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I OXIDATION METHODS FOR THE DETERMINATIGN Q3“ HYP‘OPHOSPHATE

i‘l A NEW METHOD FOR‘FHE PREPARATION OF SQDEUM DIHYDROGEN HYPOPHOSFHATE

Thesis for Eh-e Degrae M. S. MHCHIGAN STATE COLLEGE Thomas Chuiski I '37:? 4-7

This is to certify that the

thesis entitled

I Oxidation Methods for the Determination of Hypophosphate

II A New Method for the Preparation of Sodium Dihydrogen Hy ophosphate presented 9

Thomas Chulski

has been accepted towards fulfillment of the requirements for

__l_~ . S . _degree in_Qh 63:1ij (Analytical)

_6 WGth Major profeJr

Date Ivzarch 10, 19h?

M495

v—‘iwr -— — — »— w'7'* ‘- _ ___—_- — ——-——~'r— -— <- Amwwow—f

I OXIDATION METHODS FOR THE DETERMINATION

OF HYPOPHOSPHATE

II A NEW METHOD FOR THE PREPARATION OF SODIUM

DIHYDROGEN HYPOPHOSPHATE

ihomas Ohulski

A THESIS Summittea to the School of Graduate Studies of Michigan

State College of Agriculture and Applied Science

in partial fulfillment of the requirements

for the degree of

MASTER OF SCIENCE

Department of Chemistry

l947

ACKNOWLEDGEMENT

Grateful appreciation is expressed to

Doctor Elmer Leininger under whose kind and efficient direction this work was carried out. TABLE OF CONTENTS

Oxidation methods for the determination of Page hypophosphate.

A. introduction. 1

B. Historical. 2

C. Preparation of Sodium dihydrogen. 8

hypophosphate.

D. Stability of 0.05000 normal solution of 14

sodium dihydrogen hypophOSphate.

E. Quantitative determination of hypophosphate

by oxidation with potassium dichromate. 16

Quantitative determination of hypophosphate

by oxidation with SOdium hypochlorite. 25

Quantitative determination of hypophosphate

by oxidation with potassium permanganate. 33

H. Oxidizing agents that do not react

quantitatively with nypophosphate. 36

l. Potassium bromate in hydrochloric

acid.

2. bromine.

3. Chlorine.

4. Iocine.

5. Potassium bromate in sulfuric acid. 6. Potassium periocate.

1. oxidizing agents that do not react xith

hypophOSphate. 44 l. Sedium chlorite.

2. Chloramine T.

11. A new method for the preparation of souium 47

dihydrogen hypophosphate.

Ill. Conclusion. 54

IV. Literature cited. 56 INTRODUCTION

Hypophosphoric acid, H4P206, is a.member of a series of cxyacids of phosphorus.

In hypOphoSphorous acid, h3P02, phosphorus exhibits a valence number of positive one; in phosporous acid,

H3P03, positive three; in hypophOSphoric acid, H4P305, positive four; and in phOSphoric acid, H3PO4, positive five.

Such a series is a fruitful field of study. For example, phosphite is rapidly and quantitatively oxidized to phosphate by iodine in a neutral solution.

HypophOSphite does net interfere. However if the reaction sclution is acidified, hypophospnite is oxidized to a.mixture of phosphite and phosphate. It can be made quantitative by neutralizing and oxidizing the phOSphite so formed to phosphatetES). However at pH 5.7 to 4.6 phoSphite is oxidized to hypophosphate with a yield of from 37tll) to 55tl2) per cent.

before further studies of the series can be made the chemistry of the individual members must be better known. The literature contains more information on phosphite than on hypophOSphite or hypOphosphate. What has been written on hyp0ph05phite is often vague and misleading while information on hypophosphate is to a large extent entirely lacking. Hence the purpose oi this work is twofOld, to work out methods of determining hypophosphate and to more fully characterize it. HISTORICAL

Pelletiertld) identified two acids of phoSphorus.

One he called 'l'acice phosphorique", and the other

'l‘acide phoSphoreux'. The acid “l'acide phosphoreux"

was commonly known as ”Pelletier's acid" until Salzer

in le77 identified it as suhphosphoric acid\22), now

called hypophosphoric acid, H4P306.

The classical method of preparation is by the slow

oxidation of yellow phosphorus. Sticks of yellow phos-

phorus are half sucmerged in water and oxidation is

allowed to proceed for several weeks. A mixture of

phOSphorous, hypophosphoric, anu phosphoric acids is

Obtained. brom 6 to l6 percent of the phosphorus is

converted to hypophosphate. SalzerQZS) reported 6 to 7

per cent of the phosphorus oxidized to hypophosphate,

Jelyt7) reported l2 to 14 per cent, and van Name and

huff (31) reported l0 to l6 per cent conversion. In all

cases the hypophoSphate was isolated as sodium dihydrogen

hypophoSphate hexahydrate, Na2H2P206°6HgO.

Hosenheim and PinsxertZO) prepared hypophOSphate

by the oxidation of yellow phosphorus with cupric ion in

dilute nitric acid sclution. This method was later

modified by Jungte). Rosenheim and Pinsker reported a

yield of 10 per cent of the theoretical, Jung reported

l8.l per cent. Hosenheim and Pinsker (20) also prepared hypophos-

phate by the electrOlysis of a l to 2 per cent sulfuric acid sclution using anodes of copper, nickel, or silver phOSphide and the corresponding metal as the cathode.

They reported a yield of 60 per cent. With iron phosphide only phosphate was formed.

Speter(27) prepared hypophOSphate by the oxidation

of red phoSphorus with bleaching powder, calcium chloride

hypochlorite. Five g. of red phosphorus was heated

with l liter of lO per cent solution of practical grade

bleaching powder. After filtering, the acids formed

were nearly neutralized with solid calcium carbonate.

Calcium phOSphite, hypophosphate, and phOSphate precipi-

tate with standing. A solution of the pure acids was

Obtained by dissolving the precipitate in the smallest possible volume of dilute sulfuric acid. Calcium

sulfate settles out. Speter did not isolate the

hypophOSphate but used the solution of mixed acids as

a test reagent for thorium. He reported that 64 per

cent of the phOSphorus was oxidized, with hypophosphoric

acid predominating.

Probst(19) modified Speter's method by utilizing

an alkaline solution of sodium hypochlorite instead of

the bleaching powder, thus entirely eliminating the

calcium ion. He reported a yield of 25 per cent.

Kclitowska(ll) prepared hypophosphoric acid by

oxidizing the hydrolysis products of phOSphorus

tetraiodide, P314, and pnOSphoruS triiodide, P13, with

iouine at approximately pH 6.7. He reported a yield of 44.6 per cent from the tetraiodide and 37.3 per cent from the triiodide.

Kolitowska(l0) also prepared hypophosphoric acid by the oxidation of the hydrolysis products of phosphorus trichloride, PCls, with iodine at a pH of 5.7.

Kolitowska(l2) also prepared hypophosphoric acid from the hydrolysis products of phOSphorus tribromide by the oxidation with iodine in a solution buffered with sodium acetate and sodium bicarbonate. He reported a yield of 55 per cent.

Hilobedzike, Kolitowska, and Berxan(l6) prepared hypophOSphoric acid by the oxidation of red phOSphorus with iodine at a pH of 5.7 to 4.6 with a yield of 34 per cent.

Salzer(24) determined hypophoSphate by direct titration With potassium permanganate. Satisfactory results could by obtained only if the potassium permanganate was standardized against pure sodium dihydrogen hypophOSphate.

Hypophosphate can be hydrolyzed to phosphite and phosphate with hydrochloric acid and the phosphite so formed ddermined(30).

Treadwell and Schwarzenbach(29) determined hypophOSphate by titration with a standard solution of uranous sulfate, U(SO4)3, in an atmOSphere of carbon dioxide obtaining the end point electrometrically. uranous hypophOSphate, UPZOG, is insoluble. Silver hypophoSphate is insoluble over a pH range of l to 2 while silver hypophosphite, phOSphite and pnOSphate are scluble. Hence the separation of hypophosphate as silver hypophosphate can be used as a method of determination. Wolf and Jung(34) buffered a solution containing hypophosphite, phOSphite, and hypothSphate with phosphoric or formic acid. Excess Standard silver nitrate solution was added. The silver hypOphOSphate was filtered off and the excess silver determined by titrating the filtrate with standard potassium thiocyanate solution. Wolf, Jung, and Uspenskaja(35} modified the method

by titrating directly with silver nitrate solution obtaining the end point potentiometrically in a solution

buffered by dideium phoSphate. Exclusion of air was

not necessary.

A qualitative test for hypophOSphoric acid and its

salts is the formation of a precipitate with guanidine

carbonate, 2NHC(NH3)2-H2003. Mﬁller(20) assigned the

formula (CN3H6)2P05+H30 to the precipitate. Rosenheim and Pinsker(21) said it was (0N335)4H2P03+5Hbo. Bell and Sugden(l) in a more recent work determined the formula of guanidine hypophospnate to be 4(CN3H4)-g§:865-

Hypophosphoric acid and its salts form a curdy

precipitate with thorium which is insoluble in boiling hydrochloric acid. The reaction was used by Speter(27) -b—

as a qualitative test for thorium. The precipitate formed is the normal thorium hypophosphate, ThPZOG'llHBO. The doubled formula, (H2P03)2 or H4P205, for

hypophosphoric acid was established by conductance and

freezing point data. Treadwell and Schwarzenbach(29)

obtained the ionization constants of the acid by

electrometrio methods. Kl has the value of 10'2‘2, K2 of 10-2.81, K3 of 10’7'27. and K4 of 10-10.03.

However Rosenheim and Pinsker<2l) found that the

freezing points of sclutions of hypophosphoric acid

were not in accord with Treadwell and Schwarzenbach's

values for K1 and KB. Moreover, the molar conductances

of pyrophosphoric acid, H4P207, are approximately the

same as those for hypophoSphoric acid. But pyrophosphoric acid is considered to be a stronger acid with ionization constants K1 of lO-O°85. Kg 0f 10‘1'96. is of 10‘5-54, and K4 of 10‘8'44.

Although the data is not in good agreement the

determination of four ionization constants substantiates

the doubled formula.

Belliand Sugdentl) found that the hypophOSphates “32H2P206’ Na3H2P205'5H20’ Ag4P206, and 4(CN3H4)2§PEOB.

are diamagnetic indicating an even number of 6 extranuclear charges. The formula, HZPO3. W0Uld contain

an odd number, 25, of extranuclear electrons and hence

would be paramagnetic.

HypophOSphoric acid forms normal salts, acid salts and complex salts(15). The salts of most of the common metals have been prepared. There is no record of a cesium salt being prepared. Also there is no reference to tin, tantalum, gOld, mercury, or platinum hypophOSphate. PREPARATION OF SODIUM DIHYDROGEN HYROPHOSPHATE

Sodium dihydrogen hypophosphate hexahydrate,

NaZHZPZOG °6H20, was prepared by the method of Jung(9).

One hundred g. of copper turnings was placed in a 2 liter beaker. Two hundred ml. of water was added and loo ml. of concentrated nitric acid added a few ml. at a time. When the reaction subsided a piece of yellow phosphorus was added. When the reaction again subsided another piece of phOSphorus was added. Phosphorus was added in this manner until the blue color of the copper ion was disCharged.

Considerable heat was generated resulting in a loss of water. The loss was minimized by covering the beaker

with a watcn glass. Nevertheless, enough water was lost

so that the reaction mixture turned into a thick paste,

Slowing the reaction. Hence a few ml. of hot water was

added from time-to time to maintain the vclume. Cold

water could not be used as the reaction was easily

quenched.

After the blue color was discharged the beaker was

cooled under the tap. hapid cooling was necessary as with

Slow cooling, or standing, the blue color reappeared.

in such cases the prouuct was contaminated with copper

which persisted through recrystallizations.

After occling the sclution was decanted through a

ﬁdechner funnel. The residue was washed twice with

50 ml. pmﬂmions of ecld water. The residue was utilized for subsequent preparations.

The combined filtrate and washings were divided into two equal portions. Une portion was neutralized by adding anhydrous sodium carbonate until neutral to litmus. The solution was then heated to 50-60° and again neutralized to litmus with soaium carbonate. The two portions were combined and allowed to OOOl to room temperature, then placed in an ice box over night.

SOdium dihydrogen hypophosphate hexahydrate tends to form supersaturated sclutions. To prevent this formation a few crystals Of the substance were introduced immediately after mixing the two portions.

The crude prOduOt was a white crystal meal. The mother liquor was filtered off and the crystals washed once with a little cold water.

it was not possible to get a second crop of crystals from the mother liquor. The mother liquor was acid to litmus. Two hundred ml. was half neutralized as before with anhydrous socium carbonate but no crystals were

Obtained.

A portion of the mother liquor was concentrated to one quarter of its original volume, a thick syrupy

liquid being obtained. Seeding with crystals of the prOduOt and chilling did not induce crystallization.

The crude prOduct was purified by recrystallizing

twice from hot water. ~lO-

Three hundred sixty g. of yellow phOSphorus yielded

136 g. of pure sedium dihydrogen hypophosphate hexahydrate.

The percentage yield is 9.2 per cent. Jung(s) reported a yield of le.l per cent.

in addition, l6 g. of a less pure product was recovered from the liquors obtained from the process of recrystallization.

The purity of the preduct was established in two ways, first by loss of weight at lOS’ and second by oxidizing the nypopnosphate to the ortho phosphate and determining total phosphorus.

Approximately 3 g. of the pure product was accurately weighed out and heated for 24 hours at 105°. After that time it was weighed, heated again for half an hour and reweighed. The two weights were the same. After heating for 5 days no additional weight was lost.

Sodium dihydrogen hypophOSphate hexahydrate contains

34.40 per cent water; the loss of weight at 105° was

34.43 per cent.

The hYPOIhvbihﬁte was oxidized to the ortho— phosphate by modifying the procedure used by Treadwell and Hall(26) in oxidizing hypophosphite to ortho phosphate, the conditions of oxidation being made more drastic.

Approximately 0.5 g. of the pure product was accurately weighed into a 250 ml. beaker. It was dieselved in 100 ml. of 10 per cent nitric acid solution. -‘Ll-

The solution was then evaponted to dryness, but not baked, on an electric hot plate. Ten ml. of fuming nitric acid was added, the beaker covered with a watch glass and refluxed on the hot plate for three hours.

More fuming nitric acid was added when needed to maintain the original volume. After three hours the watch glass was removed and the solution again evaporated to dryness. The evaporation to dryness with fuming nitric

acid was repeated three times.

After the final evaporation the residue was taken up in l ml. of 6 normal hydrochloric acid and 50 ml. of water. Thirty five ml. of “magnesia minure“(l4) was

added. Magnesium ammonium phosphate was precipitated

by Slowly neutralizing with dilute ammonium hydroxide

solution.

The magnesium aamonium phosphate was purified by diSSOlving it in hydrochloric acid and reprecipitating.

Hiltration was made with a Gooch crucible. The magnesium ammonium phOSphate was ignited to magnesium

pyrophosphate, Mg2P207 and weighed.

Sodium dihydrogen hypophosphate hexahydrate is

l9.73 per cent phosphorus. The per cent of phOSphorus

in the preparation was also 19.73 per cent.

Both methods gave one hundred per cent purity with

an accuracy of one part per thousand.

Kalitowska(l2) prepared hypophosphoric acid by

oxidizing the hydrolysis products of phosphorus -L2-

tribromide, PBrs, with iodine in a solution buffered with sodium acetate and sodium bicarbonate using very dilute solutions. He reported a yield of 55 per cent of the theoretical. '

Kolitowska(ll) also prepared hypophosphoric acid by oxidation of the hydrolysis products of phosphorus tetraiooide, P214, and phoSphoric triiocide, P13, at pH 5.7 using very dilute solutions. He reported a yield of 44.66 per cent from the triiodide.

The hydrolysis products of phosphorus tribromide, phoSphorus tetraiocide, and phOSphorus triiodide are phosphorous acid and the reSpective hydrohalic acid.

An attempt was made to prepare hypophosphoric acid by the oxidation of pure phOSphorous acid with iodine. Ten g. of pure phoSphorous acid was dissolved in lOU ml. of a 6 per cent socium hydroxide solution. One hundred ml. of a solution containing lb g. of potassium iodide and l5 g. of iodine was added dropwise with censtant mechanical stirring. AS the iodine was added the pH decreased so l0 ml. of a 10 per cent sodium hydroxide solution was added every 30 minutes for 120 minutes at which time all the iodine solution had been

added. In this manner the pH was kept approximately 5.

The reaction solution was allowed to stand for over

night, the iooine color disappearing during that time.

The solution was half-neutralized with anhydrous

sodium carbonate at 50-60°. After cooling the solution I -13-

was seeded with a few crystals of sodium dihydrogen hypophoSpnate hexahydrate. Crystallization was not induced by so doing.

The procedure was repeated but before half- neutralizing the solution was concentrated to one quarter the original volume. ho dodium dihydrogen hypophosphate hexahydrate was isolated.

Failure to obtain hypophosphate by the above pro— cedure is in no way a criticism of the work of Kolitowska.

The conditions were purposely changed. The dilute

solutions of Koiitowska were not employed, as the object of the above was to prepare a quantity of hypophosphate.

Probst(19) prepared hypophosphate by the oxidation of red phosphorus with sedium hypochlorite using an

alkaline solution. Close contrOl of both the alkalinity

and the concentration of the hypochlorite was necessary. Probst's method was investigated. It was deemed

impraCtical because of the difficulty in closely

controlling the concentration and alkalinity of the

required solution. No hypOphOSphate was isolated. STABILITY OF 0-05000 NORMAL SOLUTION or SODIUM DIHYDROGEN HYPOPHOSPHATE

Exact information on the stability of a 0.05000 normal solution of souium dihydrogen hypOphosphate was desired.

A liter of 0.02500 molar sclution of sooium dihydrogen hypophoSphate was made by dissolving 7.8540 g. of pure NaZHZPZOS-eﬁzo in warm water and making up to

exactly l liter. The solution was stored in a brown

bottle in diffused light.

when hypophosphate is oxidized to phosphate each phOSphorus atom goes from an oxidation number of plus

4 to an oxidation number of plus 5. As there are two

atoms of phosphorus per molecule of hypophOSphortc acid

the total change in oxidation number is plus 2. Hence

the equivalent weight on oxidation is one half the molecular weight. Therefore a 0.02500 melar solution of

dihydrogen hypophosphate is also 3 0.05000 normal

sclution. In this thesis a 0.02500 mOlar solution of

sodium dihydrogen hypophoSphate will be referred to as

a 0.05000 normal sclution.

The amount of hypophosphate in the solution was

determined periodically by oxidation with potassium

dichromate according to the method develOped in this

theSiSe -15-

Table 1

Stability of a 0.05000 N solution of NagHZPZOg-bﬂzo with standing Total weeks g./l. Na2H2P205'5H30 0 7.850 l 7.838 2 7.842 6 7.833 l0 7.838

hypophoSphate hydrolyzes to phosphate and phosphite:

NaZHZPZOg 4- 320 —¢ NaH2P03 + Naﬁzpo‘. Oxidation with potassium dichromate would not give information on the hydrOlysis reaction as on oxidation one molecule of phosphite under goes the same change in oxidation units as one molecule of hypophosphate.

Hence the amount of phoSphite present was determined by oxidation with iouine in a neutral solutionk33). In no case was any iOdine reduced indicating the absence of phosphite. -l6-

QUANTITATIVE DETERMINATION OF HYPOPHOFPHATE

BY OXIDATION WITH POTASSIUM DICHROMATE

Sedium dihydrogen hypothSphate is oxidized by potassium dichromate in hot sulfuric acid solution.

The reaction is quantitative if a sufficiently strong

acid solution is used and heating continued long enough.

Yost and Russell(36) state that hypothSphoric acid

is not oxidized by boiling dichromate.

Table 2 illustrates the effect of acid concentration.

Table 2

Per cent of hypophOSphate oxidized

with varying acid concentration

Uoncentration % NagHgPZOS of 82804(normal) oxidized

2 53.15 4 80.56 6 89.26

6 96.52 9 98.02 10 99.64

l2 99.88 l3 99.76

l8 99.94 gas- 5' Time of heating: 45 minutes in'a

boiling water bath. -17-

Table 3 illustrates the effect of the time of heating upon the completeness of the reaction.

Table 3

Per cent of hypophosphate oxidized

with varying time of heating Time of heating % NagHngOB (minutes) 0x1d1zed

10 96.52 20 99.28

30 l00.l 40 100.0

50 99.76

Acid concentration: 12 normal sulfuric.

Table 4 snows that the amount of excess dichromate is not critical.

Table 4

Per cent of hypophosphate oxidized with varying amount of K20r307 “t;:: 533%)07 % ii §?§§86

20 99.82

80 100.2

140 99.94

Acid concentration: 12 normal.

Time of heating: 40 minutes.

Procedure: 50 mi. of 0.05000 normal solution of Sodium dihydrogen hypophosphate and 25 m1. 0f 0-1500 -13-

normal potassium dichromate solution are pipetted into a

500 ml. Erlenmeyer flask. Thyﬂty five ml. of concentrated sulfuric acid is added (to bring the acid concentration to l2 normal) and the flask heated in a boiling water bath for 40 minutes.

After heating the flask is cooled and the contents diluted to approximately 250 ml.. Three ml. of 80 per cent phOSphoric acid is added for each 50 ml. of vOlume.

Excess standard ferrous sulfate sclution is added and the excess back titrated with standard dichromate sclution using diphenylamine soaium sulfonatekZS) as an oxidation indicator. 0.05 ml. of dichromate is subtracted as an indicator blank.

Altegnate_pgggedu;§; instead of heating in a boiling water bath the reaction can be carried out by heating on a steam bath. ﬂith the steam bath used, the naction was complete in 30 minutes.

As a high concentration of acid and a prolonged heating is necessary, it is quite possible that the hypophosphoric acid is first hydrolyZed to phosphorous and phosphoric acids, the reactive phosphorous acid reacting with the oxidizing agent.

However the reacting ratio is the same if hydrOlysis does or does not take place as shown by the equations

below: 3H4P205-+ K20r207 +- 4H2804 -9 6H3P04 -+ Ur3{SO4)3 -19-

g; Hipzoe + H20 "" Home '*' 331’04 3H5P03 -- K20r207 + 4HBSU4“’ 3351’04 4' Cr2(604)5 + K2804 + 4H20

in both cases:

equivalent weight NagHgP306 3ggggg£gge " “ K202 0? - Kallng?

As the ferrous-dichromate end point is involved in

this method, more information was necessary on the

accuracy of the end point in stronger acid solution and in larger v01umes than recommended(25).

Fifty ml. of ferrous sulfate solution was titrated

by adding 0.1500 normal potassium dichromate solution to

it. The concentration of sulfuric acid was varied but

the velume was kept at 150 ml. The results are given in

Table 5. UOlumn A is the volume of potassium dichromate

solution used in the titration having present 10 ml. of

80 per cent phosphoric acid. Column B is the volume of

dichromate used having present a volume of 80 per cent

phosphoric acid equal to one half the volume of concen-

trated sulfuric acid used.

The same procedure was repeated but the ferrous

sulfate was added to the potassium dichromate. Twenty

ml. of potassium:ﬁchromate solution was titrated with

ferrous sulfate sclution at varying acid concentrations,

the velume being kept at 150 ml. The results are given

1 n Table 60 Table 5

Titrating FeSU4 with K3020?

varying the acid concentration

and the amount of H3PO4 present . B 32382igoiial) “ﬁgs8:“°' (ml. £20r207) (ml- K20r207) 2 8 25.07 25.05

4 16 25.08 25.06

6 26 25.04 25.07

8 34 24.84 25.04 10 42 24.75' 25.06 l2 50 24.76, 25.04

‘End point indefinite.

Table 6 Titrating K20r207with re804 varying the acid concentration

and the amount of H3PO4 present H2§32fgo$§a,, mﬁésginc° (ml. :eSO4) (m1. Eeso4; 2 8 28.24 28.24

4 16 28.22 28.22

8 34 28.38 28.36

10 42 28.85 28.36

12 50 28.92 28.39

Table 5, column B, shows that ferrous sulfate can be titrated with potassium dichromate in strong acid solution if sufficient phosphoric acid is present.

However, the titration can not be carried out in the reverse direction in strong acid as shown by Table 6.

Twenty-five ml. of ferrous sulfate solution was titrated with potassium dichromate solution varying the volume of the reaction solution, 3 ml. of 80 per cent phosphoric acid being present for each 50 ml. of volume.

The results given in Table 7 show that the volume of reaction sclution does not influence the titer.

Table 7

Titrating FeSO4 with K30r207

varying the velume of reaction solution

rinal volume ml. K§Cr207 (ml.) us d

156 17.65

213 17.64

268 17.65

324 17.64

380 17.65

Acid concentration: 2 normal sulfuric.

Hence from Table 5, 6 and 7 the end point in the

above method is substantiated. However, if 3 ml. of

phoSphoric acid per 50 ml. of reaction solution is a

smaller velume than one half the volume of concentrated

sulfuric acid used, the latter volume would be used

(Table 5) .

It was desired to know if the determination of

hypophoSphate could be carried out in other than a

sulfuric acid solution. Therefore the reaction was carried out in a perchloric acid solution. Table 8 illustrates the effect of time of heating upon the completeness of the reaction.

Table 8

Per cent of hypophosphate oxidized

with varying time of heating

Time of heating % NaZH P 06 (minutes) OXlQEZBG

10 96.64

20 98.44

30 100.2 40 100.0

Acid concentration: 30 per cent perchloric acid.

Procedure: 50 ml. of 0.05000 normal solution of sodium dihydrogen hypophosphate and 25 ml. of 0.1500 normal potassium dichromate sclution were pipetted into

a 500 ml. Erlenmeyer flask. An equal volume of perch-

loric acid (70-72 per cent hydrogen perchlorate by weight) was added and the flask placed in a boiling

water bath.

After heating, the flask was cooied and the contents diluted to approximately 250 ml. Three ml. of 80 per

cent phosphoric acid was added for each 50 ml. of volume.

Excess standard ferrous sulfate solution was added and

the excess bacx titrated with.standard potassium dichromate solution using diphenylamine sodium sulfonate as an oxidation indicator. (Substitution of percnloric acid for sulfuric acid under the conditions used did not influence the ferrous-aicnromate and point.)

If phosphite reacts rapidly and completely with dibhromate under conditions at which hypophosphate is

Stable towards dichromate, a differential titration could be made. hence some data on the oxidation of phosphite with potassium dichromate was obtained. Table 9 gives information on the oxidation of phosphorous acid ltth potassium dichromate.

Table 9

Per cent of phosphite oxidized

with varying time of heating Time of heating %‘Hapo (minutes) oxidize 10 95.32 20 99.96 30 100.0

ACid concentration: 12 normal sulfuric.

Procedure: An approximately 0.025 melar solution

of phOSphorous acid was prepared by dissolving 2.1 g.

of the pure acid in 1 liter of water. The exact

concentration was determined by oxidation with iooine

following the method of Wolf and Jung(33).

Fifty ml. of the phoSphorous acid scintion and

25 ml. of 0.1500 normal dichromate sclution Here

pipetted into a 500 ml. hrlenmeyer flask and 35 ml. of -34-

concentrated sulfuric acid was added. The prooedure used in the determination of hypophosphate was then

carried out.

As the quantitative oxidation of phosphite with

potassium dichromate requires conditions very similar

to those required by hypophoSphate a differential

titration cannot be made.

Potassium dichromate was found to be stable under

all conditions used.

A stable sclution of ferrous sulfate was easily Obtained by use of the Duke Reductor(zﬁj. -25-

QUANTITATIVE DETERMINATION OF HYPOPHOSPHATE

BY OXIDATION WITH SODIUM HYPOCHLORITE

Souium dihydrogen hypophOSphate is quantitatively oxidized by sodium hypochlorite in a bicarbonate solution.

As a Source of sedium nypochlorite a commercial bleach (Chlorox, Uhlorox Chemical Co., Oakland, Calif.) containing 5.25 per cent sedium hypochlorite by might was used. Sixty ml. was diluted to 1 liter to give an approximately 0.07 normal sclution. As shown by Table 10, such a sclution was stable for a number of days when kept in a br0wn bottle in diffused light.

Table 10

Stability of a 0.07 normal solution of NaClO

Days after ml. Na 8 0 equivalent dilution to ml. a010 used

0 20.03

2 19.98 4 19.96

6 19.96

8 19.98 l0 19.97

25 ml. of Na010 used.

Sooium dihydrogen hypophosphate was allowed to

react with sodium hypochlorite at various pH’s. —26-

Reaction solutions of various pH's were obtained by the use of buffer solutions. A M/lb solution of primary

Sodium phoSphate and a M/lb solution of secondary sodium phoSphate were mixed in various ratios. The pH's of the resulting sclutions was determined with a Beckmann pH meter. The phosphate buffer covered the pH range from

6.0 to 7.7.

In like manner buffer sclutionScovering pH‘s from

7.7 to 8.5 were prepared by mixing a M/20 solution of

s0dium tetraborate withxrarious proportions of M/5

sclution of boric acid. buffer solutions covering the range from 8.5 to

16 were made by mixing M/5 sclutions of boric acid

with M/lO sclutions of socium hydroxide.

Fifty ml. of the buffer solution was added to the

hypophosphate-pypochlorite reaction sclution to maintain

a conStant pH. The pH of the reaction mixtures was also

determined with the Leeds & Northrup Universal pH meter

as at the extremes of each buffer system the pH of the

reaction mixture was not the sane as that of the buffer

solution used.

fifty ml. of the buffer solution was introouced into

a 250 ml. iodine flask. Twenty-five ml. of 0.05000

normal solution of SOdium dihydrogen hypophOSphate and

26 ml. of sodium nypochlorite approximately 0.07 normal

were pipetted in and the flask immediately stoppered.

The gutter was filled with a saturated solution of potassium iodide and the flask allowed to stand over night.

The amount of nypochlorite remaining was determined by allowing the potassium ioaide solution to run into the reaction sclution, aCidifing with 20 ml. of glacial acetic acid(4) and titrating the liberated iooine with standard thiosulfate solution. rreshly prepared starch sclution was used as indicator.

The strength of the hypochlorite solution was determined at each pH by omiting the hypophosphate solution but maintaining the volume by the addition of

25 ml. of water. These “normality blanks“ were handled in the same manner as the determinations.

Table 11

Per cent of hypophosphate oxidized

by Na010 at various pH's pH % Nazﬂzpzoe 5.7 93.67 6.3 94.88 6.9 96.73 7.7 l00.4 8.1 100.2 2.5 92.23 9.3 75.36

Time of standing: over night. Slightly different results were obtained if the concentration of the NaClO was determined directly. That is, the sodium nypochlorite solution was pipetted from the stock bottle into an iOdine flask, solid potassium iodide immediately added, the sclution acidified with glacial acetic acid and the liberated iOdine titrated with standard sodium thiosulfate sclution. This method of standardizing the sodium hypochlorite solution will be referred to as “direct standardization".

The above procedure was repeated, but the amount of hypophosphate oxidized was also calculated on the basis of "direct standardization“ of the hypochlorite Solution.

The results are given in Table 12. Column A is the per cent of hypoanSphate oxidized with calculations based on the strength of hypochlorite obtained by carrying along “normality blanksﬂ, column B is the per cent of hypophosphate oxidized with.calculations based on the strength of nypochlorite obtained by ”direct standard- ieation". -283!-

Table 12

Per cent of hypophosphate oxidized at various pH's calculated on the basis of "normality

blanks" and on "direct standardization“ pH per cent NaEHZqusoxidized

A b (cone. of NaClO by (conc. of NaClO "normality blanks") by ”direct standardi zati on)

97.02 101.1

99.44 101.9

714 100.1 100.4

7.8 99.72 100.7

8.4 99.72 100.4

Time of standing: over night. -29..

From Table 12 the optimum pH range for the reaction is from 7.4 to 8.4. The mede of arriving at the strength of tne nypochlorite solution has but little influence over the optimum pH range although slightly better results are obtained on the basis of "direCt Standardization".

A solution with a pH of approximately 7 is conveniently obtained by buffering with SOdium bicarbonate haVing present sufficient carbon dioxide to saturate tne sclution. When l g. of s0dium bicarbonate was added to a sclution composed of

25 ml. of the hypophosphate sclution and 25 ml. of the hypochlorite solution the pH was 7.7, well within the optimum ph range.

Twentyifive ml. of 0.05000 normal solution of sedium dihydrogen hypophosphate was allowed to react with varying amounts of sodium nypochlorite solution approXimately 0.07 normal having present 1 g. of sedium bicarbonate. The determination was carried out -30-

in the same manner as when the buffer solutions were used. The acidifing was done cautiously to avoid loss of iodine from the effervescing Solution.

The results shown in Table 13.

Table 13

Per cent of hypophosphate oxidized

with varying amount of ma010

per cent of Na3H2P2% oxidized % excess A B ﬂaClO (cone. of NaClO by (conc. of NaClO "normality blanks ) by "direct standardization")

15 99.52 100.0

40 99.57 100.1

70 99.29 100.0

125 95.39 100.2

Time of standing: over night.

From Table 15, the amount of excess nypochlorite is not critical if the normality of the hypochlorite solution is determined by "direct standardization". The amount of excess nypochlorite is critical if the concentration of the hypochlorite is determined by the procedure using "normality blanks". As the shorter procedure gives better results it is the recommended one. Too large an excess of sodium bicarbonate should

be avoided as sedium hypochlorite is unstable in a

strong solution of sodium bicarbonate as shown in Tabkill. -31-

Twenty~five ml. of approximately 0.07 normal sodium hypochlorite solution and 25 ml. of water were pipetted into iodine flasks containing varying amounts of sodium bicarbonate. The flasks were stoppered, the gutters filled with potassium iodide solution and the mixtures allowed to stand over night.

The amount of sodium hypochlorite remaining was determined iodometrically.

Table 14

Stability of NaClO in varying concentration

of sodium bicarbonate Naﬂgggaﬁgesent m1. NaZSBOS used

0.0 19.98

1.0 19.59

2.0 19.80

3.0 18.60

4.0 17.52

5.0 17.28

25 ml. sedium hypochlorite solution

Procedure: 1 g. of sodium bicarbonate, 25 ml. of

0.05000 normal solution of sodium dihydrogen hypophosphate, and 25 ml. of standard sodium hypochlorite solution approximately 0.0? normal, standardized by

”direct standardization“ are introeuced into an iodine flask. The flask is immediately stoppered, the gutter filled with a saturated solution of potassium iodide and the reaction solution allowed to stand over night.

The amount of souium hypochlorite remaining is

determined iodometrically by allowing the potassium

iodide solution in the gutter to run into the flask,

acidifying with 20 m1. of glacial acetic acid(4) and

titrating the liberated iodine with a standard solution

of potassium thiosulfate. Freshly prepared starch isolution is used as indicator.

NaClO + Na2H3P306 —92NaH2PO4 4- NaCl Equivalent Might NazHgPZOS = Nagﬂzgzga_ 2 " " NaClo = NaClQ 2 -33-

QUANTITATIVE DETERMINATION OF HYPOPHOSPHATE

BY OXIDATION WITH POTASSIUM PERMANGANATE

Salzer(24) quantitatively determined hypophoSphates by oxidation with potassium permanganate, satisfactory results were Obtainable only if six conditions were met. The conditions he determined were as follows. One, titration must take place at a gently boiling temperature. Two, the sulfuric acid used must be free from potassium permanganate decomposing substances. Three, the titration must begin immediately after adding the mineral acid and proceed as rapidly as possible to the end point. Four, addition of a large amount of water must be avoided, 10 cc. being enough for the analysis of a Small amount of salts. Addition of 50 to 100 cc. cause a marked effect on the result. Five, the end point

is marked by the appearance of the pink color of manganic salts which do not deposit manganese dioxide

on prolonged heating. Six, the potassium permanganate

solution must be standardized against pure sodium

dihydrogen hypophoSphate.

Salzer found that standardizing potassium perman-

ganate against pure sedium dihydrOgen hypOphOSphate gate

results about 2 per cent lower than standardizing

against more common substances.

A further study of the permanganate method was

made. -34—

Part 1.

A hot acid solution of sodium dihydrogen hypophosphate cannot be titrated directly with potassium permanganate solution.

Procedure: Direct titration of hot acid solutions

of sedium dihydrogen hypophosphate was attempted. l, 2,

4, and 6 normal sulfuric acid solutions were used and

titration attemtpted at 90 and 100°. In all cases the

reactions were too slow. Sooium vanadate, thallous

sulfate, and osmium tetraoxide(5) were employed as

catalysts, varying the acid concentration and temperature

as before. The osmium tetraoxide catalyzed the oxidation,

but not enough to make direct titration possible.

The same procedure was followed using 1, 3, and 6

normal perchloric acid sclutions at 90 and 100°, with

and without the mentioned catalySts. In all cases the

reaction was too slow for direct titration.

Part 2.

Sedium dihydrogen hypophOSpnate is quantitatively

oxidized by excess potassium permengarste in 0.5 normal

sulfuric acid solution. However, if the reaction

solution is allowed to stand for more than four hours

manganese dioxide is precipitated by the oxidation of

manganous ion by the exdess permanganate present.

Potassium warmanganate is stable for 5 hours in 0.5

normal sulfuric acid sclution if osmium tetraoxide is

not present. Ergcedure: 25 ml. of 0.05000 normal solution of soaium dihydrogen hypophosphate and 25 ml. of standard potassium permanganate solution approximately 0.1 normal are pipetted into a 250 ml. glass stoppered Erlenmeyer flask. Four ml. of 6 normal sulfuric acid is added, the flask tightly stoppered, and allowed to stand three and a half hours at room temperature.

Five ml. of 60 per cent phOSphoriO acid, 5 ml. of

6 normal sulfuric acid, and excess standard ferrous sulfate sdution is added. The ferrous sulfate is then back titrated with standard potassium permanganate solution

Table l5 shows the effect of varying the amount of

excess potassium permanganate.

Table 15

Per cent of hypophosphate oxidized with varying amounts of KMn04

excfss KMno4 NagHngoe oxidized 40 99.l2 80 99.38 90 99.65 l00 99.82 llO 99.56 120 99.73 130 99.65 l40 99.73 150 99.73

Time of standing: three and a half hours.

Acid concentration: 0.5 normal sulfuric. -35-

OXIDIZING AGENTS THAT DO NOT REACT

QUANTITATIVELY WITH HYPOPHOSPHATE

Potassium Bromate in Hydrochloric Acid

Potassium bromate in hydrochloric acid solution partially 0x1dizes sodium dihydrogen hypophOSphate in l0 normal sulfuric acid solution and to a lesser extent

in weaker acid.

The method of Schwicker(26) for determining hypophOSphorous and phOSphorous acids was modified by adding sulfuric acid to increase the acidity.

Procedure: 20 ml. of 0.05000 normal solution of

soaium dihydrogen hypophosphate and 20 ml. of 0.1000

normal potassium bromate solution were pipetted into

250 ml. iodine flasks.

Twenty-four ml. of water and 6 ml. of concentrated

hydrocnloric acid was added to one. The flask was

immediately stoppered and the gutter filled with a

saturated solution of potassium iodide.

To the other flasks was added sufficient

concentrated sulfuric acid diluted to 30 ml. with

water to give the desired acid concentration. Six ml.

of concentrated hydrochloric acid was added to each.

The flasks were stoppered in the same manner as above.

The flasks were allowed to stand over night at

room temperature. The potassium iodide solution was

allowed to run into the flasks and the amount of -37-

potassium bromate remaining determined by titrating the liberated iodine with a standard solution.of'sodium thiOSulfate. Freshly prepared Starch solution was used as a indicator.

Total acid % . ““2233?“ 32392526

1 3.l0

5 22.22

l0 33.80

Time of standing: over night.

Bromine

The free bromine liberated by the addition of hydrochloric acid to a bromate-bromide solution partially oxidized sodium dihydrogen hypophosphate in 10 normal sulfuric acid sclution. Mellor(l5) states that there is no reaction between hypophosphoric acid and bromine.

Procedure: 20 ml. of 0.l000 normal sclution of potassium bromate, 1.0 g. of solid potassium bromide, l0 ml. of 0.05000 normal solution of sodium dihydrogen

hypophosphate, 2 ml. of l.2 normal hydrochloric acid,

and 5 ml. of concentrated sulfuric acid were introduced

into a 250 ml. iouine flask in the order named. The

flask was quickly stoppered and the gutter filled with

a saturated solution of potassium ioaide.

The amount of sulfuric acid was increased to 10 ml.

in a second trial and omitted in a third. The three flasks were allowed to stand at room temperature over night.

The potassium iodide in the gutters was allowed to run into the flasks and the amount of bromine remaining determined by titrating the liberated ioaine with a standard solution of sodium thiosulfate. Freshly prepared starch solution was used as indicator.

concentration (353331) %oxi§§z:&06 0 0.00 5 9.00 l0 24.50

Chlorine

Chlorine reacts with sedium dihydrogen hypophOSphate to some extent in l0 normal sulfuric acid sclution and to a lesser extent in weaker acid sclution. Mellor(15) states that there is no reaction between hypophosphoric acid and chlorine.

Erocedure: 25 ml. of 0.05000 normal solution of sodium dihydrogen hypophosphate was pipetted into 500 ml.

Erlenmeyer flasks. Sufficient sulfuric acid was added to give the desired acid concentration. Chlorine gas from a metal cylinder was bubbled through the solutions for two hours.

The chlorine was removed by sweeping the flasks with nitrogen for one and a half hours and then the last -39-

traces were removed by heating for ten minutes in a water bath. (A strip of filter paper moistened with a

starch-potassium iodide solution was used to test for

complete removal of chlorine. The failure of the blue

starch-iodine complex to develOp when the strip was held

over the mouth of the flask indicated the absence of

chlorine.)

The amount of sodium dihydrogen hypophosphate

remaining was determined by using the potaSSium

dichromate procedure developed in this thesis, page l6

and the amount oxidized by the chlorine determined by

difference.

concentration % N H P H2804 #2 2 206 (normal) 0x1dized

0025 1002

5 3.72

10 6.64

Satisfactory blanks could not be obtained under

less acidic conditions.

Iooine

Sodium dihydrogen hypophosphate does not react with

iodine except in Strong acid sclution and then only to a

Slight extent. Mellor(l5) states that there is no react-

ion between hypophOSphoric acid and iOdine.

Procedure: 20 ml. of 0.05000 normal solution of

souium dihydrogen hypophOSpnate and 25 ml. of standard -40-

ioaine solution approximately 0.l normal were pipetted into 250 ml. iodine flaSks, the flaSks tightly stoppered and the gutters filled with saturated potassium iodide solution. After standing over night, the iodine remaining was determined by titrating with a standard sclution of soaium thiosulfate using freshly prepared starch solution as indicator.

A pH of 7 was obtained by having present l g. of sodium bicarbonate.

When the sodium dihydrogen hypophosphate and iodine sclutions were mixed the pH of the reaction mixture was

Acid concentration % NazHZPZOB oxidized

pH 7 0.00

pH 5 0.00

4 normal 5.40

The iodine solution used was a potassium iodide-

ioaine solution. As the iocide ion is oxidized to free

iodine by atmospheric oxygen, correction blanks were

carried along. At pH 7 and at pH 5 no iOdine was formed

by atmoSpheric oxidation. In 4 normal acid an

appreciable amount of additional iodine was present in

the flaSk. This amount was taken into consideration

when calculating the amount of iocine reacted with the

hypophosphate. -41-

Pctassium Bromate in Sulfuric Acid

Jung(5) states that sodium dihydrogen hypophosphate does not react with potassium iodate in a neutral or an alkaline sclution. The rate of reaction in 10 normal sulfuric acid is ecual to the rate of hydrolysis by 10 normal sulfuric acid. Hence the oxidation of sodium dihydrogen hypophoSpnate was actually a hydrOlysis to_ phosphorous and phosphoric acids, the former being oxidized to phosPhoric acid by the potassium iodate.

As phOSphorous acid is readily oxidized by potassium bromate(26), hypOphOSphate could be quantitatively determined if hydrOlysis could be made complete. but since phOsphorous acid is unstable the hydrolysis reaction would have to be carried out in the presence of the oxidizing agent to obtain accurate

results.

when potassium bromate is used as the oxidizing

agent, a quantitative hydrolysis of hypophosphate is not possible as potassium bromate is unstable under conditions

that cause hydrolysis.

Procedure: 20 ml. of 0.05000 normal solution of

sodium dihydrogen hypophosphate and 20 ml. of 0.1000

normal potassium bromate solution were pipetted into

250 ml. glass stoppered Erlenmeyer flasks. Sufficient

concentrated sulfuric acid was added to give the desired

acid concentration. -42..

A blank was prepared to check the stability of the pctassium bromate.

The reaction solutions were allowed to stand over night. The potassium bromate remaining was determined by diluting the reaction solution, adding solid potassium iodide, and titrating the liberated bdine with a standard solution of sodium thiosulfate. Freshly prepared starcn sclution was used as indicator.

Trial Conc. ml. ml. (E35231) agtaifgg regﬁgiggaby used KBrO3 blank 10 l9.l4 l9.35 #l 5 l8.90 l9.35 #2 lo l7.70 l9.35

hydrolysis was attempted by heating in a boiling water bath for 10 minutes blank 10 18.93 l9.35 #3 10 17.38 19.35

As the potassium bromate is unstable, the amount of hypophosphate cannot be accurately calculated. In the data above, the difference between the v0lume of thio-

sulfate actually used in the titration and the vOlume of thiosulfate theoretically required by the potassium

bromate is a measure of the total reaction. Both the decomposition of the bromate and the reduction of the

bromate by the hypophOSphate are ineluded in the total

reaction. Comparison of the trials with the blanks show -4 3—

that some of the hypophOSphate was oxidized by the bromate.

Potassium Periodate

Potassium periodate does not react with scdium dihydrogen hypophosphate in neutral or moderately acid sclution. In 4 and 6 normal sulfuric acid solution there is a slight Odeation.

Procedure: 20 ml. of 0.05000 normal solution of sodium dihydrogen hypOphosphate and 25 ml. of standard pctassium periodate solution approximately 0.05 normal were pipetted into glass stoppered Erlenmeyer flasks.

Sufficient concentrated Sulfuric acid was added with cooling to give the desired acid concentration.

After standing over night the amount of periOdate remaining was determined iOdeetrically. The solutions were diluted, soid potassium iOdide was added and the liberated iodine titrated with standard socium thiosulfate solution. Freshly prepared starch solution was used as indicator.

concentration % NagHgPZOG H SO GXidized (ngrmgl) 2 0.00

4 0.93

8. 4.l8

Potassium periodate was found to be

stable under the conditions used. -44-

OXIDIZING AGENTS THAT DO NOT REACT

WITH HYPOPHOSPHATE

SOdium Chlorite

Sodium chlorite, NaCloz, cannot be used in acid solution for a slow quantitative reaction because it is not Stable under such conditions(32).

In neutral solution and in 0.2 and 0.4 normal s0dium hydroxide solution SOdium dihydrogen hypophosphate is not oxidized by sodium chlorite.

Egocedure: 25 ml. of 0.05000 normal solution of sedium dihydrogen hypOphoSphate and 20 ml. of Standard

SOdium chlorite sclution approximately 0.l normal were pipetted into 250 ml. iodine flasks, SuffiCient 0.8 normal sodium hydroxide solution was added to make the reaction sclution 0.2 normal in respect to socium hydroxide.

A neutral solution was obtained by having present l gram of sodium oicaroonate.

After Standing over night the sodium hydroxide and

sodium bicarbonate were neutralized with equivalent

amounts of hydrochloric acid. The sodium chlorite was

determined iodometrically. 1.5 g. of sclid potassium

iodide was added and the solution acidified with 15 ml.

of 30 per cent acetic acid(6). The liberated iodine was

titrated with a standard solution of sodium thiosulfate 4.5-

using fresnly prepared starch sclution as indicator.

in all cases the original amount of sodium chlorite was present.

Uhloramine T

\Sodium Salt of p-TOluene Sulfochloramine)

Chloramine T does not oxidize sooium dihydrogen hypOpnoSpnate in acidic, basic, or neutral solution.

Procedure: 25 ml. of 0.05000 normal sedium dihydrogen hyOphospnate sclution and 25 ml. of Standard chloramine I sclution approximately 0.08 normal were pipetted into iooine flasks. Sufficient sulfuric acid

«as added to give reaction solutions of 0.25, 4, and 6 normal in respect to sulfuric acid.

In like manner, reaction solutions of 0.2 and 0.4 normal in reSpect to sodium hydroxide were prepared by

adding the proper volume of 0.8 normal sodium hydroxide

solution.

A neutral solution was obtained by the addition of

1 gram of sodium bicarbonate. The flasks were tightly stoppered and the gutters

filled with a saturated solution of potassium iOdide.

After standing over night the amount of chloramine T

remaining was determined iodometrically(13). .

This was done by allowing the potassium iodide

solution to run into the reaction solution and titrating

the liberated iodine with a standard solution d sodium -46..

thiosulfate using freshly prepared starch solution as

indicator.

The neu ral and basic solutions were acidified with

sulfuric acid before titrating.

In all cases the original amount of Chloramine T

was present. A NEW METHOD FOR THE PREPARATION OF SODIUM DIHYDROGEN HYPOPHOSPHATE _48..

If l0 per cent sodium chlorite solution is drOpped

on red phOSphorus a vigorous reaction takes place. If a

small amount of red phOSphorus is added to a 10 per cent

sedium chlorite solution there is no visible reaction unless the temperature is raised to about 40°. At that

temperature the phOSphorus reacts and heat is evolved.

If a larger quantity of phosphorus is added the reaction

becomes violent with an evolution of gases. Chlorine

and phosphine can be identified in the evolved gases by

their characteristic odor.

As chlorite does not oxidize hypophosphate, it was

thought that hypophoSphoric acid might be one of the products of the oxidation of red phOSphorus by sodium

chlorite. This possibility was investigated in the

following manner.

One g. of red pnOSphorus was added to 80 ml. of

l0 per cent sedium chlorite solution. The violence of

the reaction was reduced by cooling the reaction mixture

in cold water.

After all the red phOSphorus had reacted, the

solution was neutralized to phenolphthalein with 1 normal

sodium hydroxide solution. Excess brium chloride solu-

tion was added, precipitating phOSphate and hypophosphate but not hypophoSphite or phOSphite(lz). The precipitate

was filtered and washed free of chloride ion with water,

filter paper pulp being used to aid filtration. -49-

The precipitate was dissolved in 40 ml. of 2 normal nitric acid solution and diluted to 250 ml. Eight ml. of 6 normal sulfuric acid was added to precipitate the barium ion. After Standing over night the barium sulfate was filtered off.

The filtrate from the barium sulfate was neutralized to phenolphthalein with 1 normal sooium hydroxide solution. A pH of between l and 2 was obtained by the addition of 25 ml. of a 10 percent phOSphoric acid solution. Silver hypophosphate was precipitated by adding 4.0 g. of silver nitrate dissolved in 30 ml. of water.

After standing over night the silver hypophoSphate was filtered and washed with a l per cent nitric acid sclution.

The silver hypophosphate was dried in a vacuum desiccator over concentrated sulfuric acid and wdghed as Ag4P297.

One g. of red phoSphorus yielded 2.8 g. of pure silver hypophosphate, a yield of 29 per cent based on the weight of phOSphorus used.

Contrary to information in text books(3) (15) the silver hypophosphate darkened when exposed to light.

Silver hypophosphate prepared from pure sodium dihydrogen hypophosphate darkened to the same extent as the silver hypoPhQSphate prepared above, eliminating the possibility of contamination with silver chloride. -50-

In order to utilize the reaction for the production of a quantity of hypOphosphate, the apparatus sketched in Figure l was devised. "A" is one half of the outer jacket of a 400 mm.

condenser. 'B' is the inner tube of the condenser with

the enlarged upper end removed. "C” is a 500 ml. suction flask with the side arm connected to the water outlet

tube of the condenser by means of a short piece of rubber

tubing. I

The bottom part of the condenser jacket was filled

with glass chips to above the outlet tube. A layer of

glass wool was built on the chips to support an asbestos mat. The asbestos mat was formed by pouring in an aqueous suSpension of asbestos fibers. The condenser jacket was then filled for half its length with a mixture

of glass beads and red phosphorus. A second layer of

glass woOl was then formed and the tube filled to within

1 centimeter of the top with the glass bead-phOSphorus

mixture. The surface of the mixture was protected with

a third layer of glass wool. DROPPINL’ «lb

‘A

1111/7/11; RED PHOﬁ‘PHORUs -.v.-;-

61—955 BEAD-s. GLASﬁ wOOL.

1° SUCT‘ON

, . .... A5815 SToa MAT'

(:LA 5 5 WOOL’

(”.555 CHIPS

RUBBER svoPPEK

NW \\ ' ‘ I \I‘ “it ﬁt“ It 4/

ﬁ '-l .1. L ACTIVATED CARBON F IGURE 1 I -51-

The apparatus was used in the following manner.

A 10 per cent sooium chbrite solution was allowed to drop from the dropping funnel. Sufficient suction was applied to prevent the accumulation of chlorite solution in the upper part of the apparatus. The reaction was self initiating and generated a considerable amount of heat. It was found that if the chlorite

solution was passed in at a rate of 8 to 12 drops per minute the reaction went smoothly. A faster rate

caused "gassing" and at a slower rate sufficient heat was not produced to keep the reaction going. As the

sodium chlorite solution passed through the column it

oxidized the phosphorus, the oxidation products

collecting in the suction flask.

It was necessary to mix the red phosphorus with

glass beads to obtain a.mass poéﬁs enough to permit the passage of the chlorite solution. The layers of glass wool served to minimize channeling. The asbestos mat

prevented the carrying of phosphorus into the suction

flask by the spent chlorite solution. The glass chips prevented the plugging of the outlet tube. Activated carbon (Aorite A) was placed in the

suction flask as a decolorizing agent. If the solution

was not decolorized the product was contaminated with a

yellow substance which was not removed by recrystallra-

tion. The yellow color was not eliminated by using a

very pure grade of sodium chlorite or by using all all glass apparatus. To obtain complete decolorization it

was found necessary to use 2 g. of activated carbon for

each l00 ml. of solution and allow the two to stand in

contact for an over night period.

The tube '3" was a safety precaution. If the

reaction became too vigorous, it could be easily quenched

by passing cold water through the tube. Under normal

operating conditions the tube was filled with water but

there was no flow.

In order to get quantitative data, 31 g. of red

phOSphorus was mixed with an equal volume of glass beads

and placed in the apparatus.

A 10 per cent solution of practical grade sooium

chlorite was passed through the generator, 900 ml. being

required to react with 31 g. of red phosphorus.

Twenty-five ml. of the solution in the suction flask

was pipetted out and quantitatively diluted to 250 ml.

This solution had no oxidizing properties as a potassium

iodideestarch solution was not colored by it.

The hypothSphite and phosphite were determined in

an aliquot by oxidation with standard iOdine solution according to the method of Wolf and Jung(l). The amount of hypophosphate formed was determined

by precipitating the phosphate and hypophOSphate with

barium chloride. The mixed barium phOSphate and

hypophosphate was filtered through a fine fritted

filtering crucible(Selas 200i) and washed with a small amount of water. The entire crucible was placed in a

beaker and the hypophosphate determined by oxidation with

standard potassium dichromate solution,(page 19), the

barium and phosphate ions not interfering.

Fourteen per cent of the phOSphorus was converted

_to hypophosphite, 29 per cent to phosphite, and 32 per

cent to hypophOSphate.

The bulk of the solution was half-neutralized with

anhydrous souium Carbonate at 50-60° and allowed to

stand over night in a refrigerator. The crystals were

filtered and washed with a small amount of cold water.

After being air dried, the crude product of sodium

dihydrogen hypophosphate weighed 44 g. After

recrystallizing twice from water it weighed 33 g., a

Zl.l per cent yield based on the amount of phosphorus

used. An appreciable amount of insoluble Sibstance was separated in the first recrystallization.

In addition 4 g. of a less pure product was obtained

from the mother liquors of recrystallization.

The purity of the sodium dihydrogen hypophosphate

was established by oxidation with potassium dichromate.

Sodium dihydrogen hypophosphate contains 19.73 per cent

phosphorus. The per cent of phOSphorus in the preparation

of dihydrogen hypophosphate was 19.76 per cent.

Yellow phosphorus does not react with a 10 per cent

solution of sodium chlorite unless the temperature is raised to 80’. Even at l00’, the reaction is not rapid. CONCLUSION

HypophOSphate can be quantitatively determined by

OXidation with excess potassium dichromate in 12 normal

sulfuric acid solution by heating for 40 minutes in a boiling water bath. The amount of excess dichromate, the

acid concentration, and the time of heating are not

critical.

HypophOSphate is quantitatively oxidized by excess

sodium hypochlorite in a sodium bicarbonate solution

when allowed to stand for an over night period. The per

cent of excess hypochlorite and the amount of sodium

bicarbonate must be kept within certain limits but are

not critical.

HypophOSphate cannot be titrated directly with

potassium penmanganate solution. Excess permanganate

quantitatively oxidizes hypophosphate in 0.5 normal sul-

furic acid solution in three and a half hours. The per

cent of excess permanganate and the time of standing are

rather critical.

Hypophosphate is partially oxidized by bromine

monochloride, bromine, chlorine, iodine, potassium

bromate, and potassium periodate.

Hypophosphate is not oxidized by sodium chlorite

or chloramine T.

A new methOd for the preparation of souium

dihydrogen hypophOSphate is described. The reaction

between red phOSphorus and a 10 per cent SOdium chlorite solution is used. Fourteen per cent of the phosphorus is oxidized to hypophOSphite, 29 per cent to phoSphite, and

32 per cent to hypophOSphate. Also a dekagram apparatus is described for the utilization of the reaction for the production of sodium dihydrogen hypophosphate hexahydrate. LITERATURE CITED

(l) Bell, F. And Sugden, J.,

J. Chem. 500., 48-49 (1933). (2) Duke, F. H.,

1nd. hng. Uhem. Anal. Ed., 11, 530 (1945). (3) Ephraim, F., "Inorganic Chemistry,“ Edited by Throne, P.C.L. and Roberts, E. R., p. 717, New York,

Nordeman Publishing Co., lnc., l943. (4) rurman, N. H., ”Scott's Standard Methods of Chemical

Analysis,“ vol.l, p. 278, New York, D. Van Nostrand

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(7) Joly, T., Compt. rend., 19;, 1058-1150 (1885). (8) Jung, W., Asoc. quim. Argentina., 29, 15-33 (1941). (9) Jung, W., ibid,‘§g, 99—111 (1942). (10) Kelitowska, J. H., Rocziniki Uhem., 1Q, 313-16 (1936).

(ll) Kolitowska, J. H., ibid, 11, 616-16 (1937).

(12) Kolitowska, J. h., Z. anorg. allgem. Chem., ng, 310-14 (1937).

(13) Komaroski, A. 5.; Filonove, V. F.; and Korenman, I.M., Z. anal. Chem., 35, 321-28 (1934). (14) Leininger, E., "Elementary Quantitative Analysis".

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(15) Mellor, J.“., “A Comprehensive Treatise on Inorganic

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New York, Longmans, Green and Co., 1928.

(16) Milobedzki, T.; Kolitowska, J. H.; and berxan, 2.,

Hoczniki Uhem.,.lz, 620-27 (1937). (17) Mﬁller, V. I.,

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(19) Probst, J.,

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(23) Salzer, T., ibid, 123, 28-39 (1878). (24) Salzer, T., ibid,.ﬁll, 1-36 (1882). (25) Sarver, L. A. and Kolthoff, I. M.,

J. Am. Chem. Soc., §§, 2906-10 (1931). (26) Schwicker, A., Z. anal. Chem.,‘9§, 110-61 (1937). (27) Speter, M., Hec. trav. Chem., 26, 588-89 (1927).

(28 Treadwell, b. P., and Hall, W. T., "Analytical

Chemistry“. Vol.2, p. 321, New York, John Ailey &

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(29) Treadwell, W. D. and Schwarzenbach, 0.,

helv. chim. Acta., 11, 405-16 (1928). (30} Von Name, H. G. and Huff, W. J.,

Am. J. Sci., 45, 91-100 (1918). (31) Van Name, H. G. and Huff, W. J.,

ibid, 36, 587-90 (1918).

(32) ﬂhite, J. F.; Taylor, M. 0.; and Vincent, G. P.,

Ind. Eng. Chem. Anal. hd.,‘§4, 782-94 (1942).

(33) Wolf, I... and Jung, 61.,

Z. anorg. allgem. Chem., 201, 337-46 (l931).

('64) Wolf, L. and Jung, “., ibid, 291, 347-52 (1931). (35) Wolf, L.; Jung, W.; and USpenskaja, L. P., ibid, 299, 125-28 (1932).

('66) Yost, D. M., and Russell, H., "Systematic Inorganic Chemistry", p. 209, New York, Prentice-Hall, 1944. hit- JAN 19 '53

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