efsa-opinion-chromium-food-drinking-water

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Chromium in food and drinking water Table 2: Some physico-chemical properties of elemental chromium Atomic number: 24 Boiling point: 2672 °C Atomic mass: 51.9961 amu Chemical family: Group 6, transition metals Electron shell configuration: [Ar], 3d 5 , 4s 1 Electronegativity (Pauling scale): 1.66 Melting point: 1857 (± 20) °C Vapor pressure: 990 Pa (1857 °C) Density: 7.19 g/cm 3 (20 °C) Solubility in water: insoluble Resistant to ordinary corrosive agents Dissolves fairly readily in non-oxidizing mineral acids (e.g. hydrochloric acid), but not in oxidizing acid media (e.g. nitric acid) due to passivation The solubility of chromium compounds depends in part on the oxidation state. What follows refers to observations at or around room temperature. The monohydrate acetate, hexahydrate chloride, hydroxide sulphate, and nitrate salts of Cr(III) are soluble in water and possibly in common polar organic solvents; however, Cr(III) chloride, (dichromium) iron tetraoxide, oxide, phosphate, sulphate, and picolinate exhibit a scant or no solubility in water (chromium picolinate is more soluble in polar organic solvents). Jelly-like Cr(OH) 3 (chromium(III) trihydroxide) has an amphoteric behaviour, the pH value having a strong influence on its solubility and the type of hydroxo-species that are formed following interaction with the aqueous media (Rai et al., 1987, 2004): a minimum solubility is observed between pH 7 and 10. Cr(IV) dioxide (CrO 2 ) is insoluble in water. As to Cr(VI) compounds, zinc and lead chromates are practically insoluble in water, whereas the chromates of alkaline earth metals are only slightly soluble; CrO 3 (chromium trioxide or chromic acid) and its ammonium and alkali metal salts are in general readily or quite soluble in water. Some Cr(VI) compounds also show a solubility in polar organic solvents. 1.1.4. Natural and artificial isotopes There are four naturally occurring stable chromium isotopes, with mass numbers 50 (4.3 %), 52 (83.8 %), 53 (9.5 %), and 54 (2.4 %). Several radioactive isotopes are also known, all artificial: with the exception of 51 Cr, they exhibit very short half-lives, in general much shorter than 24 hours. 51 Cr, whose decay is by electron capture with emission of 0.32-MeV gamma rays and a half-life of 27.7 days, has been used as a tracer in medical research on blood: for example, Na 2 51 CrO 4 has been employed to tag red blood cells (RBCs) and platelets in survival studies and blood volume measurements (Gray and Sterling, 1950; Najean et al., 1963; Pearson, 1963; Dever et al., 1989; Veillon et al., 1994); in addition, 51 Cr is commonly used in toxicokinetics investigations. 50 Cr is also suspected of being radioactive, but with such a long half-life (> 10 17 years) that it is regarded as a stable isotope. 1.1.5. Redox chemistry Aside from possible negative oxidation states, of no interest in this opinion, chromium can exist in oxidation states from Cr(I) (Cr 1+ ) to Cr(VI), with the trivalent and hexavalent states being largely predominant. Elemental chromium, Cr(0), seldom if ever occurs naturally. Cr(V) and Cr(IV), of which a few solid compounds are known, are observed as transient labile species in the reduction of Cr(VI) solutions; on the other hand, in solution they both can readily transform to Cr(III) and Cr(VI). As is typical of transition metals, chromium compounds are characterized by an elaborate coordination chemistry (Cotton et al., 1999), whose principal morphologic features may be summarized as follows: an octahedral geometry is associated with a coordination number of 6 and with all the oxidation states from Cr(0) to Cr(V); Cr(V) also exhibits a tetrahedral geometry with a coordination number of 4, just like Cr(VI). Clear examples of octahedral and tetrahedral geometries are exhibited in Figures 2 and 3. As discussed later in the opinion (see Section 7.1), oxidation state and molecular geometry of chromium compounds have a strong bearing on cellular uptake. Much of chromium chemistry deals with Lewis acid-base coordination complexes, in which ligands (ions or molecules) bind to the coordinating metal (atom or ion): ligands act as electron-pair donors (Lewis bases) while the metal acts as an electron-pair acceptor (Lewis acid) owing to its valence-shell orbitals that can EFSA Journal 2014;12(3):3595 14
Chromium in food and drinking water accommodate electron pairs. Therefore, ligands must have at least one pair of electrons suitable for being donated to the metal. The metal-ligand bonding can have various degrees of covalent nature even when both chromium and ligands are formally ionic species. Figure 2: Chromium hexa-carbonyl Figure 3: Chromate ion DIVALENT CHROMIUM. All Cr(II) compounds are energetic reducing agents and under environmental conditions Cr(II) (chromous state) is relatively unstable. In aqueous media, the chromous ion is readily oxidized to the stable Cr(III) species (Cr 3+ + e − → Cr 2+ ; E 0 = - 0.41 V), for instance by the dissolved molecular oxygen, O 2 (Kotaś and Stasicka, 2000). Therefore, Cr(II) solutions can only be preserved if degassed (anaerobic conditions). The only coordination number observed for Cr(II) is six, in the form of a twisted octahedral geometry. TRIVALENT CHROMIUM. Cr(III) is the most stable and important oxidation state of the element, in particular in relation to its aqueous chemistry (Kotaś and Stasicka, 2000). This state is characterized by the formation of a very large number of relatively kinetically inert complexes, in which Cr(III) is always hexacoordinate (octahedral geometry). This kinetic inertness allows many complex species to be isolated as solids and to persist for relatively long periods of time in solution, even if their thermodynamical condition is unstable. In aqueous media and in the absence of specific ligands, Cr(III) is present as Cr(H 2 O) 6 3+ (hexa-aquachromium (3+) , a moderately strong acid), Cr(OH) 3 (chromium trihydroxide), and their reaction products. Therefore, the aqueous compositions of these groups of substances are complex and depend on environmental conditions and their influence on processes such as hydrolysis, complexation, redox reactions, and adsorption. Even at naturallyoccurring concentrations and substantially neutral pHs, Cr(III) compounds in aqueous systems may be actively oxidized to Cr(VI) by strong oxidants such as chlorine or hypochlorous acid, ozone, or potassium permanganate — used, for instance, in water purification treatments (Schroeder and Lee, 1975; Lai and McNeill, 2006; Saputro et al., 2011; Lindsay et al., 2012). HEXAVALENT CHROMIUM. Cr(VI), or chromate, is the second most stable state: its compounds, whose aqueous chemistry is of particular relevance, primarily arise from anthropogenic sources (Shanker et al., 2005; Johnson et al., 2006). In addition to its occurrence in rare minerals, naturally occurring Cr(VI) has also been occasionally detected in groundwater (McNeill et al., 2012a). In its highest oxidation state, chromium forms oxy-compounds that are fairly potent oxidizing agents (Kotaś and Stasicka, 2000). In basic solutions (pH > 6.5), it exists predominantly as the yellow chromate ion (CrO 4 2− ), exhibiting a coordination number of four and a tetrahedral geometry (Figure 3). As the pH is lowered (pH < 6), the solution of chromate ions turns orange owing to the formation of dichromate ions (Cr 2 O 7 2− ). In Cr 2 O 7 2− two chromium atoms are linked by an oxygen bridge and exhibit a slightly distorted tetrahedral geometry (Figure 4). Acid solutions of dichromate are quite powerful oxidizing agents, the Cr(VI) reduction process yielding Cr(III). In basic solution, the chromate ion exhibits a much lower oxidizing power as the CrO 4 2− species undergoes a relative stabilization. EFSA Journal 2014;12(3):3595 15
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7.2.1.4. Genotoxicity Chromium in f
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7.5. Dose-response assessment Chrom
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Human observations Chromium in food
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J.1.2. mice Chromium in food and bo
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