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same constituent is being made, this device saves a lot of time since it obviates the necessity of making an individual calculation for each determination. The general method of arriving at the desired weight is shown herewith:

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1000

desired weight of sample in grams

% constituent to be represented by 1 milligram of precipitate weight in milligrams of precipitate from W

number of grams of constituent corresponding to 1 gram of precipitate

= grams or constituent in W

Also by the condition that has been set, we have

(1)

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From this we see that if W is taken equal to f, p is equal to 0.1, namely, each milligram of precipitate will correspond to 0.1% constituent in sample; if W is taken equal to one-tenth of ƒ, p is equal to 1, namely, — each milligram of precipitate will correspond to 1% constituent in sample.

Example. Suppose that it is required to weigh out such a quantity of pig iron for analysis that each milligram of barium sulphate precipitate will correspond to 0.0025% sulphur in the sample. We have

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It is easy to verify the foregoing. Thus assume the weight of BaSO4 obtained from 5.492 g. sample was 0.050 g., this multiplied by 0.0025 gives 0.125% as the percentage of S in the sample. If we compute the percentage by using the factor for S in BaSO4 0.050 X 0.1374 we have = 0.00125 g. S in 1 g. which is equal to 5.492

0.125%.

197. Examples.

1. What error, expressed in percentage of sulphur, would result if in the analysis of an ore containing 8.34% of sulphur, ten percent of the BaSO, were reduced to BaS by the filter paper and so weighed?

Ans. Error = 0.23%

2. 5.018 g. of a sample of crucible iron yielded 0.0159 g. barium sulphate. What was the content of sulphur? Ans. 0.044%

3. 4.160 g. of the same sample of crucible iron as mentioned in Ex. 2 gave 0.0640 g. magnesium pyrophosphate. What was the content of phosphorus? Ans. 0.43%

4. 10.00 g. of potash marl (soil) was dissolved and made up to 500 c.c. (method of A. O. A. C. loc. cit., § 13). 50 c.c. of this solution was used for the determination of the phosphorus by the alkalimetric method (§ 188). The "yellow precipitate" was dissolved in 20.00 c.c. of 0.1086 M NaOH; the excess of NaOH was titrated back with 0.1509 M HNO, in the presence of phenolphthalein solution as indicator, the back titration requiring 9.84 c.c. of the acid. What was the content of phosphorus calculated as P2O5?

Ans. 0.22%

5. How many grams of steel should be taken for analysis in order that each milligram of barium sulphate shall correspond to 0.004% sulphur?

Ans. 3.433 g.

6. How many grams of steel should be taken for analysis in order that each milligram of ammonium phospho-molybdate found, shall correspond to 0.001% phosphorus? Ans. 1.654 g.

7. How many grams of iron ore must be taken in order that each milligram of magnesium pyrophosphate shall correspond to 0.004% P in the ore?

Ans. 6.97 g.

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for the determination of calcium but they are all based upon the procedure of getting it in the form of calcium ion and then precipitating the latter as calcium oxalate by means of a solution of ammonium oxalate,1 according to the reaction

Ca+++ C2O → CaC204↓

The calcium oxalate may be ignited to calcium oxide and the latter compound weighed, or the calcium oxalate may be dissolved in dilute sulphuric acid and the oxalate ion titrated by means of standard permanganate solution. In the precipitation of calcium oxalate care must be taken that the solution is free of all forms of silicic acid, likewise all ions that give insoluble oxalates, namely, ions of the copper, iron and aluminum groups and also strontium ion. The following ions are permissible: magnesium, barium, sodium, potassium and ammonium, although it is to be remarked that the first four of these are always more or less co-precipitated with calcium oxalate and that double precipitation with intervening filtration is always necessary when they are present.

The presence of magnesium ion deserves special consideration because calcium and magnesium so often have to be determined from the same solution (particularly in the analysis of limestones, cements, rocks, etc.) after the removal of the iron and aluminum group by means of ammonium hydroxide and ammonium chloride. The precipitation of calcium oxalate in the presence of

1 This precipitant is selected in preference to others, such as fluoride or carbonate ion, because the calcium oxalate formed is the most insoluble salt of calcium and also because its crystalline character permits of easy and thorough washing. While the fluoride possesses about the same solubility as the carbonate, it has the objectionable feature that it is precipitated in a slimy condition. Calcium carbonate is much more soluble than the oxalate.

magnesium ion always results in the co-precipitation of magnesium oxalate. This necessarily requires re-solution of the precipitate and reprecipitation under conditions which reduce the tendency of magnesium oxalate to separate out with the calcium. If magnesium is present in excess of the calcium, the calcium is incompletely precipitated unless an excess of oxalate ion is present. Too great excess of oxalate is detrimental in that it favors the co-precipitation of magnesium oxalate. The incomplete precipitation of calcium when much magnesium is present is due, according to Noyes,2 to the slight ionization of magnesium oxalate in respect to other salts of its type. The oxalate ions, as soon as they are added, are first removed to form molecular magnesium oxalate until the necessary equilibrium conditions are satisfied. At no time during this stage will the concentration of oxalate ion be sufficient to exceed the solubility product of calcium oxalate. The amount of oxalate added might appear to the analyst as more than enough to precipitate the calcium from solution, an assumption which would probably be true if magnesium were absent, but a large part of the oxalate concentration which was designed to precipitate the calcium will have been removed to form molecular magnesium oxalate; hence it is important to add sufficient oxalate to take care of this formation of molecular magnesium oxalate.

The co-precipitation of magnesium, which amounts to about 0.1% to 0.2% of the amount of magnesium present when there is no excessive disparity in the relative proportions of magnesium and calcium, is reduced to a minimum in the second precipitation by causing the precipitation to take place from as acid a solution as possible. The solution must be acidulated with enough hydrochloric acid to repress the concentration of oxalate ion to a point where the solubility product of calcium oxalate is not reached. Oxalic acid is a comparatively weak acid and hence any increase in the hydrogen ion concentration of the solution will result in a corresponding decrease in oxalate ion concentration, according to the scheme

H2C2O4 → HC204 ↔ C204

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On slow addition of dilute ammonium hydroxide, the hydrogen ion concentration is gradually decreased, due to its removal to form H2O, while the oxalate ion concentration increases, which is the reverse of the reaction above. A point is reached where cca++ X CcOS.P.Cacio, and calcium oxalate precipitates out. The solution at this stage is distinctly acid (p1 about 4.0) and remains so until most of the calcium has been precipitated out as calcium oxalate. The remaining calcium precipitates out when the solution reacts alkaline. This removal of most of the calcium in acid solution reduces the co-precipitation of magnesium oxalate and at the same time gives a coarse crystalline precipitate which is easily filtered and washed.

199. In washing calcium oxalate it is to be noted that this precipitate shows a marked tendency to creep up the sides of the filter paper with the result that some of it reaches the top edge of the filter paper and becomes washed down on the under side of the paper. With ordinary amounts of calcium oxalate and a 12 cm. filter paper, as much as 1.0 to 2.0 mg. of the precipitate may pass into the filtrate.

200. The calcium oxalate is not weighed as such because its composition is not uniform. It is mainly CaC2O4 H2O. This compound on drying loses water incompletely and at 200° it begins to decompose into CaCO3 and CO. This decomposition proceeds slowly and is complete at about 400°. The decomposition of the oxalate into carbonate and weighing it as such are sometimes recommended, but it is far preferable to carry the decomposition to the oxide, which takes place at about 850°, and to weigh it as such. In this way a close check upon the decomposition temperature is unnecessary. The one drawback of weighing the precipitate as oxide is that the oxide takes on water and CO2 from the air very rapidly and also probably occludes gases. In order to reduce this source of error to a minimum the calcium oxide is kept covered during the cooling and weighing period, and the weighing conducted immediately after the cooling. If a platinum crucible is used for the ignition of calcium oxalate, the decomposition to oxide is quick and satisfactory. If porcelain is used the decomposition to oxide is slow and not so satisfactory, especially if the amount of calcium oxa

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