Lab #3: Spectrophotometry p. 1

Lab #3: Spectrophotometry Background One of the key functions of the homeostatic mechanisms of the human body is to maintain the chemical composition of the fluid environment in which the cells of the body live. The ability of the body to do so is really quite remarkable when one considers how the flow of different materials into and out of the body as well as the diverse array of chemical reactions that occur as a result of metabolism are constantly altering the composition of body fluids. Yet when physiological systems are functioning properly, the concentrations of various metabolites, nutrients, waste materials, etc. are maintained within very narrow ranges. Deviations in the concentrations of different substances in body fluids, therefore, suggest a problem with physiological mechanisms that normally regulate those substances. Therefore, abnormal levels of specific substances in the body fluids can be indicative of specific pathologies, and thus can be used to diagnose particular ailments.

There are numerous methods for measuring the concentrations of specific substances within body fluids. One commonly-used method is called spectrophotometry (also called “colorimetry”). Spectrophotometry works on a very basic principle—that if your solution contains a solute that absorbs light, then the more concentrated that solute is in the solution the more light should be absorbed by that solution and the less light will pass through that solution. The mathematical relationship between solute concentration and light absorbance is described in Beer’s Law*:

(1) [X] α A where [X] is the concentration of solute X and A is the amount of light absorbed. Incidentally, α means “proportional to” (i.e., that [X] multiplied by some constant number will give you the value of A). This describes a direct linear relationship between the concentration and light absorbance (i.e., per each unit increase in

concentration, the absorbance of the solution increases a specific amount) (Figure 3.1).

Beer’s Law also describes the mathematical relationship between solute concentration and light transmittance by a solution (i.e., how much light passes through the solution). (2) where T is the percentage of light transmittance (i.e. the percentage of the light entering the solution that passes through the solution). This equation describes an inverse curvilinear relationship between solute concentration and light transmittance (as concentration increases, light transmittance decreases, but there is progressively less of a decrease in transmittance per unit concentration as the concentration increases) (Fig 3.1).

Because there is a mathematical relationship between solute concentration and light absorption/transmittance, we can use spectrophotometry to determine the concentration of a substance in a sample of body fluid (or any other solution where we do not already know the substance’s concentration) by comparing its light absorbance or transmittance with a that of a solution that has a known concentration of that substance (called a reference, or a standard solution). For example, let’s say we wanted to determine the concentration of glycerol in a sample of blood plasma. We react the glycerol with a reagent that changes that absorbs light at a specific wavelength when it reacts with the glycerol (this allows us to identify light absorbance by glycerol as opposed to that by other substances in the solution). We then determine the light absorbance of the solution by placing the solution in a *referring to the German mathematician and physical scientist August Beer (1825-1863) who first described this relationship, not after the alcoholic beverage (although scientists often make great insights over a pint at the local brew pub). This mathematical relationship is also called the Beer-Lambert Law or the Boulguer-Beer Law. The mathematical proportionalities here are simplified versions of this principle, and assume a constant light path length and a constant molar extinction coefficient.

[X] α 1__ log T

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spectrophotometer (an instrument that measures light absorbance at specific wavelengths). We determine the absorbance our blood plasma sample (Ablood) to be 0.090. We also run the same procedure on a glycerol solution of known concentration ([standard]) of 250 mg/dL, and determine its absorbance (Astandard) to be 0.300.

So given these values, how can we determine the concentration of glycerol in the blood plasma sample ([blood])? Remember that according to Beer’s Law the concentration of a solution is proportional to its absorbance. That proportionality (i.e., how much light is absorbed per unit of concentration) will be the same both in the blood plasma sample and in the standard solution. So if that proportionality is the same, we can describe it through the following equation: (3)

We can then rearrange the equation to solve for [blood] by multiplying both sides of the equation by Ablood (4)

(5) Experimental Error and Standard Curves Typically, we determine the concentration of some solute of interest by reacting it with another substance to produce a substance that absorbs light at a particular wavelength, and comparing it to a second solution with a known concentration. When we do so, we are trusting that we have conducted the experiment correctly, and that the results accurately reflect the actual concentration of the substance. But is this a valid assumption?

The potential for error is always present in any measurement. Here we are referring to error as any deviation in our measured value away from the true value for that parameter. In that respect, all measurements contain some degree of error—there is no way you can perfectly measure anything down to the last molecule, to the exact temperature, etc., and conditions that you cannot control can change the reading you obtain for a particular measurement. Despite this, there are ways of minimizing experimental error: by controlling extraneous conditions that might affect your results (e.g., temperature), by carefully measuring out quantities, and by repeating the measurements when possible to see if similar results are obtained.

When determining concentrations using a spectrophotometer, there is certainly a potential for experimental error. For instance, in the example above where we measured plasma glycerol, samples of blood plasma and the standard solution were mixed with a reagent to produce a substance that absorbed light at a

[blood] Ablood = [standard]

Astandard

[blood] = [standard] × Ablood_ Astandard

We now fill in the values we have measured and solve for [blood].

[blood] = 250 mg/dL × 0.090 0.300

= 75 mg/dL

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Fig. 3.1. Examples of (left) the linear (direct proportional) relationship between solute concentration and light absorbance and of (right) the inverse logarithmic relationship between solute concentration and light transmittance as predicted via Beer’s Law.

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particular wavelength. There could be slight differences in the amount of plasma or standard solution used, the amount of reagent used, the specific concentration of the standard solution, etc., that could influence the absorbance readings. Indeed, this could be especially problematic if we do not account for possible error in the readings obtained for the standard solution. If only one standard value is used, calculations of the unknown based on the absorbance of that standard will reflect any errors made in the preparation of the standard. This could have serious effects on the results.

For example, let’s say you had very carefully conducted the glycerol assay in the example above. Thus the results should fairly accurately reflect the true glycerol levels in the plasma. Now pretend that you conducted the experiment again, but were not quite as careful with your pipetting, incubation, and other aspects of preparation. In this case you obtain an absorbance for the blood plasma solution of 0.100 (a bit higher than before, but not too terribly off). However, because you were not paying attention, you pipetted in only half as much standard solution as you were supposed to, and ended up obtaining an absorbance reading of 0.145 for the 250 mg/dL standard. The glycerol concentration for the blood plasma would be calculated to be 172 mg/dL (way too high)! Replication of an experiment is one means by which we can we account for experimental error (either due to preparation mistakes or simple random variation between one reading and another) to get accurate estimates of the concentrations.

Another potential problem is that a light-absorbing solution may adhere to the Beer’s law relationship between light absorbance and concentration only over a certain range of concentrations, beyond which there may be a curvilinear relationship. This problem can be avoided by referencing the measurements for the test solutions against multiple standards, some of which are higher in concentration than the test solution and some of which are lower.

Typically, a researcher will prepare several standard solutions of different concentrations in order to determine the solution of an unknown. The absorbance values of these solutions are measured and plotted against the concentration of the standards. A “best fit” line (e.g., a linear regression line) would be drawn through the points to produce a standard curve (Fig 3.2). Note that even in the best of experiments not all the individual points fit on the line—some are above and some are below. The line effectively balances out variation and provides an average value of absorbance for a solution of a given concentration. We can then determine the concentration of the unknown by extrapolation. First we find the value for absorbance on the vertical (y) axis, and draw a straight horizontal line from the y-axis to the best-fit line. We then draw a straight vertical line from the point where the horizontal line meets the best-fit line straight down to the horizontal (x) axis. The value on the x-axis where this vertical line intersects it is the concentration of the unknown.

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y = -0.03 + 0.027x R= 0.996

Fig. 3.2. An example of a standard curve generated by measurement of absorbance from multiple standard solutions of different concentrations. The best fit line helps to even out experimental error, and helps with identification of samples where there might have been some procedural mistake. The dashed lines provide an example of extrapolation (e.g., a solution with an absorbance of 1.000 would have a concentration of ~38). The line slope equation derived from linear regression could also be used.

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Fig 3.3. The Spectronic Spec 20 spectrophotometer. Experiment: Determining the Concentration of Glucose In today’s experiment we will be determining the concentration of monosaccharide glucose in various solutions using the spectrophotometer to measure light absorbance and comparing the absorbance of unknown solutions with those of standard solutions. The procedure used is one used in clinical settings for measurement of plasma glucose. Samples of blood plasma are mixed together with a reagent solution containing o-toluidine (a purplish substance). When glucose reacts with o-toluidine, it forms a greenish complex that absorbs light optimally at 620-650 nm. Thus the more glucose present in the sample, the more of the complex will be formed and the more light will be absorbed by the solution within that range of wavelengths.

Note of Caution! The reagent solution we will be using contains concentrated acetic acid and o-toluidine.

o The concentration of acetic acid is caustic: the vapors can irritate the eyes and mucus membranes, and it will damage the skin and eyes if they come into direct contact. Wear eye protection and gloves, and do not wear open-toed shoes. Also, I recommend wearing clothes you do not care too much about to this lab, just to be safe.

o o-toluidine is a suspected carcinogen. Again, avoid getting the reagent on you. o Clean up any spills immediately with LOTS of water

Procedures 1. Obtain four clean test tubes and label them with masking tape 2. Add 5.0 ml glucose reagent to each tube with an automatic pipetter

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3. Add the following to the tubes:

Tube #1 – 0.1 ml distilled water (will serve as a “blank” solution for calibrating the spectrophotometer)

Tube #2 – 0.1 ml blood serum (from some domesticated animal)

NOTE: Be very careful to remove only plasma from the sample. Avoid contaminating the plasma by disturbing the packed cells in the bottom of the microtube. Do not add contaminated (red tinted) plasma to your tube.

Tube #3 – 0.1 ml unknown #1 Tube #4 – 0.1 ml unknown #2 4. Immerse tubes 2, 3 and 4 in 100°C water bath for 10 min 5. While the tubes are in the hot water bath, calibrate the spectrophotometer:

o Turn on the machine with the Power Knob (front left) o Set wavelength to 620-650 nm with the Wavelength Adjustment Knob (top right) o The instrument should have nothing in the sample chamber, the sample chamber lid should be

closed, and the mode should be set to “Transmittance”. If it is not set to transmittance, push the Mode Button until the mode indicator light is next to “Transmittance”.

o Set the transmittance to read “0.0” with front left knob. o Transfer the blank solution (Tube #1) into a clean cuvette. Place the cuvette in the sample

chamber and close the lid. The readout on the instrument should read “100.0”. If it does not, use the Zero Knob (front right) to adjust the value to 100.0.

o Push the Mode Button to switch the mode to “Absorbance”. The readout should change to “.000”. The instrument has now been calibrated.

5. Remove Tubes #2-4 and immediately immerse in ice bath for 3 min. 6. Record absorbance values for remaining three tubes. Place the tube in the sample chamber, close the

lid, and record the absorbance value provided.

NOTE: The biggest mistake students make with this exercise is improperly calibrating the instrument. Usually, they forget to switch modes between absorbance and transmittance at the proper times. Make sure that your readings are for Absorbance. At the end of the calibration procedure your blank solution should give you a reading of “.000”. All other measurements should have three digits to the right of the decimal, and should not exceed 1.400. If you have only one digit to the right of the decimal point, or if your readings are between 2 and 100, your spec is set to the wrong mode.

7. Once complete, dispose of your solutions in the waste bottle under the fume hood, and clean out the

test tubes and cuvettes per the instructions posted by the sink. 8. Calculate the concentration of glucose in each of your three samples using your absorbance readings

and the concentration and absorbance of a standard solution (provided by your instructor). Record these values on your data sheet and on the table on the instructor’s station. We will see whose measurements are the closest to the actual concentrations for Tubes #3 and 4.