The role of niobium carbides in the localised corrosion initiation of 20Cr-25Ni-Nb advanced gas-cooled reactor fuel cladding. Corrosion Science, [108365].

In-Situ Imaging of Corrosion Processes in Nuclear Fuel Cladding Spent nuclear fuel in the UK is stored within ponds dosed with NaOH in order to inhibit corrosion and, to ensure the efficiency of storage regimes, there is a need to define and quantify the corrosion processes involved during immersion of fuel cladding. In this project, state-of-the-art characterisation techniques were employed to image the corroding surfaces of two nuclear fuel cladding materials: stainless steel and Magnox. Advanced gas-cooled reactor (AGR) fuel cladding consists of 20Cr-25Ni-Nb stabilised stainless steel and during irradiation the microstructure of the cladding undergoes significant changes, including grain boundary element depletion and segregation. Highspeed atomic force microscopy (HS-AFM) with nanoscale resolution, enabled precipitates and pit initiation in stainless steel to be imaged. Magnox is a magnesiumaluminium alloy and during irradiation in a reactor the outer metal surface oxidises, forming an adherent passive layer which subsequently hydrates when exposed to water. Corrosion processes encompass breakdown of passivity and filiform-like corrosion, both of which were imaged in-situ using the scanning vibrating electrode technique (SVET).

General chemistry texts usually devote a portion of one chapter to the solubility product principle (1). This is commonly in the form of K sp = f (concentrations of ions raised to some powers). For example, K sp = [Ag + ][Cl ᎑ ], where K sp is the solubility product constant for the slightly soluble substance silver chloride. [Ag + ] and [Cl ᎑ ] are the concentrations in moles per liter of the ions that are in equilibrium with the solid phase of AgCl. Tables of K sp values are given, and the usual two assumptions made are: the substances are strong electrolytes which ionize 100% in solution; and activities of the ions are close enough to the molar concentrations in these dilute solutions to allow defining the K sp in terms of concentrations. Reasoning from the equations, which presume that the only forward reaction is the production of the expected ions, the student is asked to calculate solubility from K sp and vice versa, to calculate the common ion effect on the solubility of substances, and to determine if slightly soluble solids will precipitate under certain concentration suppositions.This Journal contained an excellent article in 1966, by Meites, Pode, and Thomas called "Are Solubilities and Solubility Products Related?" (2). The authors showed clearly that, due to ion pair formation, hydrolysis, complex ion formation, and activity coefficient variations, there are only a few cases in which solubility and K sp are related in a simple way (Fig. 1). Meites et. al. concluded, "It would be better to confine illustrations of the solubility-product principle to 1-1 salts, like silver bromide and thallium iodide, in which the student's calculations will yield results close enough to the truth to permit him to feel it is worth his trouble to try to master what he is being taught."Ignoring the masculine pronoun common in the sixties, this advice is excellent. In the majority of cases solubilities are not simply related to K sp . Unfortunately, the authors of current general chemistry texts have largely ignored this advice and provide students with large K sp tables that include numerous slightly soluble substances. Further, many authors imply by example and by assigned exercises that K sp is a simple function of solubility. In effect they teach (3) that for A x B y (s) → xA z+ + yB z᎑ where solubility is S o , K sp = [xS o ] x [yS o ] y which results in the familiar K sp = S o 2 K sp = 4S o 3 K sp = 27S o 4 and K sp = 108S o 5

This paper, for chemistry teachers who have beginning students, computers, and spreadsheets, presents tables of titration data simulated using the program EQUIL. Instructors are invited to give students the titration data to enter into their spreadsheets, have them plot it in several ways, and produce an assortment of graphs. In the process, they will discover many things about pH,. the taking of derivatives, buffer capacity, and the way buffers behave upon dilution. These resulting plots show why there is pseudobuffering at high and low pH values, and show the equilibrium buffering maximum at the pKa or pKb of the electrolyte species. A convenient definition of buffer capacity with respect to dilution is beta dil, where beta dil = d(-log[conc]/d(pH). This definition has the advantage of being an intensive property of the solution, and also of being large for equilibrium buffer solutions undergoing dilution.

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