Question

Which of the following statements is true concerning the decomposition of liquid water to form hydrogen gas and oxygen gas? 2H2O(l) → 2H2(g) + O2(g)

a. ΔH is greater than ΔU because the pressure is constant.

b. ΔH is less than ΔU because of the pressure–volume work done by the gaseous products.

c. ΔH is less than ΔU because the atmosphere does pressure–volume work on the gaseous products.

d. ΔH equals ΔU because both are state functions.

e. ΔH is greater than ΔU because of the pressure–volume work done by the gaseous products.

The Answer is E. But I need help in understand why that is and why all the other answers are wrong.

Answer 1

Ok then here we are comparing the H and the U.

First lets remember the definitions.

Internal energy U is defined as the energy associated with the random, disordered motion of molecules. The internal energy U might be thought of as the energy required to create a system in the absence of changes in temperature or volume. But if the process changes the volume, as in a chemical reaction which produces a gaseous product, then work must be done to produce the change in volume.

H = U + PV

it is a useful quantity for tracking chemical reactions. If as a result of an exothermic reaction some energy is released to a system, it has to show up in some measurable form in terms of the state variables. An increase in the enthalpy H = U + PV might be associated with an increase in internal energy which could be measured by calorimetry, or with work done by the system, or a combination of the two.

Knowing all these how we apply this knowledge to the reaction above.

We know that the  H of the decomposition of water is a positive value. That tell us that the energy of the reactants are greater than the products,So we need to introduce energy because this is a nonspontaneous reaction. We know that the internal energy system wil be always the same because of the conservation of the energy law.