Atomic spectra

Introduction 

The concept of atomic spectra played a crucial role in defining Bohr’s atomic model and the behaviour of the electrons, both in the ground and the excited state, as it is one of the essential parts of molecular chemistry and physics. Considering the spectrum has become imperative for several studies further. The atomic spectrum is very different from the electromagnetic wave spectrum because the former is based on energy, and the latter considers wavelengths and frequencies. In a true sense, the concept of the elemental spectrum was coined by Gustav Kirchhoff and Robert Bunsen, two German physicists. 

What is the atomic spectrum? 

Atomic spectrum can be termed the energy pattern or signature of an atom. It helps physicists understand atomic behaviour both on the ground and excited state and the energy levels. The study of the atomic spectrum is termed spectroscopy, and it is done via a proper laboratory setup. Once the spectrum is obtained, it is compared to the hydrogen spectrum, which is the basic pattern and then calculations are done using the Rydberg constant. 

Types of atomic spectrum 

The atomic spectrum is divided into two main types- line and absorption spectrum. However, before learning about them, it’s crucial to know what ground and excited states are and how the electrons behave in the energy levels. 

Ground and excited states of an atom

When all the electrons of an atom are in the natural or zero energy levels, it is known as the ground state. Here, the negatively charged particles do not have any excess energy, and they continue to revolve with the energy of the stationary orbits. On the other hand, when the electrons get energy from an external source like heat or electric spark, they cannot stay in the dedicated orbits. And therefore, they transit to higher energy levels. This atomic state is termed an excited state.

Atomic energy levels 

According to Bohr’s atomic model, the energy levels are nothing but stationary orbits where the electrons revolve around a fixed energy amount. When it gains energy, it jumps to higher levels, becoming unstable. As a result, they instantly transit to the lower levels while emitting the exact amount of energy they have gained. It is this energy that forms lines in the atomic spectrum. 

Absorption vs line spectrum

The absorption spectrum defines the energy absorbed by the atoms to transit from the ground state to the excited state. Most of the time, this particular spectrum of different elements overlap. 

On the contrary, the line spectrum is particular to the element, which means that the sodium line spectrum cannot be compared with that of the calcium line spectrum. Therefore, it is known to be an elemental fingerprint used to define the behaviour of the electrons about the energy levels. 

The following section will draw a brief parallel study between these two absorption types.

What are the atomic spectral lines, and what do they signify?

In the emission of the line spectrum of an element, several black lines are spaced differently and also have different thicknesses. These black lines represent the amount of energy being released by an excited atom while transitioning from higher to lower energy states. After Bohr provided his theory about energy levels and the permissible orbits, it became clear how the atomic spectrum can be defined and segregated. 

Atomic spectral lines of hydrogen

From the atomic configuration hydrogen spectrum series, the five important sections were formed- 

• Lyman series determine the spectral lines emitted by electrons when they jump from En to E1 or the K shell. 
• Balmer series represents the wavelengths emitted when electrons jump from En to E2 or the L shell.
• When an electron transition is from En to E3 or the M shell, the series of spectral lines is defined as Paschen. 
• The Brackett series considers the electron emission spectra when they jump from En to E4 or the N shell.
• The PFund series represents the electron emission at the time of moving from En to E5.

Based on these lines, Bohr was able to replace the hydrogen spectrum series with the EMW spectrum. The Lyman series has the highest energy since electrons jump to the K shell. Therefore, the wavelengths emitted are within the ultraviolet region. 

The wavelengths emitted by the Balmer series falls within the visible light spectrum, while that of the Paschen and Brackett series represent the infrared radiations. 

Rydberg’s constant 

According to Balmer, the wavelengths of the spectral lines falling within the visible spectrum can be calculated based on a certain formula. Frequency is nothing but the mathematical inverse of wavelength. Therefore, the expression can be written as: 

λ = 1/f

Or, 1/ λ = f = 109678cm-1 X [(1/n2)-(1/22)]

This is known as Rydberg’s equation and is used to determine the wavelengths of the Balmer series. “n” represents the energy level from where the electron will jump to the second shell or L. The value, 109678 cm-1, is said to be a Rydberg’s constant, and it wouldn’t change, no matter what element is concerned or the stationary orbits. 

What is the importance of atomic spectroscopy? 

By studying the atomic spectrum of any element, one can know the behaviour of the atoms in both ground and excited states. Also, following the hydrogen spectrum, it becomes possible to judge whether the wavelengths being emitted falls within the light spectrum’s visible, UV, or IR sections. 

Conclusion 

The atomic spectra are one of the most important parts of physical chemistry because it helps in defining how the electrons behave when they absorb energy and jump to higher levels and again return to their destined shell upon losing energy. Moreover, by simply studying the line spectra, one can know which element it is because the number of lines, the distance between them, and the thickness of each line are unique to different elements. That’s why atomic spectroscopy is used for understanding the spectral contexts of atoms at the ground and excited levels.