Energy $E$ of a hydrogen atom with principal quantum number $n$ is given by $E=\frac{-13.6}{n^2}$ $\mathrm{eV}$. The energy of a photon ejected when the electron jumps from $\mathrm{n}=3$ state to $\mathrm{n}=2$ state of hydrogen is approximately:

Ionization potential of hydrogen atom is $13.6 \mathrm{eV}$. Hydrogen atoms in the ground state are excited by monochromatic radiation of photon energy $12.1 \mathrm{eV}$. According to Bohr's theory, the spectral lines emitted by hydrogen will be

The diagram shows the energy levels for an electron in a certain atom. Which transition shown represents the emission of a photon with the most energy?

Energy levels $A, B$ and $C$ of a certain atom corresponding to increasing values of energy, i.e., $E_{\mathrm{A}}<E_{\mathrm{B}}<E_{\mathrm{C}}$. If ${\lambda }_1,{\lambda }_2$ and ${\lambda }_3$ are the wavelengths of radiations corresponding to the transitions $C$ to $B, B$ to $A$ and $C$ to $A$ respectively, which of the following statements is correct?