Which of the following electronic transitions in a hydrogen atom will require the largest amount of energy?

If the radius of an electron in the excited state of $He^{+}$ is $0.4232 \ nm$, the energy of the electron in that excited state in $J$ is: (The radius and energy of an electron in the first orbit of a hydrogen atom are $52.9 \ pm$ and $-2.18 \times 10^{-18} \ J$ respectively)

What is the energy in joules,required to shift the electron of the hydrogen atom from the first Bohr orbit to the fifth Bohr orbit and what is the wavelength of the light emitted when the electron returns to the ground state?
The ground state electron energy is $-2.18 \times 10^{-11} \ erg$.

Which of the following represents the wavelength of a spectral line of the Balmer series of the $He^{+}$ ion? ($R=$ Rydberg constant,$n > 2$)

The potential energy of an electron present in $He^{+}$ is:

Calculate the wavelength of the emitted radiation when an electron transitions from $n = 3$ to $n = 2$ in a hydrogen atom. To which region does this radiation belong?