An Introductory Course of Quantitative Chemical Analysis / With Explanatory Notes

L DISC

d. Just as an acid solution was the principal reagent in alkalimetry, and the alkali solution used only to make certain of the end-point, the solution of the oxidizing agent is the principal reagent for the

s are potassium bichromate, potassium permangenate, potassium

ate), oxalic acid, sodium thiosulphate, stannous chloride, arsenious acid, and potassium cyanide. Other reducing agents, as su

s salts; potassium permanganate and ferrous salts; potassium permanganate and oxalic aci

SS FOR THE DETER

old solution, by the addition of potassium bichromate, provided sufficient acid

ally preferred, since it is by far the best solvent for iron and its co

7} + 14HCl -> 6FeCl_{3} +

OF OXIDIZING OR

iter which is equivalent in oxidizing power to 8 grams of oxygen; a normal reducing solution must be equivalent in reducing power to 1 gram of hydrogen. In order to determine what the amount per liter will be it is necessary to know how the reagents

}SO_{4} -> 3Fe_{2}(SO_{4}){3} + K{2}

areful inspection shows that there are three less oxygen atoms associated with chromium atoms on the right-hand side of the equation than on the left-hand, but there are three more oxygen atoms associated with iron atoms on the right than on the left. In other words, a molecule of potassium bichromate has given up three atoms of

t six molecules of FeO.SO_{3} were required to react with the t

6(FeO.SO_{3}) + 3O ->

ular weight in grams of ferrous sulphate (151.9) is equivalent to 1 gram of hydrogen. Since the ferrous sulphate crystalline f

ION OF S

ate Stren

gh to expel any oxygen. The pure salt thus prepared, may be weighed out directly, dissolved, and the solution diluted in a graduated flask to a definite volume. In this case no standardization is made, as the

lve the bichromate in distilled water, transfer the solution to a liter bottle, a

mmonium sulphate (FeSO_{4}.(NH_{4}){2}SO{4}.6H_{2}O) and dissolve in distilled water containing 5 cc. of concentrated sulphuric acid. Transfe

TOR SO

tion. Drops of the indicator solution are placed upon a glazed porcelain tile, or upon white cardboard which has been coated with paraffin to render it waterproof, and drops of the titrated solution are transferred to the indicator on the end of a stirring rod. When the oxidation is nearly completed only very small amounts of the ferrous compounds remain unoxidized and the reaction with the indicator is no longer instantane

, as needed each day, by dissolving a crystal of potassium ferricyanide about the size of a pin's head in 25 cc. of distilled water.

ified by adding to its solution a little bro

XIDIZING AND RE

al procedure with respect to cleaning and rinsing already prescribed. The b

0.5 cc. of the bichromate solution may be added before testing again. The stirring rod which has touched the indicator should be dipped in distilled water before returning it to the iron solution. As soon as the blue appears to be less intense, add the bichromate solution in small portions, finally a single drop at a time, until the point is reached at which no blue color appears after the lapse of thirty seconds from the time of mixing solution and indicator. At the close of the titration a large drop of the iron solution should be taken

rette readings, and, if need be, for the temperature of solutions, calcul

size may properly be taken for the final tests. The stirring rod should be washed to prevent transfer of indicator to the main solution. This cautious removal of solution does not seriously affect the accuracy of the determination, as it will be noted that the volume of the titrated solution is about 200 cc. and the portions r

t of other methods; and if a ferrous solution is at hand, the titration need consu

OF POTASSIUM BIC

on of a

and it must be permanent against oxidation in the air, at least for considerable periods. Such standards may take the form of pure crystalline salts, such as ferrous ammonium sulphate, or may be in the form of iron

t which answers the purpose well, and its iron content may be determined for each lot purchased by a number of gravimetric determinations. It may best be preserved in ja

ps even better to determine by gravimetric methods once for all the iron content of a large commercial sample

ARDIZ

ly for rust. It should be handled and wiped with filter paper (not touched by the fingers), should be w

e flasks covered (Note 1), then wash the sides of the flasks and the watch-glass with a little water and add stannous chloride solution to the hot liquid !from a dropper! until the solution is colorless, but avoid more than a drop or two in excess (Note 2). Dilute with 150 cc. of water and cool !completely!. When cold

to oxidize the iron actually know to be present in the wire,

he results are concordant within a

tion to insure the presence of at least sufficient free aci

f compounds of hydrogen and carbon which are formed from the small amount of ca

some oxidation may have occurred from the oxygen of the air during solution. It is also evident that any exces

ly, the completion of the reaction is most easily noted, and the excess of the reagent is most readily removed. The latter object is accomplished by oxidation to s

l_{2} + SnCl_{4} SnCl_{2} +

chloride is

e should not be present in great excess, otherwise a secondary reaction takes

2HgCl -> Sn

precipitate; and, since potassium bichromate oxidizes this mercury slowly, so

mercuric chloride to permit the complete deposition of mercurous chloride. It should then b

ON OF IRON

solvent action has ceased. If the residue is white, or known to be free from iron, it may be neglected and need not be removed by filtration. If a dark residue remains, collect it on a filter, wash free from hydrochloric acid, and ignite the filter in a platinum crucible (Note 3). Mix the ash with five times its weight of sodium carbonate and heat to fusion; cool, and disintegrate the fused mass with boiling wa

, and the known weight of the sample, calculat

at beds and contains more or less organic matter which, if brought into solution, would be acted upon by the potassium bichromate. This organic matte

-say 5 grams-and to take one tenth of it for titration. The sampl

tion of ferric chloride for any length of time, since the platinum is attacked and dissolved, and the platinic chloride is later reduced by the stannous chloride, and in the reduced

l be much larger than that added to the solution of iron wire, in which the iron was mainly already i

OF CHROMIUM IN

balances, 5 grams of dry sodium peroxide for each portion, and pour about three quarters of the peroxide upon the ore. Mix ore and flux by thorough stirring with a dry glass rod. Then cover the mixture with the remainder of the peroxide. Place the crucible on a triangle and ra

evolution of gas ceases, rinse off and remove the crucible; then heat the solution !while still alkaline! to boiling for fifteen minutes. Allow the liquid to cool for a few minutes; then acidify with dilute sulphuric acid (1:5), adding 10 cc. in excess of the

lutions, and their relations to normal solutions,

ombination of FeO and Cr_{2}O_{3}. It must be reduced to a state

bright redness and plunged into cold water. In this process oily matter is burned off and ad

tice in the rough weighing of chemicals. If paper to which the peroxide is adhering is exposed to moist air it is like

ible, as this will raise the temperature to a point at which the c

oxide is dissolved by the flux and oxidized to chromic anhydride (CrO_{3}) which co

exercised to keep the temperature as low as possible. Preference is here given to iron crucibles, because the resulting ferric hydroxide is more readily brought into solution than the n

cold before it is placed in water, otherwise scales of magnetic iron oxide may separate from the

decomposed with the evolution of oxygen. The subsequent boiling insures the complete decomposition of the peroxide. Unless this is complete,

s the sodium chromate to bichromate, which behaves

omite were pure FeO.Cr_{2}O_{3} may be weighed out and dissolved in sulphuric acid; after reduction of all the iron by stannous chloride and the addition of

CESS FOR THE DET

etely to the ferric condition, while in hot acid solution it also enters into a definit

ns involved

S_{4} -> 5Fe_{2}(SO_{4}){3} + K

_{4} +3H_{2}SO_{4} -> K_{2}SO_{4}

nd manganese (the latter only in neutral solution), by which these metals are changed from a lower to a higher state of oxidation; and it also reacts with sulphurous acid, sulphureted hydrogen, nitrous acid, f

of the reaction), but slow in cold, dilute solutions. However, the greater solubility of iron compounds in hydrochloric acid makes it desirable to use this acid as a solvent, and experiments made with this end in

of sulphuric acid until the latter fumes. This procedure is somewhat more time-consuming, but t

idation of the iron and the reduction of the permanganate is colorless, the latter becomes its own

OF A STAND

ate Stren

of KMnO_{4} and 10 molecules of FeSO_{4} on the left-hand side, and 2 molecules of MnSO_{4} and 5 molecules of Fe_{2}(SO_{4})_{5} on the right-han

{3}) -> K_{2}O.SO_{3} + 2(MnO.S

the ferrous compound. Since 8 grams of oxygen is the basis of normal oxidizing solutions and 80 grams of oxygen are supplied by 3

y tenth-normal solution of the rea

Heat slowly and when the crystals have dissolved, boil the solution for 10-15 minutes. Cover the solution with a watch-glass; allow it to stand unti

reduction of the permanganate is less complete, and, under these conditions, two gram-molecular weights of KMnO_{4} will furnish o

ed from an alkaline permanganate solution. The solutions should be protected from light and heat as far as possible, since both induce decomposition with a deposition of manganese dioxide, and it has been shown that decomposition proceeds with considerable rapidity, with the evoluti

RMANGANATE AND F

be used in burettes with rubber tips, as a reduction takes place upon contact with the rubber. The solution has so deep a color that the lower line of the meniscus cannot be detected; readings must therefore be made from the upper edge. Run out into

F A POTASSIUM PER

on of a

dize them against substances of known value. Those in most common use are iron wire, ferrous ammonium sulphate, sodium oxalate, oxalic acid, and some other derivatives of oxalic acid. With the exception of sodium oxalate, th

tho

Stan

as standards differ from those applicable in connection with oxalate standard

sence of hydrochloric acid. Since the excess of both the gaseous reducing agents can only be expelled by boiling, with consequent uncertainty regarding both the removal of the excess and the reoxidation of the iron, zinc or stannous chlorides are the most satisfactory agents. For prom

ration:

layer of asbestos, such as is used for Gooch filters, 1 mm. thick. The tube is then filled with the amalgamated zinc (Note 1) to within 50 mm. of the top, and on the zinc is placed a plug of glass wool. If the top of the tube is not already shaped like the mouth of a thistle-tube (B), a 60 mm. funnel is fitted

be amalgamated by stirring or shaking it in a mixture of 25 cc. of normal mercuric chloride solution, 25 cc. of hydrochloric acid (sp. gr. 1.12) and 250 cc. of water for two minutes. The solution should then be poured of

ARDIZ

llows: Connect the vacuum bottle with the suction pump and pour into the funnel at the top warm, dilute sulphuric acid, prepared by adding 5 cc. of concentrated sulphuric acid to 100 cc. of distilled water. See that the stopcock (C) is open far enough to allow the acid to run through slowly. Continue to pour in acid until 200 cc. have passed through, then close the stopcock !while a small qu

cc. of the warm acid and finally with 75 cc. of distilled water, leaving the funnel partially filled. Remove the filter bottle and cool the solution quickly under the water tap (Note 4), avoiding unnecessary exposure to the oxygen of th

, calculate the relation to the normal of the permanganate solution. The

artially filled with water the reductor is ready for subsequent use after a very li

allowance must be made for the iron in the zinc. !Great care! must be used to prevent the access of air to the reductor after it has be

nt may be considered to be !nascent! hydrogen, and it must be borne in mind that the visible b

n the wash-water is added. It is well to allow the iron solution to run nearly, but not entirely, out of the funnel before the wash-water

he zinc must be present in excess of the quantity necessary to reduce the iron and is finally complete

iron begins, and it is of the first importance that the volume of acid and of wash-water should be measu

ation. If the solution turns brown, it is an evidence of insufficient acid, and more should be immediately added. The results are likely to be less accurate in this case,

t to an exactly 0.1 N solution from the data here obtained. The percentag

tho

te Sta

t the liquid, if necessary, until near its boiling point (70-90°C.) and run in the standard permanganate solution drop by drop from a burette, stirring constantly until an end-point is reached (Note 2). Make a blank test with 20 cc. of manganous sulphate

relation of the permanganate solution

2}SO_{4} -> 5Na_{2}SO_{4} + K_{2}SO_

grams of MnSO_{4} in 200 cubic centimeters of water and adding 40 cc. of con

nates takes place quantitatively only in hot acid s

ON OF IRON

tho

es. Many iron ores contain titanium, and this element among others does interfere with the determination of iron by the process described. If, however, the solutions of such ores are treated with

rcelain crucibles. Roast the ore at dull redness for ten minutes (Note 1), allow the crucibles to cool, an

r the flame of the burner, holding the casserole in the hand and rotating it slowly to hasten evaporation and prevent spattering, until the heavy white fumes of sulphuric anhydride are freely evolved (Note 3). Cool the casseroles, add 100 cc. of water (measured), and boil gently until the ferric sulphate is dissolved; pour the warm solution thro

nate solution used, calculate the per

ic acid would subsequently char the carbonaceous matter, certain nitrogenous bodies are

volatilize sulphuric acid. Solutions may, therefore, be left t

ic sulphate separates at this point, since there is no water to hold it in solution and care is required to prevent bumping. The fe

n-Reinhard

hod

necessary, and reduce the iron by the addition of stannous chloride, followed by mercuric chloride, as described for the bichromate process on page 56. Dilute the solution to about 40

eady obtained calculate the perce

although somewhat fugitive end-point in the presence of manganous sulphate and phosphoric acid. The explanation of the part played by these reagents is somewhat obscure

the iron by stannous chloride, too large an amount should be avoided in order to

THE OXIDIZING P

CT OXI

of MnO_{2} in the sample. This percentage is determined by an indirect method, in which the manganese dioxide is reduced and dissolved by an excess o

pered weighing-tube, and weigh out two portions of about 0.5 gram into beakers (400-500 cc.) Read Note 2, and then calculate in each ca

H_{2}O) + H_{2}SO_{4} -> Mn

e notebook. Pour into each beaker 25 cc. of water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the beaker and its contents gently until the evolution of

below boiling, add 15 cc. of a manganese sulphate solution and while hot, titrate f

the pyrolusite; subtract this from the total quantity of acid used, and calculate the weight of mangan

pon the fineness of the powdered mineral. If properly groun

s increasing the accuracy of the process. On the other hand, the excess of oxalic acid should not be so large as to react with more of the permanganate solution than is contained in a 50 cc. burette. If the pyrolusite under examination is known to be of high grade, say 80 per cent pure, or above the calculation of the oxalic acid needed may be based upon an assumption that the mineral is all MnO_{2}. If the quality o

osed by heat alone if crystallization should occur on the sides of the vessel. Strong sulphuric acid also dec

phate, or iron wire may be substituted for the

+ 2H_{2}SO_{4} -> Fe_

h potassium bichromate, if desired. Care is required to prevent th

e determined by other volumetric processes, one

-> MnCl_{2} + C

2KI -> I_

{2}O_{3} -> Na_{2

de. The liberated iodine is then determined by titration with sodium thiosulphate, a

IME

thiosulphate solution may be used in both acid and neutral solutions to measure free iodine and the latter may, in turn, serve as a measure of any substan

n which iodometric processes

{2}O_{3} -> 2 NaI

Na_{2}S_{4}O_{6}, called sodium tetrathionate, is quantitatively exact, and differs in that respe

OF IODINE AND S

ough, therefore, the iodine contains no oxygen, the two atoms of iodine have, in effect, brought about the addition of one oxygen atoms to the sulphur atoms. That is the same thing as saying that 253.84 grams of iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8 grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97 grams of iodine should be contai

OF STANDAR

ate Stren

ith 18 grams of potassium iodide and triturate with small portions of water until all is dissolv

able quantity of the solution prepared by a laboratory attendant

een previously boiled and cooled, and dilute to 1000 cc., also with boiled

res the iodine tends to volatilize from solution. They should, therefore, be protected from light and heat. Iodine solutions are not

by recrystallization. The carbon dioxide absorbed from the air by distilled water decomposes the salt, with the separa

TOR SO

tainable which serves well, and a solution of 0.5 gram of this starch in 25 cc. of boiling w

ste, pour 150 cc. of !boiling! water over it, warm for a moment on the hot plate, and put it aside to settle. Decant

d particles of the starch, if not removed by filtration, become so colored by

. It is regarded as a "solid solution" of iodine in starch. Although it is unstable, and easily destroyed by heat, it serves as an indicator for the pr

ODINE AND THIOS

tion into a beaker, dilute with 150 cc. of water, add 1 cc. to 2 cc. of the soluble starch solution, and titrate with the iodine to the appearance of the blue

ZATION OF

as prepared. It may be purified by sublimation after mixing it with a little potassium iodide, which reacts with the iodine chloride, forming potassium chloride and setting free the iodine. The sublimed iodine is then dried by placing it in a closed container over concentrated sulphuric ac

ndardized against arsenious oxide (As_{4}O_{6}). This substanc

inable in pure condition or may be easily purified by re-crystallization. Copper wire of high grade is sufficiently pure to serve as a standard.

involved are

2}SO_{4} -> KBr + 3I_{2}

{2} + 2NO + 4H{2}O, 2Cu(NO_{3}){2

ium thiosulphate solution are here described, and one

tho

lphuric acid with 5 volumes of water), allow the solution to stand for three minutes, and dilute to 150 cc. (Note 2). Run in thiosulphate solution from a burette until the color of the liberated iodine is nearly destroyed, and then add 1 cc. or 2 cc. of starch solution, titrate to the disappearance of the iodo-st

hiosulphate solution to a normal solution, and subseque

he amount of thiosulphate which reacts with the iodine thus liberated by making a "blank test" with the iodide and acid alone. As the iodate is not always uniformly distributed throughout the iodide, it is better to make up a sufficient volume of a

ctory in concentrated solutions of the alkali salts, notably t

tho

and add strong ammonia (sp. gr. 0.90) drop by drop until a deep blue color indicates the presence of an excess. Boil the solution until the deep blue is replaced by a light bluish green, or a brown stain appears on the sides of the flask (Note 2). Add 10 cc. of strong acetic acid (sp. gr. 1.04), cool under the water tap

test" of the iodide, calculate the relation

or its use as a conductor of electricity are such that the impurities constitute

ess must be removed by boiling, which is tedious. If too much ammonia is present when acetic acid is added

the prompt liberation of iodine. While a large excess will do no har

tho

and add dilute hydrochloric acid until the solutions contain a few drops in excess, and finally add to each a concentrated solution of 5 grams of pure sodium bicarbonate (NaHCO_{3}) in water. Cover the beakers befor

nious oxide, calculate its relation to the normal. From the ratio between

stic alkali during the titration is not admissible. It is therefore destroyed by the addition of acid, and the solution i

ring titration

+ 2NaHCO_{3} -> Na_{3}

H_

presence of even the weakly alkaline bicarbonate, it is best to avoid the addition of any considerable exc

ION OF COP

arsenic, and antimony. In nearly all varieties there will be found a siliceous residue insoluble in acids. The method here given, which is a modification of that described by A.H. Low (!J. Am. Chem. Soc.! (1902),

d (sp. gr. 1.42) and heat very gently until the ore is decomposed and the acid evaporated nearly to dryness (Note 1). Add 5 cc. of concentrated hydrochloric acid (sp. gr. 1.2) and warm gently. Then add abou

chloride, and receive the filtrate in a small beaker, washing the precipitate and filter paper with warm water until the filtrate and washings amount to 75 cc. Bend a strip of aluminium foil (5 cm. x 12 cm

Erlenmeyer flask under the funnel. Pour 15 cc. of dilute nitric acid (sp. gr. 1.20) over the aluminium foil in the beaker, thus dissolving any adhering copper. Wash the foil with hot water and remove it. Warm this nitric acid solution and pour it slowly through the filter paper, thereby dissolving the copper on the paper, receiving the acid solution in the Erlenmeyer flask. Before washing the paper, pour 5 cc. of saturated bromine water (Note 3) through it and finally wash the paper carefully with hot water and transfer any particles of copper which may be left on it to the Erlenmeyer flask. Boil to exp

alculate the percentage o

per from carbonate ores. The hydrochloric acid is added to dissolve oxides of iron and to precipitate silver and lead. The sulphuric acid displa

directions carefully at this point. Lead and silver have been almost completely removed as sulphate and chloride respectively, or they too would be precipitated on the aluminium. Bismuth, though

ompounds from the original sample, and to oxidize nitrous acid f

etic acid, but in the presence of these strong acids arsenic and antimonic acids may react with the hydriod

N OF ANTIMON

many impurities, notably iron, seriously interfere with the accurate determination of the antimony by iodometric methods. It is, moreov

When the residue is white, add to each beaker 2 grams of powdered tartaric acid (Note 2). Warm the solution on the water bath for ten minutes longer, dilute the solution very cautiously by adding water in portions of 5 cc., stopping if the solution turns red. It is poss

lution (which should not occur if the directions are followed

0 cc. of water in a 500 cc. beaker, and pour the cold solution of the antimony chloride into this, avoiding loss by effervescence. Make sure that the solution contains an

n required to oxidize the antimony, calculate t

rom its concentrated solutions; hence these solutio

but in the presence of tartaric acid compounds with complex ions are formed, and these are soluble. An excess of hydrochloric acid also prevents pre

dilution at which the antimony sulphide, being no longer held in solution by the acid, separates. If the dilution is immediately stopped and the solution warmed, this sulphide is again brought into solution and at the same time more of the sulphureted hydrogen

he acid, thus lessening the amount of sodium bicarbonate to

h takes place during titra

+ I_{2} -> Na_{3}SbO_{4}

of sodium bicarbonate, leaving the solution slightly acid, or a very slight precipitation of an antimony compound which is slowly acted upon by the iodine when the latter

ORI

chlorine, hypochlorites, bromine, and hypobromites. The reagent employed is sodium arseni

sO_{3} -> Na_{3

mall quantity. To prepare the solution, dissolve about 5 grams of the powdered oxide, accurately weighed, in 10 cc. of a concentrated sodium hydroxide solution, dilute the solution to 300 cc., and make it faintly acid wi

but solution in sodium hydroxide takes place much more rapidly, and the excess of the hydroxide is readily neutral

to which 1 gram of potassium iodide has been added. These strips are allowed to drain and spread upon a watch-glass until dry. When touch

E AVAILABLE CHLORIN

or disinfecting agent, or as a source of chlorine, depends upon the amount of hypochlorous acid which it yields when treated with a stronger acid. It is customary to e

ions of water until it is well ground and wash the contents into a 500 cc. measuring flask (Note 2). Fill the flask to the mark with water and shake thoroughly. Measure off

burette until no further reaction takes place on the starch-iodide paper when touched by a d

, calculate the percentage of available chlorine in the latter, assumin

{2} + 4H_{2}O -> 2

the original weight of bleaching

clear supernatant liquid gives percentages which are below, and the sediment percentages which are above, the average. The liquid measured off should, therefore, ca

orine. The original material for analysis should be kept in a closed container and protected form the air as far as possible. It is difficult to obtain analytical samples whic