RESIDUE ON EVAPORATION.
TOTAL RESIDUE.[[16]]
Ignite and weigh a clean platinum dish, and measure into it 100 cc. of the thoroughly shaken sample. Evaporate to dryness on a water bath. Then heat the dish in an oven at 103° C. or 180° C. for one hour. Cool in a desiccator and weigh. The temperature of drying should be mentioned in the report. The increase in weight gives the total solids or residue on evaporation. If 100 cc. of the sample was taken this weight expressed in milligrams and multiplied by 10 is equal to parts per million of residue on evaporation. The residue from waters low in organic matter but relatively high in iron may be used, as a matter of convenience, for the determination of iron.
FIXED RESIDUE AND LOSS ON IGNITION.[[13]][[96]]
The residue from sewages and waters high in organic matter may be ignited to burn off the organic matter, which, with some volatile inorganic matter, constitutes the loss on ignition.
Procedure.—Ignite the residue in the platinum dish at a low red heat. If great accuracy is desired this should be done in an electric muffle furnace or in a radiator, which consists of a platinum or a nickel dish large enough to allow an air space of about half an inch between it and the dish within it, the inner dish being supported by a triangle of platinum wire laid on the bottom of the outer dish. A disc of platinum or nickel foil large enough to cover the outer dish is suspended over the inner dish to radiate the heat into it. The larger dish is heated to bright redness until the residue is white or nearly so. Allow the dish to cool, and moisten the residue with a few drops of distilled water. Dry the residue in the oven, cool in a desiccator, and weigh. The fixed residue on evaporation is the difference between this weight and the weight of the dish.
The loss on ignition is the difference between the total residue on evaporation and the fixed residue on evaporation.
If the odor and color on ignition of some residues give helpful clues to the character of the organic matter record them.
SUSPENDED MATTER.[[56]][[110]]
DETERMINATION WITH GOOCH CRUCIBLE.
Reagent.—Prepare a dilute cream of asbestos fibre which has been finely shredded, thoroughly ignited, treated with strong hydrochloric acid for at least 12 hours, and washed with distilled water till free from acid.
Procedure.—1. Prepare a mat of the asbestos fibre 1/16 inch thick in a Gooch crucible. Dry it in an oven at 103 or 180° C., cool and weigh. Filter 1,000 cc. of samples having a turbidity of 50 parts per million or less. If the turbidity is higher use sufficient water to obtain 50 to 100 mg. of suspended matter. Dry for one hour at 103 or 180° C., cool and weigh. Report the temperature at which the residue was dried. If 1,000 cc. is filtered the increase in weight expressed in milligrams is equal to parts per million of suspended matter.
DETERMINATION BY FILTRATION.
The difference between the total solids in filtered and unfiltered portions of a sample may be used as a basis for calculating suspended matter.
DETERMINATION OF VOLUME.
The determination of the volume[[9]][[69b]] of suspended matter in sewages has received considerable attention abroad. Imhoff recommends the use of conical glass vessels holding 1 liter with the lower portions graduated in cubic centimeters. Others recommend centrifuges with sediment tubes.
FIXED RESIDUE AND LOSS ON IGNITION.
Treat the total residue from a filtered sample in the same manner as described for the total residue, and obtain the loss on ignition due to dissolved matter, and by difference the loss on ignition due to suspended matter.
HARDNESS.[[94e]]
A water containing certain mineral constituents in solution, chiefly calcium and magnesium, which form insoluble compounds with soap, is said to be hard. Carbon dioxide in water increases the solubility of calcium and magnesium carbonates, forming bicarbonate. If carbon dioxide is removed from the water by boiling the bicarbonate is decomposed and calcium and magnesium are partly precipitated. The proportion of calcium or magnesium carbonate that a water can hold in solution depends on the concentration of carbon dioxide, which in turn depends on the temperature of the water and the proportion of carbon dioxide in the atmosphere with which the water has been in contact. Consequently, when the carbon dioxide is removed from the water by boiling or otherwise the carbonates of calcium and magnesium are partly, but not completely, precipitated, and the hardness of the water is thus diminished and the water is softened to the extent to which these substances are precipitated. The hardness thus removed is called temporary hardness. The hardness which still remains after boiling is due mainly to calcium and magnesium in equilibrium with sulfate, chloride, and nitrate, and residual carbonate, and it is called permanent hardness. Non-carbonate hardness is the hardness caused by sulfates, chlorides, and nitrates of calcium, magnesium, iron, and other metals that form insoluble soaps.
TOTAL HARDNESS BY CALCULATION.
The most accurate method of ascertaining total hardness is to compute it from the results of determinations of calcium and magnesium in the sample. (See methods, pp. [57]–58.) Iron and other metals must be included in the calculation if they are present in significant amounts. Total hardness as CaCO3 equals 2.5 Ca plus 4.1 Mg.
TOTAL HARDNESS BY SOAP METHOD.[[121b]]
The determination of hardness by the soap method roughly approximates the amount of calcium and magnesium in a water, though it actually measures the soap-consuming power of the water.
Reagents.—1. Standard calcium chloride solution. Dissolve 0.2 gram of pure calcite (calcium carbonate) in a little dilute hydrochloric acid, being careful to avoid loss of solution by spattering. Evaporate the solution to dryness several times with distilled water to expel excess of acid. Dissolve the residue in distilled water and dilute the solution to 1 liter. One cc. of this dilution is equivalent to 0.2 mg. of calcium carbonate.
2. Standard soap solution. Dissolve 100 grams of dry white Castile soap in 1 liter of 80 per cent alcohol, and allow this solution to stand several days before standardizing. Pure potassium oleate made from lead plaster and potassium carbonate may be used in place of Castile soap.
First method of standardization.—Dilute 20 cc. of the calcium chloride solution in a 250 cc. glass-stoppered bottle to 50 cc. with distilled water which has been recently boiled and cooled. Add soap solution from a burette, 0.2 or 0.3 cc. at a time, shaking the bottle vigorously after each addition until a lather remains unbroken for five minutes over the entire surface of the water while the bottle lies on its side. Then adjust the strength of the stock solution with 70 per cent alcohol so that the resulting diluted soap solution will give a permanent lather when 6.40 cc. of it is properly added to 20 cc. of standard calcium chloride solution diluted to 50 cc. Usually 75 to 100 cc. of the stock soap solution is required to make 1 liter of the standard soap solution. The quantity of calcium carbonate equivalent to each cubic centimeter of the standard soap solution consumed in the titration is indicated in Table 6.
| Table 6.—Total hardness in parts per million of CaCO3 for each tenth of a cubic centimeter of soap solution when 50 cc. of the sample is titrated. | ||||||||||
|---|---|---|---|---|---|---|---|---|---|---|
| Cubic centimeters of soap solution. | 0.0. | 0.1. | 0.2. | 0.3. | 0.4. | 0.5. | 0.6. | 0.7. | 0.8. | 0.9. |
| 0.0 | 0.0 | 1.6 | 3.2 | |||||||
| 1.0 | 4.8 | 6.3 | 7.9 | 9.5 | 11.1 | 12.7 | 14.3 | 15.6 | 16.9 | 18.2 |
| 2.0 | 19.5 | 20.8 | 22.1 | 23.4 | 24.7 | 26.0 | 27.3 | 28.6 | 29.9 | 31.2 |
| 3.0 | 32.5 | 33.8 | 35.1 | 36.4 | 37.7 | 38.0 | 40.3 | 41.6 | 42.9 | 44.3 |
| 4.0 | 45.7 | 47.1 | 48.6 | 50.0 | 51.4 | 52.9 | 54.3 | 55.7 | 57.1 | 58.6 |
| 5.0 | 60.0 | 61.4 | 62.9 | 64.3 | 65.7 | 67.1 | 68.6 | 70.0 | 71.4 | 72.9 |
| 6.0 | 74.3 | 75.7 | 77.1 | 78.6 | 80.0 | 81.4 | 82.9 | 84.3 | 85.7 | 87.1 |
| 7.0 | 88.6 | 90.0 | 91.4 | 92.9 | 94.3 | 95.7 | 97.1 | 98.6 | 100.0 | 101.5 |
This table does not provide for the use of so large volume of soap solution for a single determination as former ones because the end-point becomes somewhat obscured in the presence of magnesium if more than 7 cc. is used.
Second method of standardization.—Dilute 100 cc. of the stock soap solution to 1 liter with 70 per cent alcohol. This dilute solution should be of such strength that approximately 6.4 cc. of it will give a permanent lather when 20 cc. of standard calcium chloride solution diluted to 50 cc. with distilled water is titrated with it. Determine the amount of soap solution required to give a permanent lather with 50 cc. of distilled water and with 5, 10, 15, and 20 cc. of standard calcium chloride solution diluted to 50 cc. with distilled water. Finally plot on cross-section paper a curve showing the relation of various quantities of soap solution to corresponding quantities of standard calcium carbonate solution and therefore to parts per million of hardness.
Procedure.—Measure 50 cc. of the water into a 250 cc. bottle and add to it soap solution in small quantities in precisely the same manner as described under the standardization of the soap solution. From the number of cubic centimeters of soap solution used obtain from Table 6 or from the plotted curve the total hardness of the water in parts per million of calcium carbonate.
To avoid mistaking the false or magnesium end-point for the true one[[35]] when adding the soap solution to waters containing magnesium salts, read the burette after the titration is apparently finished, and add about 0.5 cc. more of soap solution. If the end-point was due to magnesium the lather will disappear. Soap solution must then be added until the true end-point is reached. Usually the false lather persists for less than five minutes.
If more than 7 cc. of soap solution is required for 50 cc. of the water take less of the sample and dilute it to 50 cc. with distilled water which has been recently boiled and cooled. This step reduces somewhat the disturbing influence of magnesium,[[107a]] which consumes more soap than an equivalent weight of calcium.
At best the soap method is not a precise test on account of the different relative amounts of calcium and magnesium in different waters. For hard waters, especially in connection with processes for purification and softening, it is advised that this method be not exclusively used. If the same water is frequently analyzed it may be of assistance to standardize the soap solution against a mixture of calcium and magnesium salts, the relative proportions of which approximate those found in the water.
The strength of the soap solution should be determined from time to time, to make sure that it has not materially changed. Record all results in parts per million of calcium carbonate.
One English degree of hardness, Clark’s scale, is equivalent to 1 grain per Imperial gallon of calcium carbonate. One French degree of hardness is equivalent to 1 part per 100,000 of calcium carbonate. One German degree of hardness is equivalent to 1 part per 100,000 of calcium oxide, and multiplied by 17.9 gives parts per million of calcium carbonate. The relations of these various scales are indicated in Table 7.
| Table 7.—Conversion table for hardness. | ||||
|---|---|---|---|---|
| Unit. | Equivalent. | |||
| Parts per million. | Clark degrees. | French degrees. | German degrees. | |
| One part per million | 1.00 | 0.07 | 0.10 | 0.056 |
| One Clark degree | 14.3 | 1.00 | 1.43 | .80 |
| One French degree | 10.0 | .70 | 1.00 | .56 |
| One German degree | 17.9 | 1.24 | 1.78 | 1.00 |
TOTAL HARDNESS BY SODA REAGENT METHOD.[[47]][[74]][[81]][[94d]]
Add standard sulfuric acid to 200 cc. of the sample until the alkalinity is neutralized. (See Procedure with methyl orange, p. [37].) Then apply the non-carbonate hardness method (pp. [34]–35). This method gives fairly satisfactory estimates of total hardness of hard waters.
TEMPORARY HARDNESS BY TITRATION WITH ACID.
Determine the alkalinity in presence of methyl orange (see p. [37]) in the original sample and also in the sample after boiling, cooling, restoring to the original volume with boiled distilled water, and filtering. The difference between the two, if any, is the temporary hardness. This is the most accurate method of determining the temporary hardness of ordinary waters. Iron bicarbonate is included as a part of the temporary hardness.
NON-CARBONATE HARDNESS BY SODA REAGENT METHOD.[[47]][[74]][[81]][[94d]]
The use of soda reagent does not avoid entirely the error due to solubility of the salts of calcium and magnesium; consequently, if much depends on the results, as in water softening, gravimetric determinations of the calcium and magnesium that remain in solution should be made and a correction should be applied for those amounts.
Reagent.—Prepare soda reagent from equal parts of sodium hydroxide and sodium carbonate. It should be approximately tenth normal.
Procedure.—Measure 200 cc. of the sample and 200 cc. of distilled water into 500 cc. Jena or similar glass Erlenmeyer flasks. Treat the contents of each flask in the following manner. Boil 15 minutes to expel free carbon dioxide. Add 25 cc. of soda reagent. Boil 10 minutes, cool, rinse into 200 cc. graduated flasks, and dilute to 200 cc. with boiled distilled water. Filter, rejecting the first 50 cc., and titrate 50 cc. of each filtrate with N/50 sulfuric acid in the presence of methyl orange or erythrosine indicator. The non-carbonate hardness in parts per million of calcium carbonate is equal to 20 times the difference between the number of cubic centimeters of sulfuric acid required for the soda reagent in distilled water and the number of cubic centimeters of N/50 sulfuric acid required for the soda reagent in the sample.
Water naturally containing bicarbonate and carbonate in excess of calcium and magnesium requires a larger amount of acid to neutralize the sample after it has been treated than is required to neutralize the volume of soda reagent originally added. (See p. [39].)
NON-CARBONATE HARDNESS BY SOAP METHOD.
Non-carbonate hardness may be calculated for waters which are soft or moderately hard in a fairly satisfactory manner by deducting the total alkalinity from the total hardness by the soap method (pp. [31]–34). For waters that are very hard, and particularly those that contain much magnesium, this method is not advised.
ALKALINITY.[[11]][[18]][[47]][[97]]
The alkalinity of a natural water represents its content of carbonate, bicarbonate, borate, silicate, phosphate, and hydroxide. Alkalinity is determined by neutralization with standard sulfuric acid or potassium bisulfate in the presence of phenolphthalein and either methyl orange, erythrosine, or lacmoid as indicators. Methyl orange may be used except in waters containing aluminium sulfate or iron sulfate. The relations between estimates in presence of these indicators and the carbonate, bicarbonate, and hydroxide radicles are indicated in Table 8. The alkalinity of carbonates in the presence of phenolphthalein is different from that in the presence of methyl orange, partly because of loss of carbon dioxide and partly because of defects in phenolphthalein as an indicator in such conditions.
| Table 8.—Relations between alkalinity to phenolphthalein and that to methyl orange, erythrosine, or lacmoid, in presence of bicarbonate, carbonate, and hydroxide. | |||
|---|---|---|---|
| Result of titration.[[C]] | Value of radicle expressed in terms of calcium carbonate. | ||
| Bicarbonate. | Carbonate. | Hydroxide. | |
| P = 0 | T | 0 | 0 |
| P < 1/2T | T − 2P | 2P | 0 |
| P = 1/2T | 0 | 2P | 0 |
| P > 1/2T | 0 | 2(T − P) | 2P − T |
| P = T | 0 | 0 | T |
[C]. T = Total alkalinity in presence of methyl orange, erythrosine, or lacmoid. P = Alkalinity in presence of phenolphthalein.
Reagents.—1. Sulfuric acid or potassium bisulfate. A N/50 solution.
2. Phenolphthalein indicator. Dissolve 5 grams of a good quality of phenolphthalein in 1 liter of 50 per cent alcohol. Neutralize with N/10 potassium hydroxide. The alcohol should be diluted with boiled distilled water.
3. Methyl orange indicator. Dissolve 0.5 gram of a good grade of methyl orange in 1 liter of distilled water. Keep the solution in the dark.
4. Lacmoid indicator. Dissolve 2.0 grams of lacmoid in 1 liter of 50 per cent alcohol. Dilute the alcohol with freshly boiled distilled water.
5. Erythrosine indicator. Dissolve 0.5 gram of erythrosine (the sodium salt) in 1 liter of freshly boiled distilled water.
PROCEDURE WITH PHENOLPHTHALEIN.
Add 4 drops of phenolphthalein indicator to 50 or 100 cc. of the sample in a white porcelain casserole or an Erlenmeyer flask over a white surface. If the solution becomes colored, hydroxide or normal carbonate is present. Add N/50 sulfuric acid from a burette until the coloration disappears.
The phenolphthalein alkalinity in parts per million of calcium carbonate is equal to the number of cubic centimeters of N/50 sulfuric acid used multiplied by 20 if 50 cc. of the sample was used, or by 10 if 100 cc. was used.
PROCEDURE WITH METHYL ORANGE.
Add 2 drops of methyl orange indicator to 50 or 100 cc. of the sample, or to the solution to which phenolphthalein has been added, in a white porcelain casserole or an Erlenmeyer flask over a white surface. If the solution becomes yellow, hydroxide, normal carbonate, or bicarbonate is present. Add N/50 sulfuric acid from a burette until the faintest pink coloration appears. The methyl orange alkalinity in parts per million of calcium carbonate is equal to the total number of cubic centimeters of N/50 sulfuric acid used multiplied by 20 if 50 cc. of the sample was used, or by 10 if 100 cc. was used.
PROCEDURE WITH LACMOID.
Add 4 drops of lacmoid indicator to 50 or 100 cc. of the sample in a porcelain casserole or an Erlenmeyer flask. Add N/50 sulfuric acid from a burette until within 1 or 2 cc. of the amount necessary for neutralization has been added. Heat the solution until bubbles of steam begin to break at the surface. Remove the dish from the source of heat and continue the titration until a drop of the acid striking the surface of the liquid and sinking to the bottom of the vessel produces no change in the uniform reddish or purple color of the solution. The calculation is the same as for phenolphthalein alkalinity.
PROCEDURE WITH ERYTHROSINE.
Add 5 cc. of neutral chloroform and 1 cc. of erythrosine indicator to 50 or 100 cc. of the sample in a 250 cc. clear glass-stoppered bottle. If the chloroform becomes rose colored on shaking, hydroxide, bicarbonate, or normal carbonate is present. Add N/50 sulfuric acid from a burette until the chloroform becomes colorless. A white surface behind the bottle facilitates detection of a trace of color as the end-point is approached. The calculation is the same as with phenolphthalein alkalinity.
BICARBONATE.
Bicarbonate is present if the alkalinity to phenolphthalein is less than one-half the alkalinity to methyl orange, erythrosine, or lacmoid. The alkalinity to methyl orange, erythrosine, or lacmoid is due entirely to bicarbonate if there is no phenolphthalein alkalinity. If there is phenolphthalein alkalinity the bicarbonate, in terms of calcium carbonate, is equal to the methyl orange, erythrosine, or lacmoid alkalinity minus twice the phenolphthalein alkalinity. Bicarbonate, carbon dioxide as bicarbonate, and half-bound carbon dioxide can be calculated as follows:
Bicarbonate (HCO3) = 1.22 times the bicarbonate expressed in terms of calcium carbonate.
Carbon dioxide (CO2) as bicarbonate = 0.88 times the bicarbonate expressed in terms of calcium carbonate.
Half-bound carbon dioxide (CO2) = 0.44 times the bicarbonate expressed in terms of calcium carbonate.
NORMAL CARBONATE.[[20]][[94]]
Normal carbonate is present if the alkalinity to phenolphthalein is greater than zero but less than the alkalinity to methyl orange, erythrosine, or lacmoid. If the phenolphthalein alkalinity is exactly equal to one-half the methyl orange, erythrosine, or lacmoid alkalinity the alkalinity is due entirely to normal carbonate. If the phenolphthalein alkalinity is less than one-half the methyl orange, erythrosine, or lacmoid alkalinity normal carbonate expressed in terms of calcium carbonate is equal to twice the phenolphthalein alkalinity. If the phenolphthalein alkalinity is greater than one-half the methyl orange, erythrosine, or lacmoid alkalinity the normal carbonate is equal to twice the difference between the methyl orange, erythrosine, or lacmoid alkalinity and the phenolphthalein alkalinity. The carbonate, carbon dioxide as carbonate, and bound carbon dioxide can be calculated as follows:
Carbonate (CO3) = 0.6 times the normal carbonate expressed in terms of calcium carbonate.
Carbon dioxide as carbonate (CO2) = 0.44 times the normal carbonate expressed in terms of calcium carbonate.
Bound carbon dioxide (CO2) is the sum of the carbon dioxide as carbonate and one-half that as bicarbonate.
HYDROXIDE.[[20]][[94]]
If hydroxide, or caustic alkalinity, is present the alkalinity to phenolphthalein is greater than one-half the alkalinity to methyl orange, erythrosine, or lacmoid; the alkalinity is due entirely to hydroxide if the phenolphthalein alkalinity is equal to the methyl orange, erythrosine, or lacmoid alkalinity. If the phenolphthalein alkalinity is more than half and less than all the methyl orange, erythrosine, or lacmoid alkalinity, hydroxide, expressed in terms of calcium carbonate, is equal to twice the phenolphthalein alkalinity minus the methyl orange, erythrosine, or lacmoid alkalinity.
ALKALI CARBONATES.
Waters which contain sodium or potassium carbonates or bicarbonates contain all of their calcium and magnesium as carbonates or bicarbonates. That is, they possess no non-carbonate hardness (sulfates, nitrates or chlorides of calcium and magnesium).
The most accurate method is to determine the total alkalinity by titration with N/50 sulfuric acid, using methyl orange, erythrosine, or lacmoid as an indicator; then determine the calcium and magnesium content; and subtract from the total alkalinity the computed alkalinity due to the calcium and magnesium expressed in terms of calcium carbonate. The remainder is the alkalinity due to carbonates and bicarbonates of sodium and potassium.
This determination may also be made by applying the method, for non-carbonate hardness with soda reagent (see p. [35]), and by noting the excess of acid required to neutralize the alkaline carbonates originally present.
With present information as to solubilities of the normal carbonates of calcium and magnesium, it is difficult in their presence to measure slight quantities of carbonates of sodium or potassium.
ACIDITY.[[24d]][[37]]
Waters may have an acid reaction because of the presence of free carbon dioxide, mineral acids, or some of their salts, especially those of iron and aluminium.
Reagents.—1. N/50 sodium carbonate. Dissolve 1.06 grams of anhydrous sodium carbonate in 1 liter of boiled distilled water that has been cooled in an atmosphere free from carbon dioxide. Preserve this solution in bottles of resistant glass protected from the air by tubes filled with soda-lime. One cc. is equivalent to 1 mg. of CaCO3.
2. N/22 sodium carbonate. Dissolve 2.41 grams of anhydrous sodium carbonate in 1 liter of boiled distilled water that has been cooled in an atmosphere free from carbon dioxide. Preserve this solution in bottles of resistant glass protected from the air by tubes filled with soda-lime. One cc. is equivalent to 1 mg. of CO2.
3. Phenolphthalein indicator (see p. [36]).
4. Methyl orange indicator (see p. [36]).
TOTAL ACIDITY.
Procedure.—Add 4 drops of phenolphthalein indicator to 50 or 100 cc. of the sample in a white porcelain casserole or an Erlenmeyer flask over a white surface. Add N/50 sodium carbonate until the solution turns pink. The total acidity in parts per million of calcium carbonate is equal to the number of cubic centimeters of N/50 sodium carbonate used multiplied by 20 if 50 cc. of the sample was used, or by 10 if 100 cc. was used.
FREE CARBON DIOXIDE.[[20]][[23]][[61]][[87]][[88]][[94a]][[118]]
Carbon dioxide may exist in water in three forms—free carbon dioxide, bicarbonate (pp. [37]–38), and carbonate (p. [38]). One-half the carbon dioxide as bicarbonate is known as the half-bound carbon dioxide. The carbon dioxide as carbonate plus one-half that as bicarbonate is known as the bound carbon dioxide.
Procedure.—Pour 100 cc. of the sample into a tall narrow vessel, preferably a 100 cc. Nessler tube. Add 10 drops of phenolphthalein indicator, and titrate rapidly with N/22 sodium carbonate, stirring gently, until a faint but permanent pink color is produced. The free carbon dioxide (CO2) in parts per million is equal to 10 times the number of cubic centimeters of N/22 sodium carbonate used.
Because of the ease with which free carbon dioxide escapes from water, particularly when the gas is present in large amount, a special sample should be collected for this determination, which should preferably be made at the time of collection. If the analysis cannot be made at the time of collection approximate results with water not too high in free carbon dioxide may be obtained on samples collected in bottles completely filled so as to leave no air space under the stopper. Bottled samples should be kept, until tested, at a temperature lower than that of the water when collected. If mineral acids or certain salts are present correction must be made. At best, the results of the titration are uncertain because the proper end-point for correct results differs in color with different types of water.
FREE MINERAL ACIDS.
Procedure.—Add 2 drops of methyl orange indicator to 50 or 100 cc. of the sample in a white porcelain casserole or an Erlenmeyer flask over a white surface. Add N/50 sodium carbonate from a burette until the pink coloration of the solution disappears. The acidity due to free mineral acids, expressed in terms of calcium carbonate, is equal to the number of cubic centimeters of N/50 sodium carbonate used multiplied by 20 if 50 cc. of the sample was used, or by 10 if 100 cc. was used.
MINERAL ACIDS AND SULFATES OF IRON AND ALUMINIUM.[[24d]][[37]]
Procedure.—Modify the method for free mineral acids by titrating the water at boiling temperature in the presence of phenolphthalein indicator. The acidity due to free mineral acids and sulfates of iron and aluminium, expressed in terms of calcium carbonate, is equal to the number of cubic centimeters of N/50 sodium carbonate used multiplied by 20 if 50 cc. of the sample was used, or by 10 if 100 cc. was used.
The acidity due to sulfates of iron and aluminium is equal to the acidity due to mineral acids and sulfates minus the acidity due to mineral acids. The acidity due to ferrous and ferric sulfate can be calculated from the determined amount of these salts (pp. [43]–48). The acidity due to aluminium sulfate is equal to the acidity due to total acid sulfates minus that due to iron sulfates.
Acidity shall be reported in parts per million of calcium carbonate (CaCO3). Sulfate (SO4) equals parts per million of calcium carbonate multiplied by 0.96.
Carbon dioxide (CO2) equals parts per million of calcium carbonate multiplied by 0.44.
CHLORIDE.[[16]]
Chloride in water and sewage has its origin in common salt, from mineral deposits in the earth, from ocean vapors carried inland by the wind, or from polluting materials like sewage and trade wastes, which contain the salt used in the household and in manufacturing. Comparison of the chloride content of a water with that of other waters in the vicinity known to be unpolluted frequently affords useful information as to its sanitary quality. If, however, the chloride normally exceeds 20 parts per million because of chloride-bearing mineral deposits the chloride content of a water has little sanitary significance.
Reagents.—1. Standard sodium chloride solution. Dissolve 16.48 grams of pure fused sodium chloride in 1 liter of distilled water. Dilute 100 cc. of this stock solution to 1 liter in order to obtain a standard solution each cubic centimeter of which contains 0.001 gram of chloride.
2. Standard silver nitrate solution. Dissolve about 2.40 grams of silver nitrate crystals in 1 liter of distilled water. Standardize this with the standard salt solution, and adjust, correcting for volume (see p. [43]), so that 1 cc. will be exactly equivalent to 0.0005 gram of chloride.
3. Potassium chromate indicator. Dissolve 50 grams of neutral potassium chromate in a little distilled water. Add enough silver nitrate to produce a slight red precipitate. Filter and dilute the filtrate to 1 liter with distilled water.
4. Aluminium hydroxide. Electrolyze ammonia-free water, using aluminium electrodes. Wash the precipitate until it is free from chloride, ammonia, and nitrite. Or dissolve 125 grams of potassium or ammonium alum in 1 liter of distilled water. Precipitate the aluminium by adding cautiously ammonium hydroxide. Wash the precipitate in a large jar by successive additions and decantations of distilled water until free from chloride, nitrite, and ammonia.
Procedure.—Add 1 cc. of potassium chromate indicator to 50 cc. of the sample in a 6–inch white porcelain evaporating dish or a 150 cc. Erlenmeyer flask over a white surface. Titrate with the silver nitrate solution under similar conditions of volume, light, and temperature as were used in standardizing the silver nitrate until a faint reddish coloration is perceptible. The detection of the end-point is facilitated by comparison of the contents of the porcelain dish with those of another dish containing the same quantity of potassium chromate indicator in 50 cc. of distilled water. Some analysts prefer to make the titration in a dark-room provided with a yellow light. The end-point is very sharp by electric light and also by daylight with photographic yellow glass. The titration may be made in Nessler tubes[[68a]] if the solutions are standardized under similar conditions.
If the amount of chloride is very high use 25 cc., or even a smaller quantity, dilute the volume taken to 50 cc. with distilled water. If the amount of chloride is very low concentrate 250 cc. of the sample to 50 cc. by evaporation. Rotate the liquid to make sure that no residue remains undissolved on the walls of the dish, and, if necessary, use a rubber-tipped glass rod to assist in this operation.
Chloride is determined by some observers by extracting with hot distilled water the residue in the platinum dish in the determination of the residue on evaporation and proceeding as just described. This is permissible if a little sodium carbonate is added before evaporation to prevent loss of chloride through decomposition of magnesium chloride in the residue.
If the sample has a color greater than 30 it should be decolorized by shaking it thoroughly with washed aluminium hydroxide (3 cc. to 500 cc. of the sample) and allowing the precipitate to settle. Make the determination on a portion of the clarified sample, filtered if necessary. If the sample is acid, neutralize it with sodium carbonate; if hydroxide is present, add dilute sulfuric acid until the cold liquid will just discharge the color of phenolphthalein. If the presence of sulfide and sulfocyanate renders it necessary, make proper corrections[[24c]][[100b]] or modifications in treatment.
Make correction for the error due to variations in the volume of the liquid and precipitate by means of the formula[[39]] X = 0.003V + 0.02, in which X = the correction in cubic centimeters of silver nitrate solution and V = cubic centimeters of liquid at the end of the titration. If 50 cc. of the sample is titrated chloride (Cl) in parts per million is equal to the number of cubic centimeters of silver nitrate solution multiplied by 10. The correction to be applied is 0.2 cc. unless unusual accuracy is required.
IRON.[[94b]][[98]]
Iron occurs in natural waters in both ferrous and ferric condition, depending on the source of the sample. In ground waters the iron is usually in an unoxidized and soluble condition, sometimes combined with carbonic or sulfuric acid, and also in combination with organic matter. Many waters, especially those that have been exposed to the air, contain the iron in the form of a colloidal hydroxide. Silt-bearing waters often contain much iron in suspension, usually in an oxidized form. Sewages and sewage effluents, particularly those receiving manufacturing wastes, contain various forms of iron of different degrees of solubility, oxidation, and coagulation.
TOTAL IRON.[[59]][[63b]]
COLORIMETRIC METHOD.
Reagents.—1. Standard iron solution. Dissolve 0.7 gram of crystallized ferrous ammonium sulfate in 50 cc. of distilled water to which 20 cc. of dilute sulfuric acid has been added. Warm the solution slightly and add potassium permanganate until the iron is completely oxidized. Dilute the solution to 1 liter. One cc. of the standard solution equals 0.1 mg. Fe.
2. Potassium sulfocyanide solution. Dissolve 20 grams of the salt in 1 liter of distilled water.
3. Dilute hydrochloric acid. One volume of acid (Sp. gr. 1.2) and one volume of distilled water. This shall be free from nitric acid.
4. N/5 potassium permanganate. Dissolve 6.30 grams of the salt in distilled water and dilute to 1 liter.
5. Hydrochloric acid. Concentrated, free from iron.
6. Nitric acid. Concentrated, free from iron.
7. Nitric acid. 5N, free from iron.
First procedure.—Evaporate 100 cc. of the water to dryness, or use the residue left after the determination of residue on evaporation (p. [29]). Ignite the residue at a low red heat taking care not to heat it hot enough to make the iron difficultly soluble. Cool the dish and add 5 cc. of concentrated hydrochloric acid. Moisten the inner surface of the dish. Warm the solution for two or three minutes, and again moisten the inner surface of the dish by permitting the hot acid to flow over it. Wash the hot solution from the dish into a 50 cc. Nessler tube, filtering if necessary through paper that has been washed with hot water. Dilute to 50 cc., and add 3 drops of potassium permanganate solution. Add 5 cc. of potassium sulfocyanide solution, mix, and compare with standards.
If it is not convenient to use the residue on evaporation and if the sample is relatively free from organic matter, boil 50 cc. of the sample with 5 cc. of 5N nitric acid for five minutes. Add a few drops of permanganate and 5 cc. of potassium sulfocyanide and compare with standards, using nitric acid in place of hydrochloric acid in the standards. This method is excellent for ground waters. The permanganate and acid liberate chlorine in water high in chloride, and produce a permanent yellow color which interferes with the determination, unless the sample is first diluted to 50 cc. An excess of permanganate, reacting with hydrochloric acid, causes similar trouble. The amounts of hydrochloric acid, 5 cc., and of sulfocyanide, 5 cc., should be approximately measured because more acid lightens the color whereas more sulfocyanide deepens it. This is especially important if permanent standards are used.
Second procedure.—For surface waters containing small amounts of organic matter, the method of Klut[[59]] is recommended. Samples containing small amounts of iron should be concentrated, if possible, until at least 0.5 mg. of iron is present in the volume tested. Boil the sample in a beaker with 2 to 3 cc. of concentrated nitric acid free from iron, adding permanganate if necessary to destroy the organic matter. To the hot liquid add ammonia in slight excess and warm until the smell of ammonia is hardly discernible. Filter and wash with water at 70° to 80° C. containing a little ammonia. Dissolve the iron in the beaker and on the filter paper in 5 cc. of concentrated hydrochloric acid, and wash with hot water until the iron is all dissolved, collecting the filtrate in a 50 cc. Nessler tube. Dilute to 50 cc. Add potassium sulfocyanide and determine the iron by comparison with standards.
Comparison with iron standards.
First procedure.—Prepare standards containing amounts of standard iron solution ranging from 0.05 to 4 cc. according to the quantity of iron in the sample. Dilute these amounts with water to about 40 cc. Add 5 cc. of dilute hydrochloric acid and 3 drops of potassium permanganate to each tube and dilute to 50 cc. Add 5 cc. of the potassium sulfocyanide to each of the standard solutions at the same time that it is added to the samples of water under examination, and compare immediately after mixing. If 100 cc. of the sample is used the iron in parts per million is equal to the number of cubic centimeters of the standard iron solution in the standard that the sample matches.
Second procedure.—For a small number of determinations it is more convenient to run the standard iron solution into a Nessler tube containing the acid, distilled water, and potassium sulfocyanide until the color matches that of the sample tested. When determining iron in three or four samples the colors may be matched in the order of their intensity and the volumes of standard iron solution required for each tube may be read from the burette.
Comparison with permanent standards.
Reagents.—1. Platinum solution. Dissolve 2 grams of potassium platinic chloride (PtCl4.2KCl) in distilled water, add 100 cc. of concentrated hydrochloric acid, and dilute to 1 liter with distilled water.
2. Cobalt solution. Dissolve 24 grams of dry cobaltous chloride crystals (CoCl2.6H2O) in a small amount of distilled water, add 100 cc. of strong hydrochloric acid, and dilute to 1 liter with distilled water.
Procedure.—Prepare a series of permanent standards by diluting to 50 cc. with distilled water the amounts of platinum and cobalt solutions, in 50 cc. Nessler tubes, indicated in Table 9. Compare the sample with these standards, and calculate the parts per million of iron.
| Table 9.—Preparation of permanent standards for the determination of iron. | ||
|---|---|---|
| Value in standard iron solution. | Platinum solution. | Cobalt solution. |
| cc. | cc. | cc. |
| 0.0 | 0 | 0.0 |
| .1 | 2 | 1.0 |
| .3 | 6 | 3.0 |
| .5 | 10 | 5.0 |
| .7 | 14 | 7.5 |
| 1.0 | 20 | 11.0 |
| 1.5 | 28 | 17.0 |
| 2.0 | 35 | 24.0 |
| 2.5 | 39 | 32.0 |
| 3.0 | 39 | 43.0 |
| 3.5 | 40 | 55.0 |
VOLUMETRIC METHOD.[[24f]]
Some samples of sewage and water mixed with trade wastes and mine drainage contain so much iron that it is preferable to use the volumetric method described on page [57] for the determination of both total and dissolved iron, rather than to work with quantities small enough to permit application of the colorimetric methods just described. If iron is present in large quantities in suspension, as in some sewages and septic tank effluents, it may be filtered off and the residue washed, ignited, and fused with potassium and sodium carbonate. The fusion is then extracted with hydrochloric acid and the iron determined as on page [57].
Samples containing much organic matter should be evaporated to dryness with 0.5 cc. of concentrated sulfuric acid and the residue then ignited before estimation of iron.
DISSOLVED IRON.
Determine, by the method described for total iron, the iron in the sample after filtration. Iron may precipitate from some samples during filtration.
SUSPENDED IRON.
The suspended iron is the difference between total iron in the unfiltered sample and dissolved iron in the filtered sample.
FERROUS IRON.[[24e]]
Determine the total ferrous iron in an unfiltered sample and the dissolved ferrous iron in a filtered sample.
Reagents.—1. Standard iron solution. Dissolve 0.7 gram of crystallized ferrous ammonium sulfate in a large volume of freshly boiled distilled water to which 10 cc. of dilute sulfuric acid has been added and dilute to 1 liter. This solution should be freshly prepared when needed. One cc. of this standard solution contains 0.1 mg. of Fe.
2. Potassium ferricyanide solution. Dissolve 5 grams of the salt in 1 liter of distilled water. Use a freshly prepared solution.
3. Dilute sulfuric acid. Dilute 1 part of sulfuric acid, specific gravity 1.84, with 5 parts of distilled water.
Procedure.—Add 10 cc. of dilute sulfuric acid to 50 cc. of the sample, remove the suspended matter by filtration if necessary, and add 15 cc. of potassium ferricyanide solution. Dilute the solution to 100 cc. with distilled water. Compare the color developed in the sample with that in standards made at the same time from the ferrous iron solution. Place in 100 cc. Nessler tubes, in the following order, 75 cc. of distilled water, 10 cc. of dilute sulfuric acid, and 15 cc. of potassium ferricyanide solution, and mix well the contents of each tube. Prepare as many tubes in this way as are needed. Add various quantities of standard ferrous iron solution to several tubes, mix well, and compare the resulting colors with the samples immediately.
FERRIC IRON.
The amount of ferric iron in solution and suspension is equal to the difference between the total iron and the ferrous iron obtained by the methods described.