Ch. 17: Fundamentals of Spectrophotometry

Outline:

• 17-1 Properties of Light

• 17-2 Absorption of Light

• 17-3 Measuring Absorbance

• 17-4 Beer’s Law in Chemical Analysis

• 17-5 Spectrophotometric Titrations

• 17-6 What Happens When a Molecule Absorbs Light?

• 17-7 Luminescence

Updated Oct. 15, 2013, new lecture, slides 22-42

SpectrophotometrySpectrophotometry is any technique that uses light to measure chemical concentrations. A procedure based on absorption of visible light is called colorimetry. The most cited article in the journal Analytical Chemistry from 1945 to 1999 describes a colorimetric method by which biochemists measure sugars.

M. Dubois, K. A. Gilles, J. K. Hamilton, P. A. Rebers, and F. Smith, “Colorimetric Method for Determination of Sugars and Related Substances,” Anal. Chem. 1956, 28, 350. (J. Riordon, E. Zubritsky, and A. Newman, “Top 10 Articles,” Anal. Chem. 2000, 72, 324A.)

The Ozone HoleOzone, formed at altitudes of 20 to 40 km by the action of solar ultraviolet radiation (hν) on O2, absorbs ultraviolet radiation that causes sunburn and skin cancer.

O2 2O· O· + O2 O3

hν

OzoneIn 1985, the British Antarctic Survey reported that the total ozone in the Antarctic stratosphere had decreased by 50% in early spring, relative to levels observed in the preceding 20 years. Ground, airborne, and satellite observations have since shown that this “ozone hole” occurs only in early spring and continued to deepen until the year 2000.

The culprits were shown to be CFCs like Freon-12 (CCl2F2), which were used in refrigerators and air conditioners. These long-lived man-made compounds diffuse into the stratosphere, and catalyze the decomposition of ozone.

Cl produced in step 4 reacts again in step 2, so a single Cl atom can destroy > 105 molecules of O3. The chain is terminated when Cl or ClO reacts with hydrocarbons or NO2 to form HCl or ClONO2.

The Ozone Hole, 2Stratospheric clouds formed during the Antarctic winter catalyze the reaction of HCl with ClONO2 to form Cl2, which is split by sunlight into Cl atoms to initiate O3 destruction:

Polar stratospheric clouds require winter cold to form. Only when the sun is rising in September and October, and clouds are still present, are conditions right for O3 destruction.

Following the discovery of the Antarctic ozone “hole” in 1985, atmospheric chemist Susan Solomon led the first expedition in 1986 specifically intended to make chemical measurements of the Antarctic atmosphere by using balloons and ground-based spectroscopy. The expedition discovered that ozone depletion occurred after polar sunrise and that the concentration of chemically active chlorine in the stratosphere was ~100 times greater than had been predicted from gas-phase chemistry. Solomon’s group identified chlorine as the culprit in ozone destruction and polar stratospheric clouds as the catalytic surface for the release of so much chlorine.

The Ozone Hole, 3

Properties of LightLight waves consist of perpendicular, oscillating electric and magnetic fields. For simplicity, a plane-polarized wave is shown below. The electric field is in the xy plane, and the magnetic field is in the xz plane. Wavelength, λ, is the crest-to-crest distance between waves. Frequency, ν, is the number of complete oscillations that the wave makes each second. The unit of frequency is 1/second. One oscillation per second is called one hertz (Hz).

Wave-like and particle-likeThe relation between frequency and wavelength is

where c is the speed of light, 2.998 × 108 m s-1 in vacuum. In a medium other than vacuum, the speed of light is c/n, where n is the refractive index of that medium. For visible wavelengths in most substances, n > 1, so visible light travels more slowly through matter than through vacuum. When light moves between media with different refractive indexes, the frequency remains constant but the wavelength changes.

It is often convenient to think of light as particles called photons. Each photon carries the energy, E, which is given by

where h is Planck’s constant, 6.626 × 10−34 J·s. We can also write:

where is equal to 1/λ, which is known as the wavenumber (usually units of cm-1). ν

EM spectrumThe names of regions of the electromagnetic spectrum are historical. There are no abrupt changes in characteristics as we go from one region to the next, such as visible to infrared.

Absorption of LightWhen a molecule absorbs a photon, the energy of the molecule increases. We say that the molecule is promoted to an excited state. If a molecule emits a photon, the energy of the molecule is lowered. The lowest energy state of a molecule is called the ground state. Microwave radiation stimulates rotation of molecules when it is absorbed. Infrared radiation stimulates vibrations. Visible and ultraviolet radiation promote electrons to higher energy orbitals. X-rays and short-wavelength ultraviolet radiation break chemical bonds and ionize molecules (hence the associated danger in medical procedures).

Photon energies

Transmittance, Absorbance & Beer’s LawWhen light is absorbed by a sample, the irradiance of the beam of light is decreased. Irradiance, P, is the energy per second per unit area of the light beam. In a rudimentary spectrophotometer, light is passed through a monochromator (a prism, a grating, or even a filter) to select one wavelength. Light with a very narrow range of wavelength is said to be monochromatic (“one color.”) The monochromatic light, with irradiance P0, strikes a sample of length b. The irradiance of the beam emerging from the other side of the sample is P. Some of the light may be absorbed by the sample, so P ≤ P0.

Transmittance, T, is defined as the fraction of the original light that passes through the sample.

Therefore, T has the range 0 to 1. The percent transmittance is simply 100T and ranges between 0 and 100%.

Transmittance, Absorbance & Beer’s Law, 2Absorbance is defined as

When no light is absorbed, P = P0 and A = 0. If 90% of the light is absorbed, 10% is transmitted and P = P0/10. This ratio gives A = 1. If only 1% of the light is transmitted, A = 2. Absorbance is sometimes called optical density.

Absorbance is directly proportional to the concentration, c, of the light-absorbing species in the sample, as described by the Beer-Lambert law, or simply Beer’s law

Absorbance is dimensionless, but some people write “absorbance units” after absorbance. The concentration of the sample, c, is usually given in units of moles per liter (M). The path length, b, is commonly expressed in centimeters. The quantity ε (epsilon) is called the molar absorptivity (or extinction coefficient in the older literature) and has the units M−1 cm−1 to make the product εbc dimensionless. Molar absorptivity is the characteristic of a substance that tells how much light is absorbed at a particular wavelength.

Transmittance, Absorbance & Beer’s Law, 3

Transmittance, Absorbance & Beer’s Law, 4The equation for Beer’s law could be written as Aλ = ελbc, because A and ε depend on the wavelength of light. The quantity ε is simply a coefficient of proportionality between absorbance and the product bc. The greater the molar absorptivity, the greater the absorbance.

The part of a molecule responsible for light absorption is called a chromophore. Any substance that absorbs visible light appears colored when white light is transmitted through it or reflected from it. The substance absorbs certain wavelengths of the white light, and our eyes detect the wavelengths that are not absorbed. The observed color is called the complement of the absorbed color. e.g., bromophenol blue has maximum absorbance at 591 nm and its observed color is blue.

Transmittance, Absorbance & Beer’s Law, 5Beer’s law applies when radiation is monochromatic and the solution is dilute (≤ 0.01 M).

Beer’s law fails in certain cases:

High concentrations: In concentrated solutions, solute molecules influence one another as a result of their proximity. When solute molecules get close to one another, their properties (including molar absorptivity) change somewhat. At very high concentration, the solute becomes the solvent. Properties of a molecule are not exactly the same in different solvents.

Interactions with other solutes: Non-absorbing solutes in a solution can also interact with the absorbing species and alter the absorptivity.

Concentration-dependent chemical equilibrium: If the absorbing molecule participates in a concentration-dependent chemical equilibrium, the absorptivity changes with concentration. For example, in concentrated solution, a weak acid, HA, may be mostly undissociated. As the solution is diluted, dissociation increases. If the absorptivity of A− is not the same as that of HA, the solution will appear not to obey Beer’s law as it is diluted.

Interference from other absorbing species: Dilute species with similar absorption maxima to the solute of interest may produce false readings.

Measuring AbsorbanceLight from a continuous source is passed through a monochromator, which selects a narrow band of wavelengths from the incident beam. This “monochromatic” light travels through a sample of path length b, and the irradiance of the emergent light is measured.

For UV/Vis spectroscopy, a liquid sample is usually contained in a cell called a cuvet that has flat, fused-silica (SiO2) faces. Glass is suitable for visible, but not UV spectroscopy, because it absorbs ultraviolet radiation (quartz!). The most common cuvets have a 1.000 cm path length and are sold in matched sets for sample and reference.

Measuring Absorbance, 2For infrared measurements, cells are commonly constructed of NaCl or KBr. For the 400 to 50 cm−1 far-infrared region, polyethylene is a transparent window. Solid samples are commonly ground to a fine powder, which can be added to mineral oil (a viscous hydrocarbon also called Nujol) to give a dispersion that is called a mull and is pressed between two KBr plates. The analyte spectrum is obscured in a few regions in which the mineral oil absorbs infrared radiation. Alternatively, a 1 wt% mixture of solid sample with KBr can be ground to a fine powder and pressed into a translucent pellet at a pressure of ~60 MPa (600 bar). Solids and powders can also be examined by diffuse reflectance, in which reflected infrared radiation, instead of transmitted infrared radiation, is observed. Wavelengths absorbed by the sample are not reflected as well. This technique is sensitive only to the surface of the sample.

Because gases are more dilute than liquids and solids, they require cells with longer path lengths, typically ranging from 10 cm to many meters. A path length of many meters is obtained by reflecting light so that it traverses the sample many times before reaching the detector.

Measuring Absorbance, 3For a single-beam instrument, which has only one beam of light, we do not measure the incident irradiance, P0, directly. Rather, the irradiance of light passing through a reference cuvet containing pure solvent (or a reagent blank) is defined as P0. This cuvet is then removed and replaced by an identical one containing sample. The irradiance of light striking the detector after passing through the sample is the quantity P. Knowing both P and P0 allows T and/or A to be determined. The reference cuvet compensates for reflection, scattering, and absorption by the cuvet and solvent.

Alternatively, one may use a dual-beam instrument:

http://chemwiki.ucdavis.edu/Physical_Chemistry

Practical Advice: Spectral Acquisition1. Record a baseline spectrum with reference solutions (pure solvent or a reagent blank) in both cuvets. The baseline usually exhibits small variations in positive and negative absorbance. The baseline absorbance is subtracted from the sample absorbance to obtain the true absorbance at each wavelength.

2. Choose the wavelength of maximum absorbance for the analyte of interest, since (a) The sensitivity of the analysis is greatest at maximum absorbance; that is, we get the maximum response for a given concentration of analyte, and (b) the absorbance curve is relatively flat at the maximum, so there is little variation in absorbance if the monochromator drifts a little or if the width of the transmitted band changes slightly.

3. Adjust conditions for optimum absorbance. Modern spectrophotometers are most precise (reproducible) at intermediate levels of absorbance (A ≈ 0.3 to 2). If too little light gets through the sample, intensity is hard to measure. If too much light gets through, it is hard to distinguish transmittance of the sample from that of the reference. Adjust sample concentration so that absorbance falls in this intermediate range. Compartments must be tightly closed to avoid stray light, which leads to false readings.

4. Avoid filth: Keep containers covered and filter solutions to exclude dust and particles, which scatter light (giving high absorbance readings). Handle cuvets with a tissue to avoid putting fingerprints on the faces or damaging them, and keep cuvets scrupulously clean.

Beer’s Law and Chemical Analysis• Sample must absorb light, and this absorption should be distinguishable from that due

to other substances in the sample.

• Because most compounds absorb ultraviolet radiation, measurements in this region of the spectrum tend to be inconclusive, and analysis is usually restricted to the visible spectrum. If there are no interfering species, however, ultraviolet absorbance is satisfactory.

• Proteins are normally assayed in the ultraviolet region at 280 nm because aromatic groups present in virtually every protein have an absorbance maximum at 280 nm.

Beer’s Law and Chemical Analysis, 2

Written in-class problems:

(a) Pure hexane has negligible ultraviolet absorbance above a wavelength of 200 nm. A solution prepared by dissolving 25.8 mg of benzene (C6H6, FM 78.11) in hexane and diluting to 250.0 mL had an absorption peak at 256 nm and an absorbance of 0.266 in a 1.000 cm cell. Find the molar absorptivity of benzene at this wavelength.

(b) A sample of hexane contaminated with benzene had an absorbance of 0.070 at 256 nm in a cuvet with a 5.000 cm path length. Find the concentration of benzene in mg/L.

This lecture will be updated with the remaining slides on Oct. 16, 2013

Spectrophotometric TitrationIn a spectrophotometric titration, changes in absorbance are monitored during a titration to tell when the equivalence point has been reached.

A solution of the iron-transport protein, transferrin, can be titrated with iron to measure the transferrin content. Transferrin without iron, apotransferrin, is colorless. Each molecule, with a molecular mass of 81,000, binds two Fe3+ ions. When Fe3+ binds to the protein, a red color with an absorbance maximum at a wavelength of 465 nm develops. The absorbance is proportional to the concentration of iron bound to the protein. Therefore, absorbance may be used to follow the course of a titration of an unknown amount of apotransferrin with a standard solution of Fe3+.

Spectrophotometric Titration, 2The titration of 2.000 mL of apotransferrin with 1.79 mM ferric nitrilotriacetate is represented in the graph below. As iron is added to the protein, red color develops and absorbance increases. When protein is saturated with iron, no more iron can bind and the curve levels off. The extrapolated intersection of the two straight portions of the titration curve at 203 μL is taken as the end point. Absorbance rises slowly after the equivalence point because ferric nitrilotriacetate has some absorbance at 465 nm.

FIGURE 17-10  Spectrophotometric titration of transferrin with ferric nitrilotriacetate. Absorbance is corrected as if no dilution had taken place. The initial absorbance of the solution, before iron is added, is due to a colored impurity.

The quantity of Fe3+ required for complete reaction (203 × 10−6 L) × (1.79 × 10−3 mol L-1) = 0.363 μmol. Each protein molecule binds 2 Fe3+ ions, so the moles of protein in the sample must be ½(0.363 μmol) = 0.182 μmol.

Spectrophotometric Titration, 3To construct the graph on the previous page, dilution must be considered because the volume is different at each point. Each point on the graph represents the absorbance that would be observed if the solution had not been diluted from its original volume of 2.000 mL.

What Happens When A Molecule Absorbs Light?When a molecule absorbs a photon (ca. 397 nm), the molecule is promoted to a more energetic excited state. Conversely, when a molecule emits a photon, the energy of the molecule falls by an amount equal to the energy of the photon.

Consider formaldehyde: In its ground state, S0, the molecule is planar (structure a), with a double bond between carbon and oxygen. From the electron dot description of formaldehyde, we expect two pairs of nonbonding electrons to be localized on the oxygen atom. The double bond consists of a sigma bond between carbon and oxygen and a pi bond made from the 2py (out-of-plane) atomic orbitals of carbon and oxygen.

Although formaldehyde is planar in its ground state, it has a pyramidal structure (structure b) in both the S1 and T1 excited states (these are defined shortly). Promotion of a nonbonding electron to an anti-bonding C―O orbital lengthens the C―O bond and changes the molecular geometry.

Electronic States of Formaldehyde

In the molecular orbital (MO) diagram for formaldehyde, one of the nonbonding AOs of oxygen is mixed with the three sigma bonding orbitals. These four orbitals, labeled σ1 through σ4, are each occupied by a pair of electrons with opposite spin (+½ and -½).

At higher energy is an occupied pi bonding orbital (π), made of the py AOs of C and O. The highest energy occupied MO (HOMO) is the nonbonding orbital (n), composed principally of the 2px AOs from O. The lowest energy unoccupied MO (LUMO) is the pi anti-bonding orbital (π*). Electrons in this orbital produce repulsion, rather than attraction, between the C and O atoms.

Electronic States of Formaldehyde, 2In an electronic transition, an electron from one MO moves to another, with a concomitant increase or decrease in the energy of the molecule. There are complex sets of guidelines for which transitions are allowed or forbidden, known as selection rules. The lowest energy electronic transition of formaldehyde promotes a nonbonding (n) electron to the anti-bonding pi orbital (π*). However, there are two possible transitions, depending on the spin quantum numbers in the excited state. The state in which the spins are antiparallel is called a singlet state (S1), and if the spins are parallel, we have a triplet state (T1).

In general, T1 has lower energy than S1. In formaldehyde, the transition n → π*(T1) requires absorption of visible light with a wavelength of 397 nm, whereas the n → π*(S1) transition occurs when ultraviolet radiation with a wavelength of 355 nm is absorbed.

Vibrational and Rotational StatesAbsorption of visible or UV radiation promotes electrons to higher energy orbitals in formaldehyde. IR and microwave radiation are not energetic enough to induce electronic transitions, but they can change the vibrational or rotational motion of the molecule.

A non-linear molecule has 3N-6 vibrational modes (linear molecules have 3N-5). Each of the atoms in formaldehyde can move in 3 different directions (x, y, z), for a total of 12 modes of motion. Three are assigned to translation, three to rotation and the remaining six to vibrations (pictured to the left).

If formaldehyde absorbs an IR photon with a wavenumber of 1251 cm−1 (14.97 kJ/mol), the asymmetric bending vibration is stimulated: Oscillations of the atoms increase in amplitude, and the energy of the molecule increases.

Spacings between rotational energy levels of a molecule are smaller than vibrational energy spacings. A molecule in the rotational ground state can absorb microwave photons with wavelengths of 4.115 or 1.372 mm (0.02907 or 0.08716 kJ/mol) to be promoted to the two lowest excited states. Absorption of microwave radiation causes the molecule to rotate faster than it does in its ground state.

Combined TransitionsIn general, when a molecule absorbs light having sufficient energy to cause an electronic transition, vibrational and rotational transitions—that is, changes in the vibrational and rotational states—occur as well (rovibronic transitions). Formaldehyde can absorb one photon with just the right energy to cause the following simultaneous changes

(1) a transition from the S0 to the S1 electronic state,

(2) a change in vibrational energy from the ground vibrational state of S0 to an excited vibrational state of S1, and

(3) a transition from one rotational state of S0 to a different rotational state of S1.

Electronic absorption bands are usually broad because many different vibrational and rotational levels are available at slightly different energies. The ground and excited electronic states are represented by Teʺ and Teʹ, respectively. The v and J are vibrational and rotational quantum numbers, respectively.

What Happens to Absorbed Energy?The fate of the absorbed energy can be described by a Jablonski diagram.

S0 is the ground electronic state. S1 and T1 are the lowest excited singlet and triplet electronic states. Straight arrows represent processes involving photons, and wavy arrows are radiationless transitions. R denotes vibrational relaxation. Absorption could terminate in any of the vibrational levels of S1, not just the one shown. Fluorescence and phosphorescence can terminate in any of the vibrational levels of S0.

What Happens to Absorbed Energy, 2After an absorption process, the system is at the S1 level, and several events can happen:

Internal conversion (IC): The molecule could enter a highly excited vibrational level of S0 having the same energy as S1. From this excited state, the molecule can relax back to the ground vibrational state and transfer its energy to neighboring molecules through collisions. This radiationless process is labeled R2. If a molecule follows the path A–R1–IC–R2, the entire energy of the photon will have been converted into heat.

Intersystem Crossing (ISC): The molecule could cross from an S1 state into an excited vibrational level of T1. After the radiationless vibrational relaxation R3, the molecule finds itself at the lowest vibrational level of T1. From here, the molecule might undergo a second intersystem crossing to S0, followed by the radiationless relaxation R4. All processes mentioned so far simply convert light into heat.

Fluorescence (F): A molecule could also relax from S1 to S0 by emitting a photon. The radiational transition S1→S0 is called fluorescence. Typical lifetimes of fluorescence processes are 10−8 to 10−4 s.

Phosphorescence (P): A molecule could also relax from T1 to S0 by emitting a photon. The radiational transition T1→S0 is called phosphorescence. Typical lifetimes of fluorescence processes are 10−4 to 102 s, because the transition is forbidden: it involves a change in spin quantum numbers (2 unpaired electrons to 0 unpaired electrons), which is improbable!

What Happens to Absorbed Energy, 3The relative rates of internal conversion, intersystem crossing, fluorescence, and phosphorescence depend on the molecule, the solvent, and conditions such as temperature and pressure. Fluorescence and phosphorescence are relatively rare processes, with molecules generally decaying from their excited states by radiationless transitions. The energy of phosphorescence is less than the energy of fluorescence, so phosphorescence comes at longer wavelengths than fluorescence.

The lifetime of a state associated with F or P is the time needed for the population of that state to decay to 1/e of its initial value, where e is the base of natural logarithms.

A few materials, such as strontium aluminate doped with europium and dysprosium (SrAl2O4:Eu:Dy) phosphoresce for hours after exposure to light. One application for such materials is in signs leading to emergency exits when power is lost.

LuminescenceFluorescence and phosphorescence are examples of luminescence, which is emission of light from an excited molecular state. Luminescence is inherently more sensitive than absorption in terms of chemical analysis, even of single molecules!

Imagine yourself in a stadium at night; the lights are off, but each of the 50 000 fans is holding a lighted candle. If 500 people blow out their candles, you will hardly notice the difference. Now, imagine that the stadium is completely dark; then 500 people suddenly light their candles. In this case, the effect is dramatic!

The first example is analogous to changing transmittance from 100% to 99%, which is equivalent to an absorbance of −log 0.99 = 0.0044. It is hard to measure such a small absorbance because the background is so bright. The second example is analogous to observing fluorescence from 1% of the molecules in a sample. Against a black background, fluorescence is striking.

Taken at the opening ceremonies of theVancouver 2010 Winter Olympics

Luminescence, 2Luminescence is sensitive enough to observe single molecules. Below are the observed tracks of two molecules of the highly fluorescent Rhodamine 6G at 0.78 s intervals in a thin layer of silica gel. These observations confirm the “random walk” of diffusing molecules postulated by Albert Einstein in 1905.

Tracks of two molecules of 20 pM Rhodamine 6G in silica gel observed by fluorescence integrated over 0.20-s periods at 0.78-s intervals. Some points are not connected, because the molecule disappeared above or below the focal plane in the 0.45-mm thick film and was not observed in a particular observation interval. In the nine periods when molecule A was in one location, it might have been adsorbed to a particle of silica. An individual molecule emits thousands of photons in 0.2 s as the molecule cycles between the ground and excited state. Only a fraction of these photons reach the detector, which generates a burst of ~10-50 electrons. [From K. S. McCain, D. C. Hanley, and J. M. Harris, “Single-Molecule Fluorescence Trajectories for Investigating Molecular Transport in Thin Silica Sol-Gel Films,” Anal. Chem. 2003, 75, 4351.]

Absorption vs. Emission SpectraSpectra arising from fluorescence and phosphorescence come at lower energies than absorption (the excitation energy). That is, molecules emit radiation at longer wavelengths than the radiation they absorb.

Absorption vs. Emission Spectra, 2

In absorption, wavelength λ0 corresponds to a transition from the ground vibrational level of S0 to the lowest vibrational level of S1. Absorption maxima at higher energy (shorter wavelength) correspond to the S0 to S1 transition accompanied by absorption of one or more quanta of vibrational energy.

After absorption, the vibrationally excited S1 molecule relaxes back to the lowest vibrational level of S1 prior to emitting any radiation. Emission from S1 can go to any vibrational level of S0. The highest energy transition comes at wavelength λ0, with a series of peaks following at longer wavelengths. Absorption and emission spectra will have an approximate mirror image relation if the spacings between vibrational levels are roughly equal and if the transition probabilities are similar.

Absorption vs. Emission Spectra, 3The λ0 transitions do not exactly overlap. In the emission spectrum, λ0 comes at slightly lower energy than in the absorption spectrum. Why? A molecule absorbing radiation is initially in its electronic ground state, S0, and it possesses a certain geometry and solvation. An electronic transition to S1 is much faster than the vibrational motion of atoms or the translational motion of solvent molecules. Hence, when radiation is first absorbed, the excited S1 molecule still possesses its S0 geometry and solvation. Shortly after excitation, the geometry and solvation change to their most favorable values for the S1 state, which lowers the energy of the excited molecule. When the S1 molecule fluoresces, it returns to the S0 state with S1 geometry and solvation. This unstable configuration must have a higher energy than that of an S0 molecule with S0 geometry and solvation. The net effect is that the λ0 emission energy is less than the λ0 excitation energy.

Excitation and Emission SpectraEmission spectra: An excitation wavelength (λex) is selected by one monochromator, and luminescence is observed through a second monochromator, usually positioned at 90° to the incident light to minimize the intensity of scattered light reaching the detector. If we hold the excitation wavelength fixed and scan through the emitted radiation, an emission spectrum is produced.

Excitation spectra: An excitation spectrum is measured by varying the excitation wavelength and measuring emitted light at one particular wavelength (λem).

Excitation and Emission Spectra, 2Emission and excitation spectra are graphs of emission/excitation intensity versus emission/excitation wavelengths. An excitation spectrum looks very much like an absorption spectrum because, the greater the absorbance at the excitation wavelength, the more molecules are promoted to the excited state and the more emission will be observed.

Excitation and emission spectra of anthracene have the same mirror image relationship as the absorption and emission spectra in Figure 17-18. An excitation spectrum is nearly the same as an absorption spectrum. [From C. M. Byron and T. C. Werner, “Experiments in Synchronous Fluorescence Spectroscopy for the Undergraduate Instrumental Chemistry Course,” J. Chem. Ed. 1991, 68, 433.]

Luminescence in Analytical ChemistrySome analytes, such as riboflavin (vitamin B2) and polycyclic aromatic compounds (an important class of carcinogens), are naturally fluorescent and can be analyzed directly. Most compounds are not luminescent.

However, coupling to a fluorescent moiety provides a route to sensitive analyses. Fluorescein is a strongly fluorescent compound that can be coupled to many molecules for analytical purposes. e.g., Fluorescent labeling of fingerprints is a powerful tool in forensic analysis.

Sensor molecules whose luminescence responds selectively to a variety of simple cations and anions are available. For example, Ca2+ concentrations can be measured from the fluorescence of a complex it forms with a derivative of fluorescein called calcein.

Luminescence in Analytical Chemistry, 2Molecular biologists use DNA microarrays (“gene chips”) to monitor gene expression and mutations and to detect and identify pathological microorganisms. A single chip can contain thousands of known single-strand DNA sequences in known locations. The chip is incubated with unknown single-strand DNA that has been tagged with fluorescent labels. After the unknown DNA has bound to its complementary strands on the chip, the amount bound to each spot on the chip is measured by fluorescence intensity.

Luminescence in Analytical Chemistry, 3Light from a firefly or light stick is an example of chemiluminescence—emission of light from a chemical reaction. Chemiluminescence detectors for sulfur and nitrogen in organic compounds are employed in gas chromatography, and nitric oxide (NO), which transmits signals between living cells, can be measured at parts per billion levels by its chemiluminescent reaction with the compound luminol. Other biological analytes measurable by chemiluminescence include Ca2+ in mitochondria and endocrine-disrupting compounds in municipal wastewater.