Question

Chlorine trifluoride is a toxic, intensely reactive gas. It was used in World War II to make incendiary bombs. It reacts with ammonia and forms nitrogen, chlorine, and hydrogen fluoride gases. When two moles of chlorine trifluoride reacts, \(1196 \mathrm{~kJ}\) of heat is evolved. (a) Write a thermochemical equation for the reaction. (b) What is \(\Delta H_{\mathrm{f}}^{\circ}\) for \(\mathrm{ClF}_{3} ?\)

Short answer

Question: Write the thermochemical equation for the reaction of chlorine trifluoride with ammonia and calculate the standard enthalpy of formation for chlorine trifluoride. Answer: The thermochemical equation for the reaction is: 2ClF₃ + 6NH₃ → 4N₂ + 3Cl₂ + 6HF + 1196 kJ The standard enthalpy of formation for chlorine trifluoride is: ΔHf°(ClF₃) = -598 kJ/mol

Step-by-step solution

(a) Write the thermochemical equation

First, we need to write the balanced chemical equation for the reaction of chlorine trifluoride with ammonia. The reaction can be written as: \(\mathrm{ClF}_{3} + \mathrm{NH}_{3} \rightarrow \mathrm{N}_{2} + \frac{3}{2}\mathrm{Cl}_{2} + 3\mathrm{HF}\) We need to balance the equation: 2\(\mathrm{ClF}_{3} + 6\mathrm{NH}_{3} \rightarrow 4\mathrm{N}_{2} + 3\mathrm{Cl}_{2} + 6\mathrm{HF}\) Now, we can include the heat evolved in the reaction. The problem states that when two moles of chlorine trifluoride reacts, \(1196 \mathrm{~kJ}\) of heat is evolved. Therefore, the thermochemical equation becomes: 2\(\mathrm{ClF}_{3} + 6\mathrm{NH}_{3} \rightarrow 4\mathrm{N}_{2} + 3\mathrm{Cl}_{2} + 6\mathrm{HF} + 1196\mathrm{~kJ}\)

(b) Calculate \(\Delta H_{\mathrm{f}}^{\circ}\) for \(\mathrm{ClF}_{3}\)

To calculate the standard enthalpy of formation for chlorine trifluoride, we can use the given heat evolved for 2 moles of chlorine trifluoride in the reaction: \(1196 \mathrm{~kJ}\) evolved for 2 moles of \(\mathrm{ClF}_{3}\) We need to find the heat evolved for 1 mole of \(\mathrm{ClF}_{3}\): \(\frac{1196 \mathrm{~kJ}}{2} = 598\mathrm{~kJ}\) evolved for 1 mole of \(\mathrm{ClF}_{3}\) Since heat is evolved in the reaction, the enthalpy change, \(\Delta H\), is negative. Thus, the standard enthalpy of formation for \(\mathrm{ClF}_{3}\), denoted by \(\Delta H_{\mathrm{f}}^{\circ}\), is: \(\Delta H_{\mathrm{f}}^{\circ} (\mathrm{ClF}_{3}) = -598 \mathrm{~kJ/mol}\)

Key concepts

Chlorine Trifluoride

Chlorine trifluoride (ClF₃) is an extremely reactive and toxic gas. Its powerful oxidizing properties make it a significant compound historically used in incendiary weapons. It's important to handle chlorine trifluoride with care due to its vigorous chemical reactivity, especially when it comes into contact with combustible materials.
This compound can react with substances like ammonia (NH₃), leading to the formation of gases such as nitrogen (N₂), chlorine (Cl₂), and hydrogen fluoride (HF). The reaction releases a considerable amount of heat, which further illustrates its intense reactivity. Understanding the behavior and characteristics of chlorine trifluoride is crucial for safely managing its applications in various industrial processes.

Enthalpy Change

Enthalpy change ( ∆H ) is a thermodynamic concept representing the heat exchange in a reaction at constant pressure. In essence, it provides insight into whether a chemical process is endothermic or exothermic.
Specifically, the enthalpy change associated with a reaction involving chlorine trifluoride and ammonia results in an evolution of 1196 kJ of heat. This indicates that the process is exothermic.
Because the reaction releases heat, ∆H is negative, reflecting the energy lost from the chemical system to the surroundings. By focusing on enthalpy change, chemists can predict the heat involved in transformations, which is crucial for calculations in larger chemical processes and for ensuring safety when handling reactive compounds.

Enthalpy of Formation

Enthalpy of formation ( ∆Hₓₓf ) refers to the energy change when one mole of a compound forms from its elements in their standard states. It provides a fundamental insight when evaluating the stability and energy associated with molecules in reactions.
For chlorine trifluoride, the standard enthalpy of formation is calculated using the energy released during its reaction with ammonia. Given that 1196 kJ of energy is evolved for two moles of chlorine trifluoride, the energy released per mole is 598 kJ, indicating the energy associated with forming chlorine trifluoride from elemental chlorine and fluorine.
In this scenario, since the heat is released when ClF₃ is formed, the value of ∆Hₓₓf is negative, specifically -598 ext{ kJ/mol} . A negative enthalpy of formation signifies that formation releases energy, suggesting that chlorine trifluoride is a thermodynamically favorable compound.