The energy of an electron in the first Bohr orbit of the $H$ atom is $-13.6 \ eV$. The magnitude of the energy value of an electron in the first excited state of $Be^{3+}$ is $.......... \ eV$ $(nearest \ integer \ value)$.

Calculate the energy released when an electron in a hydrogen atom transitions from the $3^{rd}$ energy level to the $2^{nd}$ energy level.

The ratio of the highest to the lowest wavelength of Lyman series is

The wavelength of a spectral line emitted by a hydrogen atom in the Lyman series is $\frac{16}{15 R} \ cm$. What is the value of $n_2$? $(R = \text{Rydberg constant})$

In a hydrogen atom, the minimum energy required to excite an electron from the $2^{nd}$ orbit to the $3^{rd}$ orbit is: (in $eV$)

Based on the equation $\Delta E = - 2.0 \times 10^{-18} \, J \left( \frac{1}{n_2^2} - \frac{1}{n_1^2} \right)$,the wavelength of the light that must be absorbed to excite a hydrogen electron from level $n = 1$ to level $n = 2$ will be: $(h = 6.625 \times 10^{-34} \, J \, s, c = 3 \times 10^8 \, m \, s^{-1})$