Question

Which will produce more energy per gram of hydrogen: H atoms undergoing an electronic transition from n=4 level to the n=1 level or hydrogen gas burned in the reaction below and ΔHf(H2O) = -285.8 kJ/mol? 2H2(g) + O2(g) → 2H2O(l)

Answer 1

To determine whether H atoms undergoing an electronic transition or the combustion of hydrogen gas releases more energy per gram, we need to calculate the energy released in both processes.

Step-by-step explanation:

The question asks which will produce more energy per gram of hydrogen: H atoms undergoing an electronic transition from n=4 level to the n=1 level or hydrogen gas burned in the reaction 2H2(g) + O2(g) → 2H2O(l).

To determine this, we need to calculate the energy released in both processes. The electronic transition in hydrogen atoms involves the emission of photons with specific energies. On the other hand, the combustion reaction releases energy in the form of heat. By comparing the energy per gram released, we can determine which process produces more energy.

It is important to note that the mass of one mole of hydrogen gas is 2 grams, so we will consider this when comparing the energy per gram for each process.

Answer 2

Comparing the energy produced from an H atom electronic transition with the burning of H₂(g) to form H₂O(l), the combustion reaction yields about 142.9 kJ/g. However, without a direct calculation, it is generally expected that an electronic transition would release significantly more energy per gram of hydrogen than combustion.

Step-by-step explanation:

The question asks which will produce more energy per gram of hydrogen: an H atom undergoing an electronic transition from n=4 level to the n=1 level or hydrogen gas (H₂(g)) burned in a chemical reaction to form liquid water, given that the enthalpy of formation for water (ΔHf(H₂O)) is -285.8 kJ/mol. To answer this, we need to compare the energies involved in both processes.

However, the energy released in the electronic transition of a hydrogen atom is a quantum mechanical phenomenon, best described using the Rydberg formula, and doesn't yield its value directly in kJ/mol without the necessary constants and calculations. Conversely, when hydrogen gas is burned to form liquid water, the enthalpy change can be directly inferred from the standard enthalpy of formation per mole of product formed.

Considering that the mass of hydrogen in a mole (2 grams, since dihydrogen has a molecular weight of approximately 2 g/mol) releases 285.8 kJ when burned, we can calculate the energy released per gram of hydrogen as roughly 142.9 kJ/g. Without specific quantum mechanical calculations for the electronic transition, we cannot accurately compare this value to the energy released during an electronic transition, but it is known that chemical reactions typically release much less energy compared to nuclear or atomic transitions.

Hence, we would generally expect the electronic transition of a hydrogen atom to release significantly more energy compared to the combustion reaction per gram of hydrogen.