[an error occurred while processing this directive] Volume 16, Number 2, May 1994, pp.14-18
This paper reports some of the work done by the authors in an attempt to find an alternative method for measuring acidity of carbohydrates directly and nondestructively. The title paraphrases a question asked during this work by the conservator (NL) of the scientist who came up with the idea for the project (HB). The necessary general laboratory work in preparation for the main experiments provided ample opportunity to examine various uses and properties of pH, which will serve as examples of some of its advantages and disadvantages in this article. Additionally, NL's interest in pH measurement was stimulated by her experiences in this preparatory work, and she investigated the use of pH for paper acidity determinations, after consulting with other conservators on typical uses.
To avoid misinterpretations of our statements, we will provide short, simplified definitions of the important terms used in this discussion. At the end of this article, we will discuss some connections of this work to paper permanence and durability, and then summarize the motivation behind it, and any conclusions that can be drawn from our observations.
Many of the important terms in this paper are commonly used in discussions of paper durability and permanence, as well as other conservation issues. To assure that our readers have the same mental picture of the concepts, and to avoid the necessity of hunting down definitions in the literature, we provide definitions here.
By acidity we mean a quality or state of something, namely the "system" we choose to observe. This system may be some liquid, like water used in washing paper or textile artifacts, or the paper or textile artifacts themselves, or artifacts and water together. We will provide a description of the system we refer to in every example. The quality of acidity refers to a sour taste or chemical acidity, where the latter statement will need clarification in the following paragraph.
Some long-known, commonly observed qualities led scientists to the definition of a class of compounds called acids and behaviors termed chemical acidity. Some of the important qualities and behaviors are a sour taste, turning blue plant dyes red, precipitating sulfur from sulfide solutions, and accelerating certain reactions. A technical definition that will explain most of the properties this article deals with, is provided by Bronsted and Lowry, who say that an acid is a substance whose molecules dissociate into a hydrogen ion (proton), and a conjugate base. Since a hydrogen ion cannot exist alone, it will instantly react with another molecule of another or of the same substance to form its conjugate acid (Bell 1973).
A buffer is a substance that maintains constant acidity in a solution. Buffers are mixtures of weak acids and their salts. A buffered system is different from a system having an alkaline reserve. The latter simply means that the acidity of the system is lower than some set limit, and the distance from that limit is the extent of alkaline reserve.
pH is theoretically defined as the negative decadic logarithm of the hydrogen ion activity. This definition has no practical utility, as hydrogen ion activity depends on all other substances in the system to be evaluated. Practically, the best one can do is to use the operational definition of pH, as adopted by the National Bureau of Standards (Bates 1973, Weast and Astle 1978). This definition uses the electromotive force generated by a galvanic cell consisting of a standard solution, the solution to be tested, a salt bridge of saturated KCl solution connecting these two solutions, and the hydrogen electrode. The pH is the sum of the standard solution pH, and this electromotive force.
The crucial properties of this acidity scale are the use of electromotive force of the hydrogen ions in the solution being tested, the use of a comparison (standard) solution, and the need for a quite complicated electrical circuit in the measurement. The last two properties restrict the applicability of this acidity scale to the acidity comparisons of very similar solutions. As can be seen from examples given in the literature (e.g.. Bates 1973), comparison of solutions with and without additions of organic solvents like alcohol, or the presence of other neutral additives like sugars invalidate the use of this acidity scale. Also, the definition clearly is not valid for any system that is not a dilute aqueous solution.
The problems just mentioned are exacerbated by the utilization of electromotive force (voltage) of a galvanic cell as the "yardstick." The galvanic cell used for these measurements always contains junctions between different solutions, implemented usually by placing one of the solutions into a glass vessel capped by glass filters with very fine pores. These junctions often cause problems when the solution to be measured contains particulate matter, dissolved gasses which may form bubbles at the junctions, or ions that set up potentials across the junctions.
Other acidity scales may readily be defined based on any other measurable effect of acidity. One commonly used acidity scale is based on the color changes of indicator molecules. A simple application of an acidity scale based on color changes is pH paper, the name being justified only when the paper is used to measure dilute aqueous solutions, in which case the two acidity scales coincide. If spectrophotometers are used to determine the concentrations of protonated (having taken up a proton) and unprotonated indicator molecules exactly, a very precise acidity measurement results, which leads to the so-called Hammett Acidity Function (Rochester 1970; for a brief review, see Berndt 1991).
Using the Hammett Acidity Function, an acidity scale can be based on any family of similar molecules that change their absorption characteristics towards any type of electromagnetic radiation upon protonation or deprotonation. If, for instance, a carbohydrate molecule would absorb different wavelengths of infrared radiation, depending on the acidity of its environment, that molecule can be used to establish an acidity scale. It is necessary, however, to find a baseline or reference for this scale. Such a baseline can be provided by the absorption characteristics of the molecule of interest, in a well-defined standard environment. The experiments reported in the following, were aimed at finding a suitable carbohydrate molecule, and a standard environment.
We wanted to find out whether gum arabic could be used as an acidity indicator, since it contains many glucuronic acid building blocks, whose carboxylic acid group is very visible in the infrared spectra. That functional group also has different infrared spectra, depending on whether it is present in the acid form, or in the salt form. As a standard environment, we planned to use buffer solutions of phosphoric acid and its various salts, because they are good buffers that can be adjusted over a wide range of pH values. These solutions, mixtures of 0.0667m monobasic sodium phosphate and 0.0667m dibasic sodium phosphate, have very nearly the same ionic composition and strength. Therefore, they approach the ideal situation of changing only the pH of the environment as closely as practically possible.
In preparation for the main experiments, we needed to characterize the buffering behavior, and other characteristics of the phosphate buffers, and of gum arabic. We found the results obtained instructive on the concepts of acidity, and the use of pH measurements. We hence report them here in some detail.
The buffering effectiveness of buffer solutions can be characterized in two ways, by the buffer capacity, and by the dilution value. The buffer capacity is the amount of acid needed per unit change of pH. It is obtained by titration of the buffer solution, i.e., by adding small amounts of acid or base to the buffer solution, taking the pH after every addition, and plotting the two values against each other. The results for our phosphate buffer solution is shown in Figure 1. A good buffer will change pH slowly on addition of acid or base. Figure 1 shows that phosphate buffers have regions of good buffering capacity, and regions of less good buffering capacity. The regions of good buffering correspond to mixtures of equal parts of one of the sodium phosphate salts (there are three), and its conjugate acid, e.g. 1/2 NaH2PO4 solution and 1/2 Na2HPO4 solution, corresponding to the first region of slowly increasing pH in Figure 1.
The "buffer curve" in Figure 1 is a typical example of a titration curve. Following acid-base titrations by pH measurements, is a very good and safe use of pH. At any point, one compares only very similar solutions with each other. Often, one is not interested in the absolute pH, but in the rate of change of pH. For instance, the very fast change of pH near 37 ml of added 0.1 n NaOH indicates that all NaH2PO4 has been converted to Na2HPO4. From the amount of 0.1 n NaOH needed to achieve that conversion, we can calculate the amount of NaH2PO4 initially present in the sample used for titration.
Another way of characterizing buffer solutions is the dilution value. The dilution value is the difference in pH between the buffer solution, and a sample of the buffer solution diluted with equal parts of distilled water. Table 1 shows dilution values for five different phosphate buffer solutions. On adding distilled water, one would expect the pH to move closer to 7, the nominal pH of pure water. The table's examples indicate that this expectation may not always be fulfilled, even in the present, simple system close to the "ideal" conditions required by the definitions of pH.
The buffer solution in the second row of the table may just be a good buffer, and not change its pH on dilution. The buffers in the last three rows, however, clearly violate expectations. The pH of the distilled water added in diluting had a measured pH of 6.2 to 6.6, typical for water in contact with air, and probably in equilibrium with the carbon dioxide in the air. We discussed this unusual rise in pH, which was very reproducible, with other chemists in the Forest Products Laboratory, but could not find a good explanation. Some possibilities are that, even at the low concentrations we worked with, molecular interactions take place that may have shielded some of the basic functional groups of the dibasic phosphate in the buffer solution. Another possibility is that the phosphate ions created a junction potential at the pH electrode. In any case, this example shows that even this simple buffer solution is not readily characterized by using pH measurements.
Another useful application of pH measurements is the acidity determination according to Ingruber (1958). The idea behind this method could be paraphrased as "what would the pH of this system be if it had a pH?" It is a rather laborious, but instructive, means of characterizing the acidity of any system that will interact with dilute aqueous solutions (unbuffered) of strong acids or bases. The assumption behind the method is that any system interacting with an acid or base solution, will change the acidity of that solution, depending on whether the acidity of the system is greater or smaller than the acidity of the solution. A solution with the pH corresponding to the acidity of the system being tested would not change at all. Practically, one adds equal amounts of samples of the system one wants to test, to a series of unbuffered solutions with varying pH values, and then follows the change of pH of the solutions over time.
An example of the method is given in Figure 2. We prepared eight dilute solutions of hydrochloric acid and sodium hydroxide respectively, having pH values of 1.0, 3.0, 4.5, 7.3, 7.7, 8.5, and 12.7. We then added 2.5 g of gum arabic per solution, and followed the change of pH of the solutions over time, until they remained stable. Figure 3 shows the results of this experiment plotted as final pH values vs. initial pH values. According to Ingruber, the intersection of the line connecting the final pH values with the line of final ph=initial pH, gives the pH value of the solution whose pH would not have changed at all on addition of the sample. This then corresponds to the pH the system would have.
The experimental setup for the Ingruber acidity determination provided an unexpected example of the importance of an environment's chemistry on biological processes. One of the unbuffered solutions with gum arabic was profusely growing molds. The plant pathologist at the FPL (Wilcox 1992) identified the mold tentatively as "penicillium and others." Mold spores are present almost everywhere, so the fact that mold grew on a medium containing a carbohydrate (gum arabic) isn't too surprising. The fact that only one of eight very similar environments had mold growth, is. It appears that just the right combination of nutrient content, acidity, and osmotic pressure, was provided by this particular solution.
From the operational definition of pH given above, it is clear that paper does not, strictly speaking, have a pH. Only solutions equilibrated with paper can have a pH. There are several more or less standardized ways of preparing solutions in contact with paper for pH determinations, all of which can be expected to yield different answers. To illustrate this point, NL collected several samples of historic papers, and one contemporary archival blotting paper, and applied several methods of pH measurement to these samples.
Figure 3 shows the results of these measurements. The papers used were (1) archival blotting paper, Paper Technologies, Inc.; (2) and (3) leaf of a book, ca. 1729-1742, (2) with image, (3) plain; (4) paper from Edmeade & Pine, 1804; (5) paper from a book, ca. 1700; (6) endleaf of a book published in 1654, but may be a later addition. NL obtained the leftmost set of values by following the TAPPI standard T 509 om-83. Regarding the time of equilibration, this method only specifies that "the specimen may stand for 3 to 4 h." Figure 3 shows that the time of equilibration may have a significant effect on the pH of the cold water extract. Furthermore, the type and extent of this effect depends on the type of sample, and cannot be known in advance. The change in extract pH with time cannot be explained by uptake of CO2 from the air, since in most cases, the pH increased with time.
The surface pH also measures the pH of a solution in contact with paper, hoping it will tell us something about the acidity of the paper itself. This method, too, is standardized by TAPPI. Practical concerns like water absorption by the paper, however, make it much more difficult to control the amount of water used, and hence the concentration of the solution being tested. As the rightmost values in Figure 3 show, the surface pH differs significantly from the cold extract pH values. Also, they differ from these values in an unpredictable manner. The indicator paper pH values are surface pH values, using indicator paper, rather than a pH electrode for measurement. The results reflect the limited accuracy of indicator paper.
There are reasons why conservators should be concerned with the acidity of objects, and environments they work with. One example given above is the interaction with biological processes: the right acidity can promote the growth of biological deterioration agents. Another concern is stains or color changes: from the use of acid indicators we know that the acidity of an environment can change the color of some pigments. As the authors found out in discussions with other conservators, historical papers that have been given an alkaline wash sometimes turn pink or show pink spots.
The effect of acidity most frequently and vehemently discussed presently, is deterioration of paper. The "slow fire" destroying reams of "acid paper" every year. Mass "deacidification" treatments are being tested in the hope of finding a cure for the rapid deterioration of papers manufactured in the last 130 - 150 years. Certain assumptions are necessary for the recently popular argument that deacidification will arrest paper deterioration. These can be formulated as follows:
Assumption 1: the rate of deterioration is equal to the rate of cellulose acid hydrolysis.
Assumption 2: the rate of acid hydrolysis is proportional to the acidity of the paper.
Assumption 3: the acidity of paper can be measured by pH.
Regarding the typical test of the efficacy of various deacidification treatments, one could add the assumption that the decrease of paper useability is proportional to the decrease in folding endurance.
From the definitions and examples presented here, we conclude that assumption 3 is unwarranted. Even following the appropriate TAPPI standard may not give good comparisons between different types of paper, as the measured pH changes significantly with time, and does so quite differently for different types of paper. Going back to the operational definition of pH we have to ask which solution in contact with the paper is representative for all papers we would like to compare. From our point of view, only the Ingruber acidity seems to have a wide enough applicability.
From published research, assumption 2 appears correct. For any paper of a given composition, the rate of hydrolysis will increase with increasing acidity. It is very difficult, though, to change only the pH/acidity of a particular system, as we tried to demonstrate here by presenting some preliminary experiments on designing such a system. It requires great care to separate effects of, for example, the ionic composition of the system from the effects of acidity. In a thorough investigation of cellulose hydrolysis Harris (1975) found that the rate of acid hydrolysis is proportional to the Hammett acidity function of the system. However, the proportionality constant depended on the type of acid used.
The interpretation of assumption 1 depends strongly on the way deterioration is quantified. If the amount of deterioration is measured by the amount of soluble carbohydrates created, it is most probably true. Any other measure of the rate of paper deterioration, would have to be compared to the amount of soluble carbohydrates created, to show its direct proportionality to the rate of acid hydrolysis.
To answer the title question glibly: there is nothing wrong with pH, it's just the way it's (sometimes) handled. pH is a very useful measure when comparing dilute solutions of strong acids and bases. It is also very useful for comparing very similar systems, like the various process streams in pulp and paper mills, or other manufacturing processes where conditions in general are controlled, and vary relatively little. It also works very well in acid-base titrations, because one follows the pH changes of a slowly varying system. The use of pH in the Ingruber method of measuring acidity is also a good application, because the measured systems are dominated by dilute, unbuffered solutions of strong acids and bases, and because one is mostly interested in the change of pH of any one of the solutions. Comparing the pH of papers of widely varying properties, like fiber composition, filler, additive, or sizing contents, is meaningless in our opinion.
Our attempts of finding an alternative means of measuring the acidity of carbohydrates produced mixed results. It appears that the carboxyl group in gum arabic indeed changes its infrared absorption behavior, when the acidity of the system gum arabic + phosphate buffer is changed. This infrared absorption change, however, occurs in a narrow range of acidities at the lower limit of the acidities we were able to study with our buffer. HB hopes to be able to continue this investigation.
Bates, R.G. 1973. Determination of pH, theory and practice. Second edition. New York, NY: Wiley-Interscience.
Bell, R.P. 1973. The proton in chemistry, 2nd edition. Ithaca, NY: Cornell University Press.
Berndt, H. 1991. Acidity: a review of fundamentals. The 1991 Book and Paper Group Annual, pp. 1-10.
Harris, J.F. 1975. Acid hydrolysis and dehydration reactions for utilizing plant carbohydrates. Applied Polymer Symposium 28: 131- 144.
Ingruber, O.V. 1958. The behaviour of wood and wood constituents as acid-buffering systems. Pulp and Paper Magazine of Canada 59(11): 135-141.
Rochester, C.H. 1970. Acidity functions. In: Organic Chemistry, a series of monographs, vol. 17. New York, N.Y.: Academic Press.
TAPPI. 1983. TAPPI method T 509 om-83, Hydrogen ion concentration (pH) of paper extracts (cold extraction method). Technical Association of the Pulp and Paper Industry, Atlanta, GA.
Weast, R.C. and M.J. Astle, eds. 1978. CRC handbook of chemistry and physics, 59th edition. Boca Raton, FL: CRC Press, Inc.
Wilcox, W.W. 1992. Personal communication.
Figure captions (figures not included in electronic version)
Figure 1: Buffer curve of 0.0667m monobasic sodium phosphate.
Figure 2: pH of 25 ml of unbuffered solutions of HCl and NaOH after addition of 1 g of gum arabic (for the determination of Ingruber acidity of gum arabic).
Figure 3: pH of paper, a comparison of three methods applied to six papers of different types (for identification of the papers, see text section "pH of paper").
Table captions (table not included in electronic version)
Table 1: Dilution values of phosphate buffers (mixtures of 0.0667m monobasic sodium phosphate and 0.0667m dibasic sodium phosphate).
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