The change in internal energy of a system ΔE is related to heat and work by the equation:
Where q is heat and w is work.
The sign of q changes depending on the type of reaction:
(+) q when a system absorbs heat or energy (endothermic)
(-) q when the system releases energy(exothermic)
In the given problem, 40.66 kJ of heat is released when the 1 mole of H2O(g) is condensed to 1 mole of H2O(l). Therefore, the value of q = -40.66 kJ.
Let us now determine the value for work.
One mole of H2O(g) at 1.00 atm and 100.°C occupies a volume of 30.6 L. When 1 mole of H2O(g) is condensed to 1 mole of H2O(l) at 1.00 atm and 100.°C, 40.66 kJ of heat is released. If the density of H2O(l) at this temperature and pressure is 0.996 g/cm3, calculate ΔE for the condensation of 1 mole of water at 1.00 atm and 100.°C.
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Based on our data, we think this problem is relevant for Professor McCamant's class at UR.
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Our data indicates that this problem or a close variation was asked in Chemistry: An Atoms First Approach - Zumdahl Atoms 1st 2nd Edition. You can also practice Chemistry: An Atoms First Approach - Zumdahl Atoms 1st 2nd Edition practice problems.