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# Atomic Spectra – Definition, Usage, Formula, Examples

Atoms have an equal number of negative and positive charges. Atoms were described as a spherical cloud of positive charges with embedded electrons in Thomson’s concept. In Rutherford’s model, one tiny nucleus carries the majority of the atom’s mass, as well as its positive charges, and the electrons orbit it.

Rutherford’s approach failed to account for two factors: It could not agree on the matter’s stability since it predicts that atoms are unstable because electrons revolving around the nucleus may spiral into the nucleus. It is unable to explain a line of atoms’ spectral properties. Every element’s atoms have their own unique spectra and are stable. The spectrum is made up of line spectrums, which are parallel lines that are isolated. Thomson’s model is electrostatically unstable, while Rutherford’s is electromagnetically unstable.

### Postulates of Bohr Atomic Model

The foundations of quantum mechanics were laid by Niels Bohr, and they are as follows:

• The electrons in the hydrogen atom spin around in stable orbits, generating no radiant energy.
• The angular momentum in stationary orbits is a multiple of the equation h / 2,π and L = n h / 2π, where n is known as the quantum number.
• The electron changes from a non-radiating orbit to a lower-energy orbit. When this occurs, a photon with the same energy as the difference between the final and beginning states is emitted. hv= Ei Ef  is used to calculate frequency (v).

### 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.

The atomic spectra of atoms are the collection of all these unique wavelengths of the atom in a certain set of variables such as pressure, temperature, and so on. Emission spectra, absorption spectra, and continuous spectra are the three forms of atomic spectra.

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

The study of the electromagnetic radiation received or emitted by atoms is known as atomic spectroscopy. There are three different forms of atomic spectroscopy:

1. The transfer of energy from the ground state to an excited state is the subject of atomic emission spectroscopy. Atomic emission can explain the electronic transition.
2. Atomic absorption spectroscopy: For absorption to occur, the lower and higher energy levels must have equivalent energy differences. The notion that free electrons created in an atomizer can absorb radiation at a given frequency is used in the atomic absorption spectroscopy principle. The absorption of ground-state atoms in the gaseous state is measured.
3. Atomic fluorescence spectroscopy combines atomic emission and atomic absorption since it uses both excitation and de-excitation radiation.

Uses of Atomic Spectroscopy are:

• 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 collection of wavelengths arranged in a logical order. Light, or any electromagnetic radiation released by energised atoms, has this property.

Because the hydrogen atom is the most basic atomic system found in nature, it produces the most basic series. When a slit allows a beam of light or other radiation to enter the device, each component of the light or radiation forms an image of the source. When resolved under the spectroscope, these images can be seen.

The photos will be in the shape of parallel lines with consistent spacing positioned next to each other. When moving from a higher to a lower wavelength side, the lines will be farther apart in the higher wavelength side and eventually closed. The shortest wavelength has the fewest separated spectral lines, which is referred to as the series limit.

### Line spectrum of the hydrogen atom

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

1. Pfund Series
2. Brackett Series
3. Paschen Series
4. Balmer Series
5. Lyman Series

### Spectral Series are Formed

Bohr’s atomic model models and well explains the set of energy levels/states that each atom encloses. Quantum numbers (n=1, 2, 3, 4, 5, 6,…..) are used to name energy states. A photon of energy nh – nl is released when electrons jump from higher energy states (nh) to lower energy ones (nl). Because the energy-related to each state is fixed, the difference between them is also fixed, resulting in a transition between similar energy states producing the same energy photon.

The electron transition to a lower energy state divides the spectral series into equivalent series. Within the series, the Greek alphabets are utilised to separate the spectral lines of corresponding energy. Hydrogen has the following spectral series:

• Lyman series (nl=1)

The series was discovered by Theodore Lyman between 1906 and 1914. As a result, it bears his name. When electrons transition from higher energy states (nh=2, 3, 4, 5, 6 ,…) to nl=1 energy state, according to Bohr’s model, the Lyman series appears. The Lyman series’ wavelengths are all in the Ultraviolet band.

For a list of wavelengths related to spectral lines, see the table below:

• Balmer series (nl=2)

Johann Balmer was the first to discover the series in 1885. As a result, the series is named after him. The Balmer series emerges when electrons go from higher energy levels (nh=3,4,5,6,7,…) to a lower energy state (nl=2). The wavelengths of the Balmer series are all visible in the electromagnetic spectrum (400 nm to 740 nm). The H-Alpha line of the Balmer series, which is also a part of the solar spectrum, is used in astronomy to identify hydrogen.

See the table below for a list of wavelengths associated with spectral lines.

• Paschen series (nl=3)

In 1908, a German physicist named Friedrich Paschen was the first to notice the series. As a result, the series is named after him. The Paschen series develops when electrons migrate from higher energy levels (nh=4, 5, 6, 7, 8, …) to lower energy states (nl=3). All of the wavelengths in the Paschen series are in the infrared portion of the electromagnetic spectrum. The Brackett series, which has the smallest wavelength, overlaps with the Paschen series. This series overlaps with all subsequent ones.

See the table below for a list of wavelengths associated with spectral lines.

Brackett series (nl=4)

In the year 1922, an American physicist named Friedrich Sumner Brackett spotted the series for the first time. As a result, the series is named after him. The Brackett series develops when electrons move from higher energy levels (nh=5, 6, 7, 8, 9 …) to lower energy states (nl=4). The wavelengths of the Brackett series are all in the infrared region of the electromagnetic spectrum.

See the table below for a list of wavelengths associated with spectral lines.

• Pfund series (nl=5)

In 1924, August Harman Pfund became aware of the series for the first time. As a result, the series is named after him. The Pfund series emerges when an electron transitions from a higher energy state (nh=6, 7, 8, 9,10, …) to a lower energy level (nl=5). The wavelengths of the Pfund series are all in the infrared region of the electromagnetic spectrum.

See the table below for a list of wavelengths associated with spectral lines.

• Humphreys series (nl=6)

In 1953, an American physicist called Curtis J Humphreys spotted the series for the first time, and the series is named after him. The Humphreys series develops when electrons migrate from higher energy levels (nh=7, 8, 9, 10, 11…) to a lower energy state (nl=6). The wavelengths of the Humphreys series are all in the infrared region of the electromagnetic spectrum.

For a list of wavelengths linked with spectral lines, see the table below.

### Sample Questions

Question 1: What exactly are atomic spectra?

The spectrum of electromagnetic radiation emitted or absorbed by an electron as it transitions between different energy levels within an atom is known as atomic spectra.

Question 2: In physics and chemistry, what does the term “spectrum” mean?

The meaning of spectrum is the same in physics as it is in chemistry. When white light is allowed to flow through a prism, it produces a band of colours on a screen.

Question 3: What do atoms consist of?

Atoms have an equal number of negative and positive charges. Atoms were described as a spherical cloud of positive charges with embedded electrons in Thomson’s concept.

Question 4: What Is the Number of Spectral Lines?

When electrons shift from higher energy levels to lower energy levels, spectral lines appear. The following are the two types of spectral lines:

1. Emission lines: Emission lines are a type of spectral line that can appear in a variety of colours and have a black background. Only when the particles emit the wavelength can these lines be seen.
2. Absorption lines are a type of spectral line that can be classified in two ways. These could take the form of dark coloured bands on a black background. When the particles absorb the wavelengths, these lines appear.

Question 5: In the spectrum, how many spectral lines may be seen?

When moving from higher to lower energy levels, the spiritual lines can be seen. The concept of numerous spectral lines has been generalised after several research. The elements of the fourth energy level migrate to the third level, and then two second-level elements move to the first level.

Question 6: Mention different series of the spectrum and where the lines fall on the spectrum.