Question

The conditions for producing ammonia industrially are \(500^{\circ} \mathrm{C}\) and 300 atm. What happens to the ammonia concentration if (a) the temperature increases and (b) the pressure increases? $$ \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftarrows 2 \mathrm{NH}_{3}(g)+\text { heat } $$

Short answer

Increasing the temperature decreases ammonia concentration, while increasing the pressure increases it.

Step-by-step solution

Understanding the Reaction Conditions

The balanced chemical equation for the production of ammonia is given as \( \mathrm{N}_{2}(g) + 3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g) + \text{heat} \). It is an exothermic reaction, meaning it releases heat as it proceeds to form more ammonia. The industrial conditions for this reaction are set at \(500^{\circ} \mathrm{C}\) and 300 atm to balance the efficiency and yield of ammonia. We will examine how changes in these conditions affect the ammonia concentration.

Evaluating the Effect of Increased Temperature

When the temperature of an exothermic reaction is increased, the equilibrium will shift to the left, according to Le Chatelier's Principle, as it attempts to absorb the extra heat. Thus, increasing the temperature will decrease the concentration of ammonia \(\mathrm{NH}_3\) as the equilibrium shifts backward to favor the formation of \(\mathrm{N}_2\) and \(\mathrm{H}_2\).

Evaluating the Effect of Increased Pressure

An increase in pressure shifts the equilibrium towards the side with fewer moles of gas. In this reaction, the forward reaction results in the formation of 2 moles of \(\mathrm{NH}_3\) from 4 moles of the reactants \(\mathrm{N}_2\) and \(\mathrm{H}_2\). Thus, increasing the pressure will shift the equilibrium to the right, increasing the concentration of ammonia, as it reduces the total number of gas moles.

Key concepts

Exothermic Reactions

In exothermic reactions, like the formation of ammonia, heat is released as the reaction progresses. This type of reaction can be easily understood by thinking about heat as a product alongside the chemical substances being produced. When heat is a product, it means that the system gives off energy to the surroundings. This can be observed in many familiar processes, like combustion or the setting of cement.

An important aspect of exothermic reactions is how they respond to changes in temperature. Applying Le Chatelier’s Principle, which helps predict the behavior of a system in equilibrium when subjected to changes, we know that increasing the temperature adds energy to the system. For an exothermic reaction, the system will try to absorb this added energy by shifting the equilibrium to the left, favoring the reactants instead of the products.

• This means for the ammonia formation reaction, a rise in temperature results in a decrease in ammonia production.
• This is why industrial processes must carefully manage temperature to optimize yields.

The Haber process, which produces ammonia, is designed to balance temperature to maximize ammonia without adding excess heat that would otherwise reverse the reaction.

Chemical Equilibrium

Chemical equilibrium occurs in a reversible reaction when the rates of the forward and backward reactions are the same. At this point, the concentrations of reactants and products remain constant over time, although both reactions continue to occur.

In the synthesis of ammonia, represented by the equation \[ \mathrm{N}_{2}(g) + 3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g) + \text{heat}\]the equilibrium can be shifted by changing the conditions such as temperature and pressure.

• Temperature adjustments affect which direction the equilibrium shifts due to the endothermic or exothermic nature of the reaction.
• Increasing pressure moves the equilibrium towards the side with fewer gas molecules. In this example, the ammonia side has fewer moles of gas (2 moles of \(\mathrm{NH}_{3}\)) compared to the reactants (4 moles total of \(\mathrm{N}_{2}\) and \(\mathrm{H}_{2}\)).

Understanding equilibrium is essential for optimizing industrial chemical processes, where shifting the equilibrium toward desired products can significantly enhance efficiency and yield.

Industrial Chemistry

In industrial chemistry, reactions are optimized to produce desired chemicals efficiently and on a large scale. The production of ammonia is a key example of how industrial chemistry applies scientific principles to achieve practical goals.

The Haber process, used for ammonia production, is a well-studied industrial reaction that operates under specific conditions of high pressure and moderate temperatures to achieve a balance between the rate and yield of ammonia production.

• High pressure favors the forward reaction towards ammonia because it reduces the number of moles of gas.
• The temperature is kept at moderate levels to ensure a steady production rate without driving the reaction in the favor of reactants.

Managing these conditions carefully is crucial for economic and efficient chemical manufacturing. Utilizing Le Chatelier’s Principle and a deep understanding of chemical equilibria allows industries to maximize the output and quality of the chemical products, making industrial-scale production viable and profitable.