Relevant bibliographies by topics / Diagramme potentiel ph

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Academic literature on the topic 'Diagramme potentiel ph'

Author: Grafiati
Published: 4 June 2021
Last updated: 8 February 2022

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Journal articles on the topic "Diagramme potentiel ph"

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Galicia, Laura, Yunny Meas, and Ignacio Gonzalez. "Diagramme potentiel—pH pour le systeme Fe(III)Fe(II)/H2O en presence de 1,10 phenanthroline." Electrochimica Acta 31, no. 10 (October 1986): 1333–34. http://dx.doi.org/10.1016/0013-4686(86)80156-6.

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Inoue, Hiroyuki. "The use of Potential-pH Equilibrium Diagram." Zairyo-to-Kankyo 45, no. 12 (1996): 746–48. http://dx.doi.org/10.3323/jcorr1991.45.746.

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Lee, Kyu Hwan. "Formation of Metallic Nanoparticles Using Potential-pH Diagram." Journal of the Korean institute of surface engineering 50, no. 2 (April 30, 2017): 131–39. http://dx.doi.org/10.5695/jkise.2017.50.2.131.

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Angus, John C., Bei Lu, and Michael J. Zappia. "Potential-pH diagrams for complex systems." Journal of Applied Electrochemistry 17, no. 1 (January 1987): 1–21. http://dx.doi.org/10.1007/bf01009127.

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Wood, P. M. "The potential diagram for oxygen at pH 7." Biochemical Journal 253, no. 1 (July 1, 1988): 287–89. http://dx.doi.org/10.1042/bj2530287.

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Successive one-electron reductions of molecular oxygen yield the superoxide radical (O2-) H2O2, the hydroxyl radical (OH) and water. Redox potentials at pH 7 for one-, two- and four-electron couples involving these states are presented as a potential diagram. The significance of each of these potentials is explained. The complete potential diagram enables complex systems to be rationalized, such as production of OH by H2O2 plus Fe3+.
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Reymond, Frédéric, Guillaume Steyaert, Pierre-Alain Carrupt, Bernard Testa, and Hubert Girault. "Ionic Partition Diagrams: A Potential−pH Representation." Journal of the American Chemical Society 118, no. 47 (January 1996): 11951–57. http://dx.doi.org/10.1021/ja962187t.

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Kriksunov, L. B., and D. D. Macdonald. "Potential-pH Diagrams for Iron in Supercritical Water." CORROSION 53, no. 8 (August 1997): 605–11. http://dx.doi.org/10.5006/1.3290292.

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Muñoz-Portero, M. J., T. Nachiondo, E. Blasco-Tamarit, A. Vicent-Blesa, and J. García-Antón. "Potential-pH Diagrams of Iron in Concentrated Aqueous LiBr Solutions at 25°C." Corrosion 74, no. 10 (June 29, 2018): 1102–16. http://dx.doi.org/10.5006/2865.

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Potential-pH diagrams of iron are developed in aqueous LiBr solutions with concentrations of 400 g/L, 700 g/L, 850 g/L, and 992 g/L LiBr at 25°C, which are common concentrations in different parts of absorption machines. Comparison of the potential-pH diagrams of iron in the absence and the presence of concentrated aqueous LiBr solutions shows that the corrosion area at acid, neutral, and weak alkaline pH extends to lower potentials and higher pH values with the increase of LiBr concentration, as a result of formation of the aqueous species FeBr2(aq) and FeBr3(aq) and destabilization of the solid species Fe, Fe(OH)2(s), Fe3O4, and Fe2O3.
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Minguzzi, Alessandro, Fu-Ren F. Fan, Alberto Vertova, Sandra Rondinini, and Allen J. Bard. "Dynamic potential–pH diagrams application to electrocatalysts for wateroxidation." Chem. Sci. 3, no. 1 (2012): 217–29. http://dx.doi.org/10.1039/c1sc00516b.

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Nikolaychuk, Pavel Anatolyevich. "The revised potential – pH diagram for Pb – H2O system." Ovidius University Annals of Chemistry 29, no. 2 (August 22, 2018): 55–67. http://dx.doi.org/10.2478/auoc-2018-0008.

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Abstract Thermodynamic properties of lead species in aqueous solution are collected. The chemical equilibria between various forms of Pb(II) are considered. The speciation diagrams for the equilibria 4[PbOH]+(aq) ⇄ [Pb4(OH)4]4+(aq) and 2[Pb3(OH)4]2+(aq) ⇄ [Pb6(OH)8]4+(aq), and the thermodynamic activity - pH diagram of Pb(II) species are plotted. Basic chemical and electrochemical equilibria for lead are calculated. The potential - pH diagram for Pb - H2O system is revised.
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Dissertations / Theses on the topic "Diagramme potentiel ph"

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Benayada, Abbès. "Prévision des réactions d'extraction liquide-liquide dans les solutions concentrées en acide phosphorique à partir des coefficients de solvatation établissement d'un diagramme potentiel-pH généralisé en milieu biphasique acide phosphorique-toluène." Grenoble 2 : ANRT, 1986. http://catalogue.bnf.fr/ark:/12148/cb37595878d.

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Drissi, Sidi El Hassane. "Oxydation des espèces du fer en milieu aqueux carbonate : préparation et propriétés thermodynamiques de l'hydroxyde carbonate Fe2+-Fe3+ en milieu aqueux (rouille verte 1 carbonatée) et sa formation directe à partir du fer métallique." Nancy 1, 1995. http://www.theses.fr/1995NAN10349.

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Les ions carbonate constituent l'espèce la plus abondante en solution dans les eaux naturelles continentales. Ce travail consiste à déterminer les conditions de préparation de l'hydroxyde Fe2+-Fe3+ carbonate de structure rouille verte I à partir de l'oxydation d'hydroxyde ferreux précipite en mélangeant du sulfate ferreux (mélantérite) et de la soude et en ajoutant immédiatement du carbonate de sodium Na2Co3. Les deux paliers d'oxydation observés correspondent, pour le premier à la formation de la rouille verte I, et pour le second à celle du produit final qui peut être un oxyhydroxyde ferrique ou de la magnétite. Le rapport des concentrations initiales R=init/init détermine le régime de précipitation et d'oxydation ultérieure. La caractérisation et l'étude structurale des produits obtenus se font par spectrométrie Mössbauer et diffraction des rayons X. La formule chimique de la rouille verte I carbonatée Fe2+4Fe3+2(oh)12Co3,2H2O, son potentiel chimique, le diagramme potentiel-pH sont déterminés. Une conséquence directe de ce travail concerne la formation de rouille verte I carbonatée à partir de fer métallique à la suite de l'immersion d'un échantillon d'acier. Grâce à la substitution de Fe2+ par Ni2+ et Mg2+ dans la rouille verte I carbonatée, on peut mettre en exergue l'isomorphisme de cette phase cristalline avec les minéraux naturels (pyroaurite et reevisie) et arrêter le processus d'oxydation à la première étape. Enfin, la rouille verte I carbonatée se forme de façon tout à fait similaire si les ions carbonate sont introduits en solution à partir de Co2 gazeux
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Lefebvre, Hervé. "Etudes thermodynamique et cinetique de la corrosion du fer par les melanges de nitrate et de nitrite de sodium fondus." Paris 6, 1988. http://www.theses.fr/1988PA066353.

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Caracterisation des oxydes formes lors de la corrosion du fer dans les nitrates alcalins et leurs melanges avec les nitrites et etablissement du diagramme potentiel acidite du fer dans ces milieux en fonction de la temperature (420-520**(o)c). Etude par spectroscopie d'impedance du pouvoir passivant du ferrate nafeo::(2). Modelisation du processus de corrosion
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Azoulay, Ilanith. "Corrosion des aciers à long terme : propriétés physico-chimiques des hydroxysels ferreux." Thesis, La Rochelle, 2013. http://www.theses.fr/2013LAROS410/document.

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Cette étude a porté sur différents hydroxysels ferreux, formés en milieu carbonaté ou sulfaté. Ces composés sont connus pour le rôle important qu’ils jouent lors des processus de corrosion à long terme des aciers en milieux naturels (sols, milieu marin). La chukanovite, hydroxycarbonate ferreux de composition Fe2(OH)2CO3, a été particulièrement étudiée. Son enthalpie libre standard de formation a ainsi pu être déterminée à partir d’une étude détaillée des conditions d’équilibre avec l’hydroxyde de Fe(II), réalisée sur des suspensions aqueuses vieillies jusqu’à 6 mois. Cette détermination a permis de tracer des diagrammes d’équilibre potentiel-pH du fer incluant la chukanovite et mettant en évidence son domaine de stabilité thermodynamique. Par comparaison avec des diagrammes similaires incluant la sidérite FeCO3, il a été possible de montrer que la chukanovite était métastable (à 25°C) par rapport à la sidérite. Nous avons également étudié les mécanismes de transformation de la chukanovite pour différentes conditions d’oxydation, en utilisant notamment le peroxyde d’hydrogène pour accélérer la cinétique de la réaction. Les résultats obtenus montrent, qu’à 25°C, la chukanovite se transforme en lépidocrocite et/ou goethite sans passer par un composé intermédiaire Fe(II,III) de type rouille verte. La goethite est favorisée par une augmentation du pH (excès de carbonate par exemple). Une oxydation violente par le peroxyde d’hydrogène conduit à la formation d’un oxycarbonate de Fe(III), structurellement très proche de la chukanovite. Enfin, deux hydroxysulfates ferreux ont été mis en évidence et caractérisés par diffraction des rayons X et spectroscopie Infrarouge. Ces composés n’ont cependant pas pu être obtenus seuls, mais toujours ensembles, et/ou avec Fe(OH)2, voire un 3ème hydroxysulfate ferreux. Une étude des processus d’oxydation a permis de révéler que tous ces composés se transformaient dans un premier temps en rouille verte sulfatée
This study deals with various ferrous hydroxysalts formed in carbonated or sulphated environments. These compounds are known to play an important role during the long term corrosion processes of carbon steel in natural media (soils, seawater). Chukanovite, the Fe(II) hydroxycarbonate with composition Fe2(OH)2CO3, was studied more particularly. Its standard Gibbs free energy of formation could be determined via the detailed study of the equilibrium conditions with Fe(II) hydroxide, performed with aqueous suspensions aged up to 6 months. Potential-pH equilibrium diagrams of iron could then be drawn including chukanovite and highlighting its domain of stability. A comparison with diagrams drawn with siderite FeCO3 revealed that chukanovite was metastable (at 25°C) with respect to siderite. The mechanisms of transformation of chukanovite were also studied for various conditions of oxidation. Hydrogen peroxide was for instance used to accelerate the reaction. The obtained results show that the oxidation of chukanovite leads, at 25°C, to lepidocrocite and/or goethite without the formation of an intermediate Fe(II,III) green rust-like compound. Goethite is favored by an increase of pH (i.e. excess of carbonate). The violent oxidation by hydrogen peroxide leads to a Fe(III) oxycarbonate structurally similar to chukanovite. Finally, two Fe(II) hydroxysulphates could be identified and characterised by X-ray diffraction and infrared spectroscopy. These compounds could not however be obtained alone, but always together, and/or with Fe(OH)2 or maybe a third Fe(II) hydroxysulphate. The study of their oxidation process revealed that all these compounds were first transformed to sulfated green rust
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Book chapters on the topic "Diagramme potentiel ph"

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Wranglén, Gösta. "Potential-pH-Diagramme." In WFT Werkstoff-Forschung und -Technik, 66–70. Berlin, Heidelberg: Springer Berlin Heidelberg, 1985. http://dx.doi.org/10.1007/978-3-642-82360-2_5.

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Wranglén, Gösta. "Potential-pH diagrams." In An Introduction to Corrosion and Protection of Metals, 52–56. Dordrecht: Springer Netherlands, 1985. http://dx.doi.org/10.1007/978-94-009-4850-1_4.

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Wranglén, Gösta. "Potential—pH diagrams for some technically important metals." In An Introduction to Corrosion and Protection of Metals, 251–75. Dordrecht: Springer Netherlands, 1985. http://dx.doi.org/10.1007/978-94-009-4850-1_17.

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"Potential versus pH (Pourbaix) Diagrams." In Corrosion: Fundamentals, Testing, and Protection, 17–30. ASM International, 2003. http://dx.doi.org/10.31399/asm.hb.v13a.a0003580.

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Brown, Susan A., and Paul L. Brown. "The pH-potential diagram for polonium." In The Aqueous Chemistry of Polonium and the Practical Application of its Thermochemistry, 121–26. Elsevier, 2020. http://dx.doi.org/10.1016/b978-0-12-819308-2.00004-8.

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Brown, Susan A., and Paul L. Brown. "The use of pH–potential diagrams in practical applications." In The Aqueous Chemistry of Polonium and the Practical Application of its Thermochemistry, 127–78. Elsevier, 2020. http://dx.doi.org/10.1016/b978-0-12-819308-2.00005-x.

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Yagi, Shunsuke. "Potential-pH Diagrams for Oxidation-State Control of Nanoparticles Synthesized via Chemical Reduction." In Thermodynamics - Physical Chemistry of Aqueous Systems. InTech, 2011. http://dx.doi.org/10.5772/21548.

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Bunker, Bruce C., and William H. Casey. "Nucleation and Growth of Solid Oxide and Hydroxide Phases." In The Aqueous Chemistry of Oxides. Oxford University Press, 2016. http://dx.doi.org/10.1093/oso/9780199384259.003.0013.

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In this chapter, we consider what happens when solids begin to form from solution. To grow solids from solution, solution conditions are changed from a condition in which all species are completely soluble to a condition in which they are insoluble. In the context of hydrolysis diagrams, the solution composition moves in pH and total dissolved metal concentration from a regime below a solubility or saturation limit (given by the bold solid line in Figs. 5.2 and 5.3) to a regime above this limit where the solution is supersaturated. Supersaturated solutions are inherently unstable and have the potential to generate hydroxide or oxide solids. Sometimes these solutions can be maintained in a metastable state in which precipitation does not occur immediately. However, Mother Nature eventually reduces the energy of the solution by forming a stable mixture of solids plus solution species. As solids form, soluble complexes are removed from solution until concentrations drop back to the solubility limit. The precipitation of a solid from an aqueous solution is a surprisingly complex process, involving nucleation and growth phenomena that occur at nanometer-length scales. Nucleation involves reactions between oligomers to form new clusters or particles that are sufficiently large that they do not redissolve spontaneously via the reversible reactions denoted in hydrolysis diagrams. Homogeneous and heterogeneous nucleation processes represent events that occur within the bulk solution or at the interface of another phase, respectively. Growth involves the addition of monomers to clusters in solution or oligomers to existing particles or surfaces. The combination of nucleation and growth phenomena can lead to oxides exhibiting a bewildering range of sizes, shapes, and crystal structures. How do metal complexes decide whether to form a new particle or add to an existing particle? What determines the size, shape, and crystal structure of evolving particles? Do the particles aggregate with one another in an organized fashion? Because nucleation typically involves extremely rapid (<1 millisecond) events involving objects that are extremely small (on the order of a nanometer), it is difficult to probe such phenomena at a molecular level.
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Conference papers on the topic "Diagramme potentiel ph"

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Saji, Genn. "A Root Cause Study on AOA-PWR and CDA-VVER: A Point of View of “Long-Cell Action” Corrosion Mechanism." In 14th International Conference on Nuclear Engineering. ASMEDC, 2006. http://dx.doi.org/10.1115/icone14-89658.

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The author has investigated the characteristics of boron co-deposition with crud experienced in AOA and iron ferrite deposition in CDA. Corrosion product deposits found in cores with appreciable AOA have been reported in mostly nickel-based (as NiO or elemental nickel) as opposed to nickel ferrite deposits common to non-boiling cores. Significant quantities of meta-ZrO2 and nickel iron oxyborates (bonaccordite), notably Ni2FeBO5 have also been found in deposits on cores with AOA. On the basis of this general characterization information, the author has constructed a potential-pH diagram of Ni2FeB(OH)10, which is a hydrated state of FeNi2(BO3)O2 as summarized in this paper. Although preliminary, the estimated E-pH diagram suggests some interesting observation, including: growth of bonaccordite “needles” on the clad is associated with a local anodic electrochemical reaction necessary to remove excess electrons from the system to a cathode. During the AOA cycle, the concentration of nickel and iron ions must have been unusually high as they should be for a significant amount of crud deposits. The author thinks such an acceleration of the anodic dissolution of metal cations is due to the effect of the long cell action corrosion mechanism. As early as 1949, an Italian scientist Petracchi demonstrated that electrochemical effects significantly influence the erosion rate. He constructed a flow nozzle with specimens kept under external electrical potential. Upon inducing as low as 0.1 mA/cm2 of the positive current, the erosion rates were reported drastically increased. No erosion was observed by reversing the polarity of the potential. As discussed in a companion paper also presented at this conference [1], the author discusses various mechanisms (electrochemical cell configurations) that induce potential differences, including those differences in ionic concentration, aeration, temperature, flow velocity, radiation and corrosion potentials. In this paper, the author discusses how these potential differences are related to the AOA/CDA issues in PWR/VVER plants. The author is calling for further verification experiments regarding this corrosion mechanism as a joint international project.
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Nishimura, Toshiyasu, and Junpha Dong. "Corrosion Behavior of Carbon Steel for the Overpack in the Groundwater Containing Bicarbonate Ions." In 16th International Conference on Nuclear Engineering. ASMEDC, 2008. http://dx.doi.org/10.1115/icone16-48081.

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Carbon steel is considered in Japan the candidate material for overpacks in high-level radioactive waste disposal. Effects of bicarbonate solutions on the corrosion behavior and corrosion products of carbon steel were investigated by electrochemical measurements, FT-IR and XRD analyses. The anodic polarization measurements showed that bicarbonate ions (HCO3−) accelerated the anodic dissolution and the outer layer film formation of carbon steel in the case of high concentrations, on the other hand, it inhibited these processes in the case of low concentrations. The FT-IR and XRD analyses of the anodized film showed that siderite (FeCO3) was formed in 0.5 to 1.0mol/L bicarbonate solution, and Fe2(OH)2CO3 in 0.1 to 0.2mol/L bicarbonate solution, while Fe6(OH)12CO3 was formed in 0.02 to 0.05mol/L bicarbonate solutions. The stability of these corrosion products was able to be explained by using the actual potential – pH diagrams for the Fe-H2O-CO2 system.
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Naitoh, Masanori, Shunsuke Uchida, Yasushi Uehara, Hidetoshi Okada, and Seiichi Koshizuka. "Evaluation of Wall Thinning Rate Due to Flow Accelerated Corrosion With the Coupled Models of Electrochemical Analysis and Double Oxide Layer Analysis." In ASME 2009 Pressure Vessels and Piping Conference. ASMEDC, 2009. http://dx.doi.org/10.1115/pvp2009-77583.

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Systematic approaches for evaluating flow accelerated corrosion (FAC) are desired before discussing application of countermeasures for FAC. Future FAC occurrence should be evaluated to identify locations where a higher possibility of FAC occurrence exists, and then, wall thinning rate at the identified FAC occurrence zone should be evaluated to obtain the preparation time for applying countermeasures. Wall thinning rates were calculated with the coupled models of static electrochemical analysis and dynamic double oxide layer analysis. Anodic current density and electrochemical corrosion potential (ECP) were calculated with the electrochemistry model based on an Evans diagram and ferrous ion release rate determined by the anodic current density was applied as input for the double oxide layer model. The thickness of oxide layer was calculated with the double oxide layer model. The dependences of mass transfer coefficients, oxygen concentrations ([O2]), pH and temperature on wall thinning rates were calculated with the coupled model. It was confirmed that the calculated results of the coupled models resulted good agreement with the measured ones. The effects of candidates for countermeasures, e.g., optimization of N2H4 injection point into the feed water system, on FAC mitigation was demonstrated as a result of applying the model.
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