Atomic Spectra

Atomic spectra refer to the distinct set of wavelengths of light emitted or absorbed by electrons in an atom as they transition between different energy levels. These transitions occur when an electron moves from one orbit to another, releasing or absorbing energy in the form of photons. The collection of these unique wavelengths, influenced by factors such as temperature and pressure, forms the atomic spectrum of an element. Here, we’ll explore what atomic spectra are, how they work, and look at the hydrogen atom’s specific light patterns.

What is Atomic Spectra?

An electron's spectrum of electromagnetic radiation is released or absorbed as it moves between different energy levels within an atom. When an electron moves from one energy level to the next, it emits or absorbs light of a given wavelength.

Atomic spectra are the specific wavelengths of light emitted or absorbed by electrons in an atom when they transition between energy levels.

These are the collection of unique wavelengths (light) emitted or absorbed by electrons during these transitions. The atomic spectrum is influenced by various factors like pressure and temperature.

The three main forms of atomic spectra are:

• Continuous spectra: A broad range of wavelengths emitted, typically from a hot object.
• Emission spectra: Light emitted when electrons move to lower energy levels.
• Absorption spectra: Light absorbed when electrons move to higher energy levels.

The Rydberg formula clearly divides the atomic hydrogen emission spectrum into a number of spectral lines with wavelengths. Atomic transitions between different energy levels cause the observable spectral lines in the hydrogen emission spectrum. In astronomical spectroscopy, spectral series are very important.

Atomic Spectroscopy

Atomic spectroscopy is the study of the electromagnetic radiation emitted or absorbed by atoms. It involves three main types of spectroscopy:

1. Atomic Emission Spectroscopy: This focuses on the transfer of energy from the ground state to an excited state, explaining the electronic transitions when atoms emit light.
2. Atomic Absorption Spectroscopy: This occurs when electrons absorb radiation to transition from a lower energy level to a higher one. It relies on the principle that free electrons in an atomizer can absorb radiation at specific frequencies. The absorption by ground-state atoms in the gaseous phase is measured.
3. Atomic Fluorescence Spectroscopy: This technique combines both atomic emission and absorption, utilizing both excitation and de-excitation radiation.

Uses of Atomic Spectroscopy

• It is used to identify the spectral lines of metallurgical materials.
• It is utilised in the pharmaceutical industry to detect traces of materials that have been used.
• It can be used to investigate elements with multiple dimensions.

Spectral Series

A spectral series is a sequence of wavelengths emitted or absorbed by energized atoms, arranged in a logical order. The hydrogen atom, being the simplest, produces the most basic spectral lines.

When light passes through a slit into a spectrometer, it forms an image of the source, which can be resolved under the spectroscope. The resulting image is shown as parallel lines with consistent spacing. When moving from the higher to lower wavelength side, the lines are farther apart at the higher wavelength end and eventually converge. The shortest wavelength corresponds to the fewest separated spectral lines, known as the series limit.

Line spectrum of the hydrogen atom

A hydrogen atom is made up of several line spectrum series, including:

• Pfund Series
• Brackett Series
• Paschen Series
• Balmer Series
• Lyman Series

Spectral Series Formation

Bohr's atomic model effectively explains the set of energy levels (states) within an atom. These energy states are represented by quantum numbers (n=1, 2, 3, 4, 5, 6,...). When an electron transitions from a higher energy state (nh) to a lower energy state (nl), a photon with energy E=h( 1/n²_h-1/n²_l) is emitted. As the energy associated with each state is fixed, the energy difference between them remains constant, producing a photon of the same energy during each transition.

These transitions divide the spectral series into equivalent series, with the spectral lines being separated using Greek letters to represent transitions between corresponding energy levels. Below are the key spectral series in hydrogen:

Lyman Series (nl=1)

• Discovered by Theodore Lyman between 1906 and 1914.
• The Lyman series occurs when electrons transition from higher energy levels (nh=2, 3, 4, 5, 6,…) to the lowest energy state (nl=1).
• All wavelengths in the Lyman series fall within the ultraviolet band.

Balmer Series (nl=2)

• Discovered by Johann Balmer in 1885.
• The Balmer series occurs when electrons transition from higher energy levels (nh=3, 4, 5, 6, 7,…) to the second energy level (nl=2).
• The wavelengths of the Balmer series fall within the visible range of the electromagnetic spectrum (400 nm to 740 nm).

Paschen Series (nl=3)

• Discovered by Friedrich Paschen in 1908.
• The Paschen series occurs when electrons move from higher energy levels (nh=4, 5, 6, 7, 8,…) to the third energy state (nl=3).
• All wavelengths in the Paschen series are in the infrared portion of the electromagnetic spectrum.

Brackett Series (nl=4)

• Discovered by Friedrich Sumner Brackett in 1922.
• The Brackett series occurs when electrons transition from higher energy levels (nh=5, 6, 7, 8, 9,…) to the fourth energy level (nl=4).
• The wavelengths of the Brackett series are in the infrared range of the electromagnetic spectrum.

Pfund Series (nl=5)

• Discovered by August Harman Pfund in 1924.
• The Pfund series appears when electrons transition from higher energy states (nh=6, 7, 8, 9, 10,…) to the fifth energy level (nl=5).
• The wavelengths of the Pfund series are in the infrared region of the electromagnetic spectrum.

Humphreys Series (nl=6)

• Discovered by Curtis J Humphreys in 1953.
• The Humphreys series occurs when electrons move from higher energy levels (nh=7, 8, 9, 10, 11,…) to the sixth energy level (nl=6).
• The wavelengths of the Humphreys series fall in the infrared region of the electromagnetic spectrum.

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• Bohr's Model of Atom
• Electromagnetic Spectrum