Final answer:
The electron emitted light with a wavelength of 397.2 nm to transition down to the n=2 level, indicating it originated from the n=4 level. For a hydrogen atom at the n=4 level, there is 1 electron, 16 orbitals, and the subshells are s, p, d, and f.
Step-by-step explanation:
The electron that emits light with a wavelength of 397.2 nm to reach the energy level n=2 must have transitioned from a higher energy level. Given the specific wavelength mentioned, the transition is associated with the Paschen series where electrons fall to the n=3 level, but since we're talking about a transition to n=2, it must instead be from the Balmer series (visible range of the spectrum).
We can identify the initial principal quantum level by using the Rydberg formula to calculate the transitions and match them with the wavelengths of the emitted light. Comparing the given wavelength with the known spectral lines for hydrogen, we can conclude that the initial quantum level was n=4 (option D).
In the principal quantum level n=4, we deal with the fourth shell of an atom. This shell can contain up to 32 electrons, but since we're dealing with a hydrogen atom, there is only one electron. The fourth shell consists of the s, p, d, and f subshells. Specifically, it has 1 s orbital, 3 p orbitals, 5 d orbitals, and 7 f orbitals, giving a total of 16 orbitals.
Since the hydrogen atom only has one electron, the correct answer to the number of electrons, orbitals, and the names of the subshells associated with n=4 is 1 electron, 16 orbitals, consisting of the s, p, d, and f subshells (none of the options A to D are correct).