Laboratory 1: Introduction to Spectrophotometry

Background Information

Electromagnetic Radiation

Spectrophotometry, generally, refers to the relative measurement of radiant energy or radiant flux as a function of wavelength.  The term "relative" specifies that in spectrophotometry the measurements are always made relative to some standard. What this standard may be in any case depends on the type of measurement.  Perhaps the oldest form of spectrophotometry is exemplified by the ability of the eye to distinguish different colors.  The production of colors by molecules is due to their ability to absorb or remove certain components of the electromagnetic radiation that impinges upon them.  The wavelength that are absorbed and the efficiency of absorption depend on the structure and environment of the molecule.  To appreciate how these work, we need to understand what we mean by electromagnetic radiation, how an electromagnetic wave is propagated, look at division of the electromagnetic spectrum, examine certain measurements of absorption and instrumentations commonly used, and some practical applications.

So, what does electromagnetic radiation really mean?  We can answer this questions by looking closely at the properties of light. Light has two properties, wave and particle size properties. If we assumed light as a wave phenomenon, it consists of an electric field that is mutually perpendicular to a magnetic field (hence, electromagnetic), both fields being at right angles to the direction of propagation (Fig 1.). Thus, an electromagnetic radiation can be thought of as a wave that results from periodic oscillations of either electrical or magnetic force at right angles to the direction of propagation.  The distance traveled by the wave between two repeating points is called the wavelength (l, "lambda"), while the number of such repeating units or cycles occurring per second is called the frequency (n, "nu"). The velocity of the wave (c) is given by: c = ln . Because the velocity of all electromagnetic radiation is constant in any particular medium, and since light is an electromagnetic radiation, c is equal to the velocity of light. Electromagnetic radiation can also behave as though it exists as discrete individual unit of energy called photons or quanta. The energy (E) of a photon or quantum is related to the frequency of radiation in direct way; i.e. E = hn = hc/l , in which h is Planck's constant (6.626 x 10^-34 Js, pronounced "joule-second"). Wavelengths, frequencies and energies of common waves are outlined in table 1.

Figure 1. The characteristic form of propagation of an electromagnetic wave through space. 

This page was copied from Nick Strobel's Astronomy Notes. Go to his site at www.astronomynotes.com for the updated and corrected version

See an applet containing the propagation of electromagnetic radiation. Here.

When a wave of electromagnetic radiation encounters a molecule, it can either be scattered (i.e., its direction of propagations changes) or absorbed (i.e., its energy is transferred to the molecule).  The probability of occurrence of that process is a property of the particle molecule encountered.  If the electromagnetic energy of the light is absorbed, the molecule is said to be "excited" or in an "excited state". A molecule or part of it that can be excited by absorption of electromagnetic energy is termed a chromophore.  This excitation energy is usually converted to heat by the collision of the excited molecule with another molecule (e.g., a solvent molecule), but in some molecules it is re-emitted as fluorescence. In both cases, the light intensity transmitted by a collection of chromophores is less than the intensity of the incident light. 

Electromagnetic radiation can  be obtained in a wide range of frequencies.  This continuum of frequencies is called the electromagnetic spectrum (Fig. 2). Notice that the shorter the wavelength of radiation, the greater its energy (recall the formula: E = hc/l). Hence, x-rays have more energy than do infrared waves.  It is also interesting to note that our eyes are only sensitive to a narrow range of the electromagnetic spectrum called the visible region (between 400-800 nm), which, by the way, represents only a small portion of the entire spectrum. Within the visible region, radiation at the high-wavelength (low energy) extreme appears red while those at the low-wavelength (high energy) extreme appears blue.  Our eyes cannot detect light absorbed by an object, but only the light which that object reflects.  Thus, when viewed in white light, objects which absorb red light appear blue and objects which absorb blue light appear red. This complement relationship between colors is very important when dealing with chemical solutions, for instance, a blue solution will have an absorption band in the red region of the spectrum around 700 nm. It follows, then, that certain regions of the spectrum can be studied only through different spectroscopic techniques.  For example, the UV and visible regions can be studied using standard absorption spectrophotometry (e.g., U.V., visible, fluorescence), whereas radio-waves can  be studied using nuclear magnetic resonance (NMR) spectroscopy. However, we will limit our attention on the use of standard absorption spectroscopy to measure UV-visible wavelengths.

Figure 2. Diagrammatic representation of the electromagnetic spectrum

Image: by Francis Carey

Looking back at the absorption process, we can illustrate how molecules, upon absorption of radiation, result in an excited state in a simple expression:

    M  +  hn--> M*

where hn is a photon of electromagnetic radiation, M is an atom or molecule in a native state, and M* is the atom or molecule in an excited state resulting from the absorption of a photon.  Although, the actual process of absorption of radiation is very fast, there is a point in the process when the excited state terminates or switches from the excited state back to the ground (relaxed) state.  This event is very important because if excitation of molecules is not terminated, all the molecules in a system under study would quickly reach the excited state and no further absorption could be observed.  Events in which molecules in excited states may lose energy are simply called relaxation processes. A common relaxation of an excited molecule is the conversion of the excitation energy to heat; i.e.,

     M* --> M + heat

The other relaxation events include re-emission of radiation (leading to fluorescence), and the chemical utilization of the excitation energy resulting in a photochemical reaction, i.e., the breakdown of M* to form new chemical species. This last process is regarded as a very important fate of pesticide chemicals under sunlight or the photochemical conversion of airborne emissions from industrial plants.

The lifetime of M* is very short, so that the concentration of M* is extremely small at any one point in time.  Also, the amount of thermal energy produced in the relaxation process is usually too small for detection.  For these reasons, spectrophotometric absorption methodologies create negligible chemical disturbances of the system during the period of study.  This is quite important in spectrophotometric monitoring of enzyme activities, where we want to maintain the temperature of the reaction environment. At any rate, the efficiency of such thermal relaxation will depend upon such factors as size, shape, and rigidity of molecule in which the transition has occurred (e.g., solvent), presence of other solute molecules,  etc. Thus, careful study of the relaxation processes and excited state lifetimes is very useful.

Now that we know a little bit about electromagnetic radiation, the electromagnetic spectrum, and the molecular absorption of this type of energy, it is time for us to look at the practical application of these principles.  We will examine two generalized laws commonly used to understand how these principles are employed in spectrophotometric measurements. 

Spectrophotometry Principles

First, Lambert's Law.  This law asserts the idea that the proportion of incident electromagnetic radiation absorbed by a medium is independent of the starting intensity and that each unit layer of medium absorbs an equal fraction of the electromagnetic radiation passing through.  For example, if the starting intensity (incident radiation, I_o) is 1.0, and each layer of medium absorbs 10% of the starting intensity, then the intensity will be diminished as it passes through each unit layer in a series resulting in a light intensity leaving the solution:

    1.0 (I_o) -->  |||| --> 0.90 --> |||| --> 0.81 --> |||| --> 0.73 --> |||| --> 0.66 (I)

Therefore, only 0.66 of incident radiation is transmitted across the media, the rest is absorbed. Based on this assumption, the transmittance (T) of radiation is mathematically described as:

     T = e^-ab

where e is the base of natural logarithms, a is the absorption coefficient (specific to each absorbing material), and b is the distances through which the light beam transverses the medium (i.e., path length).  In short, the fraction of light absorbed is proportional to the thickness of the absorbing solution and independent of the initial light intensity. Transmission can also be expressed in terms of starting light intensity and leaving light intensity by the following:

     T = I / I_o

The second generalization is called Beer's Law. This law assumes that absorption is proportional to the number of absorbing molecules through which radiation passes. In practical terms, if the absorbing molecules are dissolved in a non-absorbing solvent, the absorption of the solution is proportional to the concentration of absorbing molecules.  Unlike Lambert's law, Beer takes into consideration the concentration of the absorbing material and describes it mathematically as:

     T = 10^-abc

where a is the absorptivity constant analogous to a above, b is the path length, and c is the concentration of the absorbing material.  However, in recent years, transmittance (T) which is the fraction of incident radiation that is transmitted by the solution, and is usually expressed as a percentage, has been replaced by absorbance (a) as a measure of absorption.  A is related to %T (in decimal number) as follows:

     A = -log T = -log (10^-abc) = abc

    or

    A = -log T = -log (I / I_o)

The absorptivity, a, has been replaced by using the extinction coefficient or molar absorptivity, e, which is the absorbance of a 1M solution of a pure compound under standard conditions of solvent, temperature, and wavelength.  Therefore, by combining the two laws, we can state that absorbance is a function of the path length through the absorbing samples, the concentration of absorbing molecules, and the molar extinction coefficient of the particular chemical species at a given wavelength. Hence;

     A = ebc

where A is the absorption of the sample (absorbance) at a particular wavelength, e is the molar extinction coefficient, b is the path length (usually measured in cm), and c is the concentration of the absorbing material in molar units. Because A is a unit-less quantity, the unit of e must be reciprocal concentration times reciprocal path length, so the units cancel out. Commonly, we use M^-1 cm^-1, and mM^-1 cm^-1 as unit of e. Also, in most spectrophotometers, we use 1.0 cm cell (path length) cuvettes. The molar extinction coefficient is an intrinsic property of the particular species under study.  It can be obtained from published absorption spectra or from the absorption spectrum of a known concentration.  So, by reading A in a spectrophotometer, c can be determined directly.

The linear relationship between absorbance and path length, now known as either  Lambert-Beer's law or simply Beer's law (we will use Beer's law for convenience), is a generalization without known exceptions. However, there are common deviations from Beer's law (i.e., deviations from linearity between absorbance and concentration of absorbing molecules at fixed path length).  Some of these are fundamental, while others are associated with instrumentation and chemical events. First, majority of the deviations observed arise from variation in concentrations which cause changes in the distribution of chromophores in the solution.  In solution, many molecules polymerize as the concentration increases. The spectrum of the dimers are different from the monomers leading to either positive or negative deviations as indicated in Figure 3. Second, at high concentrations, molecules aggregate leading to scattering of light, thus decreasing the amount of light transmitted and causing a positive deviation.  Beer's law is based upon the assumption of the absorption of pure monochromatic light, which is produced by most spectrophotometers.  Hence, the instrument can produce a departure from Beer's law as an effect of polychromatic radiation.  Third, the denaturation of proteins at low concentrations may lead to departure from Beer's law. These include interactions with the solvent as well as dissociation of species. These events can occur in buffered solutions as a result of reactions between a product of an enzyme action under study with a buffer. What all of these points mean, in practical terms, is that accurate spectrophotometric determinations of concentration of an absorbing species depend on first determining the linear range of concentrations, and secondly on selecting an appropriate size aliquot to keep absorption values within the linear range.

In this laboratory activity, our main "toy" is the spectrophotometer.  To appreciate how a spectrophotometer works, it is very important to understand its different components.  A typical spectrophotometer includes at least 5 principal components (Fig. 4): (1) a light source adequately covering the spectrum range of interest for the measurements; (2) a monochromator for wavelength selection (e.g., prism), or spectral dispersing system by which light energy of any desired frequency or wavelength (usually a more or less narrow range of frequencies or wavelengths) may be isolated and used in the desired measurements; (3) a transparent sample holder called a cuvette, for transmission and reflection measurements; (4) a light detector, by which the desired ratio of radiant energies may be determined at the various wavelengths; (5) a suitable recorder or meter for measuring the output of the detector. In a typical operation at a single wavelength, we will measure the light transmitted by the solvent alone (which may be a buffer or a solution of small molecules as a blank or control), followed by a measurement of that transmitted by the sample when dissolved in the same solvent; the first value is then subtracted from the second to give the absorbance of the solute.  In our case, this subtraction will not be done manually because the instrument will be adjusted to read zero absorbance when the solvent is measure (i.e., zeroing the instrument). So after the instrument is adjusted, the absorbance of the sample is read directly.  We will be using a Beckman spectrophotometer, a rather sophisticated instrument compared to the old Bausch and Lomb type.  I will take you through how this spectrophotometer is operated.

Figure 4. Principal components of a spectrophotometer

Image by Ian Sutherland

Finally, we can mention some of the useful applications of absorption spectrophotometry, i.e., chemical analysis within the visible and UV regions of the spectrum.  This way, you can relate your introductory knowledge of spectrophotometry to designing studies in the future.  Basically, absorbance measurements by spectrophotometry allow: (1) the determination of concentrations of a substance0 can be done if the absorption coefficient is known and Beer's law is obeyed; (2) assay of certain chemical reactions- especially when the absorbance of one of the reactants changes during the course of the reaction; and (3) identification of substances- since most substances have characteristic spectra, measuring the ration of absorbance at different wavelengths or by measuring a complete spectrum could be used to determine the identity of a substance.  These are just a few of the many uses of absorption spectrophotometry. So far, we have covered only a superficial aspect of spectroscopy.  Our primary goal in this exercise is to gain an understanding of the applications of the subject to biological studies, particularly in insect physiology. The following sections will describe what you are going to do for the rest of the time period.  

Laboratory Activity

Pre-Lab Experience

As an introduction to the practices of spectrophotometry, this exercise allows you to practice the procedures associated with establishing wavelengths of maximal absorbance, establishing relationships between absorbance and path length, absorbance and concentration and determination of analyte concentration.  As part of this lab experience it is necessary to record data as if you were conducting an in-class experiment.  Documentation is an integral part of developing a working ideology of research techniques and outcomes from these. You will need to record the data and draw the appropriate graphs from each exercise. Go to the following site (from link below), click on the first link, spectrophotometry, read through the introduction and conduct the exercise then proceed through the remaining exercises, finishing with determination of analyte concentration.

http://www.chm.davidson.edu/ChemistryApplets/spectrophotometry/index.html

In-Class Lab Experience

Objective: 

To gain knowledge on the basic principles involved in absoption spectroscopy, its application to entomology, and experience using a spectrophotometer.

Materials:

Methyl Orange (4-[[(4-Dimethylamino)phenyl]-azo]benzenesulfonic acid sodium salt)

    = 327.33 MW, more soluble in hot water, insoluble in alcohol

    = Use: As indicator in 0.1% aqueous solution; pH: 3.1 red, 4.4 yellow

Procedure:

To gain experience using the spectrophotometer (which has countless applications in entomology as well as all other life sciences), we will study the absorbance properties of methyl orange, a chromophore.

  1. Plot of Absorbance (A) versus Wavelength (l)

We will work with 0.01 M solution of methyl orange in water.  This is probably too concentrated for our purposes, so we will make some dilutions [e.g. 1.0 ml of methyl orange + 9.0 ml of water (1:10), then 1.0 ml of methyl orange (1:10) + 9.0 ml of water (1:100)]. Then we will generate absorption spectra by plotting absorbance as a function of wavelength from 400 to 600 nm at 25 nm intervals (in the range of greatest absorbance, we will use 5nm intervals).  Take %T as raw data, which can be read more accurately, then covert this to A (absorbance) using the equations discussed previously, and plot your data.  You should be able to use your plotted information to read the wavelength of maximum absorption.

  2. Plot of Absorbance versus Concentration (Application of Beer's law)

Working with the wavelength of maximum absorption, we can then make 7 dilutions or so of the stock solution.  Read absorption as a function of concentration.

  3. Questions

Does the data follow Beer's law or not?  Based on your data, if this was a real experiment, what range of dilutions should you use to determine the concentration of your sample using a spectrophotometer?  If you are working on a protein sample, what appropriate standards (to zero out the absorbance, or as a blank) can you use in determining the protein concentration using a spectrophotometer and why?

References.

Brown, S.B. (1980) An Introduction to Spectroscopy for Biochemist. Academic Press, London.

Cooper, T.G. (1977) The Tools of Biochemistry. pp. 36-54. John Wiley and Sons, New York.

Copeland, R.A. (1993) Methods for Protein Analysis: A Practical Guide to Laboratory Protocols. Chapman and Hall, New York.

Freifelder, D.M. (1982) Physical Biochemistry. Applications to Biochemistry and Molecular Biology (2nd ed.), pp. 494-536. W.H. Freeman and Co., San Francisco, CA.

Molecular Expressions. Optical Microscopy Primer. Physics of Light and Color. http://micro.magnet.fsu.edu/primer/lightandcolor/electromagintro.html

 Astronomy 103: The Evolving Universe: Stars, Galaxies, Cosmology. Matthew A. Bershady, Univ. of Wisconsin.  

     http://www.astro.wisc.edu/~mab/education/astro103/

Instrumentation for UV-Visible Absorption Spectroscopy. Ian Sutherland. http://www.hull.ac.uk/chemistry/lectures/06217a/lecture2.html