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

M495

_ ___—_- — ——-——~'r— -— <-

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?

Amwwow—f -— — — »— w'7'* ‘- v—‘iwr

I OXIDATION METHODS FOR THE DETERMINATION
OF HYPOPHOSPHATE

II A NEW METHOD FOR THE PREPARATION OF SODIUM
DIHYDROGEN HYPOPHOSPHATE

by

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

E.

H.

1.

sodium dihydrogen hypophOSphate.

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
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.
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 phos-
phide 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 pyrophos-
phoric 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 add-
ing 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 accurate-
ly 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 main-
tain 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 dichro-

mate 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

714
7.8

8.4

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
100.1 100.4
99.72 100.7
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 hypophos-
phate, 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 tempera-
ture. 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 hypophos-
phate 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

5.
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 thio-

sulfate 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 hypophos-
phate 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 ,

-.v.-;- 1111/7/11;

RED PHOfi‘PHORUs

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 react-
ion 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.

(l)

(2)

(3)

(4)

(7)
(8)

(9)
(10)

(ll)
(12)

(13)

LITERATURE CITED

Bell, F. And Sugden, J.,

J. Chem. 500., 48-49 (1933).

Duke, F. H.,

1nd. hng. Uhem. Anal. Ed., 11, 530 (1945).

Ephraim, F., "Inorganic Chemistry,“ Edited by
Throne, P.C.L. and Roberts, E. R., p. 717, New York,
Nordeman Publishing Co., lnc., l943.

rurman, N. H., ”Scott's Standard Methods of Chemical
Analysis,“ vol.l, p. 278, New York, D. Van Nostrand
Uo., lnc., 1939.

Gleu, K., Z. anal. Chem., §§, 305 (1933).

Jackson, D. T. and Parsons, J. L.,

Ind. Eng. Chem. Anal. Ed., 24 14-37 (1937).

Joly, T., Compt. rend., 19;, 1058-1150 (1885).

Jung, W.,

Asoc. quim. Argentina., 29, 15-33 (1941).

Jung, W., ibid,‘§g, 99—111 (1942).

Kelitowska, J. H.,

Rocziniki Uhem., 1Q, 313-16 (1936).

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

Z. anorg. allgem. Chem., ng, 310-14 (1937).
Komaroski, A. 5.; Filonove, V. F.; and Korenman, I.M.,

Z. anal. Chem., 35, 321-28 (1934).

(14)

(15)

(16)

(17)

(18)
(19)

(20)

(21)

(22)
(23)
(24)

(25)

(26)
(27)

(28

Leininger, E., "Elementary Quantitative Analysis".
p. 63, Ann Arbor, Mich., Edwards brothers, lnc.,
1941.

Mellor, J.“., “A Comprehensive Treatise on Inorganic
and Theoretical Chemistry.“ Vol. 7111, p. 930, 936,
New York, Longmans, Green and Co., 1928.
Milobedzki, T.; Kolitowska, J. H.; and berxan, 2.,
Hoczniki Uhem.,.lz, 620-27 (1937).

Mfiller, V. I.,

Z. anorg. allgem. Chem., as, 29-70 (1916).
Pelletier, 8., Ann. de. Chem.g‘lg, 113 (1792).
Probst, J.,

Z. anorg. allgem. Chem., 119, 155-60 (1929).
hosenheim, A. and Pinsker, J.,

Z. anorg. Chem., 64, 327-41 (l909).

hosenheim, A. and Pinsker, J.,

ber.,.4fi, 2003-14 (1910).

Salzer, T., Ann., 1&1, 322-40 (1877).

Salzer, T., ibid, 123, 28-39 (1878).

Salzer, T., ibid,.fill, 1-36 (1882).

Sarver, L. A. and Kolthoff, I. M.,

J. Am. Chem. Soc., §§, 2906-10 (1931).

Schwicker, A., Z. anal. Chem.,‘9§, 110-61 (1937).
Speter, M., Hec. trav. Chem., 26, 588-89 (1927).
Treadwell, b. P., and Hall, W. T., "Analytical
Chemistry“. Vol.2, p. 321, New York, John Ailey &

Sons, lnc., 1942.

(29)

(30}

(31)

(32)

(33)

('64)
(35)

('66)

-58-

Treadwell, W. D. and Schwarzenbach, 0.,

helv. chim. Acta., 11, 405-16 (1928).

Von Name, H. G. and Huff, W. J.,

Am. J. Sci., 45, 91-100 (1918).

Van Name, H. G. and Huff, W. J.,

ibid, 36, 587-90 (1918).

flhite, J. F.; Taylor, M. 0.; and Vincent, G. P.,
Ind. Eng. Chem. Anal. hd.,Ԥ4, 782-94 (1942).
Wolf, I... and Jung, 61.,

Z. anorg. allgem. Chem., 201, 337-46 (l931).
Wolf, L. and Jung, “., ibid, 291, 347-52 (1931).
Wolf, L.; Jung, W.; and USpenskaja, L. P.,

ibid, 299, 125-28 (1932).

Yost, D. M., and Russell, H., "Systematic Inorganic

Chemistry", p. 209, New York, Prentice-Hall, 1944.

hit-

JAN 19 '53
MAR 2 0 ’53
MAY 2 5 '53
was" '55
“3J1 5 '5‘

 

 

 

 

{q‘llfl‘nxui

 

.00 .
3511118121 1““1’3’7

 

 

_ 3 1293 ‘024468