Ch. 17: Fundamentals of Spectrophotometry
Outline:
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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
Spectrophotometry
Spectrophotometry 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 Hole
Ozone, 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.
hν
O2
2O·
O· + O2
O3
Ozone
In 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 longlived 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, 2
Stratospheric 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 Light
Light 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-like
The 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 ν is equal to 1/λ, which is known as the wavenumber (usually units of cm-1).
EM spectrum
The 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 Light
When 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 Law
When 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, 2
Absorbance 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, 4
The 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, 5
Beer’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 Absorbance
Light 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, 2
For 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, 3
For 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 dualbeam instrument:
http://chemwiki.ucdavis.edu/Physical_Chemistry
Practical Advice: Spectral Acquisition
1. 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
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Sample must absorb light, and this absorption should be distinguishable from that due
to other substances in the sample.
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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.
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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 Titration
In 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, 2
The 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, 3
To 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 antibonding 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 antibonding orbital (π*). Electrons in this orbital
produce repulsion, rather than attraction,
between the C and O atoms.
Electronic States of Formaldehyde, 2
In 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 States
Absorption 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 Transitions
In 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, 2
After 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, 3
The 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.
Luminescence
Fluorescence 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 the
Vancouver 2010 Winter Olympics
Luminescence, 2
Luminescence 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 Spectra
Spectra 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, 3
The 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 Spectra
Emission 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, 2
Emission 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 Chemistry
Some 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, 2
Molecular 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, 3
Light 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.