Problem 10 - Chemistry in 30 days

Question

In the manufacture of ammonia by haber's process \(\mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{3}(\mathrm{~g}) \neq 2 \mathrm{NH}_{3}(\mathrm{~g})+92.3 \mathrm{~kJ}\) Which of the following conditions is favourable? (a) increasing the amount of \(\mathrm{NH}_{3}\) (b) increasing the temperature (c) reducing the temperature (d) removing nitrogen

Step-by-Step Solution

Answer

(c) reducing the temperature

Step 1: Identify Le Chatelier's Principle

To determine favorable conditions for the Haber process, we apply Le Chatelier's Principle. This principle states that if a dynamic equilibrium is disturbed by changing conditions, the position of equilibrium shifts to counteract the change. The Haber process is exothermic, as indicated by the release of 92.3 kJ.

Step 2: Analyze Temperature Effect

Since the reaction is exothermic, an increase in temperature will shift the equilibrium to the left, minimizing ammonia production. Conversely, reducing the temperature shifts the equilibrium towards the formation of more ammonia to absorb the added heat.

Step 3: Consider Concentration Changes

According to Le Chatelier's Principle, increasing the concentration of products (NH3) would shift the equilibrium to the left, reducing ammonia production. Reducing the concentration of reactants like nitrogen will also shift the equilibrium to the left, unfavorable for ammonia production.

Step 4: Choose the Favorable Condition

Considering the effects of temperature and concentration changes, reducing the temperature favors the forward reaction based on Le Chatelier's Principle, thereby increasing the production of ammonia.

Key Concepts

Haber process

The Haber process is a critical industrial method used for producing ammonia. Discovered by Fritz Haber and Carl Bosch in the early 20th century, this process synthesizes ammonia from nitrogen and hydrogen gases. It is instrumental in producing fertilizers, essential for global food production. To facilitate this conversion, a catalyst is used, often iron-based, but it requires high temperatures (around 400-500°C) and pressures (about 200 atmospheres) to effectively drive the reaction. This process can be represented as:\[\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)\]

exothermic reactions

Exothermic reactions are chemical reactions that release energy, typically in the form of heat. In the context of the Haber process, we can see that the reaction releasing 92.3 kJ signifies it is exothermic. Exothermic reactions are generally characterized by:

• A release of heat, making the surroundings warmer
• A negative enthalpy change (\(\Delta H < 0\))

For the Haber process, this exothermic nature means the formation of ammonia releases heat. Understanding whether a reaction is exothermic or endothermic is crucial in determining conditions that favor or disfavor product formation.

chemical equilibrium

Chemical equilibrium occurs in reversible reactions where the forward and reverse reactions occur at the same rate, leading to constant concentrations of reactants and products. For the Haber process, equilibrium is expressed as:\[\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) + 92.3 \text{ kJ}\] At equilibrium, the amount of nitrogen, hydrogen, and ammonia remains stable over time. The equilibrium state is dynamic, meaning that even though the concentrations are constant, molecules continue to react. Le Chatelier’s Principle helps us understand how a system at equilibrium responds to changes in concentration, pressure, or temperature to restore balance.

temperature effects in chemical reactions

Temperature plays a pivotal role in chemical reactions, profoundly affecting reaction rates and equilibrium states. In an exothermic reaction like the Haber process, increasing temperature shifts the equilibrium to the left, favoring reactants due to the reaction's heat-releasing nature. Conversely, decreasing temperature shifts the equilibrium to the right, favoring more product formation. The response of equilibrium to temperature changes is grounded in Le Chatelier's Principle:

• Increasing temperature drives the reaction towards the endothermic direction (absorbing heat).
• Decreasing temperature favors the exothermic direction, increasing ammonia yield in the Haber process.

In industrial settings, optimizing temperature is crucial for maximizing production while balancing energy costs. Making these adjustments requires a nuanced understanding of temperature's effects on reaction dynamics.