$A$ hypothetical atom has only three energy levels. The ground level has energy,$E_1 = -8 \ eV$. The two excited states have energies,$E_2 = -6 \ eV$ and $E_3 = -2 \ eV$. Which of the following wavelengths will $NOT$ be present in the emission spectrum of this atom (in $nm$)?

Which of the following transitions will have the highest emission wavelength?

The number of spectral lines emitted by atomic hydrogen that is in the $4^{\text{th}}$ energy level is

$A$ monochromatic light is incident on a hydrogen sample in the ground state. Hydrogen atoms absorb a fraction of light and subsequently emit radiation of six different wavelengths. The frequency of incident light is $x \times 10^{15} \ Hz$. The value of $x$ is (Given $h = 4.25 \times 10^{-15} \ eVs$)

If $13.6 \ eV$ energy is required to ionize the hydrogen atom,then the energy required to remove an electron from $n=2$ is (in $eV$)

The energy levels with transitions for an atom are shown in the figure. The transitions corresponding to the emission of radiation of maximum and minimum wavelength are respectively: