Explain the transition between energy levels in an atom by drawing a line spectra diagram.

(N/A) The wavelength of light emitted during a transition between different atomic energy states is given by the Rydberg formula:
$\frac{1}{\lambda_{if}} = R \left[ \frac{1}{n_f^2} - \frac{1}{n_i^2} \right]$
Where $R$ is the Rydberg constant,$n_f$ is the final energy level,and $n_i$ is the initial energy level $(n_i > n_f)$.
$1$. Lyman Series: If $n_f = 1$ and $n_i = 2, 3, 4, \ldots$,the spectral lines are obtained in the ultraviolet region.
$2$. Balmer Series: If $n_f = 2$ and $n_i = 3, 4, 5, \ldots$,the spectral lines are obtained in the visible region.
$3$. Paschen Series: If $n_f = 3$ and $n_i = 4, 5, 6, \ldots$,the spectral lines are obtained in the infrared region.
$4$. Brackett Series: If $n_f = 4$ and $n_i = 5, 6, 7, \ldots$,the spectral lines are obtained in the infrared region.
$5$. Pfund Series: If $n_f = 5$ and $n_i = 6, 7, 8, \ldots$,the spectral lines are obtained in the infrared region.
The spectral lines resulting from these transitions in a hydrogen atom are illustrated in the diagram below.