The Chemistry and Technology of the Pretreatment and Preservation of Fruit and Vegetable Products with Sulfur Dioxide and Sulfites

The Chemistry and Technology of the Pretreatment and Preservation of Fruit and Vegetable Products with Sulfur Dioxide and Sulfites BY M. A. JOSLYN

AND

J.

€3.

S. BRAVERMAN*

Ijepurtnaent of Food Technology, University of California, Berkeley, California

CONTENTS Page 97 99 113 123 128 130 134 136 141 143

I. Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11. Chemistry of Sulfur Dioxide, Sulfites, and Their Organic Compounds. . . . 111. Determination of Sulfur Dioxide in Fruit and Vegetable Products. . . . . . . IV. Sulfur Dioxide as Preservative and Sanitizing Agent.. . . . . . . . . . . . . . . . . . V. Sulfur Dioxide as an Inhibitor of Enzymic and Nonenzymic Browning.. . VI. Source and Application of Sulfur Dioxide.. . . . . . . . ... VII. Sulfur Dioxide in Fruit Juices, Syrups, Concentrates and PurBes . . . . . . . . VIII. Sulfur Dioxide in Wine and Vinegar Maki .......... IX. Sulfur Dioxide in Dehydrated and Dried F Products.. X. Sulfur Dioxide in “Brining” of Cherries and “Barrelling” of Fruit. . . . . . XI. Sulfur Dioxide in Transportation and Storage of Grapes and in Other Products . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

146 147

I. INTRODUCTION Sulfur dioxide, in its gaseous or liquid form, or dissolved in water t o form sulfurous acid, or in the form of its neutral or acid salts (sulfites, bisulfites, metabisulfites) is used widely as a chemical preservative t o reduce or prevent spoilage by microorganisms, and as a selective inhibitor of undesirable organisms in the fermentation industries. It is also used as an antioxidant and inhibitor of enzyme-catalyzed oxidative discoloration and of nonenzymic browning during preparation, storage, or distribution of many food products. As an antioxidant it is useful also in improving the retention of ascorbic acid, carotene, and other oxidizable biologically active components. The preservative value of sulfurous acid for color retention and its inhibition of microbial growth and activity have long been recognized, and i t has long been used for the control of undesirable changes in color and flavor during processing and as a relatively cheap and readily available preservative for a wide variety of products-juices, * Dr. Braverman now is Professor of Food Technology at Technion, Israel Institute of

Technology, Haifa, Israel. The authors gratefully acknowledge the assistance of Dr. C. J. B. Smit, now a t Stellenbosch University, South Africa, for preparing the graphs used in the text and for assisting in compiling the literature. 97

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M . A. JOSLYN AND J.

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S. BRAVERMAN

syrups, concentrates, purBes, fermented beverages, dehydrated and sundried foods, in preserves, jams, and marmalades, in brined fruits and pulps for subsequent processing, and in the storage and distribution of perishables (fresh fruits, vegetables, meat, and fish), Its pronounced bleaching action has been employed in the preparation of specialty products such as Maraschino cherries, citron and citrus peel, watermelon rind, and Zucca melon. The convenience with which it can be used and the ease with which its concentration can be reduced to a level low enough to be tolerated in certain products make it peculiarly attractive. Sulfur dioxide has been used in food preservation since ancient times. There is evidence that the use of the fumes of burning sulfur in wine making was known to the early Egyptians and Romans (Bioletti, 1911). The old practice of sulfuring was strictly empirical and based on secular experiences and customs. Its correct use was possible under these conditions only in a few cases and in the hands of wine makers of long experience. Its abuses were many and were recognized as early as 1738 in France, according to Brhmond (1937). It was not until the beginning of the last decennial of the nineteenth century that serious study of the effects of sulfurous acid on must and wine were undertaken in France. Shortly afterwards, in 1902, the use of sulfur dioxide in wine making in France was regulated (BrBmond, 1937). Sulfuring of fruits for drying was an established practice in the late nineteenth century. Although the practice of sulfuring fruit for drying was widely used in California in the late nineteenth century, it apparently was not recognized abroad. In 1902, Beythien and Bohrisch discovered that practically all the dried fruit imported into Germany from America was heavily sulfured and this observation was followed by publication of numerous papers in which the admissibility of sulfur dioxide as a preservative was actively debated (Beythien and Bohrisch, 1902, 1903). It was used for preserving meats in the United States as early as 1813 and somewhat later for fish. So widespread was the use of sulfur dioxide in foods in the nineteenth century that Wiley (1907) warned against its promiscuous use as a preservative (Anon., 1907). Sulfurous acid was used almost from the beginning in the purification of sugar-beet juices in the European beet sugar industry (Maxwell, 1916; McGinnis, 1951). I t was introduced by Proust in 1810 in the form of calcium sulfite, but since 1858 sulfur dioxide gas has been used for decolorization of cane and beet juices. It was first tested with cane juice in Mauritius in 1865 and introduced into Java in 1895 (Marches, 1953). In the period 1935-1940, 40% of the white sugar produced annually in Java was sulfitation sugar. In the purification of sugar-containing juices it is used for neutralization of excess alkalinity, for the decolorization of

SULFUR DIOXIDE TREATMENT OF FRUIT AND VEGETABLE PRODUCTS

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the extracted juices by reduction of naturally occurring pigments and colored decomposition products, and for prevention or inhibition of color formation in later stages of processing, evaporation, and crystallization. It has been claimed t o facilitate crystallization. The sulfitation process in cane-sugar and beet-sugar technology is described in detail in the texts on sugar technology edited by Honig (1953) and McGinnis (1951). Investigations of the scientific basis of the use of sulfur dioxide and sulfites in the preparation and preservation of foods have been carried out on wines early in the twentieth century and are still being conducted both in the United States and in Europe. Extensive investigations of sulfuring practices and problems were conducted by Long et al. (1940), during the period of 1936-1939. The widespread use of sulfur dioxide in the dehydration of vegetables during World War I1 in Britain and British Dominions and in the United States focused attention anew upon the problems involved. The intensive investigation of the fundamental nature of browning and discoloration in processed foods (Mitchell and Peterson, 1952) also resulted in support of investigations on the chemistry of the interaction of sulfite with sugar. In spite of its long history of use and the many factors influencing its applicability our knowledge of the chemistry and technology of sulfurous acid and sulfite pretreatment and preservation of foods is still incomplete. The chemistry of the products formed and the state of sulfurous acid in foods, the reactions involved in their formation, the changes in the initial products formed during storage, and the actual nature of the preservative and antibrowning effect still have t o be established. The present status of our knowledge in this field is summarized in the following review, which is restricted to foods of plant origin. Only the more important literature in the field is reviewed and the references cited are not complete. Where they are available, reference is made t o review articles and all phases of the field are not reviewed in the same detail. 11. CHEMISTRY OF SULFUR DIOXIDE, SULFITES, AND THEIRORGANIC COMPOUNDS 1. Sulfur Dioxide and the Sulfites

Sulfur dioxide formed by burning sulfur in air also contains between

6-8 % of sulfur trioxide, whose presence accounts for the ‘(foggy” appear-

ance of the gas, the fog being due t o droplets of water condensed about SO3. The sulfur trioxide present in sulfur dioxide may be determined readily by suitable filtration (Eckman, 1927). The formattion of sulfur trioxide is promoted by high oxygen content, low temperature of the air stream in sulfur furnaces, and presence of catalysts such as iron oxide.

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M. A. JOSLYN A N D J. B. S. BRAVERMAN

Complete combustion of sulfur in sulfur furnaces used in the sugar industry or in the sulfur houses used in the dried fruit industry which would theoretically produce gas with 21 volume per cent of SO2is not obtained. According to Marches (1953) the gas produced in the sulfur furnaces used in Java contains 6-14 volume per cent of SO2 and small amounts of SOa.

PH

FIG.1. Distribution of various constituents of sulfurous acid at various pH values. After Vas and Ingram (1949).

In more efficiently designed and operated sulfur burners such as those used in the paper industry and in the manufacture of sulfuric acid (Shreve, 1944), the sulfur is melted, atomized in compressed air, and burnt in a separate combustion chamber. In these burners gas of a constant composition, 19-20 volume per cent SO2, without any sublimation and with only 0.14%of the sulfur transformed into SOi, can be obtained. Sulfur dioxide is a colorless gas having a characteristic odor, a normal molecular volume of 21.89 l., and a molecular weight of 64.06 g. It is soluble to the extent of 36.4 volumes in one volume of water at 20" C. Its solubility in water decreases from 8.6% by weight a t 20" C. t o 0.1% at 100" C. At atmospheric pressure SO2liquefies a t - 10" C.; a t 20" C. liquid SO2 exerts a pressure of 3.25 atm. or 40.6 p.s.i. The solubility of sulfur dioxide in water has been determined accurately by Beuschlein and

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101

Simenson (1940), and the thermodynamic data for the system S02-Hn0 are given by Plummer (1950).* In water sulfur dioxide exists as the dissolved gas, as undissociated sulfurous acid, H2S03,as the bisulfite ion, HSOa-, and as the sulfite ion, so3--. Sulfurous acid is a rather weak dibasic acid; its first ionization constant is 1.7 X and the second is 5 X loF6.At 25" C. the corresponding pK' values are 1.8 and 5.3, respectively. The calculated distribution of the various ionic species in aqueous solution of sulfurous acid in the range of pH of 0 to 8 is shown in Fig. 1. At pH's over 9.5 only SO,- ions exist, in the range of pH 9.5 to 4.5 both SO3-- and HS03- occur, and at pH 4.5 and lower, SO3-- no longer exists in appreciable concentrations. The alkali sulfites are but slightly hydrolyzed. The sulfites of many of the heavy metals are insoluble. The sulfites occur as the normal salts, NazS03, as the acid sulfites, NaHS03, and as the metabisulfites, Na2S206. The latter is the anhydride of the acid sulfite:

The various sulfites are particularly adaptable as sources of sulfur dioxide because of the convenience of handling as dry chemicals and the ease with which their solutions can be prepared. They are best when added to acid products of fairly high acid content so that possible reduction in total acidity or increase in pH as a result of liberation of sulfurous acid is avoided. This may be avoided also when they are used as solutions by the addition of a slight excess of inorganic acid such as HC1 or HzS04. The theoretically available sulfur dioxide content and the usual range of SO2 content for various preparations is shown in Table I. During storage, particularly when exposed to moist conditions, the sulfites decrease in available SO2 content as a result of oxidation. In the dry state stored

* Plummer

(1950) made a complete survey of the literature on the equilibrium vapor pressure exerted by the system SOZ-H~O,and statistically correlated such data for SO2 concentrations of 0.03 g. to 10.0 g. of SO2 per 100 g. H20, at temperatures of 0" to 130" C. The correlated results were presented as straight lines on a n Othmer-Cox-type chart by modifying the ordinate. He found that the total pressure P, = (P, P,) could be presented as a linear equation log,, P , = b log,, P , a and tabulated values of the over-all slope b and the over-all intercept a for various concentrations of SO*. Whitney and Vivian (1949) and Whitney et al. (1953) reported data on the mechanism of absorption of sulfur dioxide gas in water, sodium carbonate, and sodium sulfite solutions. This is basic not only to the manufacture of sulfite cooking liquors used in the paper industry, but also in other industries where sulfur dioxide gas is absorbed in liquids. I n the absorption of sulfur dioxide gas in water, both gas and liquid film resistances are involved. I n strongly alkaline solution the resistance is primarily in the gas film, and in sulfite liquors it, is primarily in the liquid film.

+

+

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M. A. JOSLYN AND J. B. S. BRAVERMAN

exposed to air in order of increasing stability they are: sulfite > metabisulfite (Mason, 1928; Phillips, 1928).

> bisulfite

TABLBI Available Sulfur Dioxide Content of Various Sources Available per cent of SOa Compound

Formula

Theoretical

Actual

Liquid sulfur dioxide Acid sulfurous, 6% Calcium sulfite Potassium sulfite Sodium sulfite Potassium bisulfite Sodium bisulfite Potassium metabisulfite Sodium metabisulfite

SO2 HaSOa CaSOa.1 f 5 H z 0 KzSOa Na2SO3 KHSOa NaHSOa

100.0 6.0 23.0 33.0 50.84 53.31 61.59 67.43 57.65

100.0 6.4-6.8 43-45 36 48

KzSz06

Na2Sz06

-

55 52 61

Antioxidant and Reducing Properties. Sulfurous acid and sulfites in solution are oxidized by air to sulfates. The rate of oxidation of sulfite ions by oxygen has been shown by a number of investigators to be a chain reaction of unusual length for condensed systems and, as is frequently the situation in such cases, to be very sensitive to both positive and negative catalysis. The early investigations on oxidizability of sulfite solutions are discussed by Kolthoff and Menzel (1928). Fuller and Crist (1941) reported the reaction of sodium sulfite solutions saturated with oxygen at one atmosphere pressure to be strictly first order with respect to sulfite ion concentration with a specific reaction rate of 0.013 see.-' at 25" C. The rate of oxidation is independent of the pH between 8.8 and 8.2 but decreases in a complicated manner between 5.9 and 3.2. The rate is directly dependent upon the cupric ion concentration when this exceeds M and the catalytic constant is 2.5 X lo6 liters moles-l set.-' at 25" C. The inhibitory effect of mannitol, however, is uniform over a 106-foldchange in mannitol concentration. Mitchell et al. (1933) reported the oxidation of sulfurous acid and sulfite t o be reduced by decreasing pH and by addition of various antioxidants. The inhibitory action of sulfite in the auto-oxidation of several phenols has been established. Branch and Joslyn (1935) reported that the addition of 0.001 M of potassium sulfite to 0.1 M catechol a t pH 8.3, reduced the rate of oxidation from 3.3 to about 0.2 initially, but after about 4 minutes the rate increased t o that normal for the conditions used. James and Weissberger (1939) reported that, in the auto-oxidation of hydroquinone and its homologs in the presence of excess sulfite, these substances react with oxygen to form quinones and hydrogen peroxide. The latter then oxidizes the sulfite to sulfate. If the quinone formed has a t least one

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103

nuclear hydrogen it reacts with sulfite to form the hydroquinone monosulfonate which auto-oxidizes and finally forms the disulfonate. The occurrence of a sulfonic acid of this type is quite significant in view of the finding by Stadtman (1948) that the sulfur dioxide which disappears during the storage of sulfured dried apricots in the absence of oxygen is converted into a relatively strong polyhydroxy acid containing 10 % sulfur. Although this compound has not been identified, Stadtman postulated it t o be a sulfonic acid of partly unsaturated sugars. Sulfurous acid and sulfites also inhibit the auto-oxidation of ascorbic acid. Some of the earlier investigators, e.g., Williams and Corran (1930), were of the opinion that sulfur dioxide, when used as a preservative for fruit juices, had a definite destructive action on ascorbic acid a t normal temperatures (15 t o 18" C,).Bennett and Tarbert (1933) reported that sulfur dioxide exerted only a slight destructive effect on ascorbic acid in comparison with the marked destruction occurring in benzoated juice which they wrongly ascribed to the benzoate itself. Hamburger and Joslyn (1941), however, found no evidence for an induced oxidation of ascorbic acid in citrus juices by sulfur dioxide. They found, on the contrary, that sulfur dioxide inhibits the oxidation of ascorbic acid in citrus juices and that this inhibition was more pronounced a t 500 p.p.m. than 250 p.p.m. of SOz. Joslyn (1941) and Downer (1942) reported that sulfur dioxide definitely inhibited the auto-oxidation of ascorbic acid in citrus juices. Busing and Raabe (1938) reported that in grape juice, also, ascorbic acid was preserved by sulfur dioxide. Morgan et al. (1929, 1931, 1935) found that sulfuring improved retention of ascorbic acid in dried fruits. Sulfurous acid and sulfites, although they serve as anti-oxidants and are useful in inhibiting oxidation of ascorbic acid, have an adverse affect upon thiamine. Sulfurous acid has long been known t o destroy the antineuritic activity of rice polish extracts (Williams et al., 1935b) and of dried fruits, grape juice, and wine. Williams et al. (1935a), Buchman (1936))and Buchman et al. (1935), showed that this destruction of vitamin B1 activity was due to the cleavage of thiamine into 4-methyl-5 hydroxy ethyl This thiazole and the sulfonic acid of 2,5-dimethyl-4-amino-pyrimidine. cleavage is complete a t room temperature a t pH 5 in 24 to 48 hours, but is slower at lower pH values. I n saturated aqueous solution of sulfurous acid, 50 % yields of cleavage products were obtained after storage a t room temperature for three months. Morgan et al. (1931, 1935, 1939) showed that sulfiting of fruit for drying and sulfiting of grape juice and wine was destructive to vitamin B1. Bleaching Action. Sulfur dioxide (sulfurous acid and the acid sulfites) reduces many colored compounds to colorless derivatives and is fre-

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M. A. JOSLYN AND J. B. S. BRAVERMAN

quently used as a bleaching agent. Sulfur dioxide has a pronounced bleaching effect on the anthocyanin pigments, and red- and blue-colored fruits are rapidly bleached in brines containing sulfites. It has a slight bleaching action on yellow fruits. Sulfurous acid, however, does not bleach chlorophyll but catalyzes the conversion of chlorophyll to pheophytin. This conversion, however, is considerably slower with sulfurous acid than with sulfuric or oxalic acids of the same concentration (Joslyn and Mackinney, 1938; Mackinney and Joslyn, 1941). Mackinney (1937) found that chlorophyll was present in raisins made from sulfured grapes. Dried cut fruit, such as dried apricots, which have darkened only slightly can be almost completely restored to their original color by treating with SO2 (Jewell, 1937). The nature of this bleaching action is unknown, but it is probably due to the formation of colorless compounds as a result of combination of SO2 with some constituent. The decolorization of basic fuchsin by sulphur dioxide has been fairly well elucidated and is believed to be due to the formation of colorless sulfonates. In the sugar industry it has long been known that sulfur dioxide bleaches not only the naturally occurring pigments such as anthocyanins and other colored nonsugars but also the nonsugars and sugars which develop color in sugar manufacture (Gillett, 1953; Zerban, 1947). Lewis et al. (1949a, b) reported that colored decomposition products formed by heating glucose are decolorized by NaHS03 a t pH 5.5. Bleached syrups exposed to air for some time darken again because of oxidation of sulfite. Sulfur dioxide not only bleaches coloring matter already formed but reduces darkening during evaporation and crystallization. Zerban (1947) suggested that this protective effect was due to combination of sulfur dioxide with reducing sugars, thus blocking the carbonyl function essential for caramel and melanoidin formation. The inhibition of browning by sulfurous acid and bisulfite may be due not only to its reaction with the carbonyl group in sugars but also to its action as reducing agent in keeping reductones in the inactive reduced form rather than in the active dehydro form. Hodge (1953) in his recent review of the chemistry of browning reactions in nitrogenous model systems has stressed the important role of dehydrogenated reductones in both enzymatic and nonenzymatic browning reactions. 2. Bisulfite Addition Products: a-Hydroxysulfonic Acids

Sulfurous acid and the alkali bisulfites have long been known to react with aldehydes and ketones. The sodium and potassium bisulfites form crystalline addition products which are useful in purification but seldom used for the identification purposes, since their melting or decomposition point is not sharp and they are essentially insoluble in organic solvents.

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The bisulfites react with unsaturated organic compounds, their addition \ to -C=Cbonds being almost as rapid as to the C=O groups.

/

Amines, particularly tertiary amines, in the presence of water, with sulfur dioxide readily form amine bisulfites or complexes of amine bisulfites with sulfur dioxide. The amine bisulfites form addition products with aldehydes and ketones which are useful for their characterization as well as their resolution (Adams and Lipscomb, 1949; Adams and Garber, 1949). \ In the reaction of amines with aldehydes Schiff bases containing ‘C=N-

/

groups are formed as intermediates and these could also react with bisulfites. It has been shown recently that glucose sodium bisulfite reacts with amines to form substituted a-amino sulfonates. This reaction, as pointed out by Danehy and Pigman (1951), may be of significance in accounting for the inhibitive effect of sulfur dioxide in the discoloration of food products. Certain alkaloids as well as the primary, secondary, and tertiary amines form bisulfite addition compounds. Only comparatively recently has it been unequivocally established for a few carbonyl compounds that the bisulfite addition products are hydroxysulfonic acid derivatives rather than sulfites (Suter, 1944). The evidence for the hydroxysulfonate structure of aldehyde and ketone bisulfite compounds, the sulfonic groups of which are mobilized by the adjacent hydroxyl group, is now quite convincing. While all aldehydes form hydroxysulfonates with bisulfites, many ketones do not react. With the exception of ethyl ketone, which reacts slowly and to a limited extent, only ketones containing either a methyl group attached to the carbonyl or having the carbonyl part of a ring system of four to seven carbon atoms combine appreciably. The addition of an alkali bisulfite to an unsaturated aldehyde or ketone in which the double bond is conjugated with the carbonyl group is more complex. Citral, for example, forms, in addition to the normal hydroxysulfonate, also a compound two molecules of bisulfite. With unsaturated aldehydes three types of bisulfite addition compounds are possible: addition of -SOaNa to the C of the C=O group, to one of the C= C carbons, or to both positions. The combination of bisulfites with sugars is much slower than with aldehydes and ketones and the products formed are relatively more unstable. Only the glucose and arabinose bisulfite addition compounds have been prepared in pure or relatively pure form. The combination with other sugars is largely inferred from measurements of decrease in “free ” sulfur dioxide content. The early work of Kerp, Schmidt, and others in the laboratories of the German Public Health Department during 1904-

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M. A. JOSLYN AND J. B. 9. BRAVERMAN

1909 (Kerp, 1903, 1904,1913), is summarized by Monier-Williams (1927)) Browne (1944), Berg (1946)) and Ingram and Vas (1950a, b). The results obtained by Ingram and Vas (1950a) confirmed those obtained previously by Kerp (1904) and by Farnsteiner (1904)) that of the sugars tested galactose, mannose, and arabinose combined with bisulfites rapidly and formed compounds which dissociated less; maltose, lactose, and glucose were less active; raffinose was but slightly active and fructose and sucrose were probably inactive. Ingram and Vas (1950a) also supported the early inference of Paris (1920) that pectinic acids combine with bisulfite, but their observations were made on 100 grade citrus pectin containing about one-third its weight of glucose. Sulfurous acid has been reported to combine with dextrins, cellulose, and lignin; proteins and possibly gelatins (MonierWilliams, 1927). The lignin sulfonic acid in sulfite pulp liquor is particularly stable and gives off sulfur dioxide with extreme slowness on distillation with strong acids (Stutzer, 1910, cited by Monier-Williams, 1927). Solutions of glucose, mannose, and xylose when heated a t 135” C. in the presence of sodium-bisulfite-sulfurous acid solutions containing sulfuric acid were found by Hagglund et al. (1929, 1930) t o yield gluconic, mannonic, and xylonic acid by the transformation of the primary formed bisulfite addition complexes. Under the same conditions fructose yielded a stable sulfur compound which was believed to be fructose sulfonic acid. Sugar solutions on heating may yield decomposition products which can combine readily with sulfur dioxide. Ingram and Vas (1950a) list among such substances reported t o occur during the decomposition of sugar solutions heated alone or in presence of added alkalis or amino acids: formaldehyde, glyceraldehyde, 5-hydroxymethyl furfural, furfuralic and other aldehydes, methylglyoxal, pryuvaldehyde, dihydroxyacetone, acetal, diacetyl, and “reductone.” Convincing proof that the bisulfite addition product of formaldehyde is a hydroxysulfonic acid, containing carbon-sulfur linkage, is based on both chemical and physical evidence. The condensation product of the formaldehyde bisulfite addition product and ethyl acetoacetate has been shown to be a true sulfonate; the amino compound derived from formaldehyde, sodium bisulfite, and ammonia is converted by nitrosyl chloride into sodium chloromethanesulfonic acid ; and acetylation of the formaldehyde addition product was found to give an acetate identical with that obtained from potassium iodomethanesulfonate. The physical properties of the hydroxysulfonates are in accord with this structure : their absorption spectra contain a band a t 4992 A. characteristic of sulfonic acids and differing from that of sulfites; the conductivity of solutions of the free acids is high, indicating the presence of typical sulfonic acids; the second

SULFUR DIOXIDE TREATMENT OF FRUIT AND VEGETABLE PRODUCTS

107

hydrogen is weakly acidic having a dissociation constali t of about 7 x 10-1O.Proof of the structure of bisulfite addition products other than those of aldehydes is not as, strong, but Braverman (1953) has presented good evidence for the structure of glucose and arabiriose sulfonates and for the possible mechanism of the interaction involved. The a-hydroxysulfonic acids decompose in either acid or alkaline solutions t o give the original carbonyl compound. In acid solutions a t room temperature the reaction is slow but on heating it is more rapid and the decomposition during distillation from strongly acid solutions is almost complete. It may be represented as follows: H R-

:$I/8

- -0Na

+ HC1-t R-

8+ I

NaCl

+ S 0 2 t + HzOT

H

In alkaline solutions the decomposition of the hydroxysulfonates is quite rapid even at room temperature, largely owing to the fact that neutral sulfites do not combine with aldehydes or ketones. The decomposition in alkali may be represented as follows: H

0 0

€d-i!!-ONa

Adl

+ NaOH

0 +

R-C

II I

H

+ NaO-S-ONa II

+

H20

0

The a-hydroxysulfonates are not oxidized by iodine in neutral or acid solutions, but this does occur in alkaline solutions where it is preceded by dissociation into the free sulfite. They are also not readily oxidizable by oxygen. The formation of the a-hydroxysulfonic acids is influenced by concentration of the reactants (carbonyl compound and bisulfite), temperature, and pH. The rate of association as influenced by the relative concentration of bisulfite, aldehyde, and sugars was investigated early by Kerp (1904) and his collaborators and others, and more recently by Ingram and Vas (1950b) and by the Corn Products Refining Company.' The effect of pH and temperature on dissociation of the sulfonic acids has been investigated also for some of the compounds (see the review by Suter, 1

The Chemical Division of the Corn Products Refining Company investigated the effect of temperature, concentration, alcohol, and free sulfurous acid on the degree of association of glucose and sodium bisulfite in aqueous solution in connection with the research sponsored by the Subsistence Section of the QMC Research and Development Branch during World War I1 on browning.

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M. A. JOSLYN AND J.

B.

8. BRAVERMAN

1944). Most of the publications in the field, however, have dealt with reactions in solution or with partially purified sulfonates. Only very recently have pure sulfonates been prepared and their properties investigated. The characteristics of two typical compounds, the acetaldehyde sulfonic acid and the glucose sulfonic acid, are discussed briefly below. The literature on the sugar sulfonates is reviewed in more detail by Gehman and Osman (1954). Acetaldehyde Sulfonate. Kerp and his collaborators prepared acetaldehyde sodium hydroxysulfonate essentially by the method of Bunte (1873). This consists in the slow addition of acetaldehyde to a concentrated solution of sodium or potassium bisulfite, with continuous cooling until excess of aldehyde is evident to the nose and no further temperature rise occurs. The sulfonate is then obtained by evaporation in a desiccator over concentrated sulfuric acid, or by storing the reaction mixture over night and precipitating the complex with methanol. The impurities present, chiefly the alkali sulfates, are eliminated by recrystallization from warm methyl alcohol at 40” C. (104’ F.). With sulfonates so prepared, Kerp reported the following percentages dissociated into free bisulfite titratable with iodine: 0.17% in 1 N solution; 0.45% with 0.1 N solution; and 0.71 % with $60 N solution. The association of bisulfite with acetaldehyde was investigated by titrating the residual “free ” bisulfite with iodine. The degree of association increases as the concentration of acetaldehyde increases, but from 70 t o 95% of the bisulfite was found to be bound after 2 minutes and after 1 hour, over 99% of the bisulfite was bound. The acetaldehyde bisulfite compound was more stable than that of glucose, and the equilibrium constants for the association reaction were such that in a mixture of acetaldehyde and glucose, the former combines first with added bisulfite. If acetaldehyde is added to a glucose bisulfite solution it replaces glucose in combination. Bianconi and Bianchi (1932) investigated the extent of association of potassium bisulfite with acetaldehyde and reported complete binding after 20 minutes in 0.05 N solutions. In the presence of tartaric acid, the rate of association was slower, and this was ascribed to inhibition by tartrate. Neither Kerp nor Bianconi and Bianchi, however, determined the effect of pH on association and dissociation. This was done first by Tomoda (1927)) who found that a t pH 6 to 8 the dissociation of the acetaldehyde sulfonate was less than 5%; a t pH 8 t o 10.5, it was 50%; and above pH 12 practically complete dissociation occurred. Tomoda calculated the degree of dissociation of the aldehyde sulfonate from the mass action law to be:

SULFUR DIOXIDE TREATMENT OF FRUIT AND VEGETABLE PRODUCTS

where a K

109

degree of dissociation of the acetaldehyde sulfonate; dissociation constant for the acetaldehyde sulfonate, which according t o Kerp is 2.8 X K I = first dissociation constant for sulfurous acid, 1.7 X K O= second dissociation constant for sulfurous acid, 5 x The degree of dissociation, expressed as percentage of bisulfite or acetaldehyde sulfonate at various pH values, as calculated by Tomoda is shown in Fig. 2. Tomoda (1928) later pointed out that the pH during = =

FIG.2. Dissociation of acetaldehyde sulfonate a t various pH values. After Tomoda (1927).

glycerine fermentation (normally conducted a t not above pH 8.3) should not be much below pH 6 if undesirable concentration of bisulfite ion is to be avoided. In the course of our investigations2 on the chemistry of acetaldehyde and other bisulfite addition compounds begun in 1950, several attempts were made to prepare pure crystalline acetaldehyde hydroxysulfonate using the procedures published in the literature, but all attempts failed. With the best preparation of acetaldehyde-sulfonate available, which according to analysis was 70% pure, the effect of pH on dissociation was determined. The results obtained, yhown as curve 3 in Fig. 3, indicate the

* These investigations were begun in our laboratories with the assistance of C. J. Smit,

without knowledge of the previous unpublished data obtained by the Corn Products Refining Company research workers, but were continued after these data became available and led to Rraverman’s successful preparation of the pure glucose sulfonate (Rravcrman, 1953).

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total iodine titratable sulfite liberated in 1 minute a t the pH range of 2 to 12. At pH 2 with a preparation equivalent to 160 p.p.m. of SO2 only 2.7% was dissociated; with a preparation equivalent to 600 p.p.m. of SOz only 1.3% was dissociated in 1 minute. As the pH increased, there was a t first an increase in dissociation which had a maximum a t about 3.2 to 3.3, which then dropped to a minimum a t pH of 4.5, increased again, and was relatively constant in the range of pH 5 to 9. Above pH 9 the dissociation increased rapidly reaching a maximum a t pH 11.5.

FIG.3. The dissociation of sulfonates expressed as total free SO2 a t various p H values in presence of 0.1 M tartrate. Data for curves 1 and 3 are from Joslyn, Braverman, and Smit (1950); for curve 2 from Ponting and Johnson (1945).

Glucose Hydroxysulfonate. The preparation of sodium glucose sulfonate originally described by Kerp (1904) was modified by Neuberg (1929) and more recently by Gehman and Osman (1954). Even the latter procedure resulted in our hands in only 70% yield of 90% pure compound. These procedures have the disadvantage that the methanol used in separating the glucose sulfonate will precipitate unchanged glucose and sodium bisulfite as well as the sulfonate. Braverman (1953) successfully prepared crystalline sodium glucose hydroxysulfonate of high purity as follows : Dissolve 144 g. (0.8 mole) of anhydrous glucose in 60 ml. of distilled water and cool to 40" C. I n another beaker dissolve 88 g. of NaHSOJ (0.8 m. 10% excess) in 100 ml. H 2 0 and also cool to 40' C. At this temperature the two solutions are mixed slowly with constant slow stirring. The mixture is then placed in an incubator or in a constant

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temperature bath at 30-35" C. and constantly stirred with a slow-moving stirrer. After an hour or so, the solution becomes quite clear. After 50 to 60 hours, the liquid becomes turbid with suspended crystals. Continue stirring for 3 to 5 days in all, after which an abundant crop of beautiful crystals is obtained. Filter on a Buchner funnel with suction to separate as much as possible the mother liquor. Wash with small portions of 75% methanol. Rinse with 99% methanol. Dry in vacuum dessicator. Yield obtained: 66% of the theoretical. (See Figure 4.)M.P. 92" C.

FIQ.4. Photograph of crystals of sodium glucose hydroxysulfonate.

The equilibrium constants for the equilibrium and velocity of the reaction of glucose with sulfurous acid were first measured by Kerp and his collaborators some forty years ago, but their results are difficult to interpret and apply because pH data are lacking and because most of their data relate to rather dilute solutions of glucose (18% or less) and bisulfite (800 p.p.m. or less). Tomoda and Taguchi (1930), Bianconi and Bianchi (1932), and more recently Vas (1949) investigated this reaction.

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M. A. JOSLYN AND J. B. 6 . BRAVERMAN

Vas (1949) reported that the equilibrium constant was mainly a function of pH, being least in the range of pH 3 to 5.5 and increasing rapidly in more acid and more alkaline regions. Changes in concentration of glucose in the range of 1 to 40 g. per 100 ml. and of SO2 in the range of 130 to 3500 mg. per liter had only a minor influence on the equilibrium constant. 12

I

I

I

I

I

PH

FIQ. 5. Effect of glucose on the dissociation of a solution of glucose hydroxysulfonate at various pH’s. 1. Glucose hydroxysulfonate 2.2175 g./l. 2. Glucose hydroxysulfonate 2.2175 g!J. 10% glucose. 3. Glucose hydroxysulfonate 2.2175 g./l. 20% glucose. 4. Glucose hydroxysulfonate 2.2175 g./L 50% glucose. From Braverman (1953).

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The apparent velocity constants of decomposition and addition were found to be functions of pH only. Braverman (1953) reported that at pH 5, over 4001, of the SOz present was bound after the first 30 minutes. In a more acid medium (pH 2) the rate of combination was slow and amounted to only 32% after prolonged storage. As the concentration of glucose was increased the percentage of SO2 bound first increased and then decreased. At pH 7 he found that the pure glucose sulfonate dissociates completely. The effect of pH on the dissociation of the glucose sulfonate measured as total free SO2 is shown in Fig. 3, in comparison with that of frozen apples and acetaldehyde sulfonate. It is evident that the constituent combining sulfur dioxide in apples is glucose and that the dissociation of the glucose sulfonate increases rapidly with pH and is practically complete a t pH 6. The dissociation of glucose sulfonate in absence and presence of glucose is shown in Fig. 5. As the excess glucose present is increased in concentration, the amount of SOz dissociated a t various pH levels decreases at first, rises to a distinct maximum in the

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range of pH 4 to 4.5, drops to a minimum in the range of 5.5 to 6.5, and then rises again.

111. DETERMINATION OF SULFUR DIOXIDE IN FRUIT AND VEGETABLE PRODUCTS

It is obvious from the evidence summarized above that sulfur dioxide added to fruit and vegetable products as liquid sulfur dioxide, as gaseous sulfur dioxide, as solution of sulfmous acid in water, or as dry or dissolved sulfites will exist as the undissociated sulfurous acid, as the free bisulfite ion, as the free sulfite ion, and as combined sulfur dioxide in the form of the hydroxysulfonates. The pH, temperature, composition of the food product, conditions of sulfuring or sulfiting, aiid the subsequent treatment and storage conditions will determine the equilibrium between the various forms of sulfur dioxide present. Oxidation of the sulfurous acid and sulfites to sulfate ions and formation of organic sulfur compounds other than the hydroxysulfonates which may occur during processing and storage will also be involved. I t is obvious that the analyses for these constituents at any given time would be most useful in determining the relative efficiency of a particular type of sulfuring or sulfiting practice and in interpretation of the mechanism of sulfite inhibition of undesirable changes. Unfortunately the methods of analysis so far available are not sufficient for this and for the most part these have been based on empirically developed procedures tested either by maximum recovery of the sulfur dioxide present or in a few instances by recovery of sulfur dioxide added as such. The recovery of sulfurous acid added as such or as bisulfite is not sufficient measure of the accuracy of a given method for total sulfur dioxide. Even in cases where the bound sulfur dioxide is likely t o be present largely as hydroxysulfonate investigation of the adequacy of the procedure selected has not been checked by addition of known amounts of the sulfonate. High results, reproducibility of results, and avoidance of possible errors in determination have been relied upon generally in development of analytical methods. The available methods of analysis can be segregated into those designed to measure the free sulfur dioxide and those for total sulfur dioxide. The latter can be subdivided into two groups: (1) those in which the “bound” sulfur dioxide is liberated by distillation from acid; and (2) those in which the combined sulfur dioxide is liberated by treatment of the liquid product or extract with excess alkali a t room temperature and subsequent acidification t o prevent recombination. The free or total sulfur dioxide may be determined volumetrically, gravimetrically, or colorimetrically. Monier-Williams (1927) has reviewed critically and thoroughly the earlier investigations in this field and has reported in detail the basis of the procedure he de-

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veloped to overcome several of the difficulties with previous methods. He stressed particularly the fact that ((determinations of free and combined sulfurous acid in foods are not likely to give reliable results unless the conditions under which the analysis is conducted are such that no appreciable alteration of equilibrium can occur during the determination.” “Free ” Sulfur Dioxide: The determination of “free” sulfur dioxide in beer, white wines, and lightly colored juices, concentrates, and beverages is based on the direct iodine titration method developed by Ripper (1892) for wine. In this method a 50-ml. sample, after acidification with 5 ml. of dilute sulfuric acid (1 3), and expelling of the air present either by displacement of air present in the flask with carbon dioxide gas as suggested by Ripper or by addition of 0.5 g. of sodium carbonate (Assoc. Offic. Agric. Chemists, 1950) as in the official A.O.A.C. method, is rapidly titrated directly with 0.02 N iodine solution to a starch end point which persists for a few minutes. The added acid is relied upon to reduce the rate of dissociation of the “bound” sulfur dioxide. Kerp and Bain (see Kerp, 1913) showed that for their preparation of the acetaldehyde sulfonate the addition of N/30 hydrochloric acid markedly decreased the rate of dissociation, although it increased the actual extent of dissociation at equilibrium. With their preparation of glucose sulfonate the addition of hydrochloric acid also markedly decreased the rate of dissociation but did not influence the actual extent of dissociation. The effect of pH and temperature on the rate of dissociation of a-hydroxysulfonates in presence of other constituents likely to influence the rate of dissociation, however, has not been determined. Ripper’s method has been applied to red wines (Benvegnin and Capt, 1931; Sumuleanu et al., 1937); to beer (Butlestone, 1928); to sweet musts (Fischler and Kretzdom, 1938); to apple juices and ciders (Warcollier and Le Moal, 1929); to concentrates and syrups; and also to extracts prepared from dehydrated vegetables, dried fruits, etc. In the analysis of diluted concentrates and extracts of dried or dehydrated foods the possible change in the equilibrium between bound and free sulfur dioxide as a result of dilution during preparation of the extracts has not been controlled. Hydrochloric acid has been favored by some over sulfuric acid for acidifying the solution, e.g., Jaulmes and Espeael (1935) used a dilute hydrochloric acid (1 4). Hydrochloric acid was preferred also by Bennett and Donovan (1943), Prater et al. (1944), Ponting and Johnson (1945), Reifer and Mangan (1945), and others. In the preparation of extracts for direct titration losses of the sulfur dioxide present by oxidation as well as changes in the distribution of sulfur dioxide occur. These losses are significant in cold water extraction and in the alkali extraction procedures of Prater et al. (1944) and Potter and Hendel (1951). The oxidation losses were reduced by Reifer and Mangan (1945) by extraction of the finely ground material in boiling water containing sugar and a phosphate buffer of pH 7.6. They claimed no loss of sulfur dioxide due to evaporation or oxidation under these conditions. Ponting and Johnson (1945) controlled oxidation by blending fruit in a solution of salt and tartrate buffer a t pH 4.5. Even when errors in true free sulfur dioxide content due to dilution or oxidation are avoided, the direct iodometric determination of sulfurous acid and sulfites is subject to error. Kolthoff and Menael (1929) point out that an appreciable error occurs when sulfite solutions are titrated with iodine owing to oxidation by air and point out that correct results can be obtained only when the sulfurous acid or sulfite solution is allowed to flow into a solution of iodine. The titration of sulfurous acid or bisulfites is

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subject also to possible reduction of SO2by hydriodic acid to sulfur and iodine, and to loss of SO2 by evaporation during the titration. In strongly acid solutions volatilization of Sot, particularly when its concentration in the aliquot taken for analysis is over 0.04%, is more important as a source of error than oxidation. Mason and Walsh (1928) have investigated the magnitude of the errors occurring during the titration of dilute sulfite solutions with iodine, and concluded that loss of SO2 due to volatilization is more important than by oxidation. Reduction of sulfur dioxide by H I was not found by them. In the titration of wines, beers, juices, and other liquid food products, MonierWilliams (1927) lists as additional errors: (1)action of iodine on substances other than sulfur dioxide and (2) recombination of sulfur dioxide and acetaldehyde or other carbonyl compound. Ingram (1947a) pointed out that, in the direct iodine titration of citrus juices by iodine using starch indicator, the end point tends to drift either because excess of iodine combines with other reducing substances present or because of slow decomposition of combined SOz. He suggested potentiometric titration using a bright platinum electrode in conjunction with an Ag-AgC1 reference electrode and an electrometer. This is particularly desirable with colored juices or beverages. Using the electrometric titration, Ingram (1947b) carefully estimated the range of the errors involved in the determination of free SO2in diluted sulfited orange juice concentrate. If the concentrated juice was diluted before titrating, extra free SOZwas formed a t the rate of 1 p.p.m. per minute in a juice containing about 100 p.p.m. of free SOZ. If delay occurred during the titration at any point, more free SO1 was formed and this error was greater the greater the degree to which the titration had progressed before delay occurred. With delay a t the end point there was a relatively large increase in the titration. To correct for iodine reducing substances other than sulfur dioxide which may be present, the iodine titration is carried out both on the original wine, juice, or extract to measure the total reducing power including SO, and on an aliquot to which is added an excess of a bisulfite binding agent to measure the reducing power of the juice or extract itself. The difference in titration is taken as a measure of sulfur dioxide present. Acetone was suggested as such a binding agent by Downer (1943) and Bennett and Donovan (1943), and used by Prater et al. (1944). It was not found satisfactory by Ponting and Johnson (1945) for fruit extracts, by Reifer and Mangan (1945) for dehydrated cabbage, by Ingram (1947b) for citrus concentrates and citrus juices, and by Potter and Hendel (1951) for dehydrated white potatoes. Ponting and Johnson (1945), Reifer and Mangan (1945), and Potter and Hendel (1951) found that the rapid fading of the end point in the acetone-treated sample, as a result of rapid dissociation of the acetone-sulfite complex, can be largely avoided by the use of formaldehyde as a binding agent, since it forms a more stable complex with sulfite than does acetone. In addition to the iodometric determination, direct precipitation as barium sulfate before and after treatment with bromine was suggested for both quantitative and quaIitative test for sulfur dioxide in wine (see Monier-Williams, 1927). Precipitation of SO2 after oxidation with HeOl as the benzidine sulfate was proposed by Rothenfusser (1929); reduction of the molybdenum in phosphomolybdic acid by the sulfite ion present in an aqueous solution of the food was proposed by Sasaki (1928) as a colorimetric method; formation of a blue color from a solution of 1%methylene blue and 5 % iodine in potassium iodide was proposed by Svershkov (1939). Mathers (1949) proposed a turbidimetric method based on the distillation of wine into a dilute solution of lead acetate to form a colloidal suspension of lead sulfite whose spectral transmittance in the range of 400 to 700 mp could be used as a measure of sulfur diaxide. This is similar to turbidimetric methods based on turbidity produced by adding BaClr to a

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sulfite oxidized by HzOr but is more rapid and more sensitive in the range of 0 to 100 p.p.m. Francis el al. (1938) proposed a conductometric method based on the change in electrical conductivity of neutral HZOZwhich results from passage of SO? for the estimation of SO2 in industrial gases and a similar method. Conductometric methods were used previously by Thomas (1932), Thomas and Abersold (1929), Thomas el al. (1943, 1946) for the determination of small concentrations of SO2 in air and photoelectric determination employing dilute starch-iodine solutions was proposed by Katz (1950). The determination of sulfurous acid and sulfites by oxidation to sulfates and precipitation as barium sulfate because of its high degree of specificity in acid solutions, high degree of insolubility, and favorable gravimetric factor, commands wide acceptance. Gravimetric determination of sulfate by precipitation as BaSOr is still suggested as optional in the official Monier-Williams method of the Association of Official Agricultural Chemists (1950). To utilize the advantages of barium sulfate in the microgram range of microanalysis, titrimetric methods using tetrahydroxyquinone or rhodizonates as indicators have been introduced (see Little, 1953; Toennies and Bakay, 1953). With tetrahydroxyquinone as indicator a practical sensitivity of k 3 micrograms of sulfur per determination can be achieved in the titration of sulfate with barium chloride. Toennies and Bakay (1953) recently investigated the determination of barium sulfate in suspension by nephelometry and developed a sensitive photonephelometric micro method based on the use of a glycerol-alcohol-water system to obtain reproducible and stable suspensions of barium sulfate. Their procedure with slight modification waa successfully applied in our laboratory (Lukton and Joslyn, 1954) to the determination of sulfur dioxide in solutions of sulfurous acid and sulfites but could not be applied directly to white wines. Recovery of added sulfite and sulfate to wine was high and variable. Even when it was applied to wines after dry or wet ashing, the results were not satisfactory. Acid bleached fuchsin to which formaldehyde is added has been developed as a sensitive method for the determination of sulfur dioxide. As developed first by Grant (1947)the acidified fuchsin formaldehyde solution is added in excess to a n aliquot containing sulfite. Hoffpauir and O’Connor (1949) suggested sensitization by addition of a ketone such as acetone. Steigmann (1950) proposed that the acidified basic fuchsin and formaldehyde reagents be prepared separately and that the former be filtered after storage for 3 days. La Rosa (1950) suggested that the sulfuric acid-basic fuchsinformaldehyde reagent be decolorised with carbon to remove colored impurities in the dye solution and filtered before use. This reagent was applied by Atkin (1950) to the determination of sulfur dioxide in presence of sulfur trioxide, by Urone and Boggs (1951)to determination of sulfur dioxide in the atmosphere, and by Dupaigne (1951) as a micromethod for sulfur dioxide in grape juice and was proposed also for use in wine and beer analysis. The above substitutes for the iodometric determination with one or two exceptions have not been tested for specificity for free SOZ.Mathers (1949) on the basis of the fact that “free” sulfur dioxide is removed early in the distillation of wine proposed that 10 ml. of the distillate from the alcohol determination be mixed with 0.5 ml. of 5% neutral lead acetate and the turbidity of the suspension formed be used to correct volatile acidity for sulfur dioxide. La Rosa (1950)assumed that the color formed on mixing the fuchsin-sulfuric acid-formaldehyde reagent with white wine was a measure of free SOz but did not give any data to confirm this. The fuchsin procedure and the lead sulfite procedures are very sensitive, the former more so, and can be applied only to very small aliquots or to dilute solutions.

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Joslyn (1953) found t h at Grant’s reagent (Grant, 1947) could not be used with the Klett-Summerson colorirneter a t SO2 levels ahove 10 p.p.m. because the reading was off the scale. When 5 rnl. of the fuchsine reagent and 1 ml. of test solution were mixed and diluted to 25 ml. before transfer to t h e colorimetcr tube, fairly reproducible readings were obtained in the range of (r50 p.p.m. A high blank reading with even the freshly mixed reagent and increase in this blank and decrease in sensitivity of the reagent with storage was observed. Higher sensitivity and greater stability was obtained with the Steigmmn (1950) modification, but the free SO2 content of wine determined colorimetrically was considerably higher than given by iodometric titration, the difference increasing with increase in SO2 level. Total Sulfur Dioxide: After Alkali l’reatrnenl: The fact that neutral sodium sulfite does not combine with carbonyl compounds and t h at the hydroxysulfonic acid compounds are rapidly decomposed on treatment with alkali was used by Ripper (1892) as the basis for the determination of total sulfur dioxide in wine b y direct iodine titration. I n his method, 50 ml. of wine were pipetted into a 200-ml. flask containing 25 ml. of 1 N KOH. The mixture was shaken and allowed to stand for 10 to 15 minutes. Then 10 ml. of dilute sulfuric acid (1 3) were added, and the solution titrated rapidly with 0.02 N iodine solution t o a starch end point which persisted for some time. This method was used as the official direct titration method for wine in the first edition (1919) of the A.O.A.C. Methods of Analysis; in the third (1930) edition it was extended t o white grape juice, wine, and similar products (1 N NaOH or KOH was used and the solution during standing for 15 minutes was occasionally agitated); but it was dropped from the fourth (1935) and succeeding editions. Ripper compared his method with the Haas distillation method on ten wines whose SO2 content varied from 42 to 1488 mg. per liter and found the difference between the two to vary from 0 to 5 mg.

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TABLEI1 Collaborative Results on Sulfur Dioxide in Wine* Total Sulfur Dioxide as p.p.m. Analyst

* Joslyn

MAJ OSN WHP JHF RWM JLB DB MM MAA LQ FQ AGP T HP Average Maximum Minimum (1930).

Distillation Ripper titration 435 486 518 380 433 408 433 400 424 425 43 1 448 426 438 435 518 380

418 433 379 240 420 336 424 388 437 408 413 420 405 370 392 433 240

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per liter higher or lower. A comparison between the Ripper method and the Nichols and Reed (1932) distillation method by fourteen collaborators with the same sample of white wine gave the results shown in Table 11. As pointed out by Monier-Williams (1927), the Ripper method for total sulfur dioxide is subject, in addition to the errors inherent in the iodometric determination of sulfurous acid, to errors due to oxidation of sulfite in alkaline solution, to possible recombination of liberated sulfur dioxide with the aldehyde, and to reduction of iodine by substances other than sulfur dioxide. Under certain conditions by a compensation of errors results closely approximating those obtained by the distillation method may be obtained. Mills and Wiegand (1942) reported close agreement between the Nichols and Reed (1932) distillation method and the Ripper method for sweet wines when the end point of the iodine titration was taken as the first blue lasting approximately one minute rather than a blue color lasting several minutes. Contrary results were reported by Archinard (1937), Casanave (1910), Ferr6 (1915), Mathieu (1910), and Taylor et al. (1937). The Ripper method is considered acceptable for white wines containing less than 300 p.p.m. of SO2 by Fabre (1936), Hennig (1946), RibCrau-Gayon and Peynaud (1947), and Jaulmes (1951). Amerine (1952) modified it by decreasing the volume of wine used from 50 ml. to 20 ml. Joslyn (1953) found better correlation between the total SO2 content of wine determined iodometrically and that determined by MonierWilliams reflux distillation when the amount of alkali present was sufficient to bring the wine to p H 12.5 and the period of standing was reduced to 2 minutes. (Ten milliliters of wine pipetted into 3 ml. 1 N NaOH titrated after acidifying with 1 N HCl before and after addition of 1 ml. of formaldehyde to correct for iodine reducing substances other than SO2.) Comparative values for several wines determined by modified acid-fuchsine colorimetric method were considerably higher than by the iodometric or Monier-Williams method. The Ripper method has been subjected to several modifications largely directed to improving precision and to obtaining results in better agreement with distillation procedures accepted as being more accurate. Little or no attention, however, has been given to determination of its relative accuracy. The fundamental data on rates of dissociation of the aldehyde and sugar sulfonic acids have not been applied, nor has recovery of added pure sulfonates been used as a measure of its accuracy. Recovery of added sulfurous acid or bisulfites, as pointed out above, can not be accepted as sufficient proof of its accuracy. Advantage of the data on effect of p H on rates of association and dissociation and position of equilibrium has been used, however, in the methods developed for determination of acetaldehyde, such as that developed by Jaulmes and Espezel (1935) for acetaldehyde in wine, which was found satisfactory by Joslyn and Comar (1938). To avoid errors due to reduction of iodine by substances other than sulfur dioxide, sulfite-binding compounds (acetone or formaldehyde) were added in excess after acidification. Bennett and Donovan (1943) and Downer (1943) used acetone to combine with the liberated or free sulfur dioxide, in order to determine the iodine-reducing power of the juice itself. Prater et al. (1944) in extending the method to dehydrated vegetables determined the effect of p H on stability of the acetone complex formed with extracts of dehydrated cabbage, carrots, and potatoes in presence of large excess of acetone and found this t o be in the range of p H 2 to 3, in which the starch-iodine end point was sharper than a t p H 4. The difficulties with acetone as a binding agent for citrus juices were pointed out by Ingram (194713) and others. The Ripper method in modified form was extended to the determination of total sulfur dioxide in frozen and dehydrated sulfited fruit and vegetable products. Jensen (1928) was one of the first to suggest its application to determination of total sulfur

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dioxide in foods after alkali extraction or digestion, and Brooks (1928) applied it to dried fruit. I n Brooks’s procedure a 50-g. aliquot of finely ground dried fruit was macerated with 100 ml. of distilled water and then treated with 50 ml. of a 5.6%solution of KOH, thoroughly mixed, allowed to stand for 15 minutes, acidified with 25 ml. (1 3) sulfuric acid, and then titrated with 0.1 N iodine to a starch-iodine end point. Prater et al. (1944) extracted and liberated the sulfur dioxide present in dehydrated vegetables by suspending 8 g. in 400 ml. of water, adding 5 ml. of 5 N NaOH, and allowing the mixture to stand 20 minutes. For cabbage, carrots, and potatoes the total sulfur dioxide liberated increased during the first 10 minutes of treatment and then remained essentially constant up to 60 minutes. A similar extraction of sulfite was used by Potter and Hendel (1951) for dehydrated white potatoes. Reifer and Mangan (1945) extracted the sulfite in dehydrated fruit or vegetable by adding 2-g. portions of finely ground material to 350 ml. of boiling water containing 5 ml. of phosphate buffer a t pH 7.6 and 10 g. of sugar. The flasks containing the mixture were then stoppered and allowed to cool. Ponting and Johnson (1945) blended a 100-g. sample of fresh or frozen fruit in a Waring Blendor with 10 ml. of 0.5 M tartrate buffer at p H 4.5 and 490 ml. of 20% NaCl solution for 3 to 5 minutes. For dried fruits they used a 20-g. sample. I n the Ponting and Johnson method, the extract was then filtered and the sulfite in the filtrate was liberated by addition of 2 ml. of 1 N NaOH to 50 ml. aliquot and allowing to stand for 30 seconds. In frozen apples they found that the liberation of sulfur dioxide was complete in 30 seconds or less a t p H of 9 (see curve 2, Fig. 3). By reducing the concentration of alkali added and the p H and time of alkaline treatment, losses due to oxidation were minimized. After alkali treatment they as well as Reifer and Mangan (1945), Prater el al. (1944), and Potter and Hendel (1951) acidified the solution with hydrochloric acid in preference to sulfuric acid used b y Ripper. In the Ponting and Johnson method, 2 ml. of 6 N HC1 were added to a 50-ml. aliquot and in the Potter and Hendel (1951) procedure 7 ml. of 5 N HCl were added to 400 ml. of extract. To determine reducing material other than sulfite, 1 ml. of 40% formaldehyde and 10 ml. of 36% formaldehyde, respectively, were added to separate acidified aliquots, and after standing for 10 minutes these were titrated with 0.02 N iodine to a definite blue starch-iodine end point. Recovery of sulfur dioxide added as sulfurous acid or bisulfite and agreement with results obtained b y the Monier-Williams (1927) procedure were used as indices of accuracy. Total Sulfur Dioxide: Distillation Procedures: The early development and application of distillation procedures to various foods, beginning with Haas’s method of distilling wine in a current of carbon dioxide into iodine solution contained in a Pelegot tube immersed in water proposed in 1882, are critically reviewed b y Monier-Williams (1927). The official distillation method of the A.O.A.C. in 1919 waa the distillation of 20 to 100 g. of sample in a current of carbon dioxide after addition of 5 ml. of 20% phosphoric acid into bromine water and the determination of sulfate formed by precipitation with BaCL after expelling excess of bromine. This waa still the official method in 1925, and in 1930 but with a cautionary note that i t was not applicable if volatile organic sulfur compounds may be set free. I n 1930 Monier-Williams method both in its volumetric and gravimetric modifications was listed as tentative. I n 1935 and subsequent editions of Methods of Analysis, the Monier-Williams method was adopted as official, and the old bromine procedure (May, 1927) is no longer cited. The iodine distillation method has been investigated from time to time since its proposal by Haas (Anon., 1928; Black, 1928; Black and Warren, 1928; Nichols and Reed, 1932, 1933; H a n d and VoFigek, 1937), but because of errors involved (volatilization of iodine solution in the receiving flask, distilhtjon of iodine-reducing substances other

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than SO*, and interference by naturally occurring volatile organic sulfur, compounds in products like onions, garlic, and horse-radish, see Dyer et al., 1941), it has not been considered desirable for other than control tests. Iodine distillation, however, has been used in analysis of wine and dried fruits and has been developed as a micromethod for sulfurous acid in wine and fruit juices by Woidich (1930) and others. The method proposed by Monier-Williams (1927) is based upon the selective oxidation of sulfur dioxide by hydrogen peroxide at room temperature. When iodine or bromine are used, oxidntion of hydrogen sulfide and of volatile organic sulfur compounds to sulfuric acid occurs. Fitelson (1929) early compared the Monier-Williams method with the A.O.A.C. method and found it to be superior to the latter even when cadmium chloride or copper salts were used to remove sulfides from the distillates. The Monier-Williams method consists in distilling for 1 hour with a reflux condenser and sweeping the sulfurous acid into a cold, sulfate-free, acetanalide-free, neutral hydrogen peroxide by a current of'carbon dioxide. As a rule, volatile acids and organic sulfur compounds do not distill over when a reflux condenser is med, although when larger quantities of volatile organic sulfur compounds are present a great part of them may be carried over. The retention of the volatile organic acids in the distilling flask permits a volumetric determination of the sulfurous acid. Because of the use of a reflux, however, a longer distillation is necessary to insure sweeping of all of the sulfurous acid into the receiver. Also, the combined effects of the use of carbon dioxide and a reflux condenser make frothing a minor problem in the actual manipulation. Neutral hydrogen peroxide is used, thereby permitting titration of the sulfuric acid with 0.1 N NaOH and subsequent gravimetric determination aa BaSO4, if desired. All manipulation of the H202 is carried out in the cold, since on heating any H B or organic sulfur compounds will be oxidized. Precipitation of the sulfuric acid as Bas04 is carried out in the cold. Although the method is time-consuming, it has the following advantages: 1. The sulfurous acid is more completely liberated, owing to the use of HC1 instead of HaP04. 2. Errors due to volatile S compounds are avoided. 3. Organic acids do not pass over with the distillate, thereby allowing a volumetric determination. 4. Frothing is kept to a minimum. Since its introduction in 1927, the Monier-Williams method has been modified from time to time, chiefly in connection with design of apparatus and conditions of refluxing and distillation (Nissen and Petersen, 1943; Thompson and Toy, 1945; and others). It has been applied to a wide variety of foods, dried and dehydrated fruits (Miller, 19271, dehydrated vegetables, juices, concentrates, wines, beers, etc. (Veisman and Svereva, 1937). I n general, it is agreed that distillation is the most reliable method of determining sulfur dioxide in foods, but opinions differ considerably as to the way in which the distillation should be carried out and as to the best method of determining sulfur dioxide in the distillate. Some authors favor distillation in steam, others under reduced pressure, and the majority distill in carbon dioxide or in some other inert gas. The sulfur dioxide in the distillate may be oxidized to sulfuric acid by bromine, hydrogen peroxide, or possibly other oxidizing agents, and determined either volumetrically or gravimetrically. I n the distillation method, precautions have to be taken to assure complete separation of SO2 from the fruit product, to avoid oxidation of the liberated SO2 prior to its absorption in the receiving flask, to eliminate the effect of other reducing substrtnces, and to eliminate the effect of other sulfur compounds. Titration of the sulfur dioxide

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in the distillate, either as such or after oxidation to sulfuric acid, gives acvxrate results if precautions are taken to climinate volatile organic compounds. Gravimetric determination as Bas04 gives accurate results in nearly all cases, except, when the distillate contains appreciable amounts of volatile sulfur compounds. Most authors have used phosphoric acid for acidifying the food prior t o distillation. In certain fruit products, such as dried fruits and molasses, where the sulfurous acid is in close combination with sugars, i t is difficult to effect a complete liberation of the sulfurous acid by using phosphoric acid. The use of hydrochloric acid insures a faster and more complete liberation of sulfurous acid from its combinations. In the case of apricots, even when relatively large quantities of hydrochloric acid are used, it takes more than 1 hour for complete distillation of sulfurous acid. Preliminary treatment of the product with sodium bicarbonate does not yield appreciably different results. Increasing the acid concentrations hastens hydrolysis of the food constituents and reduces frothing, but it may favor production of volatile sulfur compounds from protein-containing foods. Dried fruit or other sugary materials change to a dark or chocolate brown color when boiled with hydrochloric acid, probably owing t o decomposition of sugars, etc. With phosphoric acid, the amount of decomposition is slight. With HC1, the change in color occurs more markedly toward the end of the distillation; some authors claim that this change, which is more or less abrupt, denotes complete evolution of SOZ,and take it as an indication of the end point. Monier-Williams has shown that reduction of sulfates owing to this decomposition does not take place, whereas others have claimed this to be the case and use it to explain the continued slow evolution of SO2 from the product. To avoid oxidation of the evolved SOzon its way to the receiving flasks, distillation in a current of COz, or other inert gas, is resorted to. However, this oxidation seems to be rather small, and similar results have been obtained with and without the use of a stream of COz. It is optional whether the system be swept out first with a stream of COzfrom an outside source and the food then distilled in this atmosphere, or whether the COz is generated in the distilling flask itself by the addition of chalk, NazCOa, or NaHC03, or a solution of NaaC03, even though in the latter case the system is not swept as clear of air as in the former case. When carbonates or bicarbonates are used to generate COz in the distilling flask, an excess of acid must be added to compensate for the acid used up in generating COz.

+ HCI + NaCl -t HzO -I- coz NaaCOa + 2 H C 1 4 2NaCl + HzO 4-COz

NaHCOs or

The carbonates are less efficient in the production of COz than the bicarbonates, for which reason, and because the latter can be obtained more readily in the pure state, they are used to a greater extent than the normal carbonates. In the distillation of solid foods, it is necessary to add water to the material in the flask. It is optional whether the water is tap, distilled, recently boiled distilled, or carbonated distilled water. The error due to the oxygen present in the water is negligible when compared to other possible errors. To avoid foaming during the initial period of distillation, it is desirable to add a small quantity of tannin or a few drops of mineral oil. Paraffin will distill over and gum up the condenser and therefore should not be used. Distillation of sulfur dioxide into standard iodine solution and titration of the excess iodine has been used with modifications by several chemists. It i s satisfactory only when volatile organic compounds capable of reducing iodine are absent, and also when

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necessary precautions are taken to avoid loss of iodine by volatilization i n the current of carbon dioxide. SO2 I2 2H20 ---t 2HI H2SO4 IZ 2NazSzOa + Na&Oa 2NaI

+ + +

+

+

The excess iodine is generally titrated with standard thiosulfate solution. To minimize the loss of Izin the receiving flask, the flask should be shielded from heat reflected by the burners used in distilling and should be equipped with a trap of some sort. Traps containing K I or NapS2Oa have been used. Cooling the receiving flask in ice water is recommended (see Fig. 6 ) . Total Sulfur Dioxide-Other Methods: The acid bleached fuchsin-formaldehyde procedure has been proposed for the determination of total sulfur dioxide in white wine and beer after alkali treatment and acidification. For red wines and colored juices this procedure can be used by distilling the sulfur dioxide with steam in a microKjeldahl still. Particular attention, however, has to be given to using an aliquot which does not contain too much sulfur dioxide. In the procedure proposed by Grant (1947), 4 ml. of the fuchsin reagent are added to 1ml. of sulfite solution and after 5 minutes the intensity of color formed is measured in a photoelectric colorirneter using a green filter. The transmittancy increases with SO2 concentration over most of the range of 0 to 10 micrograms of S02. The bound SO1 in white wine is liberated by La Rosa (1950) by treating 0.5 ml. of wine diluted with 15 ml. water with one drop of 10% NaOH and allowing the mixture to stand 1 minute. With red wine 0.5-ml. sample and 2 ml. of 25% phosphoric acid is still distilled in a micro-Kjeldahl still into 1 ml. of water containing one drop of 10% NaOH. One ml. of beer or its equivalent of distillate may be used. Mathers (1949) found that distilling 50-ml. samples of wine after addition of 50 ml. of 5% sulfuric acid and determining the sulfite in an aliquot of the distillate with neutral lead acetate gave results as low as one-half that obtained by iodometric titration or the Monier-Williams method. The neutral lead acetate solution, however, was shown to be a good absorption medium for sulfur dioxide. Lead sulfite suspensions could be acidified and the sulfur dioxide determined by iodometric titration. Low recovery in the distillation procedure was ascribed to oxidation of sulfur dioxide during distillation, since the same wines when distilled under reflux after addition of oxalic acid or sodium arsenite gave photometric values closely approximating those by iodometric titration and the Monier-Williams titrimetric or gravimetric procedure. The photometric lead sulfite method is limited to 10-ml. aliquot of distillates containing 5 to 100 mg. per liter of sulfur dioxide. The fuchsin-formaldehyde procedure is limited to determination of 0 to 10 micrograms of sulfur dioxide in 1ml. of sample with an error of less than 0.1 microgram for pure solutions (Grant, 1947). Joslyn (1953) obtained higher reproducibility3 with the Steigmann (1950) modification. Acidbleached fuchsine prepared by adding 100 ml. of concentrated sulfuric acid to about 800 ml. water, cooling the solution, and then adding 0.5 g. of National Aniline Company’s basic fuchsine (NF60) moistened with 20 ml. of alcohol, gave a stable clear solution. This was mixed with dilute formaldehyde solution (1/20) before use in the volume ratio of 10 parts of the fuchsine solution to 1 part of formaldehyde solution. To 1 ml. of wine pipetted into a 100 ml. glass-stoppered flask were added 20 ml. of mixed reagent and after mixing diluted to volume and transferred to colorimeter tube for free S02. For total SO2 in white wines, to 1 ml. of wine were added 3 ml. of 0.1 N NaOH, and after 2 minutes exactly 3 ml. of 0.1 N NaOH and 20 ml. of reagent and brought to volume. If the alkaline wine were only partially neutralized before addition of color reagent, the readings were lower than with wine that was not acidified or wine that was over acidified. Similar results were obtained with pure sulfurous acid solu-

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tions and with distillates from red wine. The values for total SOz content by this method, however, did not agree with iodometric distillation or titration values.

IV. SULFURDIOXIDEAS PRESERVATIVE AND SANITIZING AGENT The preservative value of sulfur dioxide in fruit and vegetable products and other foods and beverages is discussed by Abdulov (1938), Cruess (1948), Borgstrom (1953), Tanner (1944), von Schelhorn (1951), and Woodroof and Cecil (1945). The actual mechanism of the preservative value of sulfurous acid and its salts, however, is not known. It is believed that its strong reducing power may beinvolved either by reducing the oxygen tension in the food tissues and in beverages to a point below that a t which aerobic organisms can grow or by maintaining some enzyme system necessary for growth in a reduced state. Sulfur dioxide is only a temporary preservative when used in moderate amounts because of loss of preservative value on oxidation to sulfate, on volatilization, and on combination with chemical constituents capable of forming a-hydroxysulfonic acid or other addition products. Because of the readiness with which it can be removed to a large extent by heating under vacuum it has been favored in Europe as a temporary preservative, particularly for the bulk storage of juices and pur6es. There is some evidence that more sulfur dioxide is required to prevent fermentative activity than to inhibit growth (Bioletti and Cruess, 1912). Sulfur dioxide also has a selective antiseptic action (Cruess, 1911). Acetic acid bacteria, lactic acid bacteria, and many varieties of molds are more sensitive to sulfur dioxide than yeasts. Among the yeasts, the more strongly aerobic species are generally more sensitive than the more fermentative species. The fermentative yeasts, particularly those strains selected as being more desirable industrially, can be adapted to sulfur dioxide, according to Porchet (1931). Bioletti and Cruess (1912) did not believe that adaptation occurred but the evidence presented by Porchet was convincing. The antiseptic action of sulfur dioxide towards microorganisms, particularly yeasts, varies with stage of growth or development, microbial population, temperature, pH, and composition of product treated (acetaldehyde content, sugar content, alcohol content, etc.) Effect of p H : Muller-Thurgau and Osterwalder (1914) were the first to observe that sulfurous acid and its salts were effective as preservatives only in acid media. Perry and Beal (1920) confirmed this by finding that with more acid solutions the lower the concentration required to inhibit fermentation and mold growth. Bioletti and Cruess (1912) did not find that acidity or sugar content was involved in the decrease in antiseptic effect of SO2in ripe as compared with unripe grapes. Subsequently however, Cruess and Irish (1932), Cruess (1932), and Cruess et al. (1931) in their investigations of the effect of pH on toxicity of sulfurous acid and

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other preservatives found that at pH 3.5, two to four times as much SO2 was required to inhibit growth as a t pH 2.5, whereas at pH 7 sulfite was without effect on yeast and molds and as much as 1000 p.p.m. were required to inhibit growth of bacteria in apple juice. Cruess and his collaborators suggested that antiseptic properties might be confined only to undissociated sulfurous acid or to molecular SO2 and be absent in the neutral sulfate ions. This was also suggested by Rahn and Conn (1944), who concluded that yeast growth was checked only by undissociated sulfurous acid and not by bisulfite or sulfite ions. They reported that 4 p.p.m. of undissociated sulfurous acid inhibited growth of Saccharomyces ellipsoideus, whereas 100 p.p.rn. of bisulfite ions were required to inhibit bacterial growth. Gillespy (1946) also considered un-ionized sulfurous acid as the lethal agent. He found that the effect of SO2in increasing the heat sensitivity of the asci of Byssochlamys fulva was a function of pH. Vas and Ingram (1949) pointed out on the basis of the calculated distribution between H2S03, HSO,-, and sos- that even slight changes in pH, in the region of pH 3.5 and above, would markedly affect the proportion of undissociated sulfurous acid present. Thus the proportion of free sulfur dioxide present in the undissociated form increases from 0.5% at pH 4 to 5.5% at pH 3. With an osmophilic Zygosaccharomyces strain isolated from orange concentrate, as little as 1.5 mg. of SO. per liter completely inhibited it. They suggested addition of acid to lower the pH value as a means of obtaining better preservation with less sulfur dioxide. A t the lower pH, combination of sulfur dioxide with glucose is delayed so that this allows a longer time for a greater quantity of free sulfur dioxide to act on the microorganisms present as well as increases the proportion of the antiseptic sulfurous acid present. Bound Sulfite: It has been known for a long time that sulfurous acid when combined with aldehydes or sugars exercises practically no antiseptic action. This was first observed by Ripper in 1892 and was later observed by Harter in Germany in 1911, by Muller-Thurgau in Switzerland in 1914, and by Laborde in France in 1916 (see Monier-Williams, 1927). Bioletti and Cruess (1912) reported that free SO2 has more than thirty times the disinfecting power of bound SO2 and is sixty times as effective as bound SO2 in inhibiting fermentation. Yeasts transferred to fresh must from water suspensions containing 350 p.p.m. of free SO2 did not ferment, whereas those from grape musts containing 2400 p.p.m. of total SO2of which only 277 p.p.m. were free did ferment. Downer (1943) explained the greater susceptibility of citrus juice concentrates to fermentation at the same level of total SO2 than dilute juices, as being due to their binding a greater portion of the added SO2. Downer divided a sulfited citrus juice inoculated with yeast into two lots, to one of which

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acetone was added to bind the free SO2. This lot having a much lower concentration of free S O r fermented sooner than the untreated lot. Although i t is generally accepted that in using SOz to preserve wines, juices, and concentrates, the free SO2 is the best guide to its preservative action, the first unequivocal proof of this was made by Ingram (1918). Ingram selected a Zygosaccharomyces species which grew a t the same rate in nutrient solutions with 4 or 40% glucose, and added a total quantity of SO2 such that in the 4 % glucose most of the SO2 was free, while in the 40% glucose only a small proportion of it was free. I n this way he varied the quantity of free SOz without, also changing any other factor which might affect growth. Under these conditions the total viable yeast count was related t o the concentration of free SO2 and not to the amounts of bound SOz present. A similar correlation was found between the growth of yeast in concentrated orange juice and the concentration of free, but not the combined or total, SOZ. Ingram concluded therefore th a t the bound SO2 has little if any germicidal value. Other aspects of preservative value are discussed by Lochhead and Farrell (1936), Hamman (1951), Souci (1951), and von Schelhorn (1953). Sanitizing Action: Sulfur dioxide has long been used in winery plant sanitation for destroying undesirable microorganisms on surfaces of wood or concrete equipment, conveyors, fermenters, storage containers, wall, and floors, and for disinfecting wood and concrete platforms or crushing equipment or for spraying on grounds and pomace piles outdoors. It is also used in sanitizing bottles and corks. It has been used abroad in sanitizing equipment and bottles and closures used in sterilization-filtration of wines, fruit juices, and vinegars. I t s corrosiveness reduces its application in the fermentation industries to surfaces more resistant than the ordinary metal surfaces and its reducing action and pronounced odor and taste limit its application to other food products. It is particularly adapted t o the sterilization of wooden surfaces because of its ready penetration into the pores and its resistance to dissipation by attack upon or combination with wood. Its application in wine manufacture is described by Amerine and Joslyn (1951). Sulfurous acid and sulfite solutions may be applied continuously by sprays or by immersion for the continuous control of microorganisms on conveyors and food-processing equipment in the dried fruit industry (Vaughn et al., 1948). Efect of Sulfur Dioxide on Heat Processing: Sulfur dioxide can be combined with heat processing in preservation of juices and pur6es. In the presence of sulfur dioxide a more rapid destruction of microorganisms occurs so that lower pasteurizing temperatures and times may be used. A more rapid destruction of Byssochlamys fulva in the presence of sulfurous acid was observed by Gillespy (1946). Pederson and Tressler (1938)

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observed more rapid destruction of yeast in apple juice in the presence of sulfurous acid. This effect has long been used in the practice of preserving fruit pulps, particularly berries, by sulfiting after preheating, but it has not been investigated. Sulfur dioxide is used in small amounts t o shorten the heat treatment in the repeated two-stage pasteurization of grape juice in France (Borgstrom, 1953). Heating, however, may alter the ratio of free to total sulfur dioxide. Under some conditions it may increase the sulfur dioxide binding power, as Ingram (1949) reported for orange juice concentrate. In addition to increasing the thermal death rate of microorganisms present sulfur dioxide also increases the resistance of fruit juices and other products to change in color and flavor on heat treatment. This increase in resistance to heat damage is quite pronounced in pasteurized orange juice and is also directly observable in dehydrated vegetables. Cruess et al. (1944a, 1944b) reported that the temperature of dehydration for a number of vegetables could be increased after sulfiting by at least 18" C. (10" F.) without appreciably affecting color and flavor. The nature of this effect on reducing heat damage, however, has not been thoroughly investigated as yet. DesulJiting: The possibility of removing sulfur dioxide from fruit products preserved with sulfur dioxide has long been recognized and used in the bulk storage of various fruit products for subsequent processing. Maraschino cherries and other fruit are preserved in a sulfite brine and then, after removal of excess sulfur dioxide by a combination of leaching with water and boiling, used in preparing candied and glaced products. Cruess and Nouty (1927) were among the first to investigate the problems involved in storage and leaching of sulfur dioxide from a wide variety of fruits. It has long been a practice in British Columbia to store small fruits and fruit pulps with added sulfite for subsequent production of jams, and this method was intensively investigated by Atkinson and Strachan (1941a) in British Columbia and by Charley (1934) in England. This practice is discussed by Morris (1933) and by Cruess (1948), and will be discussed later in this review. The removal of sulfur dioxide from whole small fruits and fruit pulps, however, is never complete. The quantities of sulfur dioxide remaining are usually too small and the subsequent processing (candying, dyeing, or jam making) is such that alteration in flavor as a result of sulfuring and desulfuring is not serious. Incomplete removal, however, is a problem where the product is canned, because of interaction of the residual SO2 with tin plate. Fruit juice beverages are prepared extensively in Europe, England, Australia, New Zealand, India, Israel, and elsewhere from sulfited juices and concentrates by dilution and sweetening which tend to mask the

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undesirable sulfite odor and taste. Such products, however, are not as desirable as the fresh, frozen, or flash pasteurized juices. The possibility of removing the sulfite added as a preservative by heating to decompose the combined sulfur dioxide and remove the free sulfur dioxide by volatilization has been investigated from time to time. In the earlier methods, heating under vacuum in steam-jacketed vacuum pans was used. Cruess and Berg (1925) investigated this as a means of removal of sulfurous acid from grape syrup but did not find it to result in complete removal or to be without effect on flavor. More recently specially devised apparatus to increase removal of sulfur dioxide by mechanically agitating the juice, bubbling an inert gas through it, or by spraying the juice while under vacuum against baffles have been introduced (Fabre, 1947). I n a recently developed French process heating under vacuum is combined with a reflux condenser to recover volatile flavoring constituents. Although some French investigators claim that grape juice so desulfited is the equal in quality of the fresh, the Swiss have not accepted these claims and desulfited juice cannot be sold as fresh in Switzerland. Liithi (1950) particularly objected to the preservation of fruit juices with sulfurous acid, now used widely in many countries, particularly in warm climates, and their distribution after desulfiting. In his opinion, although sulfur dioxide is a very convenient method of preventing fermentation and changes due to oxidation, these results can be obtained without difficulty by other means and with less effect on quality. In addition to desulfiting by physical methods, chemical methods ranging from treatment with oxidizing agents like hydrogen peroxide to addition of substances to combine with free sulfur dioxide such as acetaldehyde have been proposed for the treatment of juices and dried fruits. Vacuum drying in a shelf drier proposed originally as a method of removing sulfur dioxide from dried fruits now is used for the production of apple nuggets and similar dried fruit products.

V. SULFURDIOXIDEAS

AN

INHIBITOR OF ENZYMIC AND NONENZYMIC

BROWNING The enzyme-catalyzed oxidative browning of fruit and fruit products was reviewed by Joslyn and Ponting (1951) and the nonenzymatic browning by Stadtman (1948). As pointed out by Joslyn and Ponting, sulfurous acid and sulfites have long been known to inhibit the enzymecatalyzed oxidative discoloration of fruits as a result of mechanical or physiological injury during the ’ preparation for canning, drying, or freezing. The application of sulfurous acid and sulfites to the preparation of apples for baker’s use as freshly peeled, cored, and sliced apples or as

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froaen apples is described by Joalyn and Mrak (1933), Anon. (1944a, 1945), MacArthur (1945), Joslyn and Hohl(l948) , and Tressler and Evers (1947). The preparation of prepeeled potatoes and the control of their discoloration by sulfites and other antioxidants is described by Olson and Treadway (1949). Under certain conditions when the structure of the fruit tissue and the residual oxygen content of the interior cells permits, it is sufficient to protect the exposed surfaces of the cut fruit against oxidation by dipping them into, or spraying them with, dilute sulfurous acid or sulfites. I n other cases it is necessary for the sulfur dioxide to penetrate the fruit tissue completely, and therefore the concentration of sulfite solution and duration of treatment are important (Ponting, 1944). Synergistic effects were obtained by Johnson and Johnson (1952) by using a combination of sodium chloride, ascorbic acid, and ‘sodium bisulfite by means of which it was possible to inhibit enzymatic browning by using quantities of sulfite too small t o be tasted and still completely inactivate the polyphenol oxidase in Jonathan apples. In this connection it is important to realize that in the taste appraisal of sulfited foods special attention should be taken not to offer too many samples at one time and to develop suitable scoring techniques (Boggs and Ward, 1950). Joslyn and Ponting (1951) suggested that polyphenol oxidases were the chief if not the only enzymes involved in browning of plant tissues. The role of sulfur dioxide in inhibiting this browning, however, is not known. Enzyme-catalyzed oxidation of sulfur dioxide was observed by Ponting and Johnson (1945) but not investigated. Sulfur dioxide conceivably could act by reducing oxygen and making it unavailable for oxidation or by reacting with the quinones or other intermediates in polyphenol oxidation. All naturally occurring melanins are conjugated to proteins, and Mason (1953) favors interaction of enzymically produced quinones with proteins rather than oxidation of tyrosine residues within the polypeptide chains. If interaction with proteins is also involved in enzymatic browning, then sulfur dioxide could act either upon the enzyme protein or on intermediate products of oxidation. The nonenzymatic browning of fruit and vegetable products apparently involves the interaction of sugars, organic acids, and amino acids and proteins (Stadtman, 1948). Whether decomposition of sugars t o furfural and other substances precedes condensation with nitrogenous constituents or not is not known. There is a strong possibility that the inhibition by sulfur dioxide of browning reactions involving interaction of sugars with amino acids and protein\s may be due to the stabilization of the intermediate formed. Neither the mechanism of browning in a particular instance nor the effect of conditions of sulfuring or sulfiting are known. As pointed out previously, there are little if any data on the actual

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distribution of sulfur dioxide into its various inorganic and organic states of combination. If for example inhibition of browning in dried apricots by sulfur dioxide were due to the fact that hydroxymethyl furfural and similar substances did not form from glucose hydroxysulfonic acid or that, the latter did not combine with amino acids and proteins, then control of browning would be based on conditions of sulfuring which would favor the formation and accumulation of glucose sulfonate. On the other hand, if the inhibition were due to the formation of a relatively stable sulfur dioxide compound of the intermediate amine or Schiff base then better protection would be afforded by allowing this t o be formed before adding sulfite. A third possibility is the reaction between glucose hydroxysulfonate and amines referred t o by Danehy and Pigman (1951). The inability of fructose t o form a-hydroxysulfonates and its reactivity in browning, however, would indicate that the formation of sugar bisulfite addition compounds is not the chief cause of inhibition of browning by sulfur dioxide. Hodge (1953) in his recent review of the chemistry of browning reactions in model systems pointed out some relationships between the many different types of reactions leading to the production of brown pigments a t moderate temperatures: carbonyl-amino, non-amino, oxidative, etc. He proposed a mechanism for browning in sugar-amine systems based on the Amadori rearrangement in the Maillard reaction and stressed the importance of dehydrogenated reductones in both enzymatic and nonenzymatic browning reactions. The known inhibitors for browning, cyanide, dimedon, hydroxylamine, hydrazines, mercaptans, and bisulfite, are chiefly carbonyl reagents. However, mercaptans and bisulfites, which are the best of the above inhibitors from the practical standpoint, are also reducing agents. Hodge suggested that their functions as inhibitors may be related to this property, i.e., their ability to keep reductones involved in browning in the inactive reduced form rather than the active dehydro form. The elucidation of the mechanism of inhibition of browning by sulfur dioxide thus still remains for the future and upon it will depend the more rational use of sulfur dioxide in the industry.

VI. SOURCE AND APPLICATION OF SULFUR DIOXIDE In the pretreatment and preservation of foods with sulfur dioxide, it may be applied by burning flowers of sulfur, obtained directly from underground or underwater deposits as is done in Louisiana or from iron pyrites or similar sources, in pans, specially devised burners, or as sulfur wicks or matches. Salts of sulfurous acid, particularly the alkali or acid salts (sodium or potassium bisulfite or metabisulfite), the alkali neutral

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salts (sodium or potassium sulfite) may be added as the dry chemicals t o liquid products or after solution in water. More recently liquid sulfur dioxide has been introduced commercially and this has largely supplanted other forms of sulfur dioxide. Sulfur Fumes: Sulfur wicks or matches were at one time widely used for sulfuring barrels or casks in wineries by suspending by means of a sulfur bung and burning in the cask. The undesirable melting and dropping of the sulfur from such wicks on the bottom and walls of the container was minimized by the use of sulfur cages or cups. Various devices for sulfuring empty containers of must and wine by burning the sulfur outside the cask and passing the fumes into the cask or into the wine are described in the old literature on wine making (for example, Bioletti and Cruess, 1912). The introduction of melted sulfur into the fermenter either by direct melting and deposition or by volatilization, the introduction of undesirable products of combustion from the supporting cloth, and the difficulties in measuring and regulating the amount of sulfur dioxide introduced strongly limited this method. It was believed by Bioletti and Cruess (1912) that hydrogen sulfide formed during alcoholic fermentation was derived by reduction of the sublimed or melted sulfur introduced during sulfuring. Fumes of burning sulfur were considered by Bioletti and Cruess (1912) to be the cheapest source of SO2and the best for disinfecting purposes but unsuitable for control of fermentation because of uncertainty in application and difficulty in regulating. The occurrence of hydrogen sulfide in the gases evolved during fermentation of grape and other fruit juices and of cereal beverages has been widely reported since Nessler showed in 1869 that free sulfur in a fermenting must causes the formation of hydrogen sulfide. The formation of hydrogen sulfide from sulfites and sulfates as well as from sulfur has been observed. Autolysis of yeast and the reduction of the sulfur-containing amino acids is also known to be involved. The extensive literature on hydrogen sulfide formation during fermentation has been reviewed recently by Ricketts and Coutts (1951), who have also published the results of their investigations on hydrogen sulfide formation during the fermentation of wort. They confirmed the early investigations of Osterwalder, who showed that hydrogen sulfide formation was a characteristic of certain strains of fermenting yeasts independent of the presence of free sulfur or of decomposition products, All bottom fermenting yeasts were found by Ricketts and Coutts (1951) to produce H2S during fermentation of malt worts or of sugar solutions, but the top fermenting yeasts either did not produce H2S or only in traces after a period of storage. Sulfites and sulfates stimulated H2S production. As a result of their investigations

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on the effect of various inhibitors and other substances on H2S production, they concluded that it is catalyzed by some dehydrogenase, either the triose phosphate or the alcohol dehydrogenase or both. Although sulfur fumes are used at present to but a limited extent in wine making and other alcoholic beverage industries, except for the sulfuring of small casks and barrels, the use of sulfur fumes in the dried fruit industry is still common. Cut fruits and to a limited extent grapes are treated before sun drying or dehydration with fumes of burning sulfur in specially constructed suIfur houses. Long el al. (1940) investigated the design factors involved in the construction of a suitable sulfur house to permit rapid and uniform sulfuring and the factors influencing absorption and retention of sulfur dioxide by cut fruit sulfured with fumes of burning sulfur. They reported that sulfur contaminated with petroleum oils and similar organic materials or burnt in dirty pans did not burn completely. Bisson et al. (1942) investigated this factor in more detail. Sulfur house construction and operation is described also by Mrak and Long (1941) and Phaff and Mrak (1948). The effect of time and temperature of sulfuring on absorption of sulfur dioxide by cut fruits is discussed by Fisher et al. (1942). Mrak et al. (1942) discussed the effect of certain substances and pretreatments on retention of sulfur dioxide. In beet-sugar and cane-sugar factories situated far from industrial centers where liquid SOz is produced, sulfur dioxide produced in sulfur burners still is used owing to high transport costs for the heavy steel containers necessary to store liquid SO2. Marches (1953) described the various types of sulfur furnaces used in the cane-sugar factories of Java for the generation of sulfur dioxide by burning commercial sulfur in a stream of air. He discusses in detail Blekkingh’s analysis of the process and the effect of air volume, temperature of sulfur, air, and gas, and other factors on the SO2 and SO3 content of the gas. It was found early in Java that the quality of sulfur used affected performance of the sulfur furnace. Sulfur containing impurities such as bitumen which form a slag on the burning surface hinder the evaporation of the sulfur underneath and thus decrease furnace capacity. Stirring devices by which the sulfur surface can be skimmed, thus eliminating the slag film, were introduced in Java to facilitate burning of poorer qualities of sulfur. Alkali Salts of Sulfurous Acid: The ease with which they could be used either in the dry form or as solutions has long made the alkali sulfites and bisulfites particularly attractive in sulfiting a wide variety of products. Potassium metabisulfite has long been used in wine making for disinfection, control of fermentation, and preservation. Stability, freedom from heavy metal and arsenic impurities, and cost are the chief factors that determine selection of salts. In addition their possible effect in reducing

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S. BRAVERMAN

total acidity and introducing metallic cations has to be considered in Borne products. The sulfites and bisulfites differ in the ease with which they are absorbed and penetrate into fruit and vegetable tissues. In the dehydration of vegetables, the neutral sulfites or a mixture of neutral sulfites and bisulfites is preferred and this is best applied as a spray one-third of the way down a continuous blancher (Cruess and Mackinney, 1943; Mackinney, 1945) or as a series blanch in batch blanching. Sulfite solutions also are better absorbed by apples and other fruit that is treated by dipping or immersion before freezing (Joslyn and Mrak, 1933). Sulfite solutions are useful also in sulfiting apples and other cut fruits before drying (Mrak et al., 1942). Commercial sulfiting procedures are discussed by Beavens and Bourne (1945), Woodroof and Cecil (1945), and Cruess (1948). Liquid Sulfur Dioxide: The advantages of liquid sulfur dioxide were apparent to Bioletti and Cruess (1912) even before it became available commercially. They cite the accuracy with which its doses can be measured and the absence of impurities and prefer it to solutions of sulfur dioxide in water, which they found to be quite variable in strength, corrosive, and bulky and inconvenient to handle. The use of liquid sulfur dioxide in wineries has many advantages: its stability and constancy of composition, the fact that it introduces no undesirable residues (S, HzS, impurities, etc.), and its adaptability to treating wines of low fixed acidity. Its apparently greater cost is more than offset by the ease with which it can be used and its freedom from harmful impurities. Its present wide adoption in the wine and other food industries, however, depended upon the development of methods of handling, transferring, and measuring. At first it was used in larger wineries by taking off directly from a cylinder on a platform scale through suitable connections to a distributer in the fermenting or storage tank and measured by weighing. As soon as pressure vessels fitted with calibrated gage glasses for introducing measured volumes of liquid sulfur dioxide became available, its use even in the smaller plants became widespread. The use of liquid sulfur dioxide in various food industries is described by Willson et al. (1943a, b). The effect of type of sulfur dioxide compound used and method of application on the changes occurring in color and flavor during subsequent treatment and storage is still largely unknown. Such investigations as have been made have been limited by unavailability of methods of analysis which would determine the distribution of sulfurous acid in the product. It has long been believed that fumes of burning sulfur are particularly efficacious in the sulfuring of cut fruits (Nichols and Christie, 1930). It was believed a t one time that the presence of SO3 as well as SOz improved penetration and retention, but it is now known that tempera-

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ture and concentration of SOz are the more important factors influencing absorption and retention (Long et al., 1940). In the industry it was felt that it was desirable to sulfur the fruit initially to a sufficiently high level and that fruit that was resulfured did not store as well. The data reported by Stadtman et al. (1946) for commercially dried apricots which were resulfured to various levels with liquid SO, or by adding charcoal saturated with SO2do not indicate that resulfuring had any effect. The storage life at a given moisture content increased linearly with initial SO2 content in the range of 1500 to 8000 p.p.m. Whether this would be true for other fruits and other conditions is not known. A difference in palatability has been observed between dehydrated white potatoes sulfured by sulfite sprays during blanching and those sulfured during dehydration by the sulfur dioxide present in the hot air as a result of sulphur impurities in the oil being used as a fuel. This field still needs t o be critically and objectively investigated.

VII. SULFUR DIOXIDE IN FRUITJUICES, SYRUPS, CONCENTRATES, AND PUREES Fruit juices, particularly citrus juices, are commonly preserved with added sulfurous acid or sulfites, particularly in warm countries. They may be preserved with sulfurous acid in bulk for subsequent marketing after sweetening with sugar as beverage bases or as diluted still or carbonated beverages. With proper precautions when the initial microbial infection is low the acid juices may be preserved by the addition of SUEcient sulfite or liquid SOz to bring the total SO2 content to between 350 to 600 p.p.m. When the free SO2 content is low and the infection is high or when SOz-tolerant microorganisms are present, this may not be sufficient to preserve the juices. Under these conditions flash pasteurization, or the use of a combination of sodium benzoate and sulfites, may be practiced. To avoid the clearing of citrus juices due to the activity of the naturally occurring pectic enzymes, flash pasteurization before sulfiting is desirable. The practice of sulfite preservation of citrus juices is discussed in detail by Braverman (1949), Downer (1943), Feigenbaum and Israelshvili (1949), Tressler et al. (1939), and others. Sulfur dioxide is used both to inhibit deterioration due to yeast and mold activity and to inhibit discoloration. The latter can be accomplished usually with much lower levels of SO, than the former. In less acid juices, such as apple juice, more SO2is required (see Section IV, p. 124, and Yamada and Okumura, 1950). The concentration of SO2 required for preservation depends on the extent and type of infection, source of SOz, and composition of the juice (sugar content and pH). There is some evidence that the bisulfites and

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M. A. JOSLYN AND J. B. S. BRAVERMAN

metabisulfites are somewhat more efficient in preserving citrus juices than is sulfurous acid (Feigenbaum and Israelshvili, 1949). The addition of a small initial dose followed by storage until clarification occurs, siphoning off, and resulfiting has been found superior to preservation by a single large dose of sulfur dioxide for lime juice. Fruit juice concentrates, particularly citrus juice concentrates, are also preserved with SO2. Large quantities of lemon juice concentrated t o not less than 40" but not more than 45" Brix were packed in California for shipment to Britain in World War I1 in fir barrels preserved with 500 p.p.m. of sulfur dioxide. Orange juice concentrate, testing a minimum of 65" Brix, was also prepared in California and Florida for the Lend Lease Program but this was usually flash-pasteurized in No. 10 cans. Some was preserved with sulfur dioxide. During the War and after large quantities of orange concentrate preserved with sulfur dioxide were prepared in Palestine for shipment t o Great Britain, and this is still an important outlet for Israeli citrus concentrates (orange and grapefruit). The concentrated juices, however, because of the fact that only a small part of the added sulfur dioxide is free, require much higher concentration of total S o t , usually ranging from 1000 t o 1500 p.p.m. or more. The combination of sulfur dioxide with concentrated orange juice was investigated extensively by Ingram (1949) and Ingram and Vas (1950a, b). The rate of association and equilibrium between SO2 and glucose was investigated early by Kerp (1904, 1913) and more recently by Vas (1949) and Braverman (1953). The rate of association and extent of association both vary with concentration of sugar. In a food product containing only 2% of glucose the percentage of bound SO2 will increase only to about 15% of the total SO2 added, but in a product containing 45% glucose more than three quarters of the total sulfur dioxide will combine, leaving only about 25% of free SO2. Ingram and Vas (1950a, b) compared the relation between the combining power of SO2 and glucose concentrations in pure solution and in various citrus juices. These authors confirm an essential feature, previously noted by Downer (1943), that the combining power in various citrus juices and concentrates is always greater, and often very much greater, than that of the sugar present. They do point out, however, that these data are somewhat unreliable because the concentrations of glucose present were calculated as half the total monosaccharides (presumed t o be invert sugar) , the temperature a t which combination occurred was not known, and the quantity of added SOz varied from one set of experiments to another. In making this comparison, however, the authors took into account (as shown in Fig. 6) only the amount of glucose present in the various juices tested, and paid no attention to the total concentration of the soluble solids.

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These experiments led Ingram and Vas t o conclude that the presence of glucose is not the only factor responsible for the binding of SO2 and that other aldehyde and ketone-like substances and possibly pectin are the cause of the higher combining power of SOz in these juices. However, 9

I

I

I

9 R\@

*--5

___ - - - - -- -- - -0 l @ l I

I

6

a"

ul P

3

i I

25

I

I

50 75 X GLUCOSE OR TOTAL SOL. SOLIDS

FIG.6. The relation between the power of combining with SO2 and the glucose concentration in pure solutions and various juices. The solid black line and the numbers are those of Ingram and Vas (1950a1b). The dotted line and the encircled numbers are those of Braverman (1953). 1. 7-fold Rhodesian orange.

2. 6 34-fold U. S. and Palestine orange (250 samples within rectangle). 5. 4-fold orange. 6. 4-fold orange.

7. 8. 9. 10. 13. 14. 15.

4-fold grapefruit. 4-fold lemon. %fold u. s. lemon. 3-fold Jamaica grapefruit. Natural grapefruit. Natural orange. Natural lemon.

by replotting the results obtained by these authors, not against glucose contents but against the total concentration of soluble solids, it can be shown that all the points lie on the same curve as that of pure glucose. This clearly indicates that although glucose is largely responsible for

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M. .4. JOSLYN AND J. B. S. BRAVERMAN

binding SO2 the degree of combination is entirely due to the total concentration of soluble solids in the food. (See encircled numbers in Fig. 6 . ) To prove that the combining power of SO2 is dependent largely on the total solids concentration, Braverman (1953) compared five samples each containing the same amount of glucose with varied amounts of sucrose, which does not bind SO2 a t all and which was used in this case solely t o raise the total solids in the respective solutions. The results showed that although they each contained the same amount of glucose, the percentage of combined SO2 increased progressively with increase in total sugar content. Fruit pulps and purees as well as juices and concentrates may be preserved with sulfur dioxide. This practice is quite widespread abroad, as was mentioned before, although it is not used in the United States, since preservation freezing of fruits and fruit pulps is preferable to sulfiting for subsequent use by jam and preserve industries. The preservation of fruit pulps with sulfur dioxide is described by Atkinson (1941), Atkinson and Strachan (1941), and Charley (1934). Both cold fruit pulps and hot fruit pulps may be barrelled with SO2; the latter usually require a lower concentration of SO2 for preservation. Cooked fruit pulp, preserved in barrels with sulfur dioxide, has long been used successfully by British jam manufacturers. This method was varied by research workers at the Long Ashton Research Station by the development in 1924 of the prcservation of raw fruit with a solution of sulfur dioxide in sealed containers (see Wallace and Marsh, 1953). This cold process method, however, had the disadvantage of toughening the skins of some of the fruits, progressive loss of pectin during storage, and of being limited only to the preservation of more acid fruits. In 1940 the cold process method was reinvestigated in England for the preservation of a glut of plums. For domestic use prepared tablets of sodium and potassium bisulfite were introduced in 1941. VIII. SULFURDIOXIDEIN WINE AND VINEGARMAKING Sulfur dioxide is used in wine making as a sanitizing agent to eliminate undesirable microorganisms from all surfaces of equipment and from fermenters and storage tanks which come into contact with the grape must or fruit juice that is to be fermented. It is also used to eliminate or inhibit the development of undesirable bacteria and yeast present in crushed fruit or fruit juice that is t o be fermented and to preserve the wine during storage and aging against bacterial spoilage. Bioletti and Cruess (1912) reported that very small amounts of SO2 (50 to 75 p.p.m.) are sufficient to prevent growth of mold, undesirable yeast, and acetic acid bacteria in musts and to insure pure fermentation with a starter

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of wine yeast. One hundred p.p.m. of SO2was reported by them to eliminate over 99.9% of the undesirable aerobic yeasts, Penicillium and Aspergillus spores, and acetic acid bacteria from musts. Wines made with sulfur dioxide were found to be lower in volatile acidity and of better storage quality. This observation of Cruess (191 1) confirming the European practice of sulfuring musts particularly for making dry wines in hot countries (Bioletti, 1905; Dupont and Ventre, 1906; Martinaud, 1908) was confirmed by Cruess (1935a, b). The use of sulfur dioxide in wine making was established by extensive investigations in Switzerland (Muller-Thurgau and Osterwalder, 1914) ; in France (Ribereau-Gayon, 1947; Mathieu, 1913; Bailly, 1924); in Italy (Casale, 1938; Mensio, 1917, 1930); and elsewhere (Amerine and Joslyn, 1951). The utility of sulfur dioxide in controlling growth and activity of microorganisms is based on the greater sensitivity of these organisms to SO2 and the peculiar power of adaptability possessed by yeasts. Generally yeasts are seldom destroyed by chemical preservatives; they are usually only temporarily inactivated or dormant under unfavorable conditions, awaiting a chance to become active again. Some strains of yeast can easily be “trained” to develop in the presence of SO2, especially if the amounts used are slowly increased. This important fact is used successfully in the wine industry when it is desired to ferment the must by a specially selected strain of yeast, and to avoid undesirable fermentation by other yeasts. The selected strain of yeast is “trained” to ferment in a starter containing a very small amount of S O z ;when the fermentation of the starter is in full progress, it is mixed with further quantities of must containing progressively increasing amounts of SO2. The yeasts in the starter thus become SO2-resistant and can be employed for the fermentation of the bulk of the must, to which a fair amount of SO2 has been added, and in which no microorganisms or undesirable yeast, other than the selected strain which has become SO2-resistant, will develop. Bioletti and Cruess (1912) did not believe that culture yeasts were acclimatized to SO2; on the contrary, they favored postponing their addition to a sulfited must until after the full effect of the maximum amount of free SO2 was exerted on the undesirable organisms and when the yeast starter was exposed only to the minimum amount of free SOZ. Sulfur dioxide is also used in wine making for the most important operation known as the “defecation” of must. It is customary when making white wines to allow the must to clear by settling before fermentation, and then t o draw off the supernatant clear must into the fermentation vats. The use of sulfur dioxide in this process, which may last from two t,o three days, is beneficial in several ways:

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M. A. JOSLYN AND J. B. 8. BRAVERMAN

1. It delays the free fermentation of the must until after it has cleared, 2. it assists in the hydrolysis of the pectins, which are largely responsible for maintaining the colloids and grape particles in suspension; and 3. it causes the removal of a greater quantity of living cells of wild yeast, molds, and other microorganisms, which are dragged down with the grape pulp and other particles held in suspension, in the process of defecation. Bettoli in 1911 (see Bioletti and Cruess, 1912) found that the original must, before defecation, contained over 0.5 % of suspended matter-a quantity large enough to cause injurious effect on the odor, taste, and fragrance of the wine. He also demonstrated that, by defecation with 100 p.p.m. of SO2, the number of active cells in a must can be reduced from over a million per milliliter to just a few hundred per milliliter at 22" C. (72" F.). The observations made previously on the relative preservative action of the various forms of sulfur dioxide, free and bound, apply equally well to the use of SO2 in wine making. As shown by Neuberg (1929), bound SO2has no toxic effect upon yeasts, and, therefore, it is only the remaining free part of the SO2 which should be taken into account. In the case of musts and wines, however, additional compounds besides sugars, particularly acetaldehyde, help to bind an appreciable part of the SO2.For a long time, ever since SO2came to be used generally in wine technology, the problem of bound SO2 occupied the minds of research workers. Rocques in 1897 maintained that SO2in musts and wine was bound by glucose, but Ripper (1892) and Schmidt (see Kerp, 1904) thought that sulfur dioxide in wines existed in the greater part as a combination with aldehydes. Bianconi and Bianchi (1932) tested the SO2 content before, during, and after the fermentation of grape musts, and found that during fermentation a great part of SO2 is set free (the current of C 0 2 evolved during the process assisting this liberation) , while the remaining SO2 remained combined with aldehydes. This part, in the opinion of the authors, remains quasiconstant in the must-wine substrate during the fermentation. Farnsteiner (1904), who examined the application of SO2 to agricultural products, came to the conclusion that constituents of the vegetable tissues, other than the sugars, can also interact with sulfurous acid. He thought that such complexes in wines are created not only with compounds containing an aldehydic group, but also with cellulose, proteins, organic acids, and tannin. Paris (1920), on the other hand, showed that no combinations are possible between tartaric or malic acids and SO2. Bianconi and Bianchi (1932) refuted the idea that tannin could form any

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possible combination with SO2. Paris, however, in work on white cherries, asserted that after cherry juice had fermented until the sugar content had been removed, it still contained 0.019 to 0.016% of combined SOz, or 25 to 27% of the total, and attributed this combination to pectic substances. Recent investigations with pure substances have showed that, of the sugars contained in must, only glucose combines with SO2, and that the existing aldehydes, mainly acetaldehyde, readily combine with SO2 into even more stable compounds than the glucose sulfonate. During alcoholic fermentation, however, acetaldehyde is formed in quantity, and this compound combines immediately with SO2. When SO2 is added gradually to a must in active fermentation, it is converted almost instantly into an acetaldehyde-bisulfite complex and the fermentation is not hindered. Bertin (1924) described this phenomenon as the auto-desulfiting of musts. In the course of his studies on wine, Paris observed that when increased quantities of sulfur dioxide are added to the grape musts, the percentage of free SO2 also increases. I n musts containing 18% sugar, he reported the following levels 5 minutes after the addition of the bisulfite: SO2 added, g.11. 1.6 1.7 2.0 2.5 11.5

Free SOZ, % 61 70 73 74 88

Combined

so2 39 30 27 26 12

Paris does not mention what sugars constitute the 18%. However, if one supposes about 7.5% of the above t o be glucose, Paris’s observations agree with those of other workers on different fruit juices and on pure sugar solutions. Bianconi and Bianchi (1932), on the other hand, found that 32 hours after adding 0.24 g./l. SO2 to an 18% grape must, 50.30% of the SO2 was bound, and this remained the same after 96 hours. These authors asserted that 36.2% of the SO2 was bound with glucose and the rest with fructose. These results must definitely be regarded with reserve, for fructose has been shown to be incapable of combining with SO2a t all. Moreau and Vinet (1937a, b) reported that the true antiseptic power of sulfurous acid is dependent on the proportion of free undissociated acid present and showed this to vary with acidity. During the period of 1928 to 1938 they studied the conditions in wine and must that determine the equilibrium between free and combined sulfur dioxide and developed an iodine index for determining this (Vineau and Moreau, 1937; Amerine

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M. A. JOSLYN AND J. B. S. BRAVERMAN

and Joslyn, 1951; Ribereau-Gayon, 1947). Oversulfiting of wine as a result of absorption of SO2 from sulfured casks and vats was observed by Wanner (1938). The sulfur dioxide fixing power of wine was investigated most recently by Procopio (1953). In controlling fermentation, the usual practice is to add all the SOz required at one time. In special fermentations, however, such as that of sweet table wines like Sauternes, several small additions of SOz are used to check the fermentation while fermentable sugar still remains. Quinn (1940) has recommended frequent small additions of sulfur dioxide to wine rather than an occasional large addition to take advantage of the germicidal power of free sulfur dioxide. Peynaud and Lafourcade (1952) recommend against the use of high amounts of SO2, either in settling the must or addition prior to or during fermentation, because under these conditions acetaldehyde accumulates and a larger percentage of the sulfur dioxide added is fixed. When possible sulfur dioxide should not be added in the presence of yeast. I n adding sulfur dioxide to new wines they recommend a large initial addition to cause precipitation of as much of the yeasts as possible rather than several smaller additions. In the bulk storage of wine following fermentation, it is necessary to maintain the sulfur dioxide level a t 50 to 75 p.p.m. t o prevent bacterial spoilage. Here too it is the free sulfur dioxide rather than the total that determines keeping quality, but unfortunately data on the changes in equilibrium states are too limited to allow for adequate control of the level of free SO2 present. In addition to grape wines, sulfur dioxide is used in fruit wine making. Its use in preserved cider was reported by Lindet (1922) and Durham (1909). The use of sulfur dioxide in preparing and preserving light sweet wines was investigated by Mills and Wiegand (1942) and Yang and Wiegand (1949). Yang and Wiegand (1951) described a method of maintaining free sulfur dioxide content in wine. Yang (1953) cautioned against excessive use of sulfur dioxide in fruit wines and recommended only 100 p.p.m. In spite of the extensive investigations on wine making, there are surprisingly few complete data on the changes in free and bound sulfur dioxide during fermentation and storage. The most extensive data published previously were those of Bioletti and Cruess (1912), but even these are insufficient to determine changes due to volatilization and oxidation during fermentation and storage. Complete analyses of the distribution of SO2 in wine8 and of changes in the sulfur compounds during storage are not available. Mills and Wiegand (1942) reported data on total SO2 before and after storage for six months in the bottle of 210 samples of

'

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different wines. The sulfur dioxide lost during storage in full sealed containers in wines containing initially 116 to 398 p.p.m. of SO2 varied from 37 to 176 p.p.m., or from 14 to 53% of the initial SO2 content. Alcohol content (11.8 to 20.0%) had no influence on SOz loss, but this loss was affected by the initial concentration and by the sugar content. The utility of sulfur dioxide in cider vinegar manufacture was pointed out early by Cruess et al. (1915), and it is now widely used to reduce losses by incomplete alcoholic fermentation. In modern vinegar generator practices considerable quantities of SO2 are tolerated in the vinegar stock and the sulfiting of vinegar stock to 50 p.p.m. or over has been found useful in obtaining better color retention in red wine vinegar. IX. SULFUR DIOXIDE IN DEHYDRATED AND DRIED FRUITAND VEGETABLE

PRODUCTS

In the preparation of fruits and vegetables for drying or dehydration, sulfur dioxide is added for the preservation of color and flavor during processing and subsequent storage. Sulfuring or sulfiting is used to prevent enzyme-catalyzed oxidative changes during preparation and also to prevent microbial deterioration and facilitate drying by plasmolyzing the cells. In the dehydration and drying of fruits, however, it is used primarily to inhibit the lionenzymatic browning reaction occurring during storage a t room temperatures and above. I n the dehydration of vegetables it is used to preserve both color and flavor. Dried and Dehydrated Fruit: To maintain the desired qualities in dried fruit it has been found necessary to incorporate in it an excess of sulfur dioxide to allow for losses occurring during drying, processing, and storage. The actual sulfur dioxide content needed will vary with the type of fruit, the final moisture content, and storage conditions. Long et al. (1940) recommend the following as a guide, in parts per million, of SO2 at the drying yard : apricots-2000; peaches and nectarines-2000; pears--1000 ; golden bleach raisins-800; sulfur bleach raisins-1500; and apples--800. Extensive data have been obtained on the various factors involved in SO2 absorption and retention and on the factors that determine the role of SO2in inhibiting browning. The earlier investigations in California are summarized by Nichols and Christie (1930), Nichols and Cruess (1932), Roleson and Nichols (1933), Nichols ef al. (1938), and Long et al. (1940). The earlier investigations in Australia are discussed by Jewel1 (1927, 1937) and Quinn (1926); in South Africa, by Anderssen (1929). The general principles of sulfuring cut and whole fruit are discussed by Nichols et al. (1925), Cameron et al. (1929), Chaw ct al. (1941), von h e s e c k e ( 19-23), Perry ct nl. (1 9 M ) , Morris (1 947), C'riirss (1 948), and

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M. A. JOSLYN AND J. B . 8. BRAVERMAN

Mrak and Mackinney (1951). Instructions for sulfur house operation are given by Phaff and Mrak (1948), for sun drying fruits by Mrak and Phaff (1949), and for dehydrating freestone peaches by Mrak and Perry (1948). The relation of total sulfur dioxide content and other factors to color changes during storage of apricots is discussed by Nichols and Reed (1931), Nichols et al. (1938), Chace et al. (1930, 1933), and Sorber (1944). The role of sulfuring in dehydration of cherries and small fruits is discussed by Wiegand et al. (1945), and the effects of methods of sulfuring, dehydration, and temperature of storage on ascorbic acid content and carotene content of dehydrated peaches by Eheart and Sholes (1946), as well as others mentioned previously. In previous investigations on the influence of sulfur dioxide and other factors on the storage deterioration of dried apricots, the quality of the fruit was poorly described. Stadtman et al. (1946) developed a visual index of darkening based on comparison of a 50% alcohol extract of the fruit with the color of a standard solution (containing cobaltous sulfate, potassium dichromate, and cupric sulfate) representing the color of the fruit at the limit of edibility. Using this darkening index they investigated the effect of moisture content, sulfur dioxide content, and storage conditions (temperature, oxygen content, etc.) on the storage deterioration of dried apricots. The results obtained have been reviewed already (Stadtman, 1948). Dehydrated Vegetables: The earlier investigation on the methods of dehydration of vegetables were reviewed by Cruess and Mrak (1940, 1942a, b), Cruess and Mackinney (1943), and Anon. (1944b). The treatment of certain vegetables (particularly cabbage and white potatoes) with sulfite solutions was used and recommended in England, Australia, and Canada a t the start of World War I1 but was not applied commercially in the United States until some two years after its use abroad. Dehydrated cabbage is usually sulfited to 750 t o 1500 p.p.m., and dehydrated white potatoes and carrots to not over 500 p.p.m. Sulfiting not only improves the retention of color, flavor, and carotene and ascorbic acid content but makes possible the use of higher finishing temperatures and consequently shorter drying times. With cabbage, finishing temperatures 5 to 8" C. (10 to 15' F.), higher can be employed safely without scorching (Mackinney et al., 1943; Mackinney and Howard, 1944; Friar and Van Holten, 1945). Vegetables before dehydration may be sulfured and blanched by immersion in hot water containing sodium sulfite, as was first proposed in England and also used in Canada, or by applying a sulfite solution as a spray on cabbage as it is conveyed through the steam blancher, as was done commercially in Australia, New Zealand,

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and the United States. In the United States sulfite spray was applied to partly blanched vegetables and in Australia to vegetables either a t the entrance or exit of the blancher. The sulfite solution used varied in concentration and composition depending on the method of application. Neutral salts or a combination of neutral sulfite with the metabisulfite have been used. The high retention of vitamins and palatability of sulfited dehydrated cabbage in large-scale food service was pointed out by Fenton et al. (1946). The use of SO2 in the dehydration of eastern potatoes and other vegetables was discussed by Green et al. (1946). A considerable amount of work has been carried out on the changes in sulfur dioxide content of vegetables during dehydration and storage and the relation of sulfite content and sulfite disappearance to changes in color, flavor, and nutritive value during storage. Many of these data, however, have not been published as yet and appear only in reports for restricted distribution. The behavior of sulfur dioxide in dehydrated vegetables was investigated by Mangan and Doak (1949) in New Zealand. The most extensive investigations available on the effect of sulfuring on susceptibility t o browning and on sulfite disappearance and browning of dehydrated sulfited vegetables are those of Legault and his collaborators at the Western Regional Research Laboratories (Legault et al., 1947, 1949, 1951 ; Hendel el al., 1953). Although extensive investigations have been conducted in this field (both published and unpublished), the mechanism of the protective effect of sulfite in inhibiting changes in color and palatability still remains to be elucidated. Ross (1948) has reviewed the present status of our knowledge of deterioration of processed potatoes (including dehydrated white potatoes).

x. SULFUR DIOXIDEI N

“BRINING” OF CHERRIES .4ND “BARRELLING” OF FRUIT For many years cherries were prepared for subsequent dyeing and marketing in syrup by storage in barrels in a sea-water brine containing sulfur dioxide, in Italy and France. These were imported by eastern Maraschino cherry processors. In the 1930’s, as a result of the investigations of Wiegand and his collaborators in Oregon and of Cruess in California, successful methods were developed for the production of a bleached cherry of uniform light color, firm and free from surface checks, cracks, and blemishes from Pacific Coast grown sweet cherries of Royal Anne or Napoleon variety. As a result of this the tonnage of cherries barrelled on the Pacific Coast increased from less than 2000 tons in 1925 to over 27,000 tons in 1946. In addition to the white-fleshed cherries, Bings, Lamberts, Republicans, and other varieties are barrelled. The present status of the

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M. A. JOSLYN AND J. B. 8. BRAVERMAN

brined cherry industry was discussed by Wiegand (1946). A standardized procedure of brining cherries was developed by Bullis and Wiegand (1931) which involved filling xtandard 50-gallon paraffin-lined fir barrels with from 240 to 250 pounds of fresh cherries, covering them with dilute solution of sulfur dioxide (1.5%) and calcium carbonate or calcium hydroxide (0.9%), and with some agitation allowing the barrels to stand for four to six weeks or until the cherries are cured. The storage of cherries in brine was investigated by Cruess and Henriques (1932) and Cruess (1937). The processing of cherries in British Columbia was investigated by Atkinson and Strachan (1935a, b, 1941a, b), and more recently Weast (1940) has described improved methods of preparing the solution to be used in brining cherries from liquid SO2 and lime. The cracking and splitting of cherries in brine, which at one time was a serious problem in the industry, was investigated by Cruess and Henriques (1932), Cruess (1933, 1935c), and Wiegand et al. (1939). It was found by Cruess that this was related to excessively high acidity and could be prevented by the addition of the proper amount of slaked lime and even better by the use of calcium sulfite instead of sulfurous acid and lime. Wiegand et al. (1939) compared the effect of storing cherries in a 1.18% sulfur dioxide “brine” with and without the addition of various alkalis (calcium carbonate, calcium hydroxide, sodium hydroxide, and potassium hydroxide). The original brine had a pH value of 1.56, whereas that with added alkali ranged initially in pH from 1.69 to 2.2 as the percentage of SO2 neutralized increased from 12.5 to 100%. After six weeks of storage the final pH of the cherry brines ranged from 2.6 to 5.6, depending on the type of alkali used and on the extent to which the SO2 was neutralized. The maximum per cent (87.6) of perfect cherries was obtained with CaC03 in brines having a pH of 1.96. With Ca(OH)2 a maximum of 50% perfect cherries was reached a t pH 1.94; with NaOH a maximum of 55% perfect cherries at pH 2.1; with KOH the per cent of perfect cherries increased progressively as the amount of KOH added increased until 77% was achieved a t 100% neutralization of SO2 a t pH 2.1. Orange, lemon, grapefruit, and citron peel is improved in quality if it is stored prior to candying in a salt solution containing added SO2. Uniformity and translucency of color and texture is markedly improved by sulfite-brine storage in comparison with fermentation in a dilute brine. This was found particularly true of citron (Cruess and Glickson, 1932; Fellers and Smith, 1937). I n the barrelling of fruit for use in jams and preserves, both small whole fruit and halved large fruit is preserved by storage in a “brine” containing sulfurous acid and lime. The barrelling of peaches was success-

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fully carried out by Mrak and Henrique; (1932) and extended to other fruit by Mrak et al. (1934) and Weast (1915). The use of sulfur dioxide in the processing of watermelon rinds for food is described by Woodroof and Cecil (1942b). They also investigated the factors influencing the quality of preserves made from sulfited fruit (Woodroof and Cecil, 1943). The varietal adaptability of strawberries to preservation in sulfur dioxidecalcium solutions was investigated extensively by Culpepper and Caldwell (1943). Charley ct al. (1943) reported on a large-scale preservation of plums by sulfur dioxide. The general aspects of preserving fruit for subsequent, preserve and jam making or candying is discussed by Woodroof and Cecil (1942a, 1945). The ease with which the microorganisms normally causing spoilage (yeasts, molds, and acetic acid bacteria) may be inhibited by sulfite (the concentration of sulfite as SO2 in the fruit required t o preserve the fruit varies from 1500 to 2000 p.p.m. to as low as 350 p.p.m.); the close similarity of pH of fruit tissues to th at of the sulfite solutions used as preservative, which Woodroof and Cecil believe to be a factor in texture preservation; the presence of tannins and organic acids which tend to mask the taste of the small quantities of residual sulfur dioxide remaining in the finished product; and the ease with which the sulfur dioxide absorbed by the fruit during sulfiting is lost during storage (by volatilization or oxidation) and removed by soaking arid cooking, render fruits particularly suitable for preservation with sulfite solutions. With soft fruits such as peaches or strawberries a firming agent (calcium hydroxide or calcium carbonate) must be added, but with firmer fruit such as blackberries or black currants no firming agent is needed. Woodroof and Cecil (1945) give directions for preparation of stock solution of preservative sulfur dioxide solution from the salts of sulfurous acid and its use in barrelling various fruits. The British jam manufacturers use considerable amounts of strawberries, raspberries, and similar soft fruit preserved in sulfurous acid solutions without heat treatment and have experienced considerable difficulty owing to varying degrees of softening and even complete mushiness of such fruit. Pandhi (1953) critically reviewed the various speculations and observations as to the cause of this condition. Among the factors suggested as being involved were: variety, maturity, and growing conditions of fruit; concentration and type of sulfurous acid solution; manner of packing the fruit into the barrels and method of adding the preservative ; character of the water used ; time elapsed between harvesting and preserving; and storage temperature. Pandhi concluded that the softening of strawberries is probably due to pectic enzymes secreted by contaminating microorganisms (chiefly molds) and not to the activity

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of the naturally occurring pectic enzymes present in the fruit or the action of sulfurous acid itself.

XI. SULFURDIOXIDEIN TRANSPORTATION AND STORAGE OF GRAPES AND IN OTHERPRODUCTS Grapes: California table grapes are widely fumigated with sulfur dioxide before transportation to Eastern markets. This fumigation reduces molding in transit and also saves on the number of icings required en route. The concentration of SO2added, however, must be carefully controlled as damage a t the stem end of the berry is likely to occur. The sensitivity t o damage varies with the maturity and variety of grapes shipped. The utilization of sulfur dioxide in the marketing of grapes was investigated by Winkler and Jacob (1925), Jacob (1929), Asbury and Pentzer (1931), Pentzer et al. (1932), and Rose and Pentzer (1932). Dunn (1940) reported that mixing 5 g. of potassium or sodium metabisulfite per case (30 t o 32 pounds) with granulated cork packing reduces wastage of grapes by mold. Insertion of sodium or potassium bisulfite tablets or spraying the packing with bisulfite solution was suggested in South Africa. Pentzer, however, preferred addition of the metabisulfite as a separate package to adding it to the packing. Van der Plank (1939) developed a device for regulated release of SO2 in packages of stored table grapes. Although Winkler and Jacob (1925) reported that 50 p.p.m. of SO2 would approximately double the keeping quality of grapes and up t o 100 p.p.m. would not injure the color, flavor, and quality, the amount of SO2 now used is much lower, ranging from 10 p.p.m. for varieties most susceptible t o injury t o not over 50 p.p.m. Other Perishables: I n addition to grapes, prepared fresh apples for baker’s use (Joslyn and Mrak, 1930) and prepeeled potatoes (Olson and Treadway, 1949) are treated with sulfites or SO2 to prevent discoloration during preparation, storage, and distribution. Small amounts of SO2 are added also t o prepared horse-radish to prevent discoloration (Blumenthal, 1937), cut or shredded fresh vegetables for salad use are sulfited in some packs, and coconuts were a t one time preserved with SO2 (Dybowski, 1908). Use in Manufacturing: Sulfur dioxide has been used for many years as a bleaching agent and browning inhibitor in the extraction and refining of sugar from sugar beets and in some cases sugar cane (Reed, 1918). It is also used in improving the keeping quality of syrups prepared from corn starch hydrolyzates. Although the presence of small quantities of SOz (usually less than 50 p.p.m.) in sugars and syrups used in confectionery and similar products is not objectionable, refined sugars bleached by SO2,or corn syrups containing it, cannot be used in canning. I n canned

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foods internal corrosion is increased and objectionable flavors due to formation of H,S may occur. Sulfur dioxide is used in the extraction of pectin from citrus peel (Wilson, 1925) as a depolymerizing agent and is also used in the preparation of liquid pectin extracts t o produce a light-colored extract. The SO2 added t o the water extract before filtration is easily removed during concentration in oacuo. It is also used in Europe as a preservative for liquid pectin preparations to which i t is added in amounts ranging from 500 to 2000 p.p.m., usually about 1250 p.p.m. I n the United States liquid pectin preparations are preserved by heat treatment. Sulfur dioxide has also been recommended for preservation of apple pomace t o be used in pectin manufacture, Charley et al. (1942), Burroughs et al. (1953), and Mehlitz (1941) reported rapid loss in pectin content in wet apple pomace during storage and recommended the use of sulfur dioxide to prevent or reduce loss in jelly grade. Dryden et al. (1952) more recently reported data on effect of freedom from rot, cold storage of apples, and of allowing pomace to stand before drying on pectin losses. They reported excellent retention of jelly grade in apple pomace dried in a laboratory drier a t 90” C. (194” F.).

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Asbury, C. E., and Pentzer, W. T. 1931. Investigations on the transportation of California table grapes. Calif. Grower 3 (6), 6. Assoc. Offic. Agr. Chemists 1950. Official Methods of Analysis. Published by the A.O.A.C., Washington, D. C., 7th Ed., pp. 145-146, 176, 470-472. Atkin, S. 1950. Determination of sulfur dioxide in presence of sulfur trioxide. Anal. Chem. 22, 947-948. Atkinson, F. E. 1941. Preservation of fruit pulps with sulfur dioxide. Canner 94, (2), 14-15. Atkinson, F. E., and Strachan, C. C. 1935a. Cherry processing. I. Sulphur dioxide treatment with special reference to a useful acid generator. Fruit Products J. 14 (5), 136-137. Atkinson, F. E., and Strachan, C. C. 1935b. Cherry processing. 11. Leaching of sulphured stock. Fruit Products J . 14 (6), 174-176. Atkinson, F. E., and Strachan, C. C . 1941a. Preservation of fruits with sulphur dioxide in British Columbia. Fruit Products J. 21, 5-8, 43-45, 60, 110-112, 141-144, 153. Atkinson, F. E., and Strachan, C. C. 1941b. Candying of fruit in British Columbia with special reference to cherries. Fruit Products J. 20 (5-lo), 132-135, 16f3-169, 185, 199-201, 217, 219, 229-232, 262-264, 289, 291, 310-312, 323, 324. Bailly, M. 1924. The usefulness of sulfurous acid (in wine making). Rev. viticult. 61 145- 147. Beavens, E. A., and Bourne, J. A. 1945. Commercial sulfiting practices. Food Znds. 17 (9), 1044-1045. Bennett, A. H., and Donovan, F. K. 1943. Determination of SO$ in citrus juices. Analyst 88, 140-142. Bennett, A. H., and Tarbert, D. S. 1933. Vitamin C in citrus juices. Biochem. J. 27, 1294-1301. Benvegnin, L., and Capt, E. 1931. Du dosage de l’acide sulfureux libre et total dana les vins rouges. Mitt. Gebiete Lebensm. u. Hyg. 22, 257-263, 365-368. Berg, V. A. 1946. Binding of sulfurous acid by sugars. Vinodelie i Vinogradarstuo S.S.S.R. 6 (2), 15-18. Bertin, C . 1924. Auto-desulfiting of musts. Bull. SOC.ind. Rouen 62, 427-434. Beuschlein, W. L., and Simenson, L. 0. 1940. Solubility of sulfur dioxide in water. J. Am. Chem. SOC.62, 610-612. Beythien, A. 1904. tfber die Verwendung der schwefligen Siiure als Konservierungsmittel, insbesondere den jetzigen Stand der Beurteilung geschwefelten Dorrobstes. 2. Untersuch. Nahr.-u. Genussm. 8, 36-53. Beythien, A., and Bohrisch, P. 1902. uber amerikanisches, getrocknetes Obst. 2. Untersuch. Nahr.-u. Genussm. 6, 401-409. Beythien, A., and Bohrisch, P. 1903. uber gesehwefeltes Dijrrobst. 2. Untersuch. Nahr.-u. Genussm. 6, 356-360. Bianconi, A., and Bianchi, A. 1932. Sui rapporti di indole chimica fra l’acido sulforoso ed i composti a funzione aldeidice e chetonica. Ann. chim. appl. 22, 291-316. Bioletti, F. T. 1905. The manufacture of dry wines in hot countries. Calif. Agr. Ezpt. Sta. Bull. 167, 1-66. Bioletti, F. T. 1911. Sulfurous acid in wine making. OGg. Corn. 8th Intern. Congr. Appl. Chem. 14, 31-59. Bioletti, F. T., and Cruess, W. V. 1912. Enological investigations. Calif. Agr. Expt. Sta. Bull. No. 230, 25-64. Risson, C. S., Atlinger, H. W., and Young, H. A. 1942. Some factors affecting the burning of sulfurs used in sulfuring fruits. Hilgardia 14 (6), 361-372.

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Black, J. W.et al. lY28. Determination of SO2. Analyst 63, 118-129. Black, J. W., and Warren, B. J. W. 1928. Notes on the effect of other reducing substances on the determination of sulfur dioxide. Analyst 63, 130-132. Blumenthal, S. 1937. Relationship between organic sulfur and sulfites in the prepared horseradish. Fruit Products J . 16, 239-240. Boggs, M. M., and Ward, A. C. 1950. Scoring techniques for sulfited foods. Food Technol. 4, 282-284. Borgstrom, G. 1953. Preservation of fruit juices by sterilization-filtration, chemical preservatives, etc. I n “Chemistry and Technology of Fruit and Vegetable Juice Production” (D. K. Tressler and M. A. Joslyn, eds.). Avi Publishing Co., New York, Chapter 6. Branch, G. E. K., and Joslyn, M. A. 1935. The kinetics of the auto-oxidation of catechol in the presence of several foreign substances. J. Am. Chem. SOC.67, 2388-2394.

Braverman, J. B. S. 1949. “Citrus Products. Chemical Composition and Chemical Technology.” Interscience Publishers, New York, see pp. 249-257. Braverman, J. B. 8. 1953. The mechanism of the interaction of sulphur dioxide with certain sugars. J . Sci. Food Agr. 4, 540-547. BrBmond, E. 1937. L’anhydride sulfureux en oenologie. Rev. agr. ( N . Africa) 37, 518523, 533-537, 548-552.

Brooks, R. 0. 1928. A new rapid method for SO2 in dried fruits. Chemist Analyst 17 (11, 3.

Browne, H. H. 1944. Evidence of the existence of bisulfite compounds of sugar. J . Org. Chem. 9, 477-483. Buchman, E. R. 1936. Studies of crystalline vitamin BI. XIV. Sulfite cleavage. IV. The thiazole half. J . Am. Chem. SOC.68, 1803-1805. Buchman, E. R., Williams, R. R., and Keresztesy, J. C. 1935. Studies of crystalline vitamin B1. X. Sulfite cleavage. 11. Chemistry of the basic product. J . Am. Chem. SOC. 67, 1849-1851. Bullis, D. E., and Wiegand, E. H. 1931. Bleaching and dyeing Royal Ann cherries for Maraschino or fruit salad use. Oregon Agr. Expt. Sta. Bull. No. 276, 1-30. Bunte, H. 1873. Ueber Aethylaldehyd-schweflige Salze und die Einwirkung des schwefligsauren Natrons Aethylidenchlorin. Ann. 170, 305-310. Burroughs, L. F., Kieser, M. E., Pollard, A., and Steedman, J. 1943. The treatment of apple pomace prior to drying for subsequent pectin extraction. Ann. Rep!., Agr. and Hort. Research Sta., Long Ashton, Bristol 1948, 136-140; Fruit Products J . 24, 4-6. 1944.

Busing, K. H., and Raabe, W. 1938. The vitamin C content of grape juice preserved with sulfur dioxide. Klin. Wochschr. 17, 1766. Butleetone, N. C. 1928. The determination of sulfur dioxide in beer. Brewers’ J. (London) 64, 143-147. Cameron, 5. S., Guinn, G., Savage, C. G., Jewell, W. R., and Lyon, A. V. 1929. The sulfuring of dried fruits. J . Council Sci. Ind. Research 2, 151-155. Casale, L. 1938. L’acide sulfureux en vinification. 6th Cmgr. Intern. Vigne et Vin, Lisbon 2, 86-94. Casanave, M. P. 1910. Determination of sulfurous anhydride and a new method for the estimation of the same in free and combined conditions. Ann. fals. et fraudes 3, 154-160. Chace, E. M., Church, C. G., and Sorber, D. G. 1930. Large-scale experiments in sulfuring apricots. I. Ind. Eng. Chem. 22, 1317-1320.

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Chace, E. M., Church, C. G., and Sorber, D. G. 1933. Large-scale experiments in sulfuring apricots. 11. Effect of dehydrating, shade drying and blanching. Znd. Eng. Chem. 26, 1366-1370. Chace, E. M., Noel, W. A., and Pease, V. A. 1941. Preservation of fruits and vegetables by commercial dehydration. U.S. Dept. Agr. Circ. No. 619, 1-40. Charley, V. L. S. 1934. The preservation of fruit with 902. Effect of hot and cold methods on residual SO2 content of jam. Ann. Rept. Long Ashton Research Sta. 1984, 255-260. Charley, V. L. S., Burroughs, L. F., Kieser, M. E., and Steedman, J. 1942. The treae ment of apple pomace prior to drying for subsequent pectin extraction. Ann. Rept., Agr. and Hort. Research Sta., Long Ashton, Bristol 1942, 89-100. Charley, V. L. S., Hopkins, D. P., Kieser, M. E., and Sills, V. E. 1943. The large scale preservation of plums by the sulfur dioxide cold process. Fruit Products J . 22,362-366. Cruess, W. V. 1911. The effect of sulfurous acid on fermentation organisms. J . Znd. Eng. Chem. 4, 581-585. Cruess, W. V. 1932. Hydrogen-ion concentration in preservative action. Znd. Eng. Chem. 24, 648-649. Cruess, W. V. 1933. Uee of calcium sulfite for the storage of cherries. Fruit Products J . 12 (8), 230. Cruess, W. V. 1935a. Notes on producing and keeping wines low in volatile acidity. Fruit Products J . 16, 76-77, 108-109. Cruess, W. V. 1935b. Further data on the effect of SO2 in preventing high volatile acidity in wines. Fruit Products J . 16, 324-327, 345. Cruess, W. V. 1935c. Splitting of cherries in brine. Fruit Products J . 14, 271-272. Cruess, W. V. 1937. Maraschino cherries. Fruit Products J . 16 (9), 263-265, 277, 279, 281, 283. Cruess, W. V. 1948. Commercial Fruit and Vegetable Products. McGraw-Hill Book Company, New York, 3rd Ed., see pp. 314-315, 472-596, 597-617, 69&761, 762-80 1. Cruess, W. V., Balog, E. G., Friar, H. F., and Lew, M. 1944a. Effect of sulfiting on dehydration temperature. Canner 98 (5), 18. Cruess, W. V., Balog, E. G., Friar, H. F., and Lew, M. 1944b. Sulfiting to improve vegetables for dehydration. Food Packer 26 ( l ) , 31, 62. Cruess, W. V., and Berg, E. W. 1925. Removal of sulfurous acid from grape sirup. Znd. Eng. Chem. 17, 849-850. Cruess, W. V., and Glickson, D. 1932. Observations on brining and candying of citron. Fruit Products J . 12, 17-18, 25. Cruess, W. V., and Henriques, V. de F. 1932. Experiments on the storage of cherries in “brine.” Fruit Products J . 11, 231-234. Cruess, W. V., and Irish, J. H. 1932. Further observations on the relation of pH to toxicity of preservatives to micro-organisms. J . Bacteriol. 2S, 163-166. Cruess, W. V., and Mackinney, G. 1943. The dehydration of vegetables. Calif. Agr. E q t . Sta. BUZZ. NO. 680, 1-76. Cruess, W. V., and Mrak, E. M. 1940. The dehydration of vegetables. Fruit Products J . 20, 100-103. Cruess, W. V., and Mrak, E. M. 1942a. The dehydration of vegetables. Fruit Products J . 21, 201-204, 241-242, 269-272, 302-307, 337-340. Cruess, W. V., and Mrak, E. M. 1942b. What’s known today about dehydrating vegetables. Food Znds. 14 ( l ) , 57-60; (2), 41-43, 96-97.

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Cruess, W. V., and Nouty, A. H. 1927. Studies on the preservation of fruits in sulfurous acid solutions. Fruit Products J . 6 ( l l ) , 18-19. Cruess, W. V., Richert, P. H., and Irish, J. H. 1931. The effect of pH on the toxicity of several preservatives to microorganisms. Hilgardia 6 (lo), 295-314. Cruess, W. V., Zion, J. R., and Sifredi, A. V. 1915. The utility of sulfurous acid and pure yeast in cider vinegar manufacture. J . Znd. Eng. Chem. 7 , 324-325. Culpepper, C. W., and Caldwell, J. S. 1943. Varietal adaptability of strawberries to preservation in sulfur dioxide-calcium solution. Fruit Products J . 23 (1 and 2), 5-9, 27, 29, 46-51. Danehy, J. P., and Pigman, W. W. 1951. Reactions between sugars and nitrogenous compounds and their relationship to certain food problems. Advances i n Food Research 3, 241-290. Davis, M. B., and staff, Division of Horticulture, Dominion Dept. Agr. (Canada). 1943. Dehydration of Vegetables in Canada. No. 3, 43 p., and No. 5, 27 p., of a series of booklets reprinting articles which have appeared in “Food in Canada.” Consolidated Press Ltd. Dept. of Scientific and Industrial Research and Ministry of Food. 1943. Dehydration. United Kingdom Progress Reports, printed by H. M. Stationery Office, London (Restricted Circulation). Sections 1-11. (See particularly Section VI. Vegetables; Section IX. Bacteriological Control.) Division of Food Preservation, C.S.I.R. (Australia) 1942. The Dehydration of Vegetables. Mimeo. 42 p. Downer, A. W. E. 1942. The stability of ascorbic acid in citrus fruit juice products. J. SOC.Chem. Znd. 61, 80-82. Downer, A. W. E. 1943. The preservation of citrus juices with sulfurous acid. J. SOC. Chem. Ind. 62, 124-127. Dryden, E. C., Willaman, J. J., and Hills, C. H. 1952. Some factors affecting pectin grade of apple pomace. Food Technol. 6, 474-478. Dunn, T. C. 1940. Sulfite preservation of table grapes. J . Dept. Agr. W . Australia NO.17, 439-443. Dupaigne, P. 1951. Micro-determination of sulfur dioxide in grape juice. Ann. fals. el fraudes 44, 111-121. Dupont, E., and Ventre, J. 1906. L’acide sulfureux en vinification. Progr. agr. et vit. 46, 157-161. Durham, H. E. 1909. Studies on ciders and Perry. I. Sulfite preservation. J . Hyg. 9, 17-20. Dybowski, J. 1908. The preservation of coconuts. Compt. rend. 147, 756-757. Dyer, B., Taylor, G., and Hamence, J. H. 1941. The occurrence of “apparent” sulphur dioxide in dried onions. Analyst 66, 9-12. Eckman, J. R. 1927. Determination of sulphur trioxide in the presence of sulphur dioxide, together with some analysis of commercial liquid sulphur dioxide. Natl. Bur. Standards ( U . S . ) Sci. Paper 664, 277-285. Eheart, M. S., and Sholes, M. L. 1946. Effects of methods of sulfuring, dehydration, and temperature of storage on ascorbic acid and carotene content of dehydrated peaches. Food Research 11, 332-340. Fabre, J.-Henri. 1936. “ Analyse de Vins. et interprbtation des resultats analytiques. ProcedEs modernes de vinificetion,” Vol. 2, 2nd ed. La Typo-Litho and Jules Carbonel, Alger, see pp. 157-168. Fabre, P. 1947. Nouvelle technique de deaulfitation. Intern. Congr. Grapes, Grape Juice, Wine, Istambul, Oct. 2-7.

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Hodge, J. E. 1953. Dehydrated foods. Chemistry of browning reactions in model systems. Agr. Food Chem. 1, 928-943. Hoffpauir, C. L., and O’Connor, R. T. 1949. Sensitization of fuchsin-sulfurous reagent for formaldehyde by addition of kctoncs. Anal. Chem. 21, 420. Honig, P. (ed.) 1953. “Principles of Sugar Technology.” Elsevier Publishing Co., New York, See pp. 240-241, 582-654. Ingram, M. 1947a. An electrometric indicator to replace starch in iodine titrations of sulfurous acid in fruit juices. J. SOC.Chem. Ind. 66, 50-55. Ingram, M. 1947b. Investigation of errors arising in iodometric determination of free sulphurous acid by the acetone procedure. J . SOC.Chem. Ind. 66, 105-115. Ingram, M. 1948. Germicidal effects of free and combined sulphur dioxide. J. Sac. Chem. Ind. 67, 18-21. Ingram, M. 1949. Behavior of sulphur dioxide in concentrated orange juice. Food Research 14, 54-71. Ingram, M., and Vas, K. 1950a. Combination of sulphur dioxide with concentrated orange juice. 1. Equilibrium states. J. Sci. Food Agr. 1, 21-27. Ingram, M., and Vas, K. 1950b. Combination of sulphur dioxide with concentrated orange juice. 11. The rate of the reaction. J. Sci. Food Agr. 1, 63-66. Jacob, H. E. 1929. The use of sulfur dioxide in shipping grapes. Caalif. Agr. Expt. Sla. Bull. NO.471, 1-24. James, T. H., and Weissberger, A. 1939. Oxidation processes. XIII. The inhibitory action of sulfite and other compounds in the auto-oxidation of hydroquinone and its homologs. J. Am. Chem. SOC.61, 442-450. Jaulmes, P. 1951. “Andyse des vins” 2nd ed. Librairie Poulain, Montpellier, see pp. 423-438. Jaulmes, P., and Espezel, R. 1935. The determination of acetaldehyde in wines and spirits. Ann. fals. et fraudes 28, 325-335. Jensen, H. R. 1928. Rapid determination of sulfites by alkaline liberation, or extraction and titration. Analyst 63, 133-135. Jewell, W. R. 1927. Sulfuring dried fruit. J. Dept. Agr. Victoria 26, 457-462. Jewell, W. R. 1937. “Moist Pack” processing of dried apricots. J. Dept. Agr. Victoria 36, 498-500. Johnson, G., and Johnson, D. K. 1952. Natural flavor retained in new frozen uncooked apple pulp. Food Technol. 6, 242-245. Joslyn, M. A. 1936. Unpublished data. .Joslyn, M.A. 1941. Color retention in fruit products. Znd. Eng. Chem. S3, 308-314. Joslyn, M. A. 1953. Unpublished observations. Joslyn, M. A., Braverman, J. B. S., and Smit, C. J. B. 1950. Unpublished observations. .Joslyn, M. A., and Comar, C. L. 1938. Determination of acetaldehyde in wines. Ind. Eng. Chem., Anal. Ed. 10, 364-366. Joslyn, M. A., and Hohl, L. A. 1948. The commercial freezing of fruit products. Calif. Agr. Expt. &a. Bull. No. 703, 1-108. .Joslyn, M. A., and Mackinney, G. 1938. The rate of conversion of chlorophyll to pheophytin. J. Am. Chem. SOC.60, 1132-1136. Joslyn, M. A,, and Mrak, E. M. 1930. Prepared fresh apples for baker’s use. Fruit Products J . 9, 309-311. Joslyn, M. A., and Mrak, E. M. 1933. Investigations on the use of sulfurous acid and sulfites in the preparation of fresh and frozen fruit for baker’s use. Fruit Products J . 12, 135-140. Joslyn, M. A., and Ponting, J. D. 1951. Enzyme-catalyzed oxidative browning of fruit products. Advances i n Food Research 8, 1-44.

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