Atoms, molecules and substances also have an emission spectrum due to the emission of light when they absorb the right amount of external energy to travel between two energy states. When this light passes through a prism, it breaks down into colored spectral lines with different wavelengths characteristic of each element.
The importance of the emission spectrum is that it allows the composition of unknown substances and astronomical objects to be determined by analyzing their spectral lines using emission spectroscopy techniques.
In the following, it is explained what it is and how the emission spectrum is interpreted, some examples are mentioned and the differences existing between the emission spectrum and the absorption spectrum.
What is an emission spectrum?
The atoms of an element or substance have electrons and protons that are held together by the force of electromagnetic attraction. According to Bohr’s model, electrons are arranged in such a way that the energy of the atom is as low as possible. This energy level is called the ground state of the atom.
When atoms acquire energy from the outside, electrons move to a higher energy level and the atom changes its ground state to an excited state.
In the excited state, the residence time of the electron is very short (± 10-8 s) (1), the atom is unstable and returns to the ground state, passing, if necessary, through intermediate energy levels.
In the process of transition from an excited state to a ground state, the atom emits a photon of light with energy equal to the energy difference between the two states, being directly proportional to the frequency v and inversely proportional to its wavelength λ.
Interpretation of the emission spectrum
Some atom transitions are caused by a rise in temperature or by the presence of other external sources of energy, such as a beam of light, a stream of electrons, or a chemical reaction.
If a gas such as hydrogen is placed in a chamber at low pressure and an electric current passes through the chamber, the gas will emit a light of its own color that differentiates it from other gases.
By passing the emitted light through a prism, instead of obtaining a rainbow of light, discrete units are obtained in the form of colored lines with specific wavelengths, which carry discrete amounts of energy.
The emission spectrum lines are unique in each element and its use of the spectroscopy technique allows to determine the elementary composition of an unknown substance, as well as the composition of astronomical objects, analyzing the wavelengths of the photons emitted during the atom transition.
Difference between the emission spectrum and the absorption spectrum.
In the processes of absorption and emission, the atom has transitions between two energy states, but it is during absorption that it obtains energy from the outside and reaches the state of excitation.
The spectral emission line is opposite the continuous white light spectrum. In the first, the spectral distribution is observed in the form of bright lines and in the second, a continuous band of colors.
If a beam of white light hits a gas such as hydrogen, locked in a chamber at low pressure, only a part of the light will be absorbed by the gas and the rest will be transmitted.
When transmitted light passes through a prism, it breaks down into spectral lines, each with a different wavelength, forming the gas absorption spectrum.
The absorption spectrum is completely opposite to the emission spectrum and is also element-specific. When comparing the two spectra of the same element, it is observed that the spectral emission lines are those that are absent in the absorption spectrum
Examples of chemical element emission spectra
a) The spectral lines of the hydrogen atom, in the visible region of the spectrum, are a red line at 656.3 nm, a light blue line at 486.1 nm, a dark blue line at 434 nm and a very faint violet at 410 nm. These wavelengths are obtained from the Balmer-Rydberg equation in its modern version (3).
is the wave number of the spectral line
is the Rydberg constant (109666.56 cm-1)
is the highest energy level
is the highest energy level
b) The helium emission spectrum has two series of main lines, one in the visible region and the other close to the ultraviolet. Peterson (4) used the Bohr model to calculate a series of helium emission lines in the visible part of the spectrum, as a result of several simultaneous transitions of two electrons to the n = 5 state, and obtained wavelength values consistent with the experimental results. The wavelengths obtained are 468.8 nm, 450.1 nm, 426.3 nm, 418.4 nm, 412.2 nm, 371.9 nm.
c) The sodium emission spectrum has two very bright lines of 589nm and 589.6nm called D lines (5) . The other lines are much weaker than these and for practical purposes all sodium light is considered to come from the D lines.