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) + 3 O2(g) → 2 H2O(l) + 2 SO2(g); ∆H = -1124 kJ.  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 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?

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