Atomic Spectra

Core Concepts

In this tutorial, you will learn about atomic spectra, including the meaning behind this term, why spectra appear the way they do, and applications of this unique phenomena.

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What is an Atomic Spectra?

Atomic spectra or atomic emission spectra are the unique wavelengths of light which each element emits due to an electron being relaxed from a higher energy state to a lower energy state in an atom. Every element on the periodic table contains a different number of electrons in its neutral state. Hence, each element has different gaps in energy between its highest occupied electronic energy level and subsequent higher energy levels in unoccupied orbitals. Once this highest energy valence electron is excited to a higher electronic energy level (excited state) via a photon, its relaxation back to ground state results in a photon being emitted. These photon’s wavelengths (which correspond to a specific energy) are what comprise an atomic spectra. You can view the characteristic atomic spectra for multiple elements from Boston University here.

Only the highest energy electron can be excited to an excited state, due to its intrinsic instability. Therefore, atomic spectra are the result of energy changes from valence electrons exclusively.

Quantization of Energy States

The reason only small, discrete, wavelengths of light are emitted upon relaxation of an electron is due to the inherent quantization of energy levels within electronic structure. Quantized energy levels means that each energy level an electron can be excited to is a discrete value. Just as a ladder contains multiple rungs, and it is impossible to have a height between two rungs when climbing, electrons can only be stable in higher atomic orbital energies or the ground state, but never in between. Excited electronic levels include higher order s, p, d, and f orbitals which the highest energy valence electron can be excited to fill. These all serve as metaphorical “steps” which the electron may briefly exist in, before it eventually falls down to the more stable ground state, releasing a singular photon in the process.

Visual Representation of Electronic Energy Levels

Upon excitation with a photon, the electron could jump to the next unoccupied orbital, or some which are even higher in energy. This is completely dependent upon the energy of the photon which meets the electron. The multiple lines on the spectrum demonstrate the various difference in energies between the ground electronic state and higher energy excited states. However, presented with a continuous spectrum of polychromatic light(light with most of the visible spectra wavelengths combined), a sample’s spectrum may have all these wavelengths combined.

The atomic emission spectrum of Hydrogen

Atomic Absorption Spectra

Just as every element has a unique emission spectrum, similarly each element also has a unique absorption spectrum. This is a measurement of what energies (or wavelengths) the atom absorbs based on different electronic energy levels.

Hydrogen Absorption Spectrum

The technique of atomioc absorption spectroscopy (AAS) uses this property to identify elements in samples. AAS is commonly used to analyze food and pharmaceutical samples.

Equation for Photon Energy

The difference between the electron’s ground state energy and the energy of the excited state is equivalent to the energy of the emitted photon. The equation below calculates the energy.

$\begin{gather*} { E=hc/\lambda } \end{gather*}$

$\begin{gather*} \par E=Energy \\ {h=Planck's} {constant} \\ c=Speed of Light \\ \lambda=Wavelength \end{gather*}$

Generally, photon energy is expressed in electron-volt(eV), which requires you to divide the energy in Joules by the elementary charge of an electron(1.6E-19).

Applications of Atomic Spectra

Atomic spectra are important for a variety of analytical techniques. Techniques including atomic emission spectroscopy(AES) and atomic absorption spectroscopy(AAS) rely on this phenomena specifically. Sample characterization or quantitative analysis can be performed using these techniques, even with extremely small sample volumes. For example, some elements can be detected in parts per billion(ppb) or even parts per trillion(ppt). This characteristic makes these instruments especially useful in environmental, biological, industrial and forensic applications.

Finally, atomic emission or absorption spectroscopy should not be confused with other techniques involving photon-electron interactions. Using non-ionizing radiation which is powerful enough to excite electrons and not eject electrons is critical to seeing atomic spectra. Non-ionizing radiation is considered radiation with a wavelength longer than that of UV waves. UV, X-ray and gamma rays are too high in energy to excite electrons, therefore samples exposed to these waves experience ionization. Sources of non-ionizing radiation include plasma, flames, and sparks.

Practice Problems

Problem 1

Upon bombardment with photons, a sodium atom in a bulb emits a photon with wavelength 580 nm. What is the energy between the ground state and excited state of the electron within the atom?

Problem 2

True or False: The energy of photons in an electromagnetic spectrum is continuous. Therefore, the emitted wavelengths of light from an atom will be continuous when exposed to a broad spectrum of polychromatic light.

Answers

Problem 1

$\begin{gather*} { E=\frac{hc}{\lambda }} \\ {E=\frac{(6.626 \times 10^{-34} J*s)(2.99 \times 10^8 m/s)}{(580 \times 10^{-9} m)}} \\ {E \approx 3.42 \times 10^{-19} J \\ {E \approx 2.13 eV}}\end{gather*}$

Problem 2

False, while a continuous spectrum is exposed, only discrete wavelengths will be emitted. Furthermore, these wavelengths correspond to the quantized energy level differences between electronic ground and excited states.