bk-2013-1130.ch010

Chapter 10

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Quantitative Use of Red Cabbage To Measure pH through Spectrophotometry: A Laboratory Experience for General Chemistry Students Mark B. Cannon* and Kai Li Ong Biochemistry and Physical Sciences Department, Brigham Young University-Hawaii, 55-220 Kulanui Street, Laie, Hawaii 96762 *E-mail: [email protected]

Red cabbage juice is a natural pH indicator: red, purple, blue, green and then yellow color is observed with increasing pH as the conjugated ring structures in the cabbage anthocyanins are rearranged. Classroom demonstrations and home kitchen lab procedures commonly employ cabbage juice extract to qualitatively display acid/base chemistry in a visual manner. We have developed a lab procedure that gives general chemistry-level students a more quantitative experience: by carefully measuring absorbance spectra of cabbage juice samples over a range of pH, students can develop a process for accurately measuring the pH of an unknown solution from its absorbance spectrum in the presence of a small amount of the cabbage extract. Using curve-fitting software, a mathematical function is developed from the absorbance data that can be fit to an equation connecting absorbance and pH.

Introduction Red cabbage (Brassica oleracea) is one of many foods containing natural pH indicator compounds. It has long been known that the characteristic coloring of red cabbage (along with many other fruits, vegetables, and flowers) is due to a cocktail of anthocyanins expressed in the leaves of the cabbage (1). © 2013 American Chemical Society In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.


Anthocyanins are a class of flavonoids that are glucosides of anthocyanidins (see Figure 1) and are the most common water-soluble pigments in the plant kingdom (2). The anthocyanidins of red cabbage are mostly based on cyanidin-3-O-diglucoside-5-O-glucoside (3–7). Anthocyanins also possess antioxidant properties and recent research suggests there may be significant health benefits associated with their consumption (8–10). Because the electronic structure of the anthocyanidin ring system is sensitive to [H+], visible light absorbance is dependent on the pH of the solution (11). Anthocyanidin derivatives from red cabbage appear red at acidic pH and become purple, then blue, then green, then finally yellow as the solution becomes more basic. Thus, the juice of anthocyanin-containing plants is a natural pH-indicator.

Figure 1. General structure of anthocyanidin.

Red cabbage juice is, therefore, very commonly used in qualitative demonstrations of acid/base chemistry because of its very visual connection between pH and color (for example, see references (12–15)). The fact that it is very easy to obtain and also environmentally friendly only increases its attractiveness as a staple of chemistry demonstrations and home instructional lab kits. In this chapter a quantitative method is described for employing red cabbage juice to measure aqueous pH in the framework of a second-semester general chemistry laboratory experience. First, a qualitative analysis of measured absorbance compared with observed color is described. Then, a procedure is described wherein students develop a mathematical connection between measured absorbance and measured pH. This laboratory gives students experience in spectrophotometry, acid/base chemistry, graphical analysis, and the use of curve-fitting software. 130 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.

Methods Juice Extraction


Fresh red cabbage (Brassica oleracea) was obtained from a local grocery store. Approximately 50 g of cabbage leaves were cut up and boiled in about 200 mL of deionized water for 5 minutes. The dark purple-colored juice was allowed to cool, and then was filtered through high flow-rate filter paper. Juice was stored at 4 °C and was stable at this temperature for up to 2 weeks or more. Before each use the juice was subjected to visual and olfactory analysis: if it was cloudy or stinky it was thrown out. Sample Preparation 1 M Sodium hydroxide or hydrochloric acid was added drop-wise to deionized water until the approximate desired pH was achieved, monitored using a standard pH meter. 4 mL of the prepared cabbage juice (above) was added to 21 mL of acidified/basified solution to make 25 mL of total solution for each sample. Additional buffering beyond what was naturally provided by the cabbage juice was not needed as no significant drift in pH was observed over several hours after sample preparation. Because the cabbage anthocyanins gradually (over a few hours) lost color at pH above about 9, absorbance measurements were consistently taken 5 minutes after sample preparation for all samples. pH was measured immediately before and again after absorbance readings to confirm that acid levels remained constant. Samples were prepared at every 1/10 pH unit from pH 1 through 13 for the initial data set used to construct the original model. Students in lab prepared only 8 samples at pH 4, 5, 6, 7, 8, 9, 10, and 11. Spectrophotometry Perkin Elmer Lambda 25 UV/Vis spectrometers were used to measure absorbance spectra of samples 5 minutes after sample preparation. Absorbance was measured from 380 through 700 nm. Vernier SpectroVis Plus spectrophotometers were also used for student data collection. Data Analysis All graphing, data analysis, and curve-fitting were done using Excel (Microsoft) and KaleidaGraph (Synergy Software).

Qualitative Analysis Students first prepare samples of cabbage juice at a range of pH values, from pH 4-11. Then, observing this lineup of brightly colored solutions (see Figure 2), they are asked to predict the range of visible wavelengths being transmitted and the range being absorbed by each sample. Students then measure absorbance for each sample in the visible spectrum and compare their predictions with the data. 131 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.


This leads to a discussion of the physical basis for color because most students, for example, will predict that a red solution is transmitting red light and absorbing all other wavelengths while in fact it is mostly absorbing only light in the green portion of the visible spectrum. A color wheel is useful here: the observed color of light will be found opposite the color wheel from the absorbed color.

Figure 2. Cabbage juice extract at various pH values. The label on each beaker corresponds with the approximate solution pH. (see color insert)

It is easy to see at this point how the cabbage juice can act as a “universal” pH indicator in a qualitative sense: the cabbage juice solutions are red below pH 3, purple from pH 4-6, blue at neutral pH, blue-green becoming green-blue from pH 8-11, and yellow above about pH 12.

Quantitative Analysis: pH from Aborbance Measurements Comparing Absorbance Spectra with pH Figures 3 and 4 illustrate the complex pattern of absorbance as the cabbage anthocyanins equilibrate with more and more basic solution. In very acidic solution a single major peak is visible centered at about 520 nm (causing red color to be observed). As the cabbage juice extract becomes less acidic this peak decreases until about pH 5 when it has become essentially nonexistent. At the same time another smaller peak at 600 nm begins to form (causing blue color to be observed) at about pH 5 and grows until pH 8 or 9 when it begins to steadily decrease and is no longer observed at about pH 12. The last and largest peak begins to form at about pH 7 and continues to increase up through pH 13. This peak is centered in the ultraviolet region of the spectrum and so is inaccessible to spectrophotometers without UV capabilities using UV-transparent cuvettes. Therefore, we ignore the ultraviolet absorbance and treat this final peak as if it were centered at the blue edge of the visible spectrum, at 400 nm (it causes the observed yellow color at basic pHs). 132 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.


Figure 3. Absorbance spectra for cabbage juice solutions at pH 1-13. (see color insert) As these three peaks wax and wane from acidic to basic solution the students observe the primary colors (red, blue, and yellow) as well as secondary colors (purple and green) in between caused by the simultaneous absorbance of two peaks. Figure 5 shows the three peak intensities as a function of pH.

Figure 4. Absorbance spectra for samples at selected pH values illustrating the three main peaks. 133 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.


Figure 5. Absorbance intensity as a function of pH for the three peaks.

Constructing the pH-Absorbance Equation The behavior of the three visible absorbance peaks is clearly dependent on pH. However, no simple relationship between absorbance and pH is apparent. The challenge at this point is to construct a mathematical model that will connect the absorbance intensity of the three peaks with solution pH. For this model to be of maximum value it must be a single equation. Before an appropriate equation can be found the absorbance data must be coalesced into a single function. This can be accomplished by simply adding the three peak intensities together at each pH:

This results in the sinusoidal function shown in Figure 6. However, we require a function with a unique value of f at each pH, so we discard the data below pH 4 and above pH 11 to create the sigmoidal function shown in Figure 7. Curvefitting software (we used KaleidaGraph) can now be used to “fit” the function to a mathematical 2-dimensional equation if the correct form of the equation is known. This author discovered an appropriate equation mostly by trial and error and because the sigmoidal curve in Figure 7 is reminiscent of a Hill plot used in biochemistry to describe cooperative ligand binding such as oxygen binding to hemoglobin (16, 17). 134 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.


Figure 6. The function f, constructed by summing the three peak intensities at each pH value as in Equation 1.

Figure 7. The function f has a sigmoidal shape from pH 4-11. 135 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.

Starting with the Hill equation:


the sigmoidal absorbance function (f) can be fit to the following equation:

where m1, m2, and m3 are constants that are assigned by the software as it fits the data to the equation. Figure 8 shows the original data set fit to this equation along with the assigned values for m1 through m3. Equation 3 can then be rearranged to solve for x (i.e. pH) given an experimentally measured value for y (i.e. f):

Figure 8. Kaleidagraph curve-fit of sigmoidal function f to Equation 3. 136 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.


Figure 9. The difference between measured solution pH and calculated pH from Equation 4, given as a function of pH. A linear regression line is given for the data between pH 6.5-9.5 where pH prediction is most accurate.

Figure 10. Student data fitted to Equation 3. 137 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.

Feeding the original values for f back into this equation gives good predictions of pH, at least between about pH 6.5-9.5, with predicted pH values only differing from actual pH values by an average of 0.11 pH units in this range (see Figure 9). Above and below this range the values for f have insufficient resolution to allow them to be assigned to unique pH values.


Student Data An example of student data is shown in Figure 10. The data fits well to Equation 3: the correlation coefficient is 0.997. Using the resulting values of m1 through m3 to solve for pH results in good predictions of pH, at least between about pH 6-9.

Conclusions We have developed a method whereby students can build a mathematical function that will allow them to accurately predict solution pH from measurements of visible absorbance (within a limited pH range). This laboratory module was developed to give students at the general chemistry level experience with spectrophotometry and its relationship to observed color and pH measurements while providing some valuable experience with graphical analysis and the use of curve-fitting software. Students enjoy the very visual quality of this lab and the connections they are able to make between scientific measurements and everyday observations. The use of curve-fitting software is generally not a part of most General Chemistry laboratory programs. However, this initial exposure should help to prepare them for more advanced applications of graphical analysis.

Acknowledgments We are grateful to Jason Oswald and Concordia Lo for kindly providing the student data.

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139 In Using Food To Stimulate Interest in the Chemistry Classroom; Symcox, K.; ACS Symposium Series; American Chemical Society: Washington, DC, 2013.

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