore, increasing the pressure will shift the equilibrium to the right, towards the products.

Analyze Removal of \(\mathrm{CH}_4\)

If \( \mathrm{CH}_4 \) is allowed to escape, the concentration of \( \mathrm{CH}_4 \) will decrease. According to Le Chatelier's principle, the system will shift to the right to produce more \( \mathrm{CH}_4 \) and replace what has been removed.

Key concepts

Le Chatelier's principle

Le Chatelier's principle is a fundamental concept in chemistry that explains how a system at equilibrium responds to external changes. Imagine a system at equilibrium as a balance scale; when one side is disturbed, the scale shifts to restore balance.

According to this principle, if a change is applied to a system at equilibrium, the system adjusts to counteract the change and restore a new equilibrium state. For instance, if you add more of a reactant, the balance tilts, and the system will shift to produce more product.

The principle applies to concentration changes, temperature variations, and pressure adjustments. Understanding how these factors influence equilibrium positions can help predict how systems behave under different conditions."}]},{

Exothermic reactions

In chemical terms, an exothermic reaction releases energy, usually in the form of heat, to the surroundings. This type of reaction is characterized by a negative change in enthalpy (\(\Delta H\)).

In general, you can think of exothermic reactions as a process where heat is a product of the reaction. So, when you write the equation for an exothermic reaction, you can imagine heat being released as the reaction moves forward.

Since heat is released, increasing the temperature can affect these reactions significantly. According to Le Chatelier's principle, adding heat to an exothermic reaction shifts the equilibrium towards the side that absorbs the heat, which often means reversing the reaction. When analyzing changes in conditions, keep this concept in mind to predict how such reactions will behave.

Reaction shifts

A reaction shift refers to the change in the position of equilibrium in a chemical reaction. Understanding how and why these shifts occur can help determine the direction in which a reaction will proceed.

When conditions such as temperature, pressure, or concentration change, the equilibrium shifts to minimize the effect of the change. This shift can be to the right, toward the products, or to the left, back to the reactants.

If the concentration of a product is decreased, perhaps by removing it, the reaction will shift towards the side that compensates for this loss. Conversely, adding more reactants pushes the reaction towards producing more products. Each shift is a direct response to the disturbance introduced to the system, aligning with Le Chatelier's principle.

Pressure effects on equilibrium

Pressure changes can extensively influence the position of equilibrium, especially in reactions involving gases. According to Le Chatelier's principle, changes in pressure will affect the equilibrium according to the number of gas molecules present on each side of the reaction.

Consider a scenario where the volume of the reaction vessel is decreased, leading to increased pressure. The equilibrium will shift towards the side that has fewer gas molecules. This shift helps reduce the pressure in the system by minimizing the number of gas particles.

For example, in a reaction involving two moles of hydrogen gas on the reactant side and one mole of methane on the product side, a pressure increase will push the reaction towards the product side. This helps restore equilibrium by decreasing the overall number of gas molecules in the system, thus reducing the pressure.