Question

Given the following information, how would the reaction between sulfur dioxide and oxygen to produce sulfur trioxide be best described?

2S(s) + 2O_2(g) = 2SO_3 (g)
∆ H= -790 kJ/mol
S(s) + O_2(g) = SO_2(g)
∆ H= -297 kJ/mol
The answer should be "exothermic with net energy of 196 kJ" but i need a step by step explanation of how they got that answer.

Answer 1

To find the net energy released in the formation of sulfur trioxide from sulfur dioxide and oxygen, we sum the enthalpy changes of the two steps involved in the reaction process and ensure the stoichiometry is consistent. The calculation shows an exothermic reaction with a net enthalpy change of -445.5 kJ per mole of sulfur dioxide, which indicates a discrepancy with the initially stated answer of 196 kJ.

Step-by-step explanation:

The reaction between sulfur dioxide and oxygen to produce sulfur trioxide can be described as exothermic because it releases heat. To understand why the net energy is 196 kJ, we need to look at the enthalpy changes of the steps involved in the process:

• The formation of SO2 from S(s) and O2(g) has an enthalpy change (ΔH) of -297 kJ/mol.
• The given reaction to form SO3 from SO2 and O2 has a ΔH of -790 kJ/mol.

To compute the net enthalpy change for the production of SO3 from elemental sulfur, we need to add the enthalpy changes of the two steps:

1. Formation of SO2: S(s) + O2(g) → SO2(g), ΔH = -297 kJ/mol
2. Formation of SO3 from SO2: 2SO2(g) + O2(g) → 2SO3(g), ΔH = -790 kJ for 2 moles of SO2, which is -395 kJ/mol for SO2.

Sum of ΔH: (-297 kJ) + (-395 kJ) = -692 kJ for 1 mole of S(s) reacted. However, the actual net energy released is doubled because the reaction involves 2 moles of sulfur: -692 kJ/mol x 2 = -1384 kJ. Therefore, when you compare -1384 kJ with the given ΔH for the formation of 2SO3, which is -790 kJ, the difference is the net energy released: -1384 kJ (formation from elemental sulfur) - (-790 kJ) = 594 kJ for 2 moles of SO3. Consequently, for 1 mole it would be 594 kJ / 2 = 297 kJ. Finally, adjust for the stoichiometry of the reaction (divide by 2 since it is per mole of S, not per 2 moles as in the initial reaction), giving us (-297 kJ / 2) = -148.5 kJ as the net energy release per mole of S. Adding the two steps to correct for half the reaction yields -297 kJ + (-148.5 kJ) = -445.5 kJ per mole of sulfur dioxide, which is the actual enthalpy difference for the desired reaction. Clearly, there seems to be an error in the initial answer of 196 kJ as it does not align with the provided equations and the calculations derived from them.