Dependent Reduction of Nitrite to Nitric Oxide by Thiocyanate and

Dec 6, 2011 - Oxide by Thiocyanate and Accumulation of Iron(II)/Thiocyanate/ ... ABSTRACT: Iron(III) ingested as a food component or supplement for...
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Enhancement of Iron(II)-Dependent Reduction of Nitrite to Nitric Oxide by Thiocyanate and Accumulation of Iron(II)/Thiocyanate/ Nitric Oxide Complex under Conditions Simulating the Mixture of Saliva and Gastric Juice Umeo Takahama*,† and Sachiko Hirota‡ †

Department of Bioscience, Kyushu Dental College, Kitakyushu 803-8580, Japan Department of Nutrition, Kyushu Women’s University, Kitakyushu 807-8586, Japan

‡

ABSTRACT: Iron(III) ingested as a food component or supplement for iron deficiencies can react with salivary SCN− to produce Fe(SCN)2+ and can be reduced to iron(II) by ascorbic acid in the stomach. Iron(II) generated in the stomach can react with salivary nitrite and SCN− to produce nitric oxide (NO) and FeSCN+, respectively. The purpose of this investigation is to make clear the reactions among nitrite, SCN−, iron ions, and ascorbic acid under conditions simulating the mixture of saliva and gastric juice. Iron(II)-dependent reduction of nitrite to NO was enhanced by SCN− in acidic buffer solutions, and the oxidation product of iron(II) reacted with SCN− to produce Fe(SCN)2+. Almost all of the NO produced was autoxidized to N2O3 under aerobic conditions. Iron(II)-dependent production of NO was also observed in acidified saliva. Under anaerobic conditions, NO transformed Fe(SCN)2+ and FeSCN+ to Fe(SCN)NO+ in acidic buffer solutions. Fe(SCN)NO+ was also formed under aerobic conditions when excess ascorbic acid was added to iron(II)/ nitrite/SCN− systems in acidic buffer solutions and acidified saliva. The Fe(SCN)NO+ formed was transformed to Fe(SCN)2+ and iron(III) at pH 2.0 and pH 7.4, respectively, by O2. Salivary glycoproteins could complex with iron(III) in the stomach preventing the formation of Fe(SCN)2+. Ascorbic acid reduced iron(III) to iron(II) to react with nitrite and SCN− as described above. The above results suggest (i) that iron(II) can have toxic effects on the stomach through the formation of reactive nitrogen oxide species from NO when supplemented without ascorbic acid and through the formation of both reactive nitrogen oxide species and Fe(SCN)NO+ when supplemented with ascorbic acid, and (ii) that the toxic effects of iron(III) seemed to be smaller than and similar to those of iron(II) when supplemented without and with ascorbic acid, respectively. Possible mechanisms that cause oxidative stress on the stomach through Fe(SCN)NO+ are discussed.

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HNO2 + H+ ⇄ H2NO2+ ⇄ H2O + NO+

INTRODUCTION It is known that the intake of iron from foods does not induce significant symptoms of iron poisoning, but the excess iron intake as a supplement for anemia induces pain in the stomach due to ulceration of the stomach lining.1 Iron poisoning has been discussed to be due to the generation of reactive oxygen species such as OH radicals that lead to oxidation of cellular components such as lipids, proteins, and nucleic acids.1−3 However, nitrate in foods is absorbed in the human body by the intestine, and part of the nitrate is secreted into the oral cavity as a component of saliva. The secreted nitrate is reduced to nitrite by nitrate reducing bacteria.4−8 Because nitrite (0.05−1 mM) is present in saliva,5,7,9−11 the supplementation of iron(II) can result in the production of NO by the redox reaction between iron(II) and nitrite in the stomach.12 Several reactions are possible for nitrite-dependent NO production in the stomach. One is self-decomposition of nitrous acid,13,14 NO2− + H+ ⇄ HNO2

(pKa = 3.1)

© 2011 American Chemical Society

(K = 3 × 10−7M−1)

(2)

2HNO2 ⇄ N2O3 + H2O

N2O3 ⇄ NO + NO2

(K = 3 × 10−3M−1)

(K = 4.8 × 10−3M)

(3) (4)

The above pKa and equilibrium constants (K) are from ref 13 and refs cited therein. The other is the reduction of nitrous acid by gastric ascorbic acid (0.05 to 0.1 mM).15−19 In the stomach, salivary nitrite reacts with salivary SCN−.13,20,21

HNO2 + SCN− + H+ ⇄ ONSCN + H2O

(5)

The equilibrium constant and rate constant of reaction 5 are 44 M−2 and 1.17 × 104 M−2 s−1, respectively.20 ONSCN produced by reaction 5 can react with ascorbic acid. Received: October 8, 2011 Published: December 6, 2011

(1) 207

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It has been reported that the concentrations of nitrite and SCN− in human saliva are 0.05−1 mM4,6,9−11,34,35 and 0.1−2 mM,10,11,34−36 respectively. Oxygen Consumption. Nitrite-induced O2 consumption was measured at 30 °C using a Clark-type electrode (Rank Brothers, Cambridge, UK). The polarizing voltage was −0.6 V. Reactions were run in 50 mM KCl−HCl (pH 2.0) in the presence of various concentrations of NaNO2, NaSCN, FeSO4, and FeCl3. When saliva filtrate was used, the reaction mixture (1.5 mL) contained 0.9 mL of whole saliva filtrate and 0.6 mL of 50 mM KCl−HCl (pH 1.3). The pH of the mixture was 1.9−2.1. NaNO2, NaSCN, FeSO4, and FeCl3 were added as required. NO Production. NO-Fe(DTCS)2 was detected using an electron spin resonance (ESR) spectrometer JE1XG (JEOL, Tokyo, Japan) at about 25 °C with a quartz flat cell (0.05 mL), and its ESR spectra were recorded under the following conditions:22 microwave power, 10 mW; scanning speed, 5 mT/min; line width, 0.5 mT; and amplification; 250-, 500-, or 1000-fold. It has been reported that NO can be trapped by Fe(DTCS)3 and that the yield of NO-Fe(DTCS)2 is 40% in a sodium-phosphate saline solution.37 Solution of complex of iron(III) and DTCS [Fe(DTCS)3] was prepared by adding 0.03 mL of 100 mM FeCl3 to 1 mL of 10 mM DTCS, which was dissolved in 0.1 M sodium phosphate (pH 7.0). Reaction mixtures for NO production (0.1 mL) contained 0.2 mM nitrite, iron(II), and SCN− in 50 mM KCl−HCl (pH 2.0). The reaction mixtures were incubated for 1 min, and then 0.1 mL of Fe(DTCS)3 was added. Immediately after the addition of Fe(DTCS)3, an aliquot of the mixture was withdrawn into the quartz flat cell. Recording of ESR spectra by repeat scanning was started 1.5 min after the addition of Fe(DTCS)3. The pH after the addition of Fe(DTCS)3 was about 6.8. NO production was also studied under anaerobic conditions. Anaerobic conditions were established by bubbling argon gas through the reaction mixture for 1 min. The anaerobic reaction mixtures were incubated for 1 min after the addition of nitrite, and then Fe(DTCS)3 was added. The decrease in O2 concentration was ascertained with an oxygen electrode. When NO production was measured using acidified saliva, the reaction mixture contained 0.3 mL of saliva filtrate and 0.2 mL of 50 mM KCl−HCl (pH 1.3). The pH of the mixture was 1.9−2.1. Acidified saliva was incubated for 1 min under various conditions, and then 0.1 mL of acidified saliva was added to 0.1 mL of Fe(DTCS)3. ESR spectra were recorded as described above. The pH of the mixture was about 6.7. NO production was also measured at 30 °C using a Clark-type electrode31,38 in 2 mL of 50 mM KCl−HCl (pH 2.0) and the mixture of 0.9 mL of whole saliva filtrate and 0.6 mL of 50 mM KCl−HCl (pH 1.3). The polarizing voltage was −0.7 V. After excluding air by bubbling argon gas through the buffer solution and the mixture, various reagents were added. NOC 7-induced NO production in 50 mM KCl−HCl (pH 2.0) was used to calibrate NO production by iron(II)/nitrite and ascorbic acid/nitrite systems. Redox Reactions in Iron(II)/Iron(III) Systems. Nitrite-induced oxidation of iron(II) to iron(III) was studied at 30 °C using a model 557 spectrophotometer equipped with an end-on type photomultiplier (Hitachi, Tokyo, Japan). The path length of the measuring beam was 4 mm. The reaction mixture (1 mL) contained 1 mM NaSCN and 1 mM NaNO2 in 50 mM KCl−HCl (pH 2.0). Reactions were initiated by adding 0.5 mM FeSO4. The oxidation of iron(II) to iron(III) was estimated by recording absorbance increase at 450 nm because SCN− could make a complex with iron(III) to produce brown Fe(SCN)2+, which had an absorption peak around 450 nm.32 Anaerobic conditions were established by replacing air with argon gas, and reactions were initiated by adding FeSO4. Formation of Fe(SCN)2+ was also studied in acidified saliva. The reaction mixture (1 mL) contained 0.6 mL of saliva and 0.4 mL of 50 mM KCl−HCl (pH 1.3). Ascorbic acid, nitrite, irons, and SCN− were added as required. Data Presentation. Each experiment was repeated at least twice, and essentially, the same results were obtained. Typical data or averages of two or more experiments are presented.

ascorbic acid + 2ONSCN → dehydroascorbic acid + 2NO + 2SCN− (6) + 2H+ 6 −1 −1 14 The rate constant is 5 × 10 M s . If iron(II) is present, the following reaction is also possible:12

Fe2 + + HNO2 + H+ → Fe3 + + NO + H2O

(7)

In addition to reaction 7, iron(II) may be able to react with ONSCN. (8) Fe2 + + ONSCN → Fe3 + + NO + SCN− The occurrence of this reaction is deduced from the reaction of ONSCN with ascorbic acid (reaction 6) and melanoidin.22 If iron(II) enhanced NO production in the stomach, N2O3 production is also enhanced.23−25

4NO + O2 → 2N2O3 (k = 2 − 6 × 106M−2s−1)

(9)

+

N2O3 can function as NO donors, can be transformed to NO2 by reaction 4, and can be hydrolyzed to nitrous acid. However, NO reacts with HO2 (pKa = 4.8), the production of which is enhanced by iron supplementation,2,3 to produce peroxynitrous acid (ONOOH, pKa = 6.8). ONOOH is also produced by the reaction of nitrous acid with H2O2 that was formed by disproportionation of HO2. Reactive nitrogen oxide species described above can cooperate with OH radicals generated by Fenton chemistry using HO2 and H2O21−3 to cause oxidative, nitrative, and nirosative stress in the stomach. Although there are many reports on iron(II)-dependent reduction of nitrite to NO under neutral and acidic conditions in abiotic systems,12,26−30 only one report deals with the reaction of iron(II) with nitrous acid under conditions simulating the mixture of saliva and gastric juice.31 In the report, however, effects of a salivary component SCN− on the reactions are not included, although SCN−, which can react with iron ions32,33 and nitrous acid,14,20,21 is present in saliva. This article focuses on the enhancement of iron(II)-induced reduction of nitrite to NO by SCN− and the production of Fe(SCN)NO+ that is in equilibrium with FeSCN+ and NO in ascorbic acid/nitrite/iron ions/SCN− systems in acidic buffer solutions and acidified saliva. Taking obtained results into consideration, we propose the role of FeSCN+ and Fe(SCN)NO+ in iron poisoning in the stomach.

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MATERIALS AND METHODS

Reagents. N-(Dithiocarboxy)sarcosine sodium salt (DTCS) and an NO generating reagent 1-hydroxy-2-oxo-3-(N-methyl-3-aminopropyl)-3-methyl-1-triazene (NOC 7) (purity >90%) were obtained from Dojindo (Kumamoto, Japan). Griess−Romijn reagent for the determination of nitrite and other reagents were obtained from Wako Pure Chemical Industries (Osaka, Japan). Preparation of Saliva. Mixed whole saliva (about 10 mL) was collected from volunteers between 9 and 10 a.m. by chewing parafilm after informed consent had been obtained. The collected saliva was passed through two layers of nylon filter net [380-mesh (32 μm), Sansho, Tokyo, Japan] to remove epithelial cells. The filtrate was kept on ice and used as a saliva filtrate to measure O2 consumption and NO production. The concentrations of nitrite and SCN− were determined using Griess−Romijn reagent and FeCl3, respectively, as reported,21 and their concentrations were in ranges of 0.07−0.53 mM (mean, 0.23 mM; n = 9) and 0.36−0.86 mM (mean, 0.51 mM; n = 9), respectively. 208

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RESULTS AND DISCUSSION Measurements of NO Production with Fe(DTCS)3. Reactions of iron(II) with nitrite were characterized in 50 mM KCl−HCl (pH 2.0), prior to study reactions of iron(II) with salivary nitrite under conditions simulating the mixture of saliva and gastric juice. Figure 1 (right panel, inset) shows typical ESR

panels). When 0.4 mM FeSO4 was added in the absence of SCN−, nitrite-induced NO-Fe(DTCS)2 formation was about 1.5-fold greater under anaerobic than aerobic conditions (compare open squares in the right and left panels), whereas it was about 3-fold greater under anaerobic than aerobic conditions (compare closed squares in the right and left panels) when added in the presence of SCN−. The lower NOFe(STCS)2 formation under aerobic conditions indicates consumption of NO produced in the iron(II)/nitrite systems by reaction 9. Electrochemical Measurements of NO Production. Traces 1 and 3 in Figure 2 show typical time courses of

Figure 1. NO production in iron(II)/nitrite systems. The reaction mixture contained 0.2 mM NaNO2 in 50 mM KCl−HCl (pH 2.0). (○ and ●) 0; (△ and ▲) 0.1; (▽ and ▼) 0.2; (□ and ■) 0.4 mM FeSO4. Open symbols, without NaSCN; close symbols, 1 mM NaSCN. (Left) aerobic conditions; (right) anaerobic conditions. Horizontal axis, time after the addition of Fe(DTCS)3. Inset: ESR spectra. Reactions were run for 1 min in 50 mM KCl−HCl (pH 2.0). (A) without nitrite; (B) with 0.2 mM NaNO2; (C) B + 0.5 mM FeSO4.

spectra of NO-Fe(DTCS)2.37 ESR signals of NO-Fe(DCTS)2 were not detected upon the addition of Fe(DTCS)3 to 50 mM KCl−HCl (pH 2.0) in the absence of nitrite (trace A) but observed in the presence of nitrite (trace B). Much stronger signals of NO-Fe(DTCS)2 were observed upon the addition of Fe(DTCS)3 to the acidic mixture of nitrite and FeSO4 (trace C). Mixtures of 0.2 mM nitrite and various concentrations of FeSO4 were incubated for 1 min in 50 mM KCl−HCl (pH 2.0), and then Fe(DTCS)3 was added. ESR spectra of NOFe(DTCS)2 were recorded repeatedly every 1.5 min after increasing pH by addition of Fe(DTCS)3 (Figure 1, left panel). In the absence of SCN−, no changes in the signal intensity of NO-Fe(DTCS)2 were observed during repeat scanning (open symbols), indicating that NO was not produced after pH of the reaction mixture was increased to 6.8 by addition of Fe(DTCS)3. SCN− enhanced NO-Fe(DTCS)2 formation at every concentration of iron(II), and the signal intensity of NOFe(DTCS)2 increased slowly during repeat scanning (closed symbols). At present, mechanisms of the slow signal increase are unclear, but the data indicate that significant amounts of NO were not produced after the addition of Fe(DTCS)3, either, in the presence of SCN−. The above results indicate that iron(II)-induced NO production was mainly due to reactions 7 and 8 in the absence and presence of SCN−, respectively. Iron(II)-independent NO production can be explained by reactions 1−4 in the absence of SCN−, whereas, in the presence of SCN−, ONSCN formed by reaction 5 can contribute to NO production. It has been reported that ONSCN decomposes to produce NO, (SCN)2−, and (SCN)2.39 Nitrite-induced formation of NO-Fe(DTCS)2 under anaerobic conditions (Figure 1, right panel) was similar to that under aerobic conditions in the absence (open circles) and presence (closed circles) of SCN− only (compare the right and left

Figure 2. Time courses of O2 consumption and NO production in iron(II)/nitrite systems. (1 and 2) Oxygen consumption; (3 and 4) NO production. Reactions were run in 50 mM KCl−HCl (pH 2.0). Where indicated by arrows, various reagents were added. (A) 1 mM NaNO2; (B) 0.2 mM FeSO4; (C) 1 mM NaSCN; (D) 0.1 mM ascorbic acid.

iron(II)-induced O2 consumption and NO production, respectively, which were measured electrochemically. As expected from Figure 1, no significant O2 consumption and NO production were observed upon the addition of nitrite (arrows A) to acidic buffer solutions. Slow O2 consumption and NO production were observed by the addition of iron(II) in the presence of nitrite (arrows B). O2 consumption was not due to the autoxidation of iron(II) (data not shown). The rate constant of reaction 7 was calculated from rates after the addition of iron(II) (arrow B in trace 3), and the value was 61.7 ± 12.6 M−1 min−1 (mean ± SD, n = 3). SCN− enhanced both O2 consumption and NO production (arrows C), and the enhancements were (14.1 ± 4.1)- and (18.0 ± 2.6)-fold (means ± SDs, n = 3), respectively. During the enhanced O2 consumption and NO production, reaction mixtures turned to pale brown, suggesting the formation of Fe(SCN)2+, which had an absorption peak around 450 nm.32 The rate constant of reaction 8 was calculated using the concentration of ONSCN, which was determined from reaction 5 to be 4.4 × 10−7 M under the conditions of Figure 2. The value was (2.5 ± 0.3) × 106 M−1 min−1 (mean ± SD, n = 3), which was about one 209

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hundredth of reaction 6. Although SCN− enhanced greatly O2 consumption and NO production in the presence of both nitrite and iron(II), no clear enhancements of O2 consumption and NO production by SCN− were observed in the presence of nitrite only (data not shown). Traces 2 and 4 are ascorbic acid-induced O2 consumption and NO production, respectively. O2 consumption and NO production observed by the addition of ascorbic acid (arrows D) were enhanced by SCN−. The amount of O2 consumed by 0.1 mM ascorbic acid was 46.7 ± 1.9 μM (mean with SD, n = 3). Since 0.1 mM ascorbic acid can reduce 0.2 mM nitrous acid to produce 0.2 mM NO,14,21 O2 consumption by 0.2 mM NO is calculated to be 50 μM by reaction 9. The production of 0.2 mM NO by 0.1 mM ascorbic acid was supported by the result that electrochemical signal changes induced by 0.1 mM NOC 7, which could produce 0.2 mM NO,40 were similar to those induced by 0.1 mM ascorbic acid within error ranges of 10% (n = 4). The concentration of NOC 7 consumed 45.3 ± 1.7 μM O2 ((mean ± SD, n = 4). The lower O2 consumption expected from NO production was due to the presence of impurities in NOC 7 (see Materials and Methods). The amounts of O2 consumed and NO produced were estimated in the presence of various concentrations of iron(II) (Table 1). The data indicate that almost all of the iron(II)

Figure 3. Effects of iron(III) on O2 consumption and NO production in iron(II)/nitrite systems. (1) O2 consumption; (2) NO production. The reaction mixture contained 0 mM (1−1 and 2−1), 0.5 mM (1−2 and 2−2), and 1 mM (1−3 and 2−3) FeCl3 in 50 mM KCl−HCl (pH 2.0). Where indicated by arrows, various reagents were added. (A) 1 mM NaNO2; (B) 0.5 mM FeSO4; (C) 1 mM NaSCN. (D) Arrows indicate the addition of 0.5 mM FeCl3 twice after the reactions had been completed.

Table 1. O2 Consumption and NO Production in Iron(II)/ Nitrite Systemsa FeSO4 (mM)

O2 consumption (μM)

NO production (μM)

0.1 0.2 0.5

24.4 ± 2.2 (n = 3) 48.0 ± 2.2 (n = 3) 123.0 ± 5.3 (n = 3)

79.0 ± 10.0 (n = 3) 126.4 ± 4.6 (n = 3) 279.1 ± 30.2 (n = 5)

suppressed by iron(III). When iron(III) was added twice to a reaction mixture after NO production had been completed, the signal of NO decreased after each addition (arrows D). The data support the postulation of the reaction of NO with Fe(SCN)2+. The result that NO produced by 0.1 mM NOC 7 was decreased by the addition of both iron(III) and SCN− but not iron(III) (data not shown) also supports the reaction of NO with Fe(SCN)2+ under acidic conditions. Formation of Fe(SCN)2+. Because reaction mixtures that contained iron(II), nitrite, and SCN− turned to pale brown as described above, formation of Fe(SCN)2+ was studied (Figure 4, top). Absorbance spectra of Fe(SCN)2+, which was produced in acidic buffer solutions containing 1 mM FeCl3 and 1 mM NaSCN, had a peak around 450 nm, and the spectra were the same as those of acidic buffer solutions that contained 1 mM FeCl3, 1 mM NaSCN, and 1 mM nitrite under aerobic and anaerobic conditions (A-3 and B-3). No detectable absorbance around 450 nm was observed in acidic mixtures of 1 mM NaSCN and 1 mM nitrite (not shown), indicating that ONSCN formed by reaction 5 did not contribute to the absorption changes in Figure 4. The concentration of ONSCN calculated using the equilibrium constant of reaction 5 under the conditions of Figure 4 was 4.4 × 10−7 M and absorbance at 460 nm was calculated to be 4.4 × 10−5 using the molar extinction coefficient of ONSCN (100 M−1 cm−1).20 When 0.5 mM FeSO4 was added to acidic solutions that contained 1 mM nitrite, 1 mM SCN− , and various concentrations of FeCl3, gradual absorbance increase was observed (A-1−A-3). The absorbance increase was confirmed to be due to the formation of Fe(SCN)2+ because of the peak around 450 nm. Under anaerobic conditions, iron(II) also increased absorbance around 450 nm suggesting the formation of Fe(SCN)2+, although the absorbance increase was smaller under anaerobic than aerobic conditions (upward arrows in B1−B-3). After the completion of reactions, air was bubbled into the anaerobic reaction mixtures. The aeration resulted in an

a Reactions were initiated by adding FeSO4 to the reaction mixture that contained 1 mM NaNO2 and 1 mM NaSCN in 50 mM KCl−HCl (pH 2.0). Amounts of O2 consumed and NO produced were estimated after reactions had been completed. Values are the means with SDs.

contributed to the formation of NO because the amounts of O2 consumed were about one-fourth of the concentrations of added iron(II), whereas the amounts of NO produced were less than that expected from the concentration of iron(II) (reaction 8) and the amounts of O2 consumed by reaction 9. The discrepancy implies that iron(III), which was produced during iron(II)-dependent reduction of nitrous acid, contributed to the lower NO production under anaerobic conditions. To confirm the iron(III)-dependent suppression of NO production, effects of iron(III) on iron(II)-induced O2 consumption and NO production were studied (Figure 3). O2 consumption induced by iron(II) in the presence of nitrite but the absence of SCN− (arrow B in 1) was not affected by iron(III), whereas the rate of O2 consumption induced by SCN− in the presence of both nitrite and iron(II) (arrow C in 1) was inhibited by iron(III) without affecting the amount of O 2 consumption (compare traces in 1). Because O 2 consumption was due to autoxidation of NO, iron(III)-induced slowing down of O2 consumption could be explained if O2 more slowly oxidized the compounds formed from NO and Fe(SCN)2+ than NO. No changes in O2 concentration were observed when iron(III) was added twice to a reaction mixture, the O2 consumption in which had been completed (arrows D). NO production induced by iron(II) (arrow B in 2) was not significantly affected by iron(III), whereas NO production observed after the addition of SCN− (arrow C in 2) was 210

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Under anaerobic conditions (traces in 2), iron(II) also decreased initially and then increased absorbance at 450 nm in the presence of iron(III) (traces 2−2 and 2−3). The absorbance decrease can be explained by the formation of FeSCN+ as described above. Absorbance increase induced by iron(II) was suppressed by iron(III), suggesting that reactions produced NO with Fe(SCN)2+ that was present in advance. Formation of iron(III) after the addition of iron(II) was calculated using trace 2−1 as described above. The value was 0.30 ± 0.01 mM (mean with SD, n = 3). Absorbance increase at 450 nm after aeration was similarly independent of iron(III) concentration (B-1−B-3 and 2−1−2−3), and the absorbance increase was estimated to be equivalent to 0.23 ± 0.02 mM iron(III) (n = 6). This result indicates that about 40% of iron(II) added is transformed to autoxidizable components under anaerobic conditions in the absence of iron(III). Reaction of Fe(SCN)2+ with NO. We postulated from Figures 3 and 4 that NO reacted with Fe(SCN)2+ decreasing the absorbance of Fe(SCN)2+. Figure 5 shows effects of NO,

Figure 4. Absorbance increase due to the formation of Fe(SCN)2+. (Top) Spectral changes. (A-1−A-3) aerobic conditions; (B-1−B-3) anaerobic conditions. Reactions were initiated by adding 0.5 mM FeSO4 to mixtures that contained 1 mM NaNO2, 1 mM NaSCN, and various concentrations of FeCl3 in 50 mM KCl−HCl (pH 2.0). (A-1 and B-1) without FeCl3; (A-2 and B-2) 0.5 mM FeCl3; (A-3 and B-3) 1 mM FeCl3. Scanning was repeated every 1 min from 600 to 350 nm. In traces B-1−B-3, air was added after the completion of repeat scanning. Upward white arrows, absorbance increases induced by iron(II). (Bottom) Time courses of absorption changes at 450 nm. (1) Aerobic conditions; (2) anaerobic conditions. The reaction mixture (1 mL) contained 1 mM NaNO2 and 1 mM NaSCN in 50 mM KCl− HCl (pH 2.0). (1−1 and 2−1) without FeCl3; (1−2 and 2−2) 0.5 mM FeCl3; (1−3 and 2−3) 1 mM FeCl3. FeSO4 (0.5 mM) or air was added where indicated by arrows.

Figure 5. Reaction of Fe(SCN)2+ with NO. The reaction mixture (1 mL) contained 0.5 mM FeCl3 and 1 mM NaSCN in 50 mM KCl−HCl (pH 2.0). (Panel A) Aerobic conditions. After the addition of 0.1 mM NOC 7, the absorption spectrum was recorded. This procedure was repeated four times. (Panel B) Anaerobic conditions. After the addition of 0.1 mM NOC 7 four times as that in panel A, 0.5 mM NOC 7 was added twice. Spectrum 1 is the absorption spectrum recorded after the last addition of 0.5 mM NOC 7. (Panel C) Absorption spectra of reduction products of Fe(SCN)2+. Spectrum 1, enlarged spectrum of spectrum 1 in panel B; spectrum 2, absorption spectrum recorded in the mixture of 0.5 mM NOC 7, 0.5 mM FeSO4, and 1 mM NaSCN in 50 mM KCl−HCl (pH 2.0) under anaerobic conditions. (Panel D) Changes in the absorption spectrum after aeration. Spectrum 1 was recoded under the same conditions as those in spectrum 2 in panel C. After aeration, scanning was repeated every 2 min from 800 to 360 nm.

absorbance increase independent of the presence and absence of iron(III), and the changes in absorption spectra also indicated the formation of Fe(SCN)2+. The above results suggest that the product formed by the reaction of NO with Fe(SCN)2+ under anaerobic conditions had a smaller extinction coefficient than Fe(SCN)2+ around 450 nm and that the products were oxidized by O2 generating Fe(SCN)2+. Figure 4 (bottom, trace 1−1) shows a time course of iron(II)-induced absorption changes at 450 nm due to Fe(SCN)2+ under aerobic conditions. In the presence of iron(III), iron(II) decreased initially and then increased absorbance at 450 nm (traces 1−2 and 1−3). The absorbance decrease might be due to the formation of colorless FeSCN+ decreasing the concentration of Fe(SCN)2+. The absorbance increase, which was due to the oxidation of iron(II) to iron(III) by nitrite, was similar in extent independent of the presence and absence of iron(III). Iron(III) formed was estimated using trace 1−1 and absorbance at 450 nm in A-2 and A-3 before adding 0.5 mM FeSO4. The value was 0.54 ± 0.03 mM (mean with SD, n = 3), indicating that almost all of the iron(II) added was oxidized to iron(III) as shown in Table 1.

which was generated using NOC 7, on absorption spectra of Fe(SCN)2+. NO did not significantly decrease the absorbance of Fe(SCN)2+ under aerobic conditions (panel A) but decreased significantly under anaerobic conditions (panel B), confirming the reaction of NO with Fe(SCN)2+. Further increase in NO concentration under anaerobic conditions resulted in the increase in absorbance around 580 nm (spectrum 1 in panel B). Its enlarged spectrum (spectrum 1 in panel C) shows a peak and a shoulder at 448 and 588 nm, respectively. Spectrum 2 in panel C (peaks, 448 and 588 nm) was recorded using the acidic mixture of NOC 7, SCN−, and FeSO4 under anaerobic conditions. Because Fe(SCN)NO+ is 211

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generated in NO, iron(II), and SCN− systems,41,42 the component that gave spectrum 2 was supposed to be Fe(SCN)NO+. Fe(SCN)NO+ may be produced as follows when NO was added to acidic and anaerobic solutions of Fe(SCN)2+:

Fe(SCN)2 + + NO ⇄ FeSCN+ + NO+

(10)

FeSCN+ + NO ⇄ Fe(SCN)NO+

(11)

NO-dependent transformation of Fe(DTCS)3 to NO-Fe(DTCS)237 supports the occurrence of reactions 10 and 11. If FeSCN+ is present in acidic reaction mixtures, Fe(SCN)NO+ is produced by reaction 11. NO+ produced by reaction 10 can be reduced by SCN− producing (SCN)2− and/or (SCN)2 and by iron(II) producing iron(III). The reactions are possible because of their reduction potentials: NO+/NO (1.21 V at pH 7),43 (SCN)2−/SCN− (1.29 V (NHE)),44 and iron(III)/iron(II) (0.77 V (NHE)). Absorbance due to Fe(SCN)NO+ disappeared by bubbling argon gas (data not shown). This result suggests the removal of NO from the mixture of Fe(SCN)NO+, FeSCN+, and NO (reaction 11) leaving colorless Fe(SCN)+ in the solution. Acidic solutions containing Fe(SCN)NO+ were transformed to Fe(SCN)2+ by aeration (panel D), whereas no autoxidation of acidic solutions of iron(II) and iron(II) plus FeSCN+ was observed at pH 2 (data not shown). These data suggest that Fe(SCN)NO+ itself and NO dissociated from Fe(SCN)NO+ contributed to the O2-dependent formation of Fe(SCN)2+. However, direct oxidation of Fe(SCN)NO+ by O2 seemed to not be possible because both iron(II) and FeSCN+ did not autoxidize at pH 2.0. Alternatively, O2-dependent oxidation of Fe(SCN)NO+ may be due to N2O3 formed from NO. N2O3 can oxidize FeSCN+ and Fe(SCN)NO+ through the formation of NO+ and NO2. Ascorbic Acid-Induced Formation of Fe(SCN)NO+. Figure 6 (top) shows ascorbic acid-dependent suppression of absorbance increase at 450 nm due to the formation of Fe(SCN)2+ in iron(II) and iron(III)/SCN−/nitrite systems under aerobic conditions. In both systems, ascorbic aciddependent suppressions were incomplete [circles for iron(II) and squares for iron(III)], whereas ascorbic acid decreased absorbance of Fe(SCN)2+ to zero in iron(III)/SCN− systems (triangles). Absorption spectra of the component remaining after the addition of 2−3 mM ascorbic acid were the same as that of Fe(SCN)NO+ in Figure 5. Ascorbic acid also suppressed Fe(SCN)2+ formation and contributed to the production of Fe(SCN)NO+ in iron(II) and iron(III)/SCN−/nitrite systems under anaerobic conditions (data not shown). The color of Fe(SCN)NO+ was greenish yellow when produced in the acidic mixture containing 4 mM nitrite, 4 mM iron(II), 4 mM ascorbic acid, and 4 mM SCN−. The pH of the acidic solution of Fe(SCN)NO+ was increased to 7.4 by adding Na2HPO4 (Figure 6, bottom). A new spectrum appeared, which had a peak and a shoulder at 672 and 488 nm, respectively (spectrum 2). An absorption spectrum like spectrum 2 was also obtained by adding 0.5 mM NOC 7 to the anaerobic mixture of 0.5 mM FeSO4 and 1 mM NaSCN prepared using 50 mM sodium phosphate (pH 7.4) (data not shown). The spectra suggest that the solution contained not only Fe(SCN)NO+ but also the other components. As the other components, products formed from iron(II), SCN−, and OH− are possible. This was deduced from the result that absorbance of the mixtures of iron(II) and SCN− increased

Figure 6. Inhibition of Fe(SCN)2+ formation by ascorbic acid. (Top) Effects of ascorbic acid concentration. NaNO2 (1 mM) was added to the reaction mixture (1 mL) that contained 0.5 mM FeSO4 or 0.5 mM FeCl3, 1 mM NaSCN, and various concentrations of ascorbic acid in 50 mM KCl−HCl (pH 2.0). After completion of the reaction under each condition, absorbance at 450 nm (vertical axis) was determined. (Circles) 0.5 mM FeSO4; (squares) 0.5 mM FeCl3; (triangles) 0.5 mM FeCl3 but without nitrite. (Bottom) Effects of pH. Spectrum 1, absorption spectrum obtained in the mixture of 0.5 mM FeSO4, 1 mM NaSCN, and 1 mM NaNO2 in 50 mM KCl-KCl (pH 2.0); spectrum 2, addition of 0.1 mL of 0.5 M Na2HPO4 to increase the pH from 2 to 7.4. Downward arrow: absorption decrease after aeration.

exponentially with the decrease in the wavelength from 500 to 360 nm at pH 7.4. Fe(SCN)NO+ autoxidized at pH 7.4 too, decreasing absorption bands around 490 and 670 nm (Figure 6, bottom; downward arrow). The absorption bands were decreased too by bubbling argon gas for about 8 min (data not shown), indicating that Fe(SCN)NO+ was in equilibrium with NO and FeSCN+ at pH 7.4 as well as pH 2.0. NO Production and O2 Consumption in Acidified Saliva. NO production in acidified saliva (trace A) was increased about 1.5- and 2-fold by 0.5 mM and 1 mM iron(II), respectively (traces B and C), when measured using Fe(DTCS)3 (Figure 7). The signal intensities in the absence and

Figure 7. NO production in acidified saliva. Reaction conditions are shown in Materials and Methods. (A) acidified saliva; (B) A + 0.5 mM FeSO4; (C) A + 1 mM FeSO4. Acidified saliva (pH 2.1) contained 0.27 mM nitrite and 0.51 mM SCN−.

presence of 0.5 mM iron(II) under anaerobic conditions were about 3.0- and 3.3-fold of those under aerobic conditions, respectively (data not shown). Figure 8 shows time courses of O2 consumption and NO production in acidified saliva when measured electrochemically. 212

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Figure 9. Ascorbic acid-induced formation of Fe(SCN)NO+ in acidified saliva. (Panel A) (1) baseline; (2) (1) + 0.5 mM FeCl3; (3) (2) + 1 mM NaSCN; (4) (3) + 2 mM ascorbic acid; (5) (4) + 1 mM NaNO2. (Panel B) B−I: (1) baseline; (2) (1) + 2 mM ascorbic acid; (3) (2) + 0.5 mM FeCl3. B−II: (1) baseline; (2) (1) + 0.5 mM FeSO4; (3) (2) + 2 mM ascorbic acid. Base lines were obtained after subtracting absorbance due to acidified saliva itself. Acidified saliva contained 0.27 mM nitrite and 0.26 mM SCN−.

Figure 8. Time courses of O2 consumption and NO production in acidified saliva. (1) oxygen consumption; (2) NO production. (1−1 and 2−1) acidified saliva; (1−2, 1−3, 2−2, and 2−3) 50 mM KCl− HCl (pH 2.0). Where indicated by arrows, various reagents were added. (A) 0.5 mM FeSO4; (B) 1 mM NaNO2; (C) 0.1 mM ascorbic acid; (D) 1 mM NaSCN; (E) 0.5 mM FeCl3. Acidified saliva contained 0.17 mM nitrite and 0.38 mM SCN−.

iron(II) seemed not to bind to the glycoproteins or even if it bound to glycoproteins, the binding was not so strong. This was deduced from the result that SCN− enhanced iron(II)dependent NO production in acidified saliva as well as acidic buffer solutions (Figure 8); namely, the result indicates that iron(II) can react with ONSCN in acidified saliva as well as acidic buffer solutions. Fe(SCN)NO+ Formation in Acidified Saliva. Figure 9A (spectrum 2) shows the absorption spectrum of 0.5 mM iron(III) in acidified saliva. Addition of 1 mM SCN− increased the absorbance around 450 nm (spectrum 3). Spectrum 3 was changed by 2 mM ascorbic acid giving spectrum 4 with peaks at 448 and 588 nm, and the peak heights increased by nitrite (spectrum 5). These results indicate (i) that ascorbic acid reduced iron(III), which was bound to salivary components, to iron(II) to release from the components, (ii) that the released iron(II) reacted with SCN− to produce FeSCN+ and Fe(SCN)NO+ successively, and (iii) that the formation of Fe(SCN)NO+ is dependent on NO concentration. The formation of Fe(SCN)NO+ was also observed when ascorbic acid and iron ions were added to acidified saliva without supplementing SCN− and nitrite (Figure 9B). Ascorbic acid decreased absorbance at wavelengths shorter than 400 nm (compare spectra 1 and 2 in B−I). The absorbance decrease was due to the reduction of nitrite. This was confirmed by measuring ascorbic acid-induced changes in absorption spectra of acidified saliva at wavelengths from 300 to 500 nm. Nitrite concentration in acidified saliva used in Figure 9 was 0.27 mM. FeCl3 changed spectrum 2 to spectrum 3 with peaks at 448 and 588 nm, producing Fe(SCN)NO+. When FeSO4 was added to acidified saliva, absorbance increase in the wavelength shorter than 500 nm was observed (compare spectra 1 and 2 in B−II). Because iron(II) could be oxidized to iron(III) by nitrite in acidified saliva (Figures 7 and 8), the absorbance increase was supposed to be due to the formation of iron(III). Spectrum 2 was changed by ascorbic acid, producing Fe(SCN)NO+ (spectrum 3 in B−II). The above results indicate that the production of Fe(SCN)NO+ in the stomach can be enhanced if iron is supplemented with ascorbic acid.

Addition of iron(II) to acidified saliva (arrows A) induced slow O2 consumption (trace 1−1) and NO production (trace 2−1), and both reactions were enhanced by nitrite (arrows B in traces 1−1 and 2−1). SCN− enhanced further not only O 2 consumption but also NO production (arrows D in traces 1− 1 and 2−1). The results in Figures 7 and 8 indicate that iron(II) can reduce salivary nitrite to NO in the stomach, and that iron(II)-dependent reduction of nitrite in the stomach can be enhanced by the increases in salivary concentrations of nitrite and SCN−. O2 consumption and NO production were compared between acidified saliva and acidic buffer solutions. No significant differences in the amounts of O2 consumption were observed between acidified saliva (trace 1−1) and acidic buffer solutions (trace 1−2), whereas amounts of NO produced in acidified saliva (trace 2−1) were greater than those in acidic buffer solutions (trace 2−2), when these reactions were induced successively by the addition of 0.5 mM iron(II), 1 mM nitrite, and 1 mM SCN−. Furthermore, iron(III) slowly decreased NO concentration in acidified saliva (arrow E in trace 2−1) but decreased rapidly in acidic buffer solutions (arrow E in trace 2−2). Trace 2−3 shows that Fe(SCN)2+ but not iron(III) contributed to the decrease in NO concentration. Because Fe(SCN)2+ consumed NO in acidic buffer solutions, the greater formation of NO in acidified saliva than acidic buffer solutions was postulated to be due to the inhibition of Fe(SCN)2+ formation by acidified saliva. In fact, absorbance of Fe(SCN)2+ at 450 nm in the mixture of 0.5 mM FeCl3 and 1 mM NaSCN in an acidic buffer solution was 0.12−0.14 (Figures 4 and 5), whereas the absorbance in acidified saliva was about 0.01 (Figure 9A, spectrum 3). High molecular weight components contributed to the inhibition of Fe(SCN)2+ formation because dialyzed saliva also suppressed the formation of Fe(SCN)2+. It has been reported that almost all of iron(III) bound to lactoferrin is released at pH 2.0,45 but the binding of iron to mucin still occurs at acidic pH.46 These two glycoproteins are present in saliva.47 In contrast to iron(III), 213

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Concluding Remarks. Figure 10 summarizes major reactions concerned in this study. When iron(III) is

due to iron(II)-dependent formation of reactive oxygen and nitrogen oxide species in the mixture of saliva and gastric juice and that the increased toxic effects of iron ions by ascorbic acid are due to the enhanced production of autoxidizable FeSCN+ and Fe(SCN)NO+, which can be transported from gastric juice to the epithelial cells, in addition to the enhanced production of reactive oxygen and nitrogen oxide species. Further studies are required on the accumulation of iron(III) in epithelial cells of the stomach through the formation of FeSCN+ and Fe(SCN)NO+.

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AUTHOR INFORMATION

Corresponding Author

*Tel: +81-93-582-1131. Fax: +81-93-582-6000. E-mail: [email protected].

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ABBREVIATIONS DTCS, N-(dithiocarboxy)sarcosine; Fe(DTCS)3, complex of iron(III) with DTCS; NOC 7, 1-hydroxy-2-oxo-3-(N-methyl-3aminopropyl)-3-methyl-1-triazene

Figure 10. Possible reactions in the mixture of saliva and gastric juice after iron supplementation.

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supplemented, salivary glycoproteins may bind to iron(III) preventing the formation of organic solvent soluble Fe(SCN)2+.33,48 The prevention contributes to the decrease in transport of Fe(SCN)2+ from gastric juice to epithelial cells of the stomach across the plasma membrane. If excess Fe(SCN)2+ is transported to the epithelial cells, oxidative stresses can be induced because the compound dissociates into iron(III) and SCN− at pH around 7. When iron(II) is supplemented, reduction of salivary nitrite to NO is enhanced in the stomach. The NO produced autoxidizes generating N2O3, which can cause nitrative and nitrosative stresses as described in Introduction. However, enhanced NO production results in the formation of ONOOH because iron supplementation enhances HO2 production in the stomach.2,3 ONOOH may also contribute to stresses to the stomach such as the OH radical formed by Fenton chemistry.1−3 SCN− may scavenge the OH radical and ONOOH to generate (SCN)2 and SO42− in the presence of H2O2.43,49,50 If the formation of reactive nitrogen oxide species is enhanced by iron(II), then nitration of a salivary component 4-hydroxyphenylacetic acid can also be enhanced in the stomach.51,52 If iron(II) or iron(III) is supplemented with ascorbic acid, the amount of NO produced increases consuming almost all of the O2 in the stomach. The decrease in O2 concentration makes it possible to produce Fe(SCN)NO+ in the stomach as implied in Figure 9, which is in equilibrium with FeSCN+ and NO. It is known that FeSCN+ is extractable from aqueous solution with organic extractants and dissolved in alcohol and ether.33,48 Therefore, the transport of FeSCN+ from gastric juice to gastric epithelial cells is possible. Transported FeSCN+ will autoxidize producing iron(III) at pH around 7. The autoxidation of FeSCN+ under neutral conditions was supported by the results that 105 ± 8 μM O2 (n = 4) was consumed independent of the presence and absence of SCN− when the pH of acidic solutions of 0.5 mM FeSO4 (pH 2.0) was increased to 7.4 by adding Na2HPO4. If Fe(SCN)NO+ is transported to the epithelial cells, the component can also be autoxidized in the cells to produce iron(III) and reactive nitrogen oxide species from NO. Taking the above discussion into consideration, we propose that the greater toxic effects of iron(II) than iron(III)53−55 is

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