SIMULTANEOUS STANDARDIZATION of 0.1 N HC1 and NaOH USING CALCITE A. J. HAMMER Chariton Junior College, Chariton, Iowa
Crystals of calcite are frequently of suficient purity to w r r a n t the use of this mineral as a primary standard. The method recommended uses a weighed sample of fiawdered calcite to which is added a measured volume of acid, the amount of acid being in excess. The excess acid i s then titrated with a base. This method gives
results which agree fairly closely with the silver chloride method of standardization. The ahelopment of the formula and its use in calculating the normalities of hydrochloric acid solutions have been found useful in presenting the topic of normal solutias to a class i n quantitative analysis.
T
proceed until all of the acid has reacted and the crystal is then dried and weighed. The normality of the acid is determined from the loss in weight of the crystal. The above method was suggested by Masson (3) and some authors (44) have adopted it as a means of
HE METHOD of standardizing an acid using calcite is not new (1-3). Two general methods are in use. In the first, a calcite crystal of known weight is placed in a measured volume of the acid to he standardized. The action is allowed to
standardization in their texts. Rivet (7), however, has pointed out that this method may lead to possible error because of the comparatively high solubility of some calcite samples in solutions containing carbon dioxide. The second method makes use of the powdered calcite, the crystal being ground to approximately 100 mesh immediately before use. This method has been in use for a number of years and has proved quite satisfactory as compared with the silver chloride method of standardization. Calcite has been shown to approach 100% purity within 0.2'% or within the range of analytical accuracy (2). The common impurities are Mn, Zu, Mg, or Fe which have replaced the calcium ( 8 ) . Manganese and iron are quite easily detected by the color they impart to the crystal. The following treatment of the subject has the advantage that acid and alkali are standardized simultaneously using this primary standard, and the method of calculation has proved of value a t Chariton Junior College. For student instruction, it has been found best to use HCl and NaOH solutions of from 0.3 to 0.5 normality, as the time necessary for the solution of the calcite is shortened sufficiently to permit the completion of the standardization in one two-hour laboratory period.
The problem states that 22.2 cc. of the NaOH were necessary to neutralize the excess acid. In order to determine the volume of the HCl which is equivalent to this volume (22.2 cc.) of NaOH we again substitute in formula (1) and find this volume to be: CC
=
(22) ( N ) n
(3)
Then if 50 cc. of acid was added to the calcite and this volume (equation 3) was not used up by the calcite, it is obvious that the volume of acid which actually reacted with the calcite would be :
But the normality of the HC1 (n in equation 2) is equal to:
p & @ ?
27
so if we substitute this value for (n) in equation (4) we have: 50 - (22.2) (N)
(30)O 27
or:
OUTLINE OR PROCEDURE
Suppose we have solutions of HCI and NaOH which we wish to standardize. We do not know the normality of either solution, but we have determined by titration that 30 cc. of the NaOH solution is equivalent to 27 cc. of the HCl solution. When 50 cc. of the acid is allowed to read with 0.75 g. of calcite and the excess acid is titrated with the NaOH solution, it is found that 22.2 cc. of NaOH are necessary to neutralize the excess acid. From this information we wish to determine the normalities of both the HCl and NaOH. Of course, if one of these values is known, it will be a simple matter to calculate the normality of the other from the volume relation of the two solutions. In order that the calculations of this problem may be less complicated the following symbols will be used to represent the normalities and volumes of the two solutions: cc n
= volume of HC1 solution
= normality of HCI solution CC = volume of NaOH solution N = normality of NaOH solution
As 30 cc. of the NaOH solution was found by titration to be equivalent to 27 cc. of the HC1 solution, we may determine the normality of the HCl in terms of the normality of the NaOH by the following equation:
(4 ( d
=
(CC) (N)
(1)
Substituting the known values in this equation we have :
From which we may cancel out the normality of the NaOH, leaving: 50
(27) - (22.2) 30
(5)
Now by substituting in the formula for the calculation of the percentage purity of calcium carbonate by titration of a weighed sample with an acid of known concentration : (mil. equiv. of CaCOd (100) = Weight of sample
(cc) (n)
%
and assuming a purity of 100% for calcite we have:
The equation then contains only one unknown which when determined gives 0.45 as the normality of the NaOH. Substituting this value with the volume relations of the solutions in equation (1) we find the normality of the acid to be 0.5. EXPERIMENTAL PROCEDURE
Grind up a piece of pure calcite with a mortar and pestle until it has the fineness of flour. Weigh out three 0.75-g. samples into three 250-cc. Erlenmeyer flasks and add 50 cc. of the unknown acid to each. Warm until the calcite has completely dissolved. While the HCl and calcite are reacting, titrate 25 cc. of the acid with the NaOH solution, using methyl red
as the indicator. As soon as solution of the calcite is complete, titrate the excess acid with the NaOH. Calculate the normalities from the above formulas. COMPARISON WITH THE SILVER CHLORIDE METHOD
by the method described above and the silver chloride method. The normality of the HC1 by the AgCl method was determined by titrating i t against the HCl using the average of the AgCl normakes (0.1001). Phenolphthalein was used as the indicator.
The following table shows the comparison of the normality values of the same HC1 and NaOH solutions
LITERATURE CITED
(1) By Cdcilc
HCL 0.0994 0.0992 0.0992
NaOH
0.1141 0.1143 0.1138
BY AgCl
GRANDEAU, 2. anal. Chem., 2, 426 426 (1863).
(1863);
PINCUS, ibid.. 2,
HCI 0.1001 o,loo4
NaOH 0.1136 0,1136
(2) THIELE AND'R~CATER, Z. a ~ g Chem., ~ . 13,486 (1900). (3) MASSON, Chem. News, (Feb. 16, 1900). (4) KOLTHOPP AND FURMAN, 'Volumetric analysis, 11: Prac-
0.1000
0.1132
0.09~9 0.0996 0.1001 0.0996 0.0999
0.1134 0.1129 0.1131 0.1130 0.1131 0,1133 0.1141 0.1129 0.0012
tical -. principles," . - - - ~. John Wiley & Sons, Inc., New York City. 1929, p. Ya. (5) PoPom, "Quantitative analysis." P. Blakiston's Son & Co., Inc.. Philadelphia, 1927, p. 98. (6) FARNSWORTH, "The theory and technique of quantitative analysis," John Wiley & Sons,Inc., New York City. 1928, p. 94. (7) RNET,"A possible error in the calc-spar estimation of hydrochloric acid." Chem. News. 132, 309-10 (1926). (8) FORD, "Dana's manual of mineralogy," John Wiley & Sons. Inc., New Ynrk City. 1912.
o,lool o.lo04 0.0996 0.0008
