Spectrum of Hydrogen Atom

An electric discharge produces excited hydrogen atoms, which dissociate hydrogen molecules and excite the atoms. These atoms emit light at specific wavelengths, some of which fall within the visible spectrum. In a spectrograph, visible light is dispersed by a diffraction grating and detected by photographic film.

The wavelength controls the position of the lines in the spectrum of the hydrogen atom, and it is calibrated using the Hg spectrum, which has well-known and intense lines. The Rydberg constant can then be calculated using measurements of the H atom lines.

Bohr’s Model

Niels Bohr (1885–1962; Nobel Prize in Physics, 1922), a Danish physicist, created a theoretical model for the hydrogen atom in 1913 that described its emission spectrum. Bohr’s model requires one assumption: the electron goes around the nucleus in circular orbits with only a few permitted radii.

Rutherford’s former atomic model similarly assumed that electrons orbited the nucleus in circular orbits, and the atom was held together by electrostatic interaction between the positively charged nucleus and the negatively charged electron. Although we now know that circular orbits are erroneous, Bohr’s breakthrough argued that the electron could only occupy particular regions of space.

Bohr’s model on spectral lines of hydrogen

Bohr’s hypothesis explains the spectrums of the hydrogen atom. Under the spectrums of hydrogen atom study material, you will learn that the energy of an atom’s electron remains unaltered while it is in the ground state. The electron jumps from the ground state orbit to a more distant excited state orbit when the atom absorbs one or more quanta of energy.

The variable n is used to identify energy levels. n = 1 is the ground state, n = 2 is the first excited state and so on. The difference in energy between the two energy levels is equal to the energy obtained by the atom. When the atom relaxes back to a lower energy state, energy equal to the difference in energy between the two orbits is released.

According to the equation E = hv, the change in energy, ΔE, causes the light of a specific frequency to be emitted. Remember that the atomic emission spectrum of hydrogen has four different frequency spectral lines. The Bohr model explains that electron orbits are not evenly spaced. The spacing between the levels gets narrower and smaller as the energy travels farther away from the nucleus.

Bohr calculated the energies that the hydrogen electron would have in each of its permissible energy levels using the wavelengths of the spectral lines. The energy level transitions corresponding to the spectral lines in the atomic emission spectrum were then mathematically demonstrated.

The outcome of Bohr’s model on hydrogen spectrum lines

He discovered that the four visible spectral lines correlated to transitions from higher to lower energy levels (n = 2); these are called the Balmer series. The Lyman series refers to transitions that finish in the ground state (n = 1), but the energy released is so high that the spectral lines are all in the ultraviolet part of the spectrum. Because the energies are too small, the transitions known as the Paschen series and the Brackett series both produce spectral lines in the infrared region.

Bohr’s concept was a massive achievement in understanding the study material notes on the spectrums of the hydrogen atom. Unfortunately, applying the model’s mathematics to atoms containing more than one electron predicted the frequencies of the spectral lines incorrectly. Bohr’s model was a significant step forward in atomic theory, and the concept of electron transitions between energy levels is still relevant. However, further work was required to comprehend all atoms and their chemical activity completely.

Limitations of Bohr model of an atom

• The Bohr model was unable to explain the multi-electron atoms’ line spectra.
• He could not explain why lines split in the magnetic and electric fields (Zeeman effect) (stark effect).
• He was unable to explain the three-dimensional atom model.
• He was unable to explain molecule shapes.
• He could not explain the Heisenberg uncertainty principle or the dual nature of matter.