Method for determining calcium by color reaction. Download the book "Analytical chemistry of calcium" (2.28Mb)

Magnesium and calcium are main or by-product components of numerous natural or artificial products. Classical methods for the analysis of these two cations are time-consuming, while complexometric titration provides the researcher with the opportunity to elegantly determine both metals, which greatly contributed to the rapid introduction of this method into analytical practice.

We think it is useful to discuss both metals at the same time, since they are almost always present together, and therefore it is important to know the behavior of a mixture of Ca and Mg, even if only one of these elements is required to be determined.

The analysis of biological fluids, due to its great practical importance, is discussed in a separate section. The given literary references represent only a part of all publications related to this topic, which seems to us completely fair, since most of the works, from the point of view of complexometric titration itself, do not contain anything new.

The cited works still provide a complete picture of the existing capabilities of the method and of the problems that have not yet been resolved.

The determination of Mg using EDTA has long been described by Schwarzenbach et al. . The indicator they used, eriochrome black T, is one of the most commonly used at present. Performing microscale titrations and even determining microgram quantities is straightforward. The accuracy of complexometric determinations and titration stoichiometry were thoroughly studied.

The stability of the EDTA and indicator complexes with Mg is quite high so that titrations can be carried out with sufficient accuracy; The color change at the equivalence point (from wine red to blue) is somewhat less distinct than in other complexometric titrations. It should be titrated until the red tint disappears completely, which, however, is not difficult to recognize. The reaction at the equivalence point proceeds somewhat slowly, so the solution should be slightly heated.

Eriochrome black T and many similar dyes are blocked by traces of heavy metals, primarily copper, which, however, are not difficult to remove using appropriate masking agents. Potassium cyanide eliminates interference from Cu, Ni, Co, Fe, etc. The same function is performed by Na2S (in this case, heavy metal impurities are precipitated in the form of sulfides) and Mn - titration of Mg in the presence of a large amount of Mn, see. Aluminum can be masked using triethanol-amine, and the titration should be carried out at 5 ° C, since otherwise the transition of Al from the complex with the masking substance to the complex with the indicator is possible.

Interference caused by the presence of trace heavy metals can often be eliminated by using the back titration method. In this case, interfering impurities are bound into a complex with EDTA and react with the indicator only slowly or not at all; thus, the back titration can be completed before the indicator is blocked. If, for example, a back titration is carried out with a Zn solution, then a Cu content of up to 20 mg per liter of solution does not have a harmful effect. The method of protective titration proposed by Hahn is based on the same principle, based on the relative absence of interference and consisting in the fact that a known amount of a titrated EDTA solution is titrated with the analyzed solution.

In addition to eriochrome black T, a large number of other indicators are used, for example aluminone, which allows for sequential titration of the Fe-Al-Ca-Mg mixture, lacquer scarlet C, acid chromium blue dyes, chromoxane green, pyrocatechol violet, arsenazo I. Deal et al., on the one hand, and a group of researchers led by Belcher, on the other, examined a large number of dyes from the point of view of their suitability as indicators. Recently, Kalmagite has enjoyed great success; in terms of the stability of complexes with metals and in terms of color changes, it is almost identical to eriochrome black T, but its solution is more stable.

Indication of the titration endpoint using instrumental methods involves predominantly photometric titration, which is performed either with self-indication in the UV region, or with eriochrome black T, or with other indicators, for example with chromasurol S or calmagite. When determining Mg and when sequentially titrating mixtures of Ni-Mg, Zn-Mg or Bi-Mg, potentiometric titration with a mercury cathode or amperometric titration is also used. Below, conductometric and thermometric definitions will also be described.

The interfering effect of Mg on the titration of other metals appears only in an alkaline medium, so its presence is hardly a problem in the determination of other metals since it became possible to carry out titrations in an acidic solution. Mg can be masked by precipitating it in the form of hydroxide in a strongly alkaline solution (caustic soda) or using fluoride ions.

The titration of Mg in the presence of phosphate ions was carried out by Collier, who advised removing large quantities of these ions by extraction. Ion exchange resins are also good for removing phosphate ions. A strong dilution of the test solution is often sufficient to slow down the formation of MgNFLjPO.), since this compound easily forms supersaturated solutions. In addition, Mg can be determined in the presence of phosphate ions by back titration. The titration of Mg in the presence of Ca will be discussed below. Here we can also note the possibility of separating Ca in the form of molybdate and titrating Mg in the filtrate if the determination of Mg- alone is required.

Magnesium can be determined complexometrically in pharmaceuticals, in aluminum alloys, in electron alloy, in cast iron and cast iron, in titanium, nickel sulfate, gunpowders, in soil and plant materials, rocks and uranium slags.

Calcium is one of the first metals for which the complexometric titration method was described. Titration can be carried out in highly dilute solutions, as well as in the presence of small quantities of Ca. The indicator murexide used in this case has been studied in detail and is often used today. In a highly alkaline environment (pH = 12), the red color of murexide changes to blue-violet, which is not as sharp as that of many other metallochromic indicators. The murexide solution is stable for only a few hours, so it is advisable to add the indicator in solid form, ground with 100 parts of NaCl. The oxidative or hydrolytic decomposition of murexide in the test solution should also be taken into account, especially during photometric titrations, where decomposition sometimes becomes noticeable due to a slow decrease in light absorption. To improve the recognition of the equivalence point, mixed indicators have been proposed, for example 0.2 g murexide with 0.5 g naphthol green B, well mixed with 100 g NaCl.

Many other substances have been proposed as Ca indicators, which, however, are not always superior to murexide. Here are some of them: tracing paper, CAL-Red, eriochrome blue-black SE (Erio SE), acid chrome blue-black and others. All these substances are o, o"-azo compounds, similar to eriochrome black T.

A systematic study of the indicator properties of such substances belongs to Diehl et al. . Numerous compounds have also been studied by Belcher et al. . Later, the following were tested as indicators for Ca: lac scarlet C, omega chromium blue-green BL, ftclein complex son, glyoxal-bis-(2-hydroxyanyl), chromazurol S, H-acid, acid alizarin black SN and pyrogallol carboxylic acid . With aluminone, sequential titration of the Fe-Al-Ca-Mg mixture is possible.

The calcichrome synthesized by West mentioned below is apparently identical to the hydron proposed by Russian authors. Methyl thymol blue and pyrocatechol violet are also suitable for determining Ca.

Calcein can be used both as a color and as a fluorescent indicator (UV rays). Fluoresce-incomplexon has residual fluorescence caused by contaminants beyond the equivalence point, which overlaps when phenolphthalein is added (0.25 g of phenolphthalein per 1 g of indicator). The situation is similar with calcein (calcein W), for which acridine was proposed to cover the residual fluorescence. Thymolphthalexone is also recommended as a fluorescent indicator for Ca. To ensure recognition of the equivalence point without interference, Toft et al. proposed a simple device that worked well in titrations with calcein and also served well in titrations with other fluorescent indicators.

Almost all Ca indicators give a sharp color transition only at a high pH value of the solution. However, there are some indicating systems that operate at pH<11, например комплекс Mg с ЭДТА (его дббавляют по меньшей мере в количестве 5% от содержания присутствующего Са) или ZnY в комбинации с эриохромом черным Т, а также комбинации ZnY с цинконом и CuY с ПАН . При этом одновременно титруется присутствующий в растворе Mg.

Indicators that operate at high pH values are usually preferred, since magnesium, which often accompanies calcium, precipitates in the form of hydroxide (see below). It should be noted that the alkali used for alkalization should not contain carbonates, and should not absorb them from air, water or other reagents, since otherwise a CaCO3 precipitate will form. The precipitate will dissolve again during the titration if done slowly.

However, it is more advantageous and time-efficient to avoid the formation of a precipitate, for which carbonate ions are removed and titrated in fairly dilute solutions in order to prevent possible precipitation of Ca(OH)g. The formation of turbidity can also be avoided by using the back titration method.

Factors that interfere with calcium titration have been studied in detail. Fe and Al, present in most natural and artificial products, can be isolated using a variety of methods. Separation by precipitation with ammonia solution is always possible, but it is often time consuming since double precipitation may be necessary. The masking of Fe, Al and Mn can be read in the sections dealing with the identification of the corresponding elements.

If only Al is present in the solution, then there is no need to do anything to determine Ca, since during normal titrations at high pH Al is present in the form of aluminate ions, which do not react with the complexone. However, you should pay attention to the choice of indicator, since some dyes are blocked by aluminum under these conditions. For very high Al contents please refer to references and for high Mn concentrations refer to references.

Titanium can be masked with hydrogen peroxide (see definition of titanium). The use of potassium cyanide and ion exchangers opens up wide possibilities for masking. The possibility of interference from anions should also be taken into account. The interference caused by hydroxyl and carbonate ions has already been discussed. Hexacyanoferrate (II) ion, which was initially present in the solution or formed during masking of Fe, can form turbidity due to the low solubility of its calcium salt; The turbidity disappears again during the titration process. The interference associated with the presence of phosphate ions has been studied in particular detail. Small amounts of the latter do not interfere with the titration of Ca. The maximum permissible ratio P: Ca = 4: 1, but it strongly depends on the dilution of the solution.

Large amounts of PO4 ions will not interfere with the determination if you use the back titration method. Zimmerman suggests a titrated solution for the determination of Ca in the presence of phosphate ions, 0.1 M for EDTA and 0.05 M for ZnY. In extreme cases, when the content of phosphate ions is exceptionally high, they are separated by ion exchange or extraction.

Since the use of titration in an acidic medium, Ca is no longer a strong interference in the determination of other metals. In some alkaline titrations (but not Mg titrations), Ca can be masked with fluoride ions.

The accuracy and precision of complexometric Ca determinations is good, as confirmed by numerous studies, for example.

There are many instrumental methods for determining Ca. Photometric titration is most often preferred, since it is difficult to recognize the color transition of murexide with the naked eye. Other indicators are also used, for example tracing paper, CuY-PAN, metalphthalein.

Photometric titration can be carried out with self-indication in the UV region (228 nm); it can be automated using various indicators. The indication of the equivalence point by the slope of the titration curve when adding Cu2+ ions is described. Amperometric indication with a mercury drop electrode makes it possible to carry out sequential titration of mixtures such as, for example, Ni-Ca or Cu-Zn-Ca, and the “complexon wave” is used for indication. In a highly ammonia solution, Ca can be determined by an indirect amperometric method: from zinc complexonate, Ca2+ ions displace Zn2+-, which are then titrated.

When performing potentiometric titration with a mercury drop as an electrode, it is advisable to use a HEDTA solution, since Mg does not interfere. Ghazlam et al. carry out potentiometric automatic titration with a silver electrode; This method can be used to titrate a Ca-Mg mixture sequentially. Radiometric and conductometric titrations are described. The thermometric indication of the equivalence point is especially interesting in relation to the analysis of the C a-Mg mixture, since the heats of formation of complexonates of both metals are not only different, but even opposite in sign.

The number of practical applications of complexometric determination of Ca is enormous. Below are just some of the possible cases. Since the determination of Ca is often related to the determination of Mg, we recommend that the reader refer to the sections concerning the determination of a mixture of Ca and Mg and water hardness. Using visual indication, analyzes of stearates, sugar juices, casein, water, rainwater, pharmaceuticals, tricalcium phosphate, technical phosphates, plant materials, photographic materials, rosin are carried out, as well as the determination of free lime in silicates and Ca in caustic soda, and in the latter case, Ca concentration is used on a chelating ion-exchange resin - Dauex A-I.

Photometric titration with murexide is used to determine the water-soluble part of gypsum and water analysis. Calcein is used as a photometric indicator in the determination of Ca in lithium salts. When analyzing forage, Ca is titrated with a mercury drop electrode with a solution of HEDTA.

Mixtures of calcium and magnesium. The separation of calcium from magnesium can be carried out in various ways. Separation is always possible, but takes a lot of time. For separation purposes, it is advisable to use ion exchange resins. Gehrke suggests separating Ca in the form of sulfite. Ca can be precipitated in the classical way in the form of oxalate and, after ashing and dissolving the precipitate, it can be titrated complexometrically.

In the case of a very low Ca content, the calcium oxalate precipitate can be dissolved in acid, EDTA can be added, and after making the solution alkaline, the excess EDTA can be titrated. However, after precipitation of calcium in the form of oxalate, the color change of eriochrome black T during titration of Mg in the filtrate is not sharp enough, so the amount of oxalate ions used is limited to a minimum.

More elegant methods are those that avoid separating both metals. The most commonly used method consists of titrating Ca in a strongly alkaline solution in the presence of a precipitate of magnesium hydroxide and determining the sum of Ca and Mg in a second aliquot of the solution (taking everything said above about titrating Mg into account), followed by calculating the Mg content from the difference. If there is a lot of Ca and little Mg in the mixture, it is unlikely that difficulties will arise during the analysis. If the situation is less favorable, attention should be paid to a number of circumstances, a discussion of which can be found in the original literature.

The presence of Mg(OH)2 can interfere, firstly, because there is the possibility of co-precipitation of Ca, and secondly, because the change in color of the indicator may not be sharp due to the adsorption of the dye by the flocculent sediment.

The addition of sugar should prevent the coprecipitation of Ca, but this is not confirmed by other authors. Co-precipitation can, according to Flaschka and Gooditz, be reduced to a minimum if an amount of EDTA not much greater than the amount equivalent to calcium is first added to the neutral or acidic solution being analyzed, and only then is it made alkaline. The alkali should always be added slowly drop by drop and the solution should be mixed well. According to Lewis et al. , in this case, a small amount of EDTA is deposited, which, when standing, due to the recrystallization of Mg(OH)2, goes back into solution.

To better recognize the color change of an indicator (for example, murexide), it is advisable, but certainly not necessary, to carry out precipitation in a volumetric flask; the volume of the solution should be brought to the mark and, after the precipitate has settled, a clear aliquot of the filtrate should be used to back-titrate a small excess of EDTA.

Bauch et al. obtained good results with very high Mg contents (determination of about 0.5% Ca contained in MgO) by slowly precipitating Mg(OH)2 with 0.5 M NaOH solution (a small amount of KCN and NH2OH HC1 was added to the alkali) with vigorous stirring and titrating Ca directly into the suspension with an EDTA solution with CaL-Red as an indicator. The importance of slow precipitation with vigorous stirring is also emphasized by Lewis and Melnik.

As shown by a study conducted by Kenya et al. , the final pH value of the solution, the indicator used and its quantity also affect the titration results. The results obtained by Belcher et al. are important in this regard. Of the numerous indicators tested, tracing paper turned out to be the most suitable. The equivalence point in the presence of precipitated Mg(OH)2 was clearer than in pure Ca solutions, and the presence of Mg did not produce low Ca results as occurs with other indicators (e.g. murexide, methylthymol blue or calcein).

Indistinct color transitions caused by the adsorption of the indicator onto the Mg(OH)2 precipitate can be improved if the indicator is added after the magnesium has precipitated, and, in addition, if the precipitate is waited until the precipitate becomes crystalline before adding the dye. As Lott and Cheng point out, adding a few drops of polyvinyl alcohol prevents the indicator's color transition from becoming less clear. A similar effect of acetylacetone was observed by Bourget et al.

Summarizing the above, it can be noted that there are various possibilities for improving the determination conditions, but it is hardly possible to propose a determination methodology that is satisfactory for all cases; For each specific case, optimal conditions should be selected in order to achieve the greatest accuracy. Therefore, it is not surprising that there are many reports of experiments carried out in order to avoid the precipitation of Mg(OH)2. For this purpose, it is proposed to add tartaric acid to the solution. According to our experiments and in accordance with the data of other authors, tartaric acid is suitable for preventing the precipitation of magnesium, but inflated results for Ca are obtained if an EDTA solution is used as a titrant. If we take HEDTA instead of EDTA, the results of determining Ca are correct, since the magnesium complex with this complexone is less stable than the calcium complex. In this regard, it is interesting to note that the equivalence point for titration with calcone is sharp only when the Mg:Ca ratio is at least 1.

If we compare this information with the data of Belcher et al. mentioned above, we must admit that it is still not clear how the precipitation of magnesium and its complexation influence the formation of the Ca-calcone complex at the equivalence point.

One of the main problems in determining Ca in the presence of Mg is the lack of a simple Ca indicator for visual determinations that operates at pH values when Mg still remains in solution. Ringbom resolved this difficulty by using an indirect indication of the equivalence point using the Zn-HEDTA-zincone system. The solution is adjusted to pH = 9.5-10 using a buffer solution containing 25 g of borax, 2.5 g of NH4C1 and 5.7 g of NaOH in 1 liter.

In pure solutions, very sharp color transitions and correct Ca content values are obtained. But for this it is necessary that, firstly, the ammonium concentration is maintained very precisely and, secondly, the Ca: Zn ratio is approximately 10; Unfortunately, meeting these optimal conditions during practical analysis is not always possible. Another way is described by Flaschka and Ganchof: they titrate with a solution of HEDTA with murexide as an indicator at a pH of about 10. With photometric indication, Ca can be determined in the presence of more than 100-fold excess of Mg. Calcium in the presence of magnesium can also be titrated potentiometrically with a HEDTA solution at pH = 10.

We should also refer to the Strafeld method, in which Mg is precipitated with phosphate ions at pH = 9 and then, in the presence of the precipitate, Ca is determined by reverse potentiometric titration of excess EDTA with a titrated solution of calcium salt with a mercury drop electrode. The amount of phosphate added must be very precise. On the one hand, this amount should be enough to reduce the solubility of MgNH4P04 so that it does not react with EDTA, on the other hand, the amount of phosphate should not be too large, since otherwise a precipitate of Ca3(P04)2 will form. There are no published data regarding the coprecipitation of Ca.

After all that has been said, we emphasize once again that it is hardly possible to give a universal method of work, however, there are satisfactory modifications of standard methods, based on which it is possible to choose a method suitable for work for all cases encountered in practice. It should not be forgotten that most studies are carried out on pure solutions, and in practical analysis the determination conditions are complicated due to the high concentration of salts, the presence of interfering elements and masking substances added to eliminate them.

The most elegant are sequential titrations, since, on the one hand, they save time, and on the other hand, they require a smaller amount of the analyzed solution, which

practical definitions are often very important. Such experiments were carried out and gave very good results, at least on artificial solutions. Karesh first titrates Ca with murexide at pH = 13, then acidifies the solution, and murexide, hydrolyzing, is destroyed, brings the pH to 10 and titrates Mg with eriochrome black T. The difficulties described above that arise when determining Ca in the presence of Mg (OH) 2, Naturally, they matter here too.

Lott and Cheng first titrate Ca with calcone at a high pH, then lower the pH of the solution by adding acid and ammonium chloride, and continue titrating with eriochrome black T to determine Mg. Schmidt and Reilly exclude the error caused by the precipitation of magnesium, for which they first titrate Ca with a HEDTA solution in a transparent solution at pH = 9.5-10 in the presence of the Ringbom indicator system, which is a mixture of Zn - HEDTA - zincone, then add KCN to mask the Zn and titrate Mg with a solution of EDTA with eriochrome black T. Flaschka and Ganchof use photometric indication of the equivalence point. First, they titrate Ca with murexide with a HEDTA solution at pH = 10, then add eriochrome black T, change the wavelength of light and determine Mg by titration with an EDTA solution. Submicrogram amounts of Ca and Mg can be determined from a single photometric titration curve; in this case, the Mg-cal-magit complex is used as a self-indicating system to establish the end point of Ca titration based on the slope of the titration curve.

The determination of Ca and Mg by the methods mentioned above is used in the analysis of a variety of materials, for example insect lymph, limestone, dolomite, magnesite, calcareous and silicate rocks, soils, glass powders, glass, ores and slags, cement, steel and similar materials; rock salt, brines, sea water and other solutions with a high alkali content, as well as welding wire containing Mn, pulp, coal mining wastewater, ordinary water and special mineral waters, milk, canned fruit juices, pharmaceuticals, plant materials after ashing , in particular tobacco ash, animal tissues and biological materials in general.

Calcium and magnesium in biological fluids. Complexometric determination of Ca and (or) Mg in blood, serum, urine and cerebrospinal fluid is currently a standard titrimetric method used in almost all laboratories. The number of publications related to this area exceeded one hundred.

Since many of the proposed methods differ only slightly in detail, only some of the published work will be reviewed here to explain the principles of the definitions.

Calcium in serum was first determined by Greenblatt and Hartmann by titrating with murexide in a strongly alkaline solution. Other authors describe the same method with only minor modifications or with photometric indication.

Other indicators are also used, for example calcein, mainly in the UV region, and a titration curve can be drawn; This method can analyze very small quantities of serum (20 µl); Photometric indication can be used. In addition, CAL-Red, tracing paper, phthalein complexone, acid alizarin black SN and fluorescent indicators are used. A thorough comparison (for example) of these methods with the classical oxalate method clearly showed the advantages of the complexometric method.

Calcium in urine can be determined by the standard EDTA method in the same way as it was determined in other materials, or by photometric titration, or with the addition of fluorexone. Due to the increased content of phosphates in urine, when analyzing it, it is often useful to strongly dilute the analyzed solution or, to avoid precipitation of poorly soluble compounds, to use back titration.

In addition to these methods developed specifically for the determination of Ca, suitable methods for the determination of Ca may also be found among the methods described below for the determination of Ca and Mg, since many Ca determinations are associated with the determination of Mg.

The first determination of magnesium in serum was described by Golasek and Flaschka. Calcium is precipitated as oxalate and titrated after dissolution of the precipitate, while Mg is determined in the filtrate after centrifugation. The advantage of this method is that both metals can be determined in the same solution. The method proposed by Gjessing, in which sequential titration is carried out, has a similar advantage. Ca is first titrated photometrically with murexide in an alkaline solution (NaOH), with small amounts of Mg(OH)2 remaining in solution, apparently in colloidal form, not interfering. Then glycine is added and boiled. In this case, murexide is destroyed and magnesium hydroxide dissolves; after this, Mg is titrated with eriochrome black T. However, most methods are based on the use of two aliquots of samples. In one sample, Ca is titrated in a solution with a high pH value with murexide (see above) or with another indicator, for example Erio SE, and in the other, the sum of Ca and Mg is titrated.

In the last titration, eriochrome black T is usually used. The method is suitable for working with ultra-microquantities and is superior in accuracy if photometric titration is used. The titration process can be automated.

Calcium and magnesium in urine can be determined in the same way as in serum, but with slight modifications.

Calcium and magnesium in plasma and cerebrospinal fluid are determined in exactly the same way as in serum.

Determination of water hardness. The determination of water hardness has long been described by Schwarzenbach et al. and is the first complexometric titration method used in practice. Numerous methods for determining water hardness, including microdeterminations, can be found in the literature.

Two groups of methods should be distinguished: determination of total hardness and separate determination of calcium and magnesium hardness. When determining total hardness, the sum of Ca and Mg is titrated. Titration is usually carried out in a solution with pH = 10 with eriochrome black T as an indicator. In order for the color transition of the indicator to be sharp, the presence of at least 5% Mg (relative to the Ca content) must be present.

Since this condition is not always met for waters of different origins, a known amount of Mg should be added and taken into account in the calculations or, even better, introduced into the analyzed solution in the form of a magnesium complex with EDTA. When performing serial analyses, it is much easier to use a titrated solution, which, along with EDTA (H2Y2~), contains the required amount of MgY2~.

When studying the factors that interfere with these titrations, it was found that they mainly include small impurities of heavy metals, which either cause excessive titrant consumption or block the indicator. Their removal is not difficult if you add a mixture of KCN with ascorbic acid or triethanolamine as masking substances. Na2S is also a good masking agent for most metals, except Al. Masking agents are often added to the buffer solution.

Hahn avoids or reduces interference by titrating a known amount of a standard EDTA solution with the water being analyzed. However, this technique is difficult for practical application. When titrating with chromasurol S, interference is less dangerous, since this dye is less susceptible to blocking. However, the color transition in this case is less abrupt than when using eriochrome black T.

When separately determining calcium and magnesium hardness, two aliquot portions of the solution are usually used. In one part of the solution, Ca is titrated at a high pH value, in the other, at pH = 10, the sum of Ca and Mg is titrated. Magnesium is calculated by its content.

Titration of calcium, as a rule, does not cause difficulties, since in all normal waters the Ca content greatly exceeds the Mg content.

To analyze waters containing polyphosphates, Brook proposed separating the Ca titration using the ion exchange method. Schneider et al. used eriochrome blue-black B as an indicator when determining the hardness of sugar syrup.

The complexometric determination of permanent hardness can be preceded by an acid-base determination of temporary hardness, after which complexometric titration can be carried out directly in the same solution. Photometric titrations of interest for the analysis of colored waters are reported. Photometric indication allows titration to be automated.

Laci describes a semi-automatic method in which the titration curve obtained in the presence of eriochrome black T is plotted with a chart recorder. The curve has two inflections, the first of which corresponds to the end of the titration of Ca. Thus, simultaneous determination of calcium and magnesium hardness is possible. Erdey et al. also obtained two inflections in the curve during high-frequency titration.

Conductometric titration has proven itself in the analysis of turbid and colored waters. Since the concentration of salts in natural waters is usually insignificant, the conductometric method is very suitable for their analysis, due to the absence of a background that interferes with the determination of electrical conductivity.

Direct determination of magnesium with eriochrome black T

Reagents EDTA, 0.01 M solution. Eriochrome black T.

Buffer solution, pH = 10.

Progress of determination. The Mg concentration in the test solution should not exceed 10 -2 M. Acidic test solutions are first neutralized with sodium hydroxide. Then, 2 ml of buffer solution and a few drops of eriochrome black T are added to each 100 ml of solution and titrated until the red color turns blue.

With the last drop of the titrant solution, the reddish tint of the indicator should disappear. Since complexation reactions do not occur instantaneously, the titration is slowed down near the end point.

Notes. The curves shown in Fig. 32 and obtained by a combination of the curves shown in Fig. 4 and 23 show that during the titration process a pH value of 10 should be maintained fairly accurately. Both too low and too high a pH value impairs the recognition of the equivalence point. Therefore, acidic test solutions before adding buffer

do not introduce additional ammonium ions into the solution. With the right choice of titration conditions, the equivalence point is so sharp that even a 0.001 M EDTA solution can be titrated.

Determination of calcium with eriochrome black T using the displacement method

Reagents

EDTA, 0.01 M solution.

Eriochrome black T.

Buffer solution, pH = 10.

Magnesium complex with EDTA, 0.1 M solution.

Progress of determination. The concentration of calcium ions should not exceed 10 -2 M. If the analyzed solution is acidic, it is neutralized with sodium hydroxide. To each 100 ml of the analyzed solution, add 2 ml of a buffer solution, 1 ml of a 0.1 M MgY solution, 2-4 drops of eriochrome black T and titrate until the red color turns blue. With the last drop of the titrant solution, the reddish tint should completely disappear. Near the end point, the titration is slowed down.

Notes. The curves shown in Fig. 33 and obtained by a combination of the curves shown in Fig. 5 and 24 show how the color of eriochrome black T changes if Ca2+ ions are titrated without adding magnesium complexonate. In this case, even at pH = 11, a sharp color transition does not occur; in addition, in such a strongly alkaline solution a pure blue color is not obtained, since in this pH region eriochrome black T behaves as an acid-base indicator.

The curves shown in Fig. 34 show the improvements achieved by the addition of magnesium complexonate. Since calcium complexonate is more stable than magnesium complexonate, Mg is displaced and, as a consequence, simultaneous titration of Ca and Mg occurs (see Fig. 11).

The curves shown in Fig. 34, obtained by a combination of Fig. 11 and 23. They show that the addition of only 1% Mg already significantly improves the recognition of the equivalence point. By adding 10% Mg, almost the maximum possible effect is obtained. Further addition of MgY2 would only lead to an unnecessary increase in the ionic strength of the solution and a decrease in the pMg jump. When titration is carried out correctly, the color change is so dramatic that microdeterminations can be made using even a 0.001 M EDTA solution.

With photometric indication of the equivalence point, titration results are noticeably improved.

Direct determination of calcium with calcon

Reagents

EDTA, 0.01 M solution. Tracing paper.

Caustic potash, 2 M solution. Diethylamine.

Progress of determination. The calcium concentration in the titrated solution should be about 10 -2 M. Acidic solutions are first neutralized with sodium hydroxide or potassium hydroxide. To each 100 ml of neutralized test solution add 5-7 ml of diethylamine. This amount is quite sufficient to establish the pH value of the solution around 12.5. Then add the indicator using tracing paper and titrate (immediately to prevent precipitation of CaCO3) with an EDTA solution until a stable pure blue color occurs.

Notes. The required pH value of the titrated solution can also be set using KOH or NaOH.

Some observers have noted that the equivalence point in calcone titrations is sharper if small amounts of magnesium are present. In this case, if there is no Mg in the analyzed solution, add 1-2 ml of a 0.1 M solution of magnesium salt. Then slowly, with strong stirring, the solution is made alkaline. The amount of diethylamine indicated above is sufficient to establish the appropriate pH in the presence of Mg. When titrating in the presence of Mg, sometimes after the end point the solution becomes discolored again when standing; then add another 1-2 drops of EDTA titrating solution to obtain a stable blue color. Therefore, if magnesium is present in the solution, you should wait about half a minute before counting on the burette.

HEDTA can be used as a titrant instead of EDTA, especially when the determination of Ca is carried out in the presence of a large amount of Mg and tartaric acid is added to prevent the precipitation of Mg.

Methods for quantitative determination of calcium. There are various methods for determining calcium.

a) Weight methods.

1) Precipitation as oxalate and weighing as or (see “Weight analysis”).

2) Precipitation in the form of sulfate from an alcohol solution.

3) Precipitation as picrolonate.

b) Volumetric methods.

1) Precipitation as calcium oxalate and subsequent determination of the calcium-bound oxalate ion by permanganatometry or cerimetry.

2) Precipitation in the form of molybdate, reduction of molybdenum and titration with ammonium vanadate.

3) Complexometric method.

The gravimetric method for determining calcium has very significant disadvantages.

1. Determining the calcium content in various technical objects by the gravimetric method is a very lengthy operation.

2. Precipitation of calcium ions in the form is associated with great difficulties due to the impossibility of achieving quantitative separation of calcium oxalate;

3. The calcium oxalate precipitate is often contaminated with foreign impurities and is difficult to isolate in a chemically pure form.

4. Obtaining a weight form involves the use of a relatively high temperature necessary for the thermal decomposition of calcium oxalate.

5. The resulting weight form is unstable and is exposed to moisture and carbon dioxide in the air, as a result of which its mass changes depending on the conditions of production and storage.

Therefore, at present, the gravimetric method for determining calcium has lost its former significance and has been replaced by more progressive volumetric methods of analysis.

One of these methods is described above (see Chapter II.I, § 28). The permanganatometric method for determining calcium has a number of advantages compared to gravimetric methods of analysis. One such benefit is faster completion of the definition operation. However, the permanganatometric method for determining calcium, based on the precipitation of calcium ions in the form of oxalate and subsequent titration of oxalate ions with permanganate, has many disadvantages of gravimetric analysis associated with the impossibility of complete quantitative precipitation and separation of calcium oxalate.

Of the volumetric methods of analysis, the most accurate and rapid method for determining calcium is undoubtedly the complexometric titration of calcium ions with complexon III.

Complexometric method for calcium determination. Complexometric determination of calcium is based on the direct method of titrating its ions with a standard solution of complexone III in the presence of murexide or acid chromium dark blue. The indicator forms a red complex compound with calcium ions. When the solution is titrated with complexone III at the equivalence point, the red color turns into the color characteristic of the free indicator.

As a result of the titration of calcium salts with complexon III, a complex and acid are formed:

The resulting complex is relatively unstable:

Therefore, the formation of free acid during the reaction or its addition to the titrated solution before titration shifts the indicated equilibrium to the left, i.e., towards the destruction of the complex.

EDTA is a tetrabasic acid, characterized by the following constants: and is a relatively weak acid, therefore the solution of its complex with should not be lower than 10.3. If it is less, it forms the corresponding hydroanions: and acid. In this case, the complex is destroyed or not formed at all.

Thus, the stability of the intracomplex salt formed by calcium ions with complexone III depends on the size of the solution.

Therefore, to ensure the optimal course of the complex formation reaction, the titration of calcium salts with an EDTA solution must be carried out in a highly alkaline medium at . In this case, complete neutralization of the free acid formed during the nitration process is achieved and a maximum jump in the titration curve is observed (Fig. 61).

Rice. 61. Titration curves of calcium ions using the complexometric method at different solution values: 1 - ; 2 - ; 3 - ; 4 - .

The method is based on the property of Trilon B (disodium salt of ethylenedi-aminoacetic acid) to produce extremely stable complex compounds with divalent metal ions, including calcium and magnesium.
For this method, absorbed bases from the soil must be displaced with 1.0 N. ammonium acetate solution at pH 6.5 or 1 N. NaCl solution. The trilonometric method is more convenient to carry out under conditions of low salt concentrations. Therefore, after displacing the absorbed bases, ammonium acetate is destroyed by evaporating the solution, then the resulting residue is calcined on a heating mantle or in a muffle at 400-600°, and calcium and magnesium are obtained in the form of carbonates or oxides. The organic matter burns out. The resulting precipitate is dissolved with 10% hydrochloric acid and, after making sure that it is completely dissolved (no crystals are visible at the bottom of the cup), the hydrochloric acid solution is diluted with hot water, filtered into a 200 ml volumetric flask and diluted to the mark with water.
Ammonium acetate displaces small amounts of sesquioxides from the soil, so in many cases they do not have to be separated from solution. High concentrations of iron interfere with titration with Trilon - the color transition loses clarity; in addition, you can get slightly overestimated data. In such cases, it is recommended to further dilute the solution with water to reduce the concentration of iron, or to isolate iron if there is a lot of it in relation to calcium and magnesium. It is better to do this before preparing the solution for final evaporation - isolate the sesquioxides in the usual way with ammonia, and then complete the evaporation and calcinate the precipitate.
The harmful effects of manganese are destroyed by adding hydroxylamine hydrochloride (1-2 ml of a 5% solution), which prevents the formation of manganese peroxide, which interferes with titration. It is also necessary to eliminate the harmful effects of copper. All reagents for this purpose are prepared using distilled water that does not contain copper. Distilled water must be obtained using a device that does not have copper parts. The harmful effects of traces of copper are destroyed by adding 1-2 ml of 2% Na2S to the test solution, which converts it into insoluble copper sulfide.
The determination is made by titrating the test sample with a solution of Trilon B in the presence of the black chromogen indicator, and calcium ions are first bound into the complex, and then magnesium ions. Magnesium ions cause a particularly sharp change in the color of the indicator, while calcium ions do not give a clear change in the color of the solution, and therefore calcium can only be determined in the presence of magnesium, i.e., determine the sum of calcium and magnesium.
Determination of the amount of calcium and magnesium. A certain part of the analyzed solution of absorbed bases (it is convenient to take 50 ml) is placed in a conical flask with a capacity of 250 ml, diluted with water to approximately 100 ml. The solution is heated to 60-70°, 5 ml of ammonia buffer solution is added to create an alkaline reaction, then 0.5 ml of Na2S and 0.5 ml of hydroxylamine, 10-15 mg of black chromogen indicator (or blue-black chromium) and titrated to 0. 01 - 0.05 n. with a solution of Trilon B with vigorous stirring until the color of the solution changes from cherry-red through violet-blue to pure blue at the equivalence point. When adding excess trilon, the color does not change. Therefore, it is recommended to carry out titration by comparing the color of the solution with a “witness” - a sample that has been obviously titrated.
The sum of calcium and magnesium (in mEq per 100 g of soil) is equal to:

Determination of calcium by the trilonometric method in the presence of the indicator murexide (ammonium salt of monobasic purple acid). With calcium ions, the purple acid anion forms a red-colored complex in an alkaline medium. This complex is less stable than the calcium compound with Trilon, and upon titration there is a sharp change in color from red to purple at the equivalent point. The harmful effects of copper and manganese are destroyed in the same way as when titrating the amount of calcium and magnesium.
Progress of the analysis. A certain volume of solution is placed in a 250 ml conical flask, and the solution is diluted with water to approximately 100 ml.
To prevent coprecipitation of calcium with magnesium during the direct determination of calcium with murexide, 2 ml of 0.5 N is first introduced into the sample (before adding NaOH). Na2CO3 solution. In this case, calcium precipitates in the form of CaCO3, forming a separate phase, which dissolves during subsequent titration. This eliminates the possibility of coprecipitation of calcium with Mg(OH)2 and ensures the completeness of calcium determination. Add 2 ml of 2.0 N. NaOH, 0.5 ml of Na2S solution and 0.5 ml of hydroxylamine solution, then dry murexide on the tip of a knife and titrate with 0.05 or 0.01 N. with a solution of Trilon B with vigorous stirring until the bright purple color of the solution turns purple.
Subsequently, the addition of Trilon does not change the color, so titration is best carried out in the presence of a “witness” - a sample that has been obviously overtitrated.

From the sum of calcium and magnesium per 100 g of soil, subtract the amount of calcium and obtain the amount of magnesium (in mEq per 100 g of soil).
The obtained data on the content of absorbed calcium and magnesium are recalculated per 100 g of dried soil.
Reagents. 1. Trilon B solution. To prepare 0.05 p. solution, dissolve 9.3 g of Trilon in 1 liter of distilled water. 0.01 n. the solution is prepared by diluting 0.05 N. solution. The titer of the Trilon solution is determined using magnesium sulfate. The commercially available chemically pure salt MgSO4 7H20 is recrystallized, dried for 24 hours between sheets of filter paper and kept in a desiccator over a mixture of 5 parts MgSO4 7H2O and 1 part water until it stays dry. 0.01 n. the solution contains 1.232 g of MgSO4 7H2O in 1 liter of water. It is recommended to check the amount of magnesium in the solution prepared to check the Trilon titer using the gravimetric pyrophosphate method and make the necessary correction.
2. Buffer solution. 20 g of ammonium chloride are dissolved in 500 ml of distilled water, 100 ml of a 25% ammonia solution is added and the volume is adjusted to 1 liter.
3. Indicator for titrating the amount of calcium and magnesium. 0.2 g of chromogen black is dissolved in 10 ml of ammonia buffer and diluted with water to 100 ml. The indicator solution is stable for 1 month. It is convenient to prepare this indicator for analysis by rubbing it with NaCl until it is uniformly colored (5 g of indicator and 95 g of NaCl), and store it in a jar with a ground stopper in a dark place. When titrating, add 10-15 mg for each determination.
To check the titer of Trilon for magnesium, 20 ml of the prepared solution of magnesium sulfate is pipetted into a 250 ml conical flask, 100 ml of distilled water, 5 ml of ammonia buffer, 10-15 ml of black chromogen are added and the cherry-red solution is titrated with 0.01 N. with Trilon solution until the color of the solution turns blue.
4. Ready-made commercial murexide is prepared for analysis by rubbing it with NaCl until the color is uniform (5 g of indicator and 95 g of NaCl). Store in a jar with a ground stopper in a dark place. 10-15 mg of the resulting salt is taken for analysis.
To create the necessary alkaline reaction when titrating with Trilon with a murexide indicator, use 2 N. caustic soda solution Small amounts of Na2S are added to hydroxylamine in case of displacement of manganese and copper from the soil.
Determination of mobile magnesium in soils is carried out in 1 N. KCl hood. 100 g of soil, sifted through a sieve with 1 mm holes, is placed in a bottle, filled with 250 ml of 1 N. KCl, shake on a rotator for 1 hour and filter through a pleated filter.
To determine the sum of Mg, Ca and Mn, place 50 ml of the extract in a 150 ml beaker, add 5 ml of an ammonia buffer mixture (20 g of chemically pure NH4Cl and 100 ml of a 25% NH4OH solution in 1 liter of water, 2 ml of 1% hydroxylamine hydrochloric acid solution, 50 ml of distilled water and on the tip of a knife a dry indicator - dark blue acid chromium, sour cream with NaCl in a ratio of 1: 99. Then the extract is titrated with a 0.02 N solution of Trilon B until the color of the solution changes from cherry red to pure blue. When using a photoelectric titrimeter of the FET-UNIZ type, titration is carried out until the ammeter needle stops.
To determine the sum of Ca and Mn, take 50 ml of the extract, add 2 ml of a 1% solution of hydroxylamine hydrochloride, 10 ml of a borate buffer mixture (6 ml of 0.05 N borax solution and 4 ml of 0.02 N boric acid solution) , 10 ml of ammonia buffer mixture and dry murexide on the tip of a knife. 0.02 N is added to the extract from a burette. Trilon B solution until the color changes from orange to crimson. Then add 2 ml of a 20% NaOH solution and continue titration until the titrimeter needle stops or the bright purple color changes to purple during visual titration.
Magnesium is determined by the formula:

Direct titration method. The analyzed solution containing cations of the metal being determined is diluted in a volumetric flask and an aliquot of the solution is taken for titration.

Titration is carried out with a standard EDTA solution in an alkaline medium with eriochrome black T or in an acidic medium with xylene orange.

To do this, the titrated solution is first adjusted to a certain pH value using a buffer solution before titration. Along with the buffer solution, an auxiliary complexing agent (tartrate, citrate, etc.) is sometimes added, which binds some cations and keeps them in a soluble state to avoid precipitation of hydroxides in the alkaline solution.

During direct titration, the concentration of the cation being determined first decreases gradually, then drops sharply near the equivalence point. This moment is noticed by the change in color of the introduced indicator, which instantly reacts to changes in the concentration of complexing metal cations.

The direct complexometric titration method is used to determine Cu 2+ , Cd 2+ , Pb 2+ , Ni 2+ , Co 2+ , Fe 3+ , Zn 2+ , Th IV , Al 3+ , Ba 2+ , Sr 2+ , Ca 2 + , Mg 2+ and some other cations. The determination is hampered by complexing substances that retain the ions being determined in the form of complex ions that are not destroyed by complexons.

Back titration method. In cases where, for one reason or another, it is impossible to carry out direct titration of the cation being determined, the reverse titration method is used. A precisely measured volume of a standard complexone solution is added to the analyzed solution, heated to boiling to complete the complexation reaction, and then the excess complexone is titrated in the cold with a titrated solution of MgSO 4 or ZnSO 4 . To establish the equivalence point, an indicator metal is used that reacts to magnesium or zinc ions.

The back titration method is used in cases where there is no suitable indicator for cations of the metal being determined, when cations form a precipitate in a buffer solution, and when the complexation reaction proceeds slowly. The back titration method is also used to determine the content of cations in water-insoluble sediments (Ca 2+ in CaC 2 O 4, Mg 2+ in MgNH 4 PO 4, Pb 2+ in PbSO 4, etc.).

Substituent titration method. In some cases, instead of the methods described above, the substituent titration method is used. The method of complexometric titration of a substituent is based on the fact that Mg 2+ ions form a less stable complex compound with a complexone (pK = 8.7) than the vast majority of other cations. Therefore, if you mix cations of the metal being determined with a magnesium complex, an exchange reaction will occur.

For example, this reaction is used to determine thorium ions when magnesium complexonate MgY 2 - is first introduced into the analyzed solution, and then the released Mg 2+ ions are titrated with a standard EDTA solution (b);

Th 4+ + MgY 2 ‑
Mg 2+ + H 2 Y 2 ‑		MgY 2 ‑ +2H +

Due to the fact that Th IV forms a more stable complex compound with the complexone than Mg 2+, the equilibrium of reaction (a) shifts to the right.

If, at the end of the displacement reaction, Mg 2+ is titrated with a standard EDTA solution in the presence of eriochrome black T, then the content of Th IV ions in the test solution can be calculated.

Methodacid-base titration. During the interaction of the complexon with certain metal cations, a certain amount of equivalents of hydrogen ions is released.

The hydrogen ions formed in equivalent quantities are titrated by the usual alkalimetric method in the presence of an acid-base indicator or by other methods.

There are other methods of complexometric titration, the description of which is beyond our scope.

Setting the titer of the EDTA solution

To prepare a standard (titrated) EDTA solution, disodium salt of ethylenediaminetetraacetic acid is used, which crystallizes with two water molecules; its composition corresponds to the formula Na 2 C 10 N 14 O 8 N 2 2H 2 O.

If a disodium salt containing water of crystallization is dried at 120-140°C, then an anhydrous salt is obtained, the composition of which corresponds to the formula Na 2 C 10 H 14 O 8 N 2.

Both salts can serve as starting materials for preparing a standard EDTA solution.

To prepare 1 liter 0.1 n. EDTA solution you need to take:

M Na 2 C 10 H 14 O 8 N 2 2H 2 O╱2 10 = 372.24╱ 2 10 = 18.61 g

M Na 2 C 10 H 14 O 8 N 2 ╱2 10 = 336.21╱ 2 10 = 16.81 g

To set the EDTA titer, use x. including calcium carbonate, x. including ZnO or x. including metallic zinc, a calculated portion of which is dissolved in x. including hydrochloric or sulfuric acid, neutralized with sodium hydroxide or ammonia, diluted with an ammonia buffer solution and titrated with a standard EDTA solution in the presence of the required indicator. Towards the end, titrate slowly.

The titer of a solution can also be determined using the magnesium salt fixation (0.01 and 0.05 N solutions of magnesium sulfate are commercially available).

Based on the titration results, T is calculated, N And TO EDTA solution.

Determination of calcium content

Methods for quantitative determination of calcium. There are various methods for determining calcium.

Gravimetric methods.

1. Precipitation in the form of CaC 2 O 4 -H 2 O oxalate and suspension in the form of CaCO 3 or CaO (see “Gravimetric analysis”).

2. Precipitation in the form of CaSO 4 sulfate from an alcohol solution.

3. Precipitation in the form of picrolonate Ca(C 10 H 7 O 5 N 4) 2 8H 2 O.

Titrimetric methods.

1. Precipitation as calcium oxalate and subsequent determination of the calcium-bound oxalate ion by permanganatometry or cerimetry.

2. Precipitation in the form of molybdate CaMoO 4, reduction of molybdenum and titration with ammonium vanadate.

3. Complexometric method.

The gravimetric method for determining calcium has very significant disadvantages.

1. Determining the calcium content in various technical objects by the gravimetric method is a very lengthy operation.

2. The precipitation of calcium ions in the form of CaC 2 O 4 is associated with great difficulties due to the impossibility of achieving quantitative separation of calcium oxalate;

3. The calcium oxalate precipitate is often contaminated with foreign impurities, and it is difficult to isolate it in a chemically pure form.

4. Obtaining the weight form (CaO) involves the use of a relatively high temperature necessary for the thermal decomposition of calcium oxalate.

5. The resulting weight form (CaO) is unstable and is exposed to moisture and carbon dioxide in the air, as a result of which its mass changes depending on the conditions of production and storage.

Therefore, at present, the gravimetric method for determining calcium has lost its former significance and has been replaced by more progressive titrimetric methods of analysis.

The permanganatometric method for determining calcium has a number of advantages compared to the gravimetric method of analysis. One such benefit is faster completion of the definition operation. However, the permanganatometric method for determining calcium, based on the precipitation of calcium ions in the form of oxalate and subsequent titration of oxalate ions with permanganate, has many analytical disadvantages associated with the impossibility of complete quantitative precipitation and separation of calcium oxalate.

Of the titrimetric methods of analysis, the most accurate and fastest method for determining calcium is undoubtedly the complexometric titration of calcium ions with EDTA.

Complexometric method for calcium determination. Complexometric determination of calcium is based on the direct method of titrating its ions with a standard EDTA solution in the presence of murexide or acid chromium dark blue. The indicator forms a red complex compound with calcium ions. When titrating an EDTA solution at the equivalence point, the red color turns into the color characteristic of a free indicator.

As a result of the titration of calcium salts with EDTA, the formation of a complex of CaY 2 - and acid occurs:

Ca 2+ + H 2 Y 2 ‑ ⇄ CaY 2 ‑ + 2H +

The resulting CaY 2 complex is relatively unstable:

╱ =310 ‑11

The formation of free acid during the reaction or its addition to the titrated solution before titration shifts the indicated equilibrium to the left, i.e., towards the destruction of the complex.

EDTA is a tetrabasic acid characterized by the following constants: pK 1 = 2; rK 2 = 2,7; rK 3 = 6,2; rK 4 = 10.3 and is a relatively weak acid, therefore the pH of the solution of its complex with Ca 2+ should not be lower than 10.3. If the pH is lower, then Y 4 ‑ with H + forms the corresponding hydroanions: HY 3 ‑, H 2 Y 2 ‑, H 3 Y ‑ and acid H 4 Y. In this case, the CaY 2 ‑ complex is destroyed or not formed at all.

Thus, the stability of the intracomplex salt formed by calcium ions with EDTA depends on the pH of the solution. To ensure the optimal course of the formation reaction of the CaY 2 complex, the titration of calcium salts with an EDTA solution must be carried out in a strongly alkaline medium at pH > 12. In this case, complete neutralization of the free acid formed during the titration is achieved and a maximum jump in the titration curve is observed.

Volume of EDTA, ml

Rice. 6.1 Titration curves of calcium ions using the complexometric method at different pH values of the solution:

1 – pH =6; 2 – pH = 8; 3 – pH = 10; 4 – pH = 12

Methods for quantitative determination of calcium. There are various methods for determining calcium.

Gravimetric methods.

1. Precipitation in the form of CaC 2 O 4 -H 2 O oxalate and suspension in the form of CaCO 3 or CaO (see “Gravimetric analysis”).

2. Precipitation in the form of CaSO 4 sulfate from an alcohol solution.

3. Precipitation in the form of picrolonate Ca(C 10 H 7 O 5 N 4) 2 8H 2 O.