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An Analysis Of Emission Spectra Environmental Sciences Essay

Paper Type: Free Essay Subject: Environmental Sciences
Wordcount: 2707 words Published: 1st Jan 2015

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Emission spectra are the radiation emitted by the atoms when their electrons jump from higher energy level to lower energy level. The emission spectrum of a chemical element or chemical compound is the relative intensity of each frequency of electromagnetic radiation emitted by the element’s atoms or the compound’s molecules when they are returned to a ground state.

The subatomic particles that comprise the atom can absorb various kinds of energy and then emit that energy as a photon of a specific energy and corresponding wavelength and frequency. This emitted energy is called an emission spectrum. Electrons in particular release electromagnetic radiation in the visible range as well as in wavelengths surrounding the visible range. The particular wavelength that an electron releases depends on the difference between its ground state energy and the energy level that it jumps to. The amount of energy required for an electron to jump to a higher energy level depends on where it is starting from (its ground state). So the specific visible wavelengths (colors) released by an atom that has absorbed energy depend on the arrangement of its electrons. All the various elements and molecules that exist have their own unique arrangement of electrons, and so the particular wavelengths (colors) produced will always be unique to any one element or molecule. This “spectrum” of specific electromagnetic waves can therefore identify the substance. Note that Bohr used discreet emission spectra to show the discreet energies possessed by electrons in atoms.

Because the electrons of different atoms so closely arranged in solid substances influence each other, the spectrum of a solid is different from that of the substance’s gas state, where the electron arrangement of individual atoms or molecules are not interfered with by neighboring atoms or molecules. Normally, therefore, substances are identified by their gas phase spectrum.

A plot of the brightness of an object versus wavelength is called a spectrum, (even called spectra), and is observed using a spectrograph. By spreading out the light by wavelength, we can gain insight into what’s happening to photons of particular wavelengths (or energies), which in turn tells us what’s happening with particular types of atoms. There are three components of a spectrum: continuum emission (or blackbody radiation), emission lines, and absorption lines.

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Continuum emission is a wide, smooth (continuous!) band of colors like a rainbow. This type of emission is caused by an opaque material which emits radiation because of its temperature. Hotter objects are brighter and bluer than cooler objects. All objects have continuum radiation. (Even you; although in your case, since it’s in the infrared, we usually call it ‘heat’.)

An absorption line is characterized by a lack of radiation at specific wavelength. Absorption lines are created by viewing a hot opaque object through a cooler, thin gas. The cool gas in front absorbs some of the continuum emission from the background source, and re-emits it in another direction, or at another frequency. Absorption lines are subtracted from the continuum emission, so that they appear fainter.

An emission line is characterized by excessive radiation at specific wavelengths. You can observe emission lines by looking through a spectrometer at an energized gas. They are created by the photons that are released by the “falling” electrons.

The important thing to know about absorption and emission lines is that every atom of a particular element (hydrogen, say) will have the same pattern of lines all the time. And the spacing of the lines is the same in both absorption and emission, only emission lines are added to the continuum, while absorption lines are subtracted.


When a sample of gaseous atoms of an element at low pressure is subjected to an input of energy, such as from an electric discharge, the atoms are themselves found to emit electromagnetic radiation.

On passing through a very thin slit and then through a prism the light (electromagnetic radiation) emitted by the excited atoms is separated into its component frequencies.

The familiar dispersion of white light is illustrated below:

Solids, liquids and dense gases glow at high temperatures. The emitted light, examined using a spectroscope, consists of a continuous band of colours as in a rainbow. A continuous spectrum is observed. This is typical of matter in which the atoms are packed closely together. Gases at low pressure behave quite differently.

The excited atoms emit only certain frequencies, and when these are placed as discreet lines along a frequency scale an atomic emission spectrum is formed.

The spectral lines in the visible region of the atomic emission spectrum of barium are shown below.

Spectral lines exist in series in the different regions (infra-red, visible and ultra-violet) of the spectrum of electromagnetic radiation.

The spectral lines in a series get closer together with increasing frequency.

Each element has its own unique atomic emission spectrum.


It was necessary to explain how electrons are situated in atoms and why atoms are stable. Much of the following discussion refers to hydrogen atoms as these contain only one proton and one electron making them convenient to study.

In the early 1913, the famous scientist Neils Bohr solved many problems in chemistry of the time by proposing his view that the electron revolves around the nucleus of the atom with a definite fixed energy in a fixed path, without emitting or absorbing energy. The electron in the hydrogen atom exists only in certain definite energy levels. These energy levels are called Principal Quantum Levels, denoted by the Principal Quantum Number, n. Principal Quantum Level n = 1 is closest to the nucleus of the atom and of lowest energy. When the electron occupies the energy level of lowest energy the atom is said to be in its ground state. An atom can have only one ground state. If the electron occupies one of the higher energy levels then the atom is in an excited state. An atom has many excited states.

When a gaseous hydrogen atom in its ground state is excited by an input of energy, its electron is ‘promoted’ from the lowest energy level to one of higher energy. The atom does not remain excited but re-emits energy as electromagnetic radiation. This is as a result of an electron ‘falling’ from a higher energy level to one of lower energy. This electron transition results in the release of a photon from the atom of an amount of energy (E = h®) equal to the difference in energy of the electronic energy levels involved in the transition. In a sample of gaseous hydrogen where there are many trillions of atoms all of the possible electron transitions from higher to lower energy levels will take place many times. A prism can now be used to separate the emitted electromagnetic radiation into its component frequencies (wavelengths or energies). These are then represented as spectral lines along an increasing frequency scale to form an atomic emission spectrum.

Principal Quantum Levels (n)

for the hydrogen atom.


A hydrogen atom in its Ground State.

The electron occupies the lowest possible energy level which in the case of hydrogen is the Principal Quantum Level n = 1.

The Bohr Theory was a marvelous success in explaining the spectrum of the hydrogen atom. He calculated wavelengths agreed perfectly with the experimentally measured wavelengths of the spectral lines. Bohr knew that he was on to something; matching theory with experimental data is successful science. More recent theories about the electronic structure of atoms have refined these ideas, but Bohr’s ‘model’ is still very helpful to us.

For clarity, it is normal to consider electron transitions from higher energy levels to the same Principal Quantum Level. The image given below illustrates the formation of spectral lines in visible region of the spectrum of electromagnetic radiation for hydrogen, called the Balmer Series.

The Spectral Lines are in Series…

As referred to above for hydrogen atoms, electron transitions form higher energy levels all to the n = 2 level produce a series of lines in the visible region of the electromagnetic spectrum, called the Balmer Series. The series of lines in the ultra-violet region, called the Lyman Series, are due to electron transitions from higher energy levels all to the n = 1 level, and these were discovered after Bohr predicted their existence.

Within each series, the spectral lines get closer together with increasing frequency. This suggests that the electronic energy levels get closer the more distant they become from the nucleus of the atom.

No two elements have the same atomic emission spectrum; the atomic emission spectrum of an element is like a fingerprint.

The diagram to the right illustrates the formation of three series of spectral lines in the atomic emission spectrum of hydrogen.


As we know light from a mercury discharge tube was composed of only three colors, or three distinct wavelengths of light. This feature, that an element emits light of specific colors, is an enormously useful probe of how individual atoms of that element behave. Indeed, the science of spectroscopy was developed around the discovery that each element of the periodic table emits light with its own set characteristic wavelengths, or “emission spectrum.” of light. If one has a collection of several elements, all emitting light, and the spectra of the different elements combine or overlap. By comparing the combined spectra to the known spectra of individual elements, we can discover which elements are present. It is amusing to note that the element helium was first discovered in this manner through the spectroscopic analysis of light from the sun in 1868 and was only later discovered in terrestrial minerals in 1895.

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But why do we see distinct wavelengths in emission spectra? And why are the spectra different for particular elements? There is nothing distinct about the light from an incandescent source such as the ordinary light bulb. In an empirical study of the spectrum of hydrogen, Balmer discovered that the precise frequencies and wavelengths of the light produced could be described by a simple equation involving a constant and an integer. Balmer’s equation was then expanded to describe the entire spectrum of hydrogen, including the ultra-violet and the infrared spectral lines. This equation is called the Rydberg equation:

= R (€­ ),

Where R is the “Rydberg” constant, and n1 and n2 are integers.

The presence of integers in this equation created a real problem for physicists until the development of the quantum theory of the atom by Neils Bohr. Bohr’s theory suggested that the electron orbiting the nucleus could have only certain quantized angular momenta. The implication of this idea is that the electron can orbit only at certain fixed distances and velocities around the nucleus and subsequently can possess only certain discrete energies. Individual electron orbits are associated with specific energy levels. Integer numbers uniquely identify these levels and these integers, “quantum numbers,” are the ones that show up in the Rydberg equation and that are labeled

n1 and n2.

The integers in Rydberg’s equations identify electron orbits of specific radius. In general, the larger the value of the integer, the larger the size of the orbit. Rydberg’s equation says that the wavelength of the light emitted from an atom depends on two electron orbits. The interpretation is that an electron makes a transition from the initial orbit identified by the integer n1 to a final orbit identified by the integer n2. Furthermore, since there is a unique energy associated with each electron orbit, these integers n1 and n2 also identify or tag the energy of the electron. Hence, a discrete amount of energy is released or absorbed when an electron makes a transition between two orbits. In the case of the atom, when an electron makes a transition from one orbit to another with a lesser value of its identifying integer, energy is released from the atom and takes the form of emitted light of a distinct wavelength, or equivalently, of distinct frequency.

So the picture we have is that electron transitions between different orbits produce different wavelengths of light and that the actual wavelength value of the light depends on the energy difference between the two orbits. Furthermore, since the energies of the different orbits and the energies of the transitions are determined by the atomic number (the number of protons in the nucleus), each atom has its own characteristic spectrum.

distances and velocities around the nucleus and subsequently can possess only certain discrete energies. Individual electron orbits are associated with specific energy levels. Integer numbers uniquely identify these levels and these integers, “quantum numbers,” are the ones that show up in the Rydberg equation and that are labeled n1 and n2.

Emission Line Spectra of Various Elements


Explanation of the above Image:

First spectrum is hydrogen, typical of a hydrogen spectrum tube.

Second spectrum is helium, typical of a helium spectrum tube.

Third spectrum is lithium, as typically from a flame or an electric arc.

Fourth spectrum is neon.

Fifth spectrum is low pressure sodium, but with secondary lines exaggerated.

Sixth spectrum is argon, typical of an argon glow lamp or spectrum tube.

Next spectrum is copper, drawn using a wavelength table and Ioannis Galidakis’ photos of a copper arc spectrum (see link below). Oxide lines which may appear in the flame spectrum are not shown.

Next spectrum is zinc, drawn using a wavelength table and a photo by Ioannis Galidakis of a zinc arc spectrum. Intensity of the red line is shown for the slightly greenish light blue usual zinc arc, but Ioannis reports getting a pinkish zinc arc and shows the red line to be brighter.

Next spectrum is barium. Oxide lines are not included.

Next spectrum is krypton. Ion lines typical of flashlamp use are not included.

Next spectrum is that of the most common variety of metal halide lamp, which is basically a mercury vapor lamp enhanced with iodides of sodium and scandium.

Next spectrum is that of a xenon flashtube of lower-than-usual pressure, operated with a higher than usual voltage and a lower than usual energy level to favor a line spectrum. An actual typical xenon spectrum generally has a strong continuous spectrum, which I show more dimly than actually occurs in order to show the lines. The lines are mainly those of excited xenon ions, rather than excited neutral xenon atoms. At lower current, the most distinct visible spectral lines are two close together in the blue and the brightness is usually low.

Next spectrum is high pressure mercury vapor, typical of a mercury vapor lamp. Low pressure mercury vapor has a similar spectrum except the green line is slightly dimmer and the yellow lines are significantly dimmer.

Next one after that is a mercury lamp with the common Deluxe White phosphor.

Next one after that is a compact fluorescent lamp of the 2700K color.

Emission line spectra of various other elements is given below


Emission Spectroscopic techniques are used in “Flame Emission Spectroscopy”

Energy spectra are used in astrophysical spectroscopy.

Energy Spectra are used in Optical Spectroscopy


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