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.
2H2S(g) + 3O2(g) → 2H2O(l ) + 2SO2(g);
∆H° = –1124 kJ
The density of sulfur dioxide at 25oC and 1.00 atm is 2.62 g/L, and the molar heat capacity is 30.2 J/(mol•oC). (a) How much heat would be evolved in producing 1.00 L of SO2 at 25oC and 1.00 atm? (b) Suppose heat from this reaction is used to heat 1.00 L of SO2 from 25oC and 1.00 atm to 500oC for its use in the next step of the process. What percentage of the heat evolved is required for this?