Theemission spectrum of achemical element orchemical compound is thespectrum offrequencies ofelectromagnetic radiation emitted due toelectrons making atransition from a high energy state to a lower energy state. Thephoton energy of the emittedphotons is equal to the energy difference between the two states. There are many possible electron transitions for each atom, and each transition has a specific energy difference. This collection of different transitions, leading to different radiatedwavelengths, make up an emission spectrum. Each element's emission spectrum is unique. Therefore,spectroscopy can be used to identify elements in matter of unknown composition. Similarly, the emission spectra of molecules can be used in chemical analysis of substances.
Inphysics, emission is the process by which a higher energy quantum mechanical state of a particle becomes converted to a lower one through the emission of a photon, resulting in the production oflight. The frequency of light emitted is a function of the energy of the transition.
Since energy must be conserved, the energy difference between the two states equals the energy carried off by the photon. The energy states of the transitions can lead to emissions over a very large range of frequencies. For example,visible light is emitted by the coupling of electronic states in atoms and molecules (then the phenomenon is calledfluorescence orphosphorescence). On the other hand, nuclear shell transitions can emit high energygamma rays, while nuclear spin transitions emit low energyradio waves.
Theemittance of an object quantifies how much light is emitted by it. This may be related to other properties of the object through theStefan–Boltzmann law.For most substances, the amount of emission varies with thetemperature and the spectroscopic composition of the object, leading to the appearance ofcolor temperature andemission lines. Precise measurements at many wavelengths allow the identification of a substance viaemission spectroscopy.
Emission of radiation is typically described using semi-classical quantum mechanics: the particle's energy levels and spacings are determined fromquantum mechanics, and light is treated as an oscillating electric field that can drive a transition if it is in resonance with the system's natural frequency. The quantum mechanics problem is treated using time-dependentperturbation theory and leads to the general result known asFermi's golden rule. The description has been superseded byquantum electrodynamics, although the semi-classical version continues to be more useful in most practical computations.
When theelectrons in the atom are excited, for example by being heated, the additionalenergy pushes the electrons to higher energy orbitals. When the electrons fall back down and leave the excited state, energy is re-emitted in the form of a photon. The wavelength (or equivalently, frequency) of the photon is determined by the difference in energy between the two states. These emitted photons form the element's spectrum.
The fact that only certain colors appear in an element's atomic emission spectrum means that only certain frequencies of light are emitted. Each of these frequencies are related to energy by the formula:where is the energy of the photon, is itsfrequency, and is thePlanck constant.This concludes that onlyphotons with specific energies are emitted by the atom. The principle of the atomic emission spectrum explains the varied colors inneon signs, as well as chemicalflame test results (described below).
The frequencies of light that an atom can emit are dependent on states the electrons can be in. When excited, an electron moves to a higher energy level or orbital. When the electron falls back to its ground level the light is emitted.
The above picture shows the visible lightemission spectrum for hydrogen. If only a single atom of hydrogen were present, then only a single wavelength would be observed at a given instant. Several of the possible emissions are observed because the sample contains many hydrogen atoms that are in different initial energy states and reach different final energy states. These different combinations lead to simultaneous emissions at different wavelengths.
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As well as the electronic transitions discussed above, the energy of a molecule can also change viarotational,vibrational, andvibronic (combined vibrational and electronic) transitions. These energy transitions often lead to closely spaced groups of many differentspectral lines, known asspectral bands. Unresolved band spectra may appear as a spectral continuum.
Light consists of electromagnetic radiation of different wavelengths. Therefore, when the elements or their compounds are heated either on a flame or by an electric arc they emit energy in the form of light. Analysis of this light, with the help of aspectroscope gives us a discontinuous spectrum. A spectroscope or a spectrometer is an instrument which is used for separating the components of light, which have different wavelengths. The spectrum appears in a series of lines called the line spectrum. This line spectrum is called an atomic spectrum when it originates from an atom in elemental form. Each element has a different atomic spectrum. The production of line spectra by the atoms of an element indicate that an atom can radiate only a certain amount of energy. This leads to the conclusion that bound electrons cannot have just any amount of energy but only a certain, specific amount of energy.
The emission spectrum can be used to determine the composition of a material, since it is different for eachelement of theperiodic table. One example isastronomical spectroscopy: identifying the composition ofstars by analysing the received light.The emission spectrum characteristics of some elements are plainly visible to the naked eye when these elements are heated. For example, when platinum wire is dipped into a sodium nitrate solution and then inserted into a flame, the sodium atoms emit an amber yellow color. Similarly, when indium is inserted into a flame, the flame becomes blue. These definite characteristics allow elements to be identified by their atomic emission spectrum. Not all emitted lights are perceptible to the naked eye, as the spectrum also includes ultraviolet rays and infrared radiation.An emission spectrum is formed when an excited gas is viewed directly through a spectroscope.
Emission spectroscopy is aspectroscopic technique which examines the wavelengths of photons emitted by atoms or molecules during their transition from anexcited state to a lower energy state. Each element emits a characteristic set of discrete wavelengths according to itselectronic structure, and by observing these wavelengths the elemental composition of the sample can be determined. Emission spectroscopy developed in the late 19th century and efforts in theoretical explanation of atomic emission spectra eventually led to quantum mechanics.
There are many ways in which atoms can be brought to an excited state. Interaction with electromagnetic radiation is used influorescence spectroscopy, protons or other heavier particles inparticle-induced X-ray emission and electrons or X-ray photons inenergy-dispersive X-ray spectroscopy orX-ray fluorescence. The simplest method is to heat the sample to a high temperature, after which the excitations are produced by collisions between the sample atoms. This method is used inflame emission spectroscopy, and it was also the method used byAnders Jonas Ångström when he discovered the phenomenon of discrete emission lines in the 1850s.[1]
Although the emission lines are caused by a transition between quantized energy states and may at first look very sharp, they do have a finite width, i.e. they are composed of more than one wavelength of light. Thisspectral line broadening has many different causes.
Emission spectroscopy is often referred to as optical emission spectroscopy because of the light nature of what is being emitted.
In 1756 Thomas Melvill observed the emission of distinct patterns of colour whensalts were added toalcohol flames.[2] By 1785James Gregory discovered the principles of diffraction grating and American astronomerDavid Rittenhouse made the first engineereddiffraction grating.[3][4] In 1821Joseph von Fraunhofer solidified this significant experimental leap of replacing a prism as the source of wavelengthdispersion improving thespectral resolution and allowing for the dispersed wavelengths to be quantified.[citation needed]
In 1835,Charles Wheatstone reported that different metals could be distinguished by bright lines in the emission spectra of theirsparks, thereby introducing an alternative to flame spectroscopy.[5][6]In 1849,J. B. L. Foucault experimentally demonstrated thatabsorption and emission lines at the same wavelength are both due to the same material, with the difference between the two originating from the temperature of the light source.[7][8]In 1853, theSwedish physicistAnders Jonas Ångström presented observations and theories about gas spectra.[9] Ångström postulated that an incandescent gas emits luminous rays of the same wavelength as those it can absorb. At the same timeGeorge Stokes andWilliam Thomson (Kelvin) were discussing similar postulates.[7] Ångström also measured the emission spectrum from hydrogen later labeled theBalmer lines.[10][11]In 1854 and 1855,David Alter published observations on the spectra of metals and gases, including an independent observation of theBalmer lines of hydrogen.[12][13]
By 1859,Gustav Kirchhoff andRobert Bunsen noticed that severalFraunhofer lines (lines in the solar spectrum) coincide with characteristic emission lines identified in the spectra of heated elements.[14][15] It was correctly deduced that dark lines in the solar spectrum are caused by absorption by chemical elements in thesolar atmosphere.[16]
The solution containing the relevant substance to be analysed is drawn into the burner and dispersed into the flame as a fine spray. The solvent evaporates first, leaving finely divided solid particles which move to the hottest region of the flame where gaseousatoms andions are produced through the dissociation of molecules. Hereelectrons are excited as described above, and the spontaneously emit photon to decay to lower energy states. It is common for amonochromator to be used to allow for easy detection.
On a simple level, flame emission spectroscopy can be observed using just a flame and samples of metal salts. This method of qualitative analysis is called aflame test. For example,sodium salts placed in the flame will glow yellow from sodium ions, whilestrontium (used in road flares) ions color it red.Copper wire will create a green colored flame, however in the presence ofchloride gives blue (molecular contribution by CuCl).[17][18]
Emission coefficient is a coefficient in the power output per unit time of anelectromagnetic source, a calculated value inphysics. The emission coefficient of a gas varies with thewavelength of the light. It has unit m⋅s−3⋅sr−1.[19] It is also used as a measure ofenvironmental emissions (by mass) per MW⋅h ofelectricity generated, see:Emission factor.
InThomson scattering a charged particle emits radiation under incident light. The particle may be an ordinary atomic electron, so emission coefficients have practical applications.
IfXdVdΩdλ is the energy scattered by a volume elementdV into solid angledΩ between wavelengthsλ andλ +dλ per unit time then the emissioncoefficient isX.
The values ofX in Thomson scattering can bepredicted from incident flux, the density of the charged particles and their Thomson differential cross section (area/solid angle).
A warm body emitting photons has amonochromatic emission coefficient relating to its temperature and total power radiation. This is sometimes called the secondEinstein coefficient, and can be deduced fromquantum mechanical theory.
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