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Spectroscopy - A More Detailed Exploration

Spectral fingerprints

Just as the fingers on our hands have a unique set of fingerprints, each chemical element has a unique emission fingerprint.

An “emission” spectrum is a plot of the amount of light emitted vs the wavelength (which is directly related to the perceived colour).

Atoms emit light at certain discrete wavelengths. The reason for the appearance of discrete wavelengths is that each atom has characteristic energy levels, and it is the jumps, or “transitions” between these levels that give rise to the emission of light.

When we watch a fireworks display, we are seeing the characteristic colours emitted by the various elements that are contained in the salts that are used in making the fireworks.

The colours that our eyes see are actually our brain’s response to the visual signals that our eyes produce when observing light of different wavelengths.

Spectral identification (i)

The presence of many of the elements well known to us on earth can also be inferred on extremely distant celestial bodies such as stars and nebulae by observing their spectra.

The image at right shows a patch of light that is emitted from the constellation Orion. The glowing gas clouds, or nebulae that are seen in this image confirm the presence of excited hydrogen atoms. The spectrum of this light shows peaks, or emission "lines" at exactly the same wavelengths as are observed in a hydrogen discharge lamp in a lab on earth.
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Gas Discharge Lamps

The spectrum of atomic hydrogen can be observed using a discharge tube like the one shown opposite.

The discharge tube has a glass envelope with an electrode at each end. It is filled with a low pressure of H2 gas. A high voltage is applied to the electrodes and collisions of energetic electrons with the hydrogen produces excited hydrogen atoms, H*. These emit a characteristic spectrum, the "Balmer series" lines shown below.

The pink colour of the discharge represents our eyes response to the various wavelengths/intensities that are emitted by the lamp. Note the similarity of the lamp’s colour to that of the Orion nebula previously. An electronically recorded emission spectrum of hydrogen can be found here.
A nice website here has downloadable .jpg images of the emission spectra of many of the chemical elements.
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Emission Spectrum of Iron (Fe)

Spectral identification (ii)

The auroral emission that is commonly seen in the earth's polar regions is due to the interaction of energetic electrons from the solar wind with oxygen atoms, approximately 100 km above the earth's surface.

The green light has a characteristic wavelength of 557.7 nm. Red emission is also sometimes seen in auroral displays - again from excited oxygen.

Flame Tests

If some table salt is placed into a bunsen burner flame we observe a very characteristic yellow emission (589 nm). Salt (NaCl) contains sodium – atoms of this element become excited in the flame and these excited atoms emit yellow light. Visual flame tests like this are commonly used to confirm the presence of particular elements in a sample.

What would happen if we were to shine white light through a vapour of gaseous atomic sodium and then disperse the light with a prism? If we looked at the spectrum we would see a dark “line” at 589 nm, because ground state Na atoms absorb strongly at the same wavelength. Consequently, this wavelength will be missing in the spectrum.

Absorption is the opposite process to emission – if a sample absorbs radiation then the intensity of the radiation from another (usually white light) source becomes depleted as it passes through the sample.

Spectrum of Sunlight

The spectrum of sunlight combines the concepts we have just been discussing - it is an example of both emission and absorption spectra.

The continuous coloured background spanning the visible colours through from violet to red is the visible portion of the Sun’s emission spectrum.

Superimposed on this "continuum" are the dark Fraunhofer lines that are due to absorption of the Sun’s radiation by elements in its photosphere. Every single one of these dark lines tells us something about the elements that are present in the sun's photosphere.

For example, prominent dark patches in the yellow part of the spectrum are due to absorption of radiation by sodium atoms in the Sun’s atmosphere.

Image Source :

The Electromagnetic Spectrum

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The diagram above is the electromagnetic spectrum, showing the wavelengths and corresponding frequencies of radiation in the various spectral regions. The tiny portion that we see - visible radiation - extends from 400 nm (violet) up to 700 nm (red). Visible light is just one form of radiation - spanning just a small region of the entire electromagnetic spectrum. Other regions of this spectrum are important for the transmission of radio and TV signals, mobile phone communications, the operation of microwave ovens, causing sunburns and ozone depletion, and for use in the X-ray machines essential in many medical and dental procedures.

Colour and Absorption

Our eyes contain rods and cones that respond to visible light spanning wavelengths from 700 nm down to 400 nm. When “white light” passes through a prism, the beam becomes dispersed into a continuous spectrum, with the colours red, orange, yellow, green, blue, indigo and violet all being visible. Passing through this series of colours the wavelength successively gets shorter until at ~400 nm the light is no longer visible – we enter the near-ultraviolet region. Our eyes are most sensitive to yellowish-green light at about 560 nm.

Molecules are collections of bonded atoms and these have much more complicated energy levels than atoms. Molecular spectra are therefore much broader than atomic spectra and instead of seeing narrow spectral lines these spectra show “bands” in both absorption and emission. A “band” is a fairly broad region of wavelengths over which the absorption or emission process takes place.

Many molecules have a single broad absorption band in the visible region. In this case it is possible to use the “colour wheel” opposite to predict the colour (and hence the range of wavelengths) that the sample actually absorbs. To do this one simply locates the perceived colour of the sample on the wheel and traverses from there across the centre of the wheel to locate the wavelengths the sample absorbs most strongly.

Why is a leaf green ?

The perceived colour of an absorbing sample is said to be "complementary" to the wavelengths that the sample actually absorbs.

Suppose we have two bottles, one containing a red jelly and the other a green jelly. If we take a red laser pointer and shine the laser into the two samples the beam will pass through the red sample but not through the green one.

Why? – the red sample appears red because it absorbs at wavelengths other than red – so the laser beam passes straight through. The green sample appears green because this is the wavelength that it does not absorb – red light however will not pass through this sample.

The same concepts explain why leaves appear green. Leaves contain chlorophyll, an important pigment in the process of photosynthesis. The absorption spectrum of chlorophyll shows two strong peaks (bands) in the visible region, as seen in the spectra below.

Absorption Spectrum of Chlorophyll

The traces at right show the wavelength dependence of the light absorbed by solutions of two similar chlorophyll molecules. Here, a high absorption means that little light passes through, while a low reading means that most of the incident light is transmitted.

The regions of strongest absorption by the chlorophyll in a leaf (and those which it uses to harvest energy from sunlight) are in the blue and red parts of the spectrum while we see that there is only a very weak absorption in the green (500-550 nm) – hence the leaf’s colour.

Measuring an Absorption Spectrum

To record a sample’s absorption spectrum, chemists use an instrument called a spectrophotometer. This instrument contains a dispersive device (here a prism) to separate a broadband source (e.g. white light) into its component colours. A slit is placed before the sample to ensure that just a single wavelength passes through the sample for each measurement. The instrument contains a means to rotate the dispersive device, thereby allowing the incident wavelength to be changed.

To measure an absorption spectrum we first place a "blank" in the sample cuvette and measure intensity as a function of wavelength I0(λ). Next, we repeat the measurement, now with the sample present. A new set of intensities is obtained, I(λ), whose values are generally lower than those measured previously because at some wavelengths the light source intensity is reduced by absorption.

The absorption spectrum of the sample is found by computing

A(λ) = log10(I0(λ)/I(λ)

The above definition of absorbance is in terms of physical quantities that are measured in the experiment just described. There is another important chemical definition of absorbance known as Beer’s law, which states that

A(λ) = ε(λ)cl

This relationship states that the absorbance of any sample is wavelength-dependent and increases in proportion to the thickness of the sample (l) and its concentration (c). The constant of proportionality is called the sample’s molar extinction coefficient (ε) – this too is wavelength dependent.

The normal units used in the Beer’s law expression are ε (M-1cm-1), c (M) and l (cm). Absorbance itself is a dimensionless quantity.


Colorimetry is an instrumental technique that can accurately measure the concentrations of a wide range of analyte species, including cations, anions and neutrals. If the sample is coloured, the method can be applied directly, but in most cases it is necessary to first convert the analyte into a coloured species by performing some chemistry.

In order to maximize the sensitivity of this method, the wavelength of the measurement should be as close as possible to the absorption maximum for the coloured form of the analyte. At this wavelength the slope of the absorbance versus concentration plot is zero since ε(λ) is a maximum.

To determine the concentration of an unknown, it is necessary to measure a calibration curve of absorbance versus concentration (a so-called Beer’s law plot). Ideally 4-5 such measurements should be recorded. The concentration of an unknown is then determined by interpolation, using its measured absorbance.

Certain conditions must be met in order to observe a linear Beer’s law plot. As long as the light used to carry out the measurement is monochromatic or nearly so, the plot should be linear. Curvature in the plot can be expected if the source emits a much broader spread of wavelengths than the absorption spectrum of the analyte species. This will typically be the case, for example, if a light-emitting diode (LED) is used for the light source. Further discussion of this more subtle point can be found elsewhere.

A Summary of the Important Concepts…

• The electromagnetic spectrum has several regions; the visible region represents just a very small slice of the complete spectrum.
• The visible region covers the wavelength range from 400 nm (violet) to 700 nm (red).
• Emission spectra provide experimental evidence for the existence of quantized energy levels in atoms.
• A spectrum is a trace of light intensity versus wavelength and provides a unique fingerprint of an atom or molecule.
• Absorption and emission are upward/downward transitions between two energy levels.
• All spectroscopic absorption experiments requires a source of electromagnetic radiation, a sample and a detector.
• A coloured sample has absorption band(s) in the visible region.
• The colour of a sample is complementary to the wavelengths of light that are absorbed.
• Colorimetry is an essentially single wavelength technique to determine the concentration of a coloured substance.