According to the Bohr model of the atom (and confirmed more rigorously by quantum mechanics), electrons in an atom can only exist in certain allowed, discrete energy levels or orbitals — they cannot possess arbitrary, continuous values of energy. Each of these levels corresponds to a specific, well-defined energy value, and the atom is said to have a quantized energy structure.

Spectral lines arise because electrons can transition between these discrete energy levels, and each such transition involves the emission or absorption of a photon whose energy exactly matches the energy difference between the two levels:

• Emission (downward transition): when an electron falls from a higher energy level E_i to a lower one E_f, the atom emits a photon of energy ΔE = E_i − E_f = hν, where h is Planck's constant and ν is the frequency of the emitted radiation.
• Absorption (upward transition): when an electron is excited from a lower energy level to a higher one, it must absorb a photon carrying exactly the required energy ΔE = hν.

Because each pair of energy levels in an atom has its own fixed energy gap, each transition produces radiation of one specific frequency (and therefore one specific wavelength). When such radiation is passed through a spectrometer, it appears as a sharp, discrete spectral line at that wavelength rather than as a continuous smear of colours — this is why atomic emission and absorption spectra consist of distinct lines unique to each element (the atomic "fingerprint" used in spectroscopy).

The remaining statements are inconsistent with this picture: transitions are never energy-neutral — every transition necessarily involves an exchange of a definite quantum of energy with the surroundings (via a photon), so the idea of "no exchange of energy" is incorrect. Spectral lines are also strongly dependent on the electronic configuration of the atom, since the specific set of allowed energy levels (and hence the specific spectral lines produced) differs from element to element — this dependence is precisely what allows spectroscopy to identify elements. Finally, electrons do not move continuously between levels; if energy levels were not quantized, atoms would emit a continuous spectrum rather than the sharp line spectra that are actually observed experimentally. Therefore, spectral lines originate from electrons jumping between quantized energy levels while emitting or absorbing photons of specific, well-defined energies.