Oxalate ions reactions

Oxidation states can be used to establish the stoichiometry for an equation. Consider the reaction between the manganate(VII) (permanganate) and ethanedioate (oxalate) ions in acidic solution. Under these conditions the MnO faq) ion acts as an oxidising agent and it is reduced to Mn (aq), i.e. 

In the laboratory you have observed the reaction of ferrous ion, Fe+i(aq), with permanganate ion, MnOiYaqJ, and also the reaction of oxalate ion, CiOi2(aq), with permanganate ion, MnO (aq). These studies show that the rate of a reaction depends upon the nature of the reacting substances. In Experiment 14, the reaction between IO and HSO3" shows that the rate of a reaction depends upon concentrations of reactants and on the temperature. Let us examine these factors one at a time. 

Both ferrous ion, Fe+2(aq) and oxalate ion, Cf)t2(aq), have the capability of decolorizing a solution containing permanganate ion at room temperature. Yet, there is a great contrast in the time required for the decoloration. The difference lies in specific characteristics of Ft+2(aq) and Cf 2(aq). On the other hand, Fe+2(aq) is also changed to Fe+ (aq) by reacting either with MnO fagJ or with ceric ion, Ce+i(aq). One of these reactions is simple and the other involves... 

An example of catalytic action is provided by the titration of oxalates with potassium permanganate solution referred to above. It is found that even though the oxalate solution is heated, the first few drops of permanganate solution are only slowly decolorised, but as more permanganate solution is added the decoloration becomes instantaneous. This is because the reaction between oxalate ions and permanganate ions is catalysed by the Mn2+ ions formed by the reduction of permanganate ions ... 

The activation energy found for the decomposition of an individual oxalate ion in a KBr matrix (270 15 kJ mole-1) [292,294] is regarded as the energy requirement for C—C bond rupture. The generally lower values of E observed for many oxalates ( 165—175 kJ mole-1) are attributed to the facilitation of reaction at the reactant—product interface. 

Desorption of the reduced metal ion is the rate determining step and is assisted by protons and oxalate ions. The reoxidized surface complex also desorbs owing to its altered molecular structure and is thus available for further reaction. The reductive dissolution step is faster than the initial complexation process. Photochemical dissolution of hematite in acidic oxalate solution is faster when air is excluded from the system (by purging with N2) than when air is present (Taxiarchou et al. 1997). 

In contrast with the sensors described elsewhere in this Chapter, the device proposed by the authors group uses no reagent, but photons, to induce a photochemical reaction, and involves electrochemical detection of the photochemical product, which allows one to continuously monitor the formation of the electroactive product. Kinetic monitoring increases the selectivity of determinations by eliminating matrix effects and the contribution of side reactions, whether slower or faster than the main reaction. The electrochemical system chosen for implementation of this special sensor was the Fe(II)/C204 couple, which was used for the kinetic determination of oxalate ion based on the following reaction ... 

Consider the parallel reactions between zinc ion and acetate ion and between zinc ion and oxalate ion. 

Murmann (18) has investigated the reaction of some amineoximatonickel (II) complexes with EDTA as well as isotopic ligand exchange with the amineoximes. These systems are rather complicated, but the replacement rate of the chelate ligand does show the typical two term rate law (22). It was also observed that the reaction rate is catalyzed by the addition of other substances such as ammonia, ethylenediamine and oxalate ion. 

In this reaction, oxalate ion may be oxidized intramolecularly by cobalt(III) ion, but it is interesting to compare the three different systems in w hich there are three, two, or one oxalate ions with the cobalt(III) cation. The last one can be boiled in l.OM add for an hour and nothing happens. In the first one, decomposition will occur very readily in aqueous solution at 50°C., so that oxalate exchange can t be measured, for instance. The middle one has not been studied in any detail yet, as far as I know, but there is oxidation-reduction in this too, though much slower than in the first. I wonder if this inhibiting effect of the nonreacting ligand, the diamine, on the oxidation has any simple explanation. 

Example of Application Large-Scale Actinometry. Neural network modelling was applied to large-scale actinometry in a continuous elliptical photochemical reactor with a concentric annular reaction chamber [2, 3,108, 148], Uranyl oxalate was used as an actinometer, which is based on the photosensitized decomposition of oxalate ions (Eq. 89) [2, 3] the experimental data were taken from the literature [108],... 

Identify the Lewis acid and the Lewis base in the reaction of oxalate ions (C2042-) with Fe3+ to give [Fe(C204)3]3-. 

The first reaction is much faster than the second. As Mn04 ion is common in both the reactions, the difference clearly lies in the nature of ferrous and oxalate ions. Fe2+ ion is a simple ion, whereas C2042- ion is a polyatomic ion and contains a number of covalent bonds which have to be broken in the oxidation reaction. 

A similar but, perhaps, more interesting case is that of the ion pairs between Co(sep)3+ and oxalate ions. Excitation of the ion pair in the IPCT band leads to the formation of the Co(II) complex and of an oxalate radical which undergoes a fast decomposition reaction. Thus, the back electron transfer reaction is prevented and Co(sep)2+, which is a good reductant, can accumulate in the system. When colloidal... 

Oxalate Destruction in Waste Stream. Although the reactions are more rapid at 8M HNO3, manganous ion catalyzes the oxidation of oxalate at lower acidities. Approximately 1M HNO3 seems to be required. As the major solids producing reagent in the waste stream is oxalic acid, it is much cheaper to oxidize the oxalic acid to CO2 gas than to store it as radioactive waste. Waste streams are therefore acidified with nitric acid, Mn(N03>2 added and the solutions evaporated. During the evaporation, the oxalate ion is oxidized to CO2 gas. 

Fig. 3.1. Schematic diagnam to illustrate the difference in the way potential electrolytes and true electrolytes dissolve to give ionb solutions (a) Oxalic acid (a potential electrolyte) undergoes a proton-transfer chemical reaction with waterto give rise to hydrogen ions and oxalate ions, (b) Sodium chloride (a true electrolyte) dissolves by the solvation of the Na and Cl bns in the crystal.

Reaction of (I) with tctra-n-butylammonium oxalate (disalt) gave the unsaturated lactone (3) as the sole product, isolated in 82 % yield. Oxalate ion thus appears to be an excellent nucleophile for effecting elimination under very mild conditions. 

In the classical case, R is sulphite and Ox sulphate. Three classes of related reactions have been recognised. To the first belong the sulphite, thiosulphate and stannous ion reactions, and with these (4) is always faster than (3) so that the starch-iodine colour emerges very suddenly when all the reductant is exhausted (by excess iodate). The second type can attain equal rates of iodine production, through (2) and (3), and decomposition (4). Starch-iodine colour is seen at about that point, with partial removal of the reductant e.g. arsenite, ferrocyanide, Fe(II) complexed with oxalate or EDTA). In the third type, reaction (3) is so much faster than (4) that the necessary iodide concentration to give starch-iodine colour is only attained late in reaction. Iodine is then present early but the blue colouration only develops later. A number of organic reductants fall into this class. The rates of colour development in the normal reaction system have been treated in semiquantitative fashion . ... 

One can easily summarize the requirements that should be fulfilled before the titanium-based ceramic materials can be obtained through an aqueous solution synthesis method. First, the precursor compound should possess good solubility in water, and preferably it should be stable over a wide pH range. In ideal case such compound should be a weighing form for titanium however, from the practical considerations it is sufficient to have a stock solution stable for a reasonably long period of time. Second, the reagent should be non-toxic, relatively cheap and its impact on the environment should be small. Its composition and chemistry should be simple and the reactions with other cations that will be introduced to the system must be well-predictable. The tendency to form precipitates with many cations, like in the case of oxalate ions, must be avoided. Finally, from an industrial point of view, the overall process should be cost effective and environmentally benign. 

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