Chapter 6, Problem 6.158QP

An industrial process for manufacturing sulfuric acid, H₂SO₄, uses hydrogen sulfide, H₂S, from the purification of natural gas. In the first step of this process, the hydrogen sulfide is burned to obtain sulfur dioxide, SO₂.

$\begin{array}{c}2{\text{H}}_2\text{S}(g)+3{\text{O}}_2(g)\to 2{\text{H}}_2\text{O}(l)+2{\text{SO}}_2(g);\\ \Delta H°=-1124\text{kJ}\end{array}$

The density of sulfur dioxide at 25°C and 1.00 atm is 2.62 g/L, and the molar heat capacity is 30.2 J/(mol · °C). (a) How much heat would be evolved in producing 1.00 L of SO₂ at 25°C and 1.00 atm? (b) Suppose heat from this reaction is used to heat 1.00 L of the SO₂ from 25°C to 500°C for its use in the next step of the process. What percentage of the heat evolved is required for this?