Question

Consider the following reaction, which occurs at room temperature and pressure: $$ 2 \mathrm{Cl}(g) \longrightarrow \mathrm{Cl}_{2}(g) \quad \Delta H=-243.4 \mathrm{~kJ} $$ Which has the higher enthalpy under these conditions, \(2 \mathrm{Cl}(g)\) or \(\mathrm{Cl}_{2}(g) ?\)

Short answer

Under the given conditions, 2 Cl(g) has a higher enthalpy than Cl₂(g), as the reaction is exothermic with a ∆H of -243.4 kJ.

Step-by-step solution

Understand the given reaction and enthalpy change

The reaction we are dealing with is: \( 2 \,\mathrm{Cl}(g) \longrightarrow \mathrm{Cl}_{2}(g)\), where ∆H = -243.4 kJ. A negative value of ∆H indicates that the reaction is exothermic, meaning that energy is released as the reaction occurs.

Relate the enthalpy change to the reactants and products

Let's use the given enthalpy change to relate the enthalpy of the reactants and products. For an exothermic reaction, the enthalpy of the products is lower than the enthalpy of the reactants. Mathematically, we can express this as: Enthalpy of reactants - Enthalpy of products = ∆H

Determine which species has the higher enthalpy

Since the reaction is exothermic, the enthalpy of the reactants (2 Cl(g)) is higher than the enthalpy of the products (Cl₂(g)). Therefore, under these conditions, 2 Cl(g) has higher enthalpy than Cl₂(g).