What is Atomic Spectra?
The Atomic spectra are defined as the Spectrum of frequencies of electromagnetic radiation emitted or absorbed during transitions of Electrons between Energy levels within an Atom. Each element has a characteristic Spectrum through which it can easily be recognized. In an Atom, Electrons have discrete and some specific energies. There are more Energy states in a tom than there are Electrons. When an Electron transitions from one Energy level to another, it emits light or photon with a specific wavelength. When an Electron gets excited from one Energy level to another, it emits or absorbs light of a specific wavelength.
As an Electron moves between different Energy levels within an Atom, its Spectrum of Electromagnetic radiation is released or absorbed. An Electron emits or absorbs light of a specific wavelength as it jumps from one Energy level to the next.
The Rydberg formula clearly splits the Atomic Hydrogen emission Spectrum into a number of wavelength-dependent spectral lines. The visible spectral lines in the hydrogen emission Spectrum are caused by Atomic transitions between distinct Energy levels. Spectral series are crucial in Astronomical Spectroscopy.
Characteristics of Atomic Spectrum
The characteristics of the Atomic Spectrum are observed as:
The Atomic Spectrum should be a pure line Spectrum.
The Atomic Spectrum should be the emission band Spectrum.
The Atomic Spectrum should be an absorption line Spectrum.
The Atomic Spectrum should be the absorption band Spectrum.
Atomic Spectrum Overview
In any given set of conditions like pressure, temperature, etc., the collection of all these specific wavelengths is what constitutes the Atomic Spectrum. Hence, Atomic spectra are the spectra of Atoms. Below we will be looking at Atomic spectra more in detail along with the Rydberg formula and the spectral series of the hydrogen Atom. There are three types of Atomic spectra: emission spectra, absorption spectra, and continuous spectra.
Spectral Series
Light frequencies emitted by a specific element follow a predictable pattern. For example, because hydrogen is the most basic Atom, it has the most basic Spectrum. At first glance, spectral lines appear to lack order or regularity, but the spacing between lines within certain sets of the hydrogen Spectrum decreases on a regular basis, and each of these sets is known as a spectral series. The first spectral series, known as the Balmer series, was discovered in the visible region of the Hydrogen Spectrum by a Swedish schoolteacher named Johann Jakob Balmer.
H is the red line with the longest wavelength (656.3 nm). The wavelength of the next line in the blue-green Spectrum is 486. 1 nm. The wavelength of the violet Spectrum's third line is 434.1 nm, and so on.
The Balmer series is as follows in the hydrogen emission Spectrum:
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As the wavelength decreases, the lines become closer together and less intense.
Rydberg’s Formula
Rydberg's equation estimates the wavelength of a spectral line in a wide variety of chemical elements in Atomic physics. For all Atomic hydrogen transitions, the equation is a generalisation of the Balmer series.
It is an Energy unit defined in terms of the Electron's ground-state Energy in the Bohr model of the hydrogen atom. In cgs, "me" is Electron mass, "e" is the Electron charge, h-bar, "Z" is the Atomic number, and "n" is the Electron state's primary quantum number. The Rydberg formula makes measuring spectral lines simple.
In Atomic physics, Rydberg's formula calculates the wavelength of a spectral line in many chemical elements. The formula was primarily presented as a generalization of the Balmer series for all Atomic transitions of hydrogen. It is a unit of Energy defined in terms of the ground-state Energy of an Electron in the Bohr model for the hydrogen Atom, in cgs, where is the Electron mass, e is the charge on the Electron, is h-bar, Z is the Atomic number, and n is the principal quantum number for a given Electron state. It becomes easy to calculate the spectral lines by the Rydberg formula. Following is the formula:
\[\frac{1}{\lambda} = RZ^{2} (\frac{1}{n'^{2}} - \frac{1}{n^{2}})\]
R = Rydberg constant (1.09737*107 m-1)
λ = wavelength of light
Z = the Atomic number
n = upper Energy level
n’ = lower Energy level
The spectral series of single-Electron Atoms like hydrogen is Z = 1.
Atomic Spectroscopy
Atomic spectroscopy studies the electromagnetic radiation absorbed or emitted by the Atoms. There are three types of Atomic spectroscopy:
Atomic Emission Spectroscopy: This includes the transfer of Energy from the ground state to an excited state. The Electronic transition can be discussed in Atomic emission.
Atomic Absorption Spectroscopy: Absorption to take place, there should be an identical Energy difference between the lower and higher Energy levels. The Atomic absorption spectroscopy principle uses the fact that generating free Electrons in an Atomizer can absorb radiation at specific frequencies. It quantifies the absorption of ground-state Atoms in the gaseous state.
Atomic Fluorescence Spectroscopy: This is a combination of Atomic emission and Atomic absorption, as it involves radiation of both excitation and de-excitation as well.
Uses of Atomic Spectroscopy
It is used to identify the spectral lines of materials used in metallurgy.
It is used in pharmaceutical industries to find the traces of materials used.
It can be used to study multidimensional elements.
It is used as a tool for studying the structures of Atoms and molecules.
It provides a precise analytical method for finding the constituents in a material having unknown chemical composition.
Atomic spectroscopy is used in occupational and environmental monitoring.
Solved Examples
Question: An Electron excites an Atom to the fourth orbit, so when it jumps back to the Energy levels, a Spectrum is formed. A total number of spectra is formed. What would be the total number of spectral lines in this Spectrum?
Answer: An Electron excites in an Atom to the fourth orbit, n=4.
The total number of spectral lines in the Spectrum is,
\[\frac{n(n - 1)}{2} = \frac{4(4 - 1)}{2} = \frac{4\times 3}{2} = 6\]
Fun Facts
When Atoms get excited, they emit certain specific wavelengths that correspond to different colors. The emitted light can be noted as a series of colored lines with dark spaces in between, this colored lines series is called a line of Atomic spectra. A unique set of spectral lines is produced through each element. Rainbow is an example of a continuous Spectrum.
FAQs on Atomic Spectra
1. What is an atomic spectrum and why is it considered an element's 'fingerprint'?
An atomic spectrum is the unique pattern of electromagnetic radiation frequencies that is either emitted or absorbed by the atoms of an element. It is considered a 'fingerprint' because every element has a distinct arrangement of electron energy levels. The transitions of electrons between these unique levels produce a set of spectral lines at specific wavelengths that is different for every element, allowing for precise identification.
2. What is the fundamental principle behind the formation of atomic spectra?
The fundamental principle is the quantisation of energy levels within an atom, as described by Bohr's model. Electrons can only exist in discrete energy orbits, not in between. A spectrum is formed when:
- An electron absorbs a photon of a specific energy to jump to a higher, allowed energy level, creating an absorption line.
- An excited electron falls to a lower energy level, releasing a photon of specific energy, creating an emission line.
3. What is the main difference between an emission spectrum and an absorption spectrum?
An emission spectrum is produced by a hot, excited gas and appears as a series of bright, coloured lines on a dark background. Each line corresponds to a photon emitted by an electron falling to a lower energy state. In contrast, an absorption spectrum is formed when white light passes through a cool gas. It appears as a continuous rainbow-like spectrum with specific dark lines where atoms have absorbed photons to move electrons to higher energy states.
4. How does quantum theory explain why atomic spectra consist of discrete lines instead of a continuous band?
Quantum theory explains that energy within an atom is quantised, meaning electrons are restricted to specific, discrete energy levels. A spectral line is formed from an electron transition between two such allowed levels. The energy of the emitted or absorbed photon is precisely equal to the energy difference between these two levels (ΔE = hf). Since only specific energy transitions are possible, only photons of corresponding discrete frequencies are involved, resulting in sharp, distinct lines rather than a continuous spectrum.
5. How is the Rydberg formula used to calculate the wavelength of spectral lines in the hydrogen atom?
The Rydberg formula is a key equation in physics used to predict the wavelength of any spectral line in the hydrogen atom. The formula is 1/λ = R (1/n₁² - 1/n₂²), where λ is the wavelength, R is the Rydberg constant, and n₁ and n₂ are the principal quantum numbers of the lower and upper energy levels, respectively. By substituting the integer values for the energy levels of a specific electron transition, one can calculate the exact wavelength of the emitted or absorbed photon.
6. What are some important real-world applications of atomic spectroscopy?
Atomic spectroscopy is a powerful analytical tool with several important applications, including:
- Astronomy: Identifying the chemical composition of distant stars and galaxies by analysing the light they emit.
- Forensic Science: Detecting trace amounts of specific elements, like lead or arsenic, in evidence samples.
- Environmental Monitoring: Measuring pollutants and heavy metals in water, soil, and air.
- Pharmaceutical Industry: Ensuring the purity of drugs by checking for trace contaminants.
7. Which spectral series of hydrogen is visible to the human eye, and why are the others invisible?
The Balmer series is the spectral series for hydrogen that falls within the visible light portion of the electromagnetic spectrum. This occurs because the electron transitions in this series all end at the second energy level (n=2), and the energy differences correspond to wavelengths our eyes can detect. Other series, like the Lyman (ending at n=1) and Paschen (ending at n=3), involve larger or smaller energy transitions, respectively, resulting in photons in the ultraviolet and infrared regions, which are invisible to the human eye.
