The maximum wavelength of incident radiation required to ionize a hydrogen atom in its ground state is nearly

The diagram shows different transitions across the energy levels for an electron in a certain atom. Among these,which transition represents the emission of a photon with the most energy?

In the $n^{th}$ orbit,the energy of an electron is given by $E_n = -\frac{13.6}{n^2} \text{ eV}$ for a hydrogen atom. The energy required to excite the electron from the first orbit $(n=1)$ to the second orbit $(n=2)$ will be..... eV.

In a hydrogen atom,when an electron jumps from the second to the first orbit,the energy emitted is:

An electron collides elastically with an $H$-like atom and excites it from the ground state to $n = 3$. Find the energy transferred to the $H$-like atom in $eV$.

The energy levels of an atom are shown in the figure. Which one of these transitions will result in the emission of a photon of wavelength $124.1 \, nm$? Given $(h = 6.62 \times 10^{-34} \, Js)$.