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. 12k DownloadsAbstractMarcel Pourbaix has developed a unique and concise method of summarizing the corrosion thermodynamic information for a given metal in a useful potential–pH diagram. These diagrams indicate certain regions of potential and pH where the metal undergoes corrosion and other regions of potential and pH where the metal is protected from corrosion. Such diagrams are usually called “Pourbaix diagrams” but are sometimes called “equilibrium diagrams” because these diagrams apply to conditions where the metal is in equilibrium with its environment. Pourbaix diagrams are available for over 70 different metals 1.
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Learning ObjectivesMake sure you thoroughly understand the following essential ideas. It is especially important that you know the precise meanings of all the highlighted terms in the context of this topic. The Nernst equation relates the effective concentrations ( activities) of the components of a cell reaction to the standard cell potential.
Of a Pourbaix diagram, which plots the ∆φ. For redox reactions vs pH of the chemical environment. Taking the first reaction (1.12) as an example, at voltages below -0.356V, any lead species will be reduced to lead metal. At voltages higher, some other species will be stable; this depends on whether.
For a simple reduction of the form M n+ + ne – → M, it tells us that a half-cell potential will change by 59/ n mV per 10-fold change in the activity of the ion. Ionic concentrations can usually be used in place of activities when the total concentration of ions in the solution does not exceed about about 0.001 M.
In those reactions in which H + or OH – ions take part, the cell potential will also depend on the pH. Plots of E vs. PH showing the stability regions of related species are very useful means of summarizing the redox chemistry of an element.The standard cell potentials we discussed in a refer to cells in which all dissolved substances are at unit activity, which essentially means an 'effective concentration' of 1 M.
Similarly, any gases that take part in an electrode reaction are at an effective pressure (known as the ) of 1 atm. If these concentrations or pressures have other values, the cell potential will change in a manner that can be predicted from the principles you already know. Electrodes with poiseThe equation just above for the Cu/Cu 2 + half-cell raises an interesting question: suppose you immerse a piece of copper in a solution of pure water. With Q = Cu 2 + = 0, the potential difference between the electrode and the solution should be infinite!
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Are you in danger of being electrocuted? You need not worry; without any electron transfer, there is no charge to zap you with. Of course it won't be very long before some Cu 2+ ions appear in the solution, and if there are only a few such ions per liter, the potential reduces to only about 20 volts. More to the point, however, the system is so far from equilibrium (for example, there are not enough ions to populate the electric double layer) that the Nernst equation doesn't really give meaningful results. Such an electrode is said to be un poised. What ionic concentration is needed to poise an electrode? I don't really know, but I would be suspicious of anything much below 10 –6 M.
The Nernst Equation works only in Dilute Ionic SolutionsIons of opposite charge tend to associate into loosely-bound ion pairs in more concentrated solutions, thus reducing the number of ions that are free to donate or accept electrons at an electrode. For this reason, the Nernst equation cannot accurately predict half-cell potentials for solutions in which the total ionic concentration exceeds about 10 –3 M.How the cell potential really depends on concentration! The Nernst equation accurately predicts cell potentials only when the equilibrium quotient term Q is expressed in. Ionic activities depart increasingly from concentrations when the latter exceed 10 –4 to 10 –5 M, depending on the sizes and charges of the ions.If the Nernst equation is applied to more concentrated solutions, the terms in the reaction quotient Q must be expressed in 'effective concentrations' or of the electroactive ionic species. The activity coefficient (gamma)) relates the concentration of an ion to its activity a in a given solution through the relationa = gamma cSince electrode potentials measure activities directly, activity coefficients can be determined by carrying out appropriate EMF measurements on cells in which the concentration of the ion of interest is known. The resulting Es can then be used to convert concentrations into activities for use in other calculations involving equilibrium constants. Chlorine in waterBecause chlorine is widely used as a disinfectant for drinking water, swimming pools, and sewage treatment, it is worth looking at its stability diagram.
Note that the effective bactericidal agent is not Cl 2 itself, but its oxidation product hypochlorous acid HOCl which predominates at pH values below its pK a of 7.3. Note also that.
Cl 2 is unstable in water except at very low pH; it decomposes into HOCl and Cl –. Hypochlorous acid and its anion are stronger oxidants than O 2 and thus subject to decomposition in water. The only stable chlorine species in water is Cl –. Decomposition of HOCl occurs very slowly in the dark, but is catalyzed by sunlight. For this reason the chlorine in outside swimming pools must be frequently renewed. Decomposition of Cl 2 and HOCl by reaction with organic material in municipal water supply systems sometimes makes it necessary to inject additional chlorine at outlying locations.Each solid line represents a combination of E and pH at which the two species on either side of it can coexist; at all other points, only a single species is stable.
Note that equilibria between species separated by diagonal lines are dependent on both E and pH, while those separated by horizontal or vertical lines are affected by pH only or E only, respectively. IronStability diagrams are able to condense a great amount of information into a compact representation, and are widely employed in geochemistry and corrosion engineering. The Pourbaix diagram for iron is one of the more commonly seen examples.Pourbaix diagram for iron.
Three oxidation states of iron (0, +2 and +3) are represented on this diagram. The stability regions for the oxidized iron states are shown only within the stability region of H 2O.
Equilibria between species separated by vertical lines are dependent on pH only.The +3 oxidation state is the only stable one in environments in which the oxidation level is controlled by atmospheric O 2. This is the reason the Earth’s crust contains iron oxides, which developed only after the appearance of green plants which are the source of O 2. Iron is attacked by H + to form H 2 and Fe(II); the latter then reacts with O 2 to form the various colored Fe(III) oxides that constitute “rust”. Numerous other species such as oxides and hydrous oxides are not shown. A really “complete” diagram for iron would need to have at least two additional dimensions showing the partial pressures of O 2 and CO 2. The LibreTexts libraries are and are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739.
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