In the Laboratory
Spectrophotometric Determination of Thiocyanate in Human Saliva Markku Lahti* Department of Chemistry, University of Turku, FIN-20014 Turku, Finland;
[email protected] Juhani Vilpo Department of Clinical Chemistry, Tampere University Hospital and University of Tampere Medical School, P.O. Box 2000, FIN-30521 Tampere, Finland Jari Hovinen* Wallac Oy, P.O. Box 10 FIN-20101 Turku, Finland;
[email protected] The experiment for determination of an equilibrium constant for a chemical reaction presented in many laboratory manuals (1) is the equilibrium between iron(III) ion, Fe3+, and thiocyanate ion, SCN{, to form thiocyanatoiron(III) ion, FeSCN2+ (reaction 1). Fe3+(aq) + SCN{(aq)
FeSCN2+(aq)
(1)
When c(Fe ) ≥ c(SCN ), the presence of the thiocyanato complexes Fe(SCN)2+ and Fe(SCN)3 can be excluded. Since the FeSCN2+ solutions of are deep blood-red (λmax = 447 nm) the equilibrium constant of the reaction 1 can be conveniently measured with visible spectrophotometry. The standard curve, from which the concentrations of FeSCN2+ are determined, can be built using standard solutions where [Fe3+] >> [SCN{]. In these solutions all thiocyanate ions are assumed to be converted to the thiocyanato complex ions, FeSCN2+, according to reaction 1. The same calibration curve can be exploited in the quantitative analysis of SCN{ ions in solution. Students seem to enjoy experiments that apply their skills and knowledge to solving “real world” problems, of which their own body and its health and welfare are the closest. Hence the above experiment can be easily made more attractive to students if the thiocyanate ion concentration measured is from human saliva. The present experiment is divided into two stages. For qualitative analysis few drops of saliva (each student is using his or her own saliva) is treated with a drop of acidic Fe(NO3)3 solution. The deep blood-red color of FeSCN2+ complex is clearly demonstrated. Then each student measures his or her saliva thiocyanate ion concentration with visible spectrophotometry. 3+
{
Experimental Procedures
Stock Solutions Prepare the following solutions in 100-mL volumetric flasks: 0.200 M Fe(NO3) 3 in 1 M aqueous HNO3 and 2.00 × 10{4 M KSCN in water. Qualitative Analysis Collect a few milliliters of saliva in a 25-mL beaker. Transfer a few drops of it to a test tube and add a drop of Fe(NO3)3 stock solution. Observe the deep blood-red color of FeSCN2+ complex.
Quantitative Analysis STANDARD CURVE. For the determination of the molar absorption coefficient of FeSCN2+ complex prepare six solutions in test tubes according to Table 1. Measure the absorbance of each solution at 447 nm against solution #0. The FeSCN2+ concentration in each solution can be calculated by assuming that in these conditions ([Fe3+] >> [SCN{]) all SCN{ ions react to form FeSCN2+ complexes. Plot absorbance as a function of FeSCN2+ concentration. Calculate the molar absorption coefficient of the FeSCN2+ complex from the slope of the standard curve (fitting of least squares). SAMPLE PREPARATION FROM HUMAN SALIVA. Collect about 5 mL of saliva in a 25-mL beaker. Transfer it to a centrifuge tube, and centrifuge at 6000 rpm for at least 5 min. Pipet a known amount of the resulting clear solution (e.g., 200 µL) into a 10-mL volumetric flask and fill the flask with the Fe(NO 3)3 stock solution. Measure the absorbance of the solution at 447 nm against Fe(NO3)3 stock solution. Calculate the FeSCN2+ concentration of the saliva solution according to Lambert and Beer’s law using the above determined ε. Calculate the real saliva thiocyanate concentration taking the dilution into account. PRECAUTIONS AGAINST BIOHAZARDS. Do the experiment with your own saliva. Although the viruses and bacteria in your saliva are probably not hazardous, you should remember the importance of washing the glassware with detergent and hot water and keeping lab areas clean. Whenever it is possible, use disposable centrifuge tubes and pipet tips. It is advisable that saliva of students having contagious infections is not used in the study. If there is any doubt that infected saliva has been used all glassware should be washed by 2 h incubation with 1% phenol or 500 ppm chlorine. Discussion Thiocyanate ion is derived endogenously as a detoxification product of the reaction between cyanide and thiosulfates in the liver. It is abundantly present in blood and especially in saliva (2, 3). Its concentration varies widely depending on diet and smoking habits. Normal nonsmokers have saliva thiocyanate concentrations of 0.5–2 mM; some smokers may have [SCN{] up to 6 mM (2, 3). Hence, a high saliva thiocyanate concentration can be qualitatively used as an indicator for the exposure to tobacco smoke. However, metabolism of vitamin B12 and certain foods containing
JChemEd.chem.wisc.edu • Vol. 76 No. 9 September 1999 • Journal of Chemical Education
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In the Laboratory
cyanide or cyanogenic glucosides Table 2. Measurements of Saliva Table 1. Preparation of Standard (such as nuts—especially almonds— Thiocyanate Concentration Solutions for Determination of e of and cruciferous vegetables) also inGroup FeSCN2+ Complex Range/ creases saliva SCN{ concentration. In N o n s m o k e r s Smokers All V o l u m e / m L T e s t mM addition to detoxification of CN{, (No.) (No.) (No.) Tube Fe(NO3)3 KSCN H2O thiocyanate ion plays an important No. 0.4–1.0 25 0 25 0.200 M 2.00 × 10{4 M role in antimicrobial systems by act1.1–2.0 71 5 76 0 5 0 5 ing as a substrate for peroxidases (4 ). 2 . 1 – 3 . 0 3 1 3 34 1 5 1 4 Also, its capacity to detoxify ingested 3 . 1 – 4 . 0 7 1 8 2 5 2 3 carcinogens has been demonstrated recently (5). All these aspects make the 4.1–5.0 1 2 3 3 5 3 2 discussion of individual saliva SCN{ 5.1–5.6 0 1 1 4 5 4 1 concentration interesting to our stuTotal 13 5 12 147 5 5 5 0 dents. The qualitative and quantitative analysis of thiocyanate ion in human saliva was performed Literature Cited as a part of the exercise to determine the chemical equilib1. Scaife, C. W. J.; Beachley, O. T. Jr. Chemistry in the Laboratory; rium between Fe3+ and SCN{ ions in the laboratory course Saunders: Philadelphia, 1987; p 305. Roberts, J. L. Jr.; Hollenberg, of general chemistry for university freshmen. Naturally, the J. L.; Postma, J. L. General Chemistry in the Laboratory, 3rd ed.; experiment can be done independently as well. Since students Freeman: New York, 1991; p 275. Beran, J. A. Laboratory Manual were using their own saliva, no repugnance to saliva handling for Principles of General Chemistry, 5th ed.; Wiley: New York, 1994; was observed. The results are collected in Table 2. The saliva p 303. Slowinski, E. J.; Wolsey, W. C.; Masterton, W. L. Chemical thiocyanate ion concentration of 147 subjects varied from Principles in the Laboratory, 6th ed.; Saunders: Fort Worth, TX, 0.4 to 5.6 mM. It is clearly demonstrated that cigarette smok1996; p 185. Bishop, C. B.; Bishop, M., B.; Whitten, K. W. Standard and Microscale Experiments in General Chemistry, 3rd ed.; ers (ca. 10% of students) tend to have higher SCN{ concenSaunders: Forth Worth, TX, 1996; p 295. tration than the nonsmokers, although the number of smok2. Tenovuo, J. In Human Saliva: Clinical Chemistry and Microbiology, ers was too small to make a meaningful statistical compariVol II; Tenovuo, J., Ed.; CRC: Boca Raton, FL, 1989; pp 55–91. son.
3. Tenovuo, J. In The Lactoperoxidase System. Chemical and Biological Significance; Pruitt, K. M.; Tenovuo, J., Eds.; Dekker: New York, 1985; pp 101–122. 4. Thomas, E. L.; Milligan. T. W.; Joyner, R. E.; Jefferson, M. M. Infect. Immun. 1994, 62, 529. 5. Hovinen, J.; Silvennoinen, R.; Vilpo, J. Chem. Res. Toxicol. 1998, 11, 91.
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