An Introductory Course of Quantitative Chemical Analysis / With Explanatory Notes

TRY AND

L DISC

nnot be boiled without danger of loss of strength; sulphuric acid solutions may be boiled without loss, but the acid forms insoluble sulphates with three of the alkaline earths; oxalic acid can be accuratel

ttles, thereby losing strength; sodium carbonate may be weighed directly if its purity is assured, but the presence of carbonic acid from the carbonate is a disadvantage with many indicators; barium hydroxide solutions may be prepared which are entirely free from carbon dioxide, and such solutions immediately show by pr

mal solution of hydrochloric acid (HCl) should contain 36.46 grams of gaseous hydrogen chloride, since that amount furnishes the requisite 1 gram of replaceable

tity is represented by the molecular weight in grams (40.01) of sodium hydroxide (NaOH), while a sodium carbonate solution

case of the less soluble barium hydroxide). Solutions of the latter strength yield

ICA

s intended to bring about. In the neutralization processes which are employed in the measurement of alkalies (!alkalimetry!) or acids (!acidimetry!) the end-point of the reaction should, in principle, be that of complete

+}, OH^{-} -> Na^{+}

own which, in solution, undergo a sharp change of color as soon as even a minute excess of H^{+} or OH^{-} ions are present. Some, as will be seen, react sharply in the presence of H

plied to aqueous solutions of electrolytes is assumed. A brief outline of

F ORGANIC

acid and alkali titrations are methy

the changes which they undergo have been carefully studied by Stieglitz (!J. Am. Chem. Soc.!, 25, 1112) and others, and it appears that the changes involved are of two sorts: First, a rearrangement of the atoms within the molecule, such as often occurs in organic compounds; and,

color which it imparts to solutions is ascribed to the presence of the undissociated base. If an acid, such as HCl, is adde

^{+}, Cl^{-} -> (M.o.)

ts a characteristic red color to the solution. As these changes occur in the presence of even a very small excess of acid (that is, of H^{+} ions), it serves as the desired index of their presence in the solution. If, now, an alkali, such as NaOH, is added to this reddened solution, the reverse series of c

)^{+} -> [M'.o'.

ce of free H^{+} ions. When an alkali is added to such a solution, even in slight excess, the anion of the salt which has formed from the ac

^{-} -> (H_{2}O) + Na^{+}, (

the other hand, occasions first the reversion to the colorless ion

} -> H^{+}, (Ph

is a weak base and, therefore, but little dissociated. It should, then, be formed in the undissociated condition as soon as even a slight excess of OH^{-} ions is present in the solution, and there should be a prompt change from red to yellow as outlined above. On the othe

_{2}^{-} + H^{+}, OH^{

_{3}O_

, and methyl orange cannot be used in the titration of such acids, while with the very weak acids, such as carbonic acid or hydrogen sulphide (hydrosulphuric acid), the salts formed with methyl orange are, in effect, completely hydrolyzed (i.e., no neutral

ee acid (that is, free H^{+} ions) in the solution. This indicator cannot, however, be successfully used with weak bases, even ammonium hydroxide; for, since it is weak acid, the salts which it forms with weak alkalies are easily hydrolyzed, and as a consequence of this hydrolysi

hich must be removed by boiling the solution before titration. It is t

icators may be used, since very little hydrolysis ensues. It has been noted above that the color change does not occur exactly at theoretical neutrality, from which it follows that no two indicators will show exactly the

al sodium hydroxide) required to produce an alkaline end-point when run into 10 cc. of tenth-normal sul

_______________|____________|__________|_____________|______________ | cc. | cc. | cc. | Methyl orange | 10 | 9.90 | Red | Yellow Lacmoid | 10 | 10.00 | Red | Blue Litmus | 10

cator. It is important, therefore, to take pains to use approximately the same volume of solution when standardizing that is likely to be employed in analysis; and when it is necessary, as is often the case, to titrate the solution at boiling temperature, the stan

fully studied to determine the exact concentration of H^{+} ions at which the color change of each occurs. It is thus possible to select an indicator for a particular purpose with considerab

ples of Physical Chemi

Edition, 1921)

Theory and Use of Indi

), (

icators, !Z. physik. Che

Indicators, !J. Am. Chem. S

ve Applications of the

!J. Am. Chem. Soc.!

ssion, !Z. Anal. Chem.!, 6

Chemistry of Indica

10 (1912)

ikatoren der Acidimet

aden,

OF INDICAT

ulphurous acids, and is particularly useful in the determination of bases, such as sodium, potassium, barium, calcium, and ammonium hydroxides, and even many of the weak organic bases. It can also be used for the determination, by titration with a standard solution of a strong acid, of the salts of very weak acids, such as carbonates, sulphides, arsenit

bases, even ammonia. It is affected by carbonic acid, which must, therefore, be removed by boiling when other acids are to be measured. It can be used in hot solutions. Some care is necessary to keep the volume of the solutions to be titrated app

tituents which cause a troublesome intermediate color if not removed. The alcohol is decanted and drained off, after which the litmus is extracted with hot water until exhausted. The solution is allowed to settle for some time, the clear liquid siphoned off, concentrated to one-third its volume and acetic acid added in slight excess. It is then concentrated to a sirup, and a large excess of 95 per cent. alcohol added to it. This precipitates the blue coloring m

ugh the salts of the latter influence the indicator unfavorably if present in considerable concentration. It

OF STANDAR

Sodium Hydroxide. Appro

to an approximate volume of 1000 cc. Shake the solution vigorously for a full minute to insure uniformity. Be sure that the bottl

a liter bottle and dilute, as above, to approximately 1000 cc. This bottle should preferably have a rubber stopper, as the hydroxide solutio

ength greater than 0.5 N, as they are more readily di

re standardization, and thoroughness in this respect will, as stated, often

ate; an allowance is therefore made for this impurity by placing the weight taken at 23 grams

ACID AND ALK

and allowing the liquid to run out through the tip to displace all water and air from that part of the burette. Then fill

alkaline solution should never be allowed to remain long in a glass-stoppered burette, as it tends to cement

her burette, stirring constantly, until the pink has given place to a yellow. Wash down the sides of the beaker with a little distilled water if the solution has spattered upon them, return the beaker to the acid burette, and add acid to restore the pink; continue these alternations until the point is accurately fixed at which a single drop of either solutions served to produce a distinct ch

acid used by the number of cubic centimeters of alkali required for neutralization. The check results of the two titrations should not vary by more than two parts in one thousand (Note 2). If the variation in results is greater than this, refill the burettes and repeat the titration until satisfactor

e; any darker tint is unsatisfactory, since it is impossible to carry

ate work to be !generally applied!. In many cases, after experience is gained, the allowable error is less than this proportion. In a few cases a larger variation is permissible, but these ar

ION OF HYDRO

D PREPARATIO

), Ferguson (!J. Soc. Chem. Ind.! (1905), 24, 784), and others, seems to indicate that the best standard is sodium carbonate prepared from sodium bicarbonate by heating the latte

Na_{2}CO_{3} +

ium carbonate, such as may occur at higher temperat

e, supporting it with a clamp so that the bulb does not rest on the bottom of the crucible. Heat the outside crucible, using a rather small flame, and raise the temperature of the bicarbonate fairly rapidly to 270°C. Then regulate the heat in such a way that the temperature rises !slowly! to 300°C. in the course of a half-hour. The bicarbonate should be frequently stirred with a clean, dry, glass rod, and after stirring, should be heaped

ARDIZ

sure that no particles fall from it or from the tube elsewhere than in the beaker. Pour out from the tube a portion of the carbonate, replace the stopper and determine approximately how much has been removed. Continue this procedure until 1.00 to 1.10 grams has been taken from the tube. Then weigh the tube accurately and record the weight under the fir

m the burette, stirring and avoiding loss by effervescence, until the solution has become pink. Wash down the sides of the beaker with a !little! water from a wash-bottle, and then run in alkali from the other burette until the pink is replaced by yellow; then finish the titration as described o

tion. Subtract this volume from the volume of hydrochloric acid. The difference represents the volume of acid used to react with the sodium carbonate. Divide th

ponding weight of HCl in each cubic centimeter of the a

e weight which would be contained in the same volume of a normal solution of sodium carbonate. A normal solution of sodium carbonate contains 53.0 grams per liter, or 0.0530 gram per cc. (see page 29). The relation of the acid solution to the norm

late, from the known ratio of the two solutions, the relation of

in a common use for this purpose are purified oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O), potassium acid oxalate (KHC_{2}O_{4}.H_{2}O or KHC_{2}O_{4}), potassium tetroxalate (KHC_{2}O_{4}.H_{2}C_{2}O_{4}.2H_{2}O), or potassium ac

able by a base, while the tetroxalate contains three such atoms and the oxalic acid two. Each of the

is of sodium chloride to be described later. Sulphuric acid solutions may be standardized by precipitation of the sulphate ions as barium sulphate and weighing the ignited precipitate, bu

HE TOTAL ALKALINE

rocess it may contain sodium sulphide, silicate, and aluminate, and other impurities. Some of these, notably the hydroxide, combine with acids and contribut

olve the ash in 75 cc. of water, warming gently, and filter off the insoluble residue; wash the filter by filling it at least three times with distilled water, and allowing it to drain, adding the washings to the main filtrate. Cool the filtrate to approximately the standard temperature of the laboratory, and transfer it to a 250 cc. measuring flask, washing out the beaker thoroughly. Add distilled water of laboratory temperature until the lowest point of the meniscus is level with the g

uated for !delivery!, using small quantities of water, which are added to the liquid in the beaker. A second 50 cc. portion from the main solution should be measured off into a second beaker. Dilute the solutions in each

erature changes, and the data derived from the standardization, calculate the perce

arbon dioxide, most of which escapes from the solution. Carbonic acid is a weak acid and, as such, does n

s involved may be s

3}^{-} -> 2Na^{+}, 2Cl^{-} + [H

termined by precipitating the carbonate by the addition of barium chloride, remov

eutralized and the carbonate has been converted to bicarbonate, and then adding methyl orange and completing the titration. The amount of acid necessary to change the methyl orange to pink is

THE ACID STRENG

more) in excess. Heat the solution to boiling for three minutes. If the pink returns during the boiling, discharge it with acid and again add 0.3 cc. in excess and repeat the boiling (Note 1). If the color does not then reappear, add alkali until it does, and a !drop or two! of acid in excess and boil aga

the oxalic acid, calculate the percentage of the crystall

s made contains some sodium carbonate. This reacts with the oxalic acid, setting free

Na_{2}CO_{3}

f the indicator is discharged in cold solutions at the point at which bicarbonate is formed. If, however, the solution is heated to boi

_{2}O + CO_{2}

hloric acid and boiling, the carbonate

ence in behavior of the two indicators, meth

ery dilute. If the directions in the procedure are strictly followed, no loss of acid

a given indicator are, therefore, comparable. As a matter of fact, a small error is involved in the procedure as outlined above. The comparison of the acid and alkali solutions was made, using methyl orange as an indicator, while the titration of the oxalic acid is made with the use of phenolphthalein. For our present purposes the small error may be negl