The total energy of an electron in the ground state of a hydrogen atom is $-13.6 \text{ eV}$. The kinetic energy of an electron in the first excited state is.....$\text{ eV}$

$A$ difference of $5.4 \ eV$ separates two energy levels in an atom. What is the frequency of radiation emitted when the atom makes a transition from the upper level to the lower level? (Take $1 \ eV = 1.6 \times 10^{-19} \ J$,$h = 6.625 \times 10^{-34} \ Js$)

In Bohr's atomic model of hydrogen,let $K$,$P$ and $E$ be the kinetic energy,potential energy and total energy of the electron,respectively. Choose the correct option when the electron undergoes transitions to a higher level.

The energy levels for an electron in a certain atom are shown in the figure. Which of the following transitions will emit a photon with the highest energy?

Transitions between three energy levels in a particular atom give rise to three spectral lines of wavelengths $\lambda_1, \lambda_2$ and $\lambda_3$. Given that the wavelengths are in increasing order of magnitude,which one of the following equations correctly relates $\lambda_1, \lambda_2$ and $\lambda_3$?

Energy levels $A, B$ and $C$ of a certain atom correspond to increasing values of energy, i.e., $E_A < E_B < E_C$. If $\lambda_1, \lambda_2$ and $\lambda_3$ are the wavelengths of radiations corresponding to transitions $C$ to $B, B$ to $A$ and $C$ to $A$ respectively, which of the following relations is correct?