Trivalent manganese as an oxidimetric reagent

ANALYTICA

322

TRIVALENT

MANGANESE

CHIMICA

AS

AN

ACTA

OXIDIMETRIC

VOL. 6

(1952)

REAGENT

Manganic sulphate was first suggested as an oxidimctric reagent by UBBELOIIUE*. It has sornc advantages over potassium pcrmanganate, for ferrous salts may be dctermincd in the prcsencc of chlorides, oxalates may be titrated at much lower tcmpcrntures and nitrites, hydrogen peroxide and divalent vanadium are much UBBELOIIDE considered that this greater efficiency more smoothly oxidised. was due to the fact that only a simple valency change occurred. Fe+” + Mnmb3+ Fe+3 + Mn+z. The solution has a cleep red tint and may bc useJ as its own indicator. Unfortunately the reagent is not stable and deteriorates rapidly. Furthermore, when prepared by UBBELOHI)E’S method, (oxidation of manganous sulphatc with pcrmanganate at a high acid concentration) large quantities of manganese dioxide precipitate. After removal of this manganese dioxide by filtration, the resulting solution is too weak to be used for most purposes. In any case it is only possible to prepare So ml batches by this method, because of the difficulty of keeping the solution sufficiently cool during the various additions of concentrated sulphuric acid. Dropwise addition of reagents from burettes, allowing a long time interval between each addition, and replacement of the 36N sulphuric acid specified by UBBELOEIDE by ISN sulphuric acid to reduce the heating effects, had no significant cffcct on inhibiting the precipitation of manganese dioside. In an attempt to prepare a more stable form of trivalent manganese, ~~ELCHER AND TOWNEND~ examined the double sulphatcs of trivalent manganese with aluminium, chromium and potassium, for it is well-known that a double sulphate containing a reactive ion is often more stable than the single sulphate. The double sulphatc with rubidium is generally regarded as the most stable salt containing trivalent manganese, but its properties were not investigated, as its cost would prohibit its use as an analytical reagent. The double salts of trivalent manganese with aluminium and chromium were prepared by ETARD’S~ method; the former being light-blue and the latter dark l?i?fCr0lUX.S

p.

332.

VOL. 6 (1952)

I

* , ‘i 7 If 4’ i,l._dLI I1bW1 d - rL;‘rL 2 II bar’.ilJ TRIVALENT M’iiGANESE AS AN OXIDIMETRIC REAGENT

323

green in colour. Potassium manganic sulphate W;IS prepared by the method of OBEYER AND BES+. It was obtained as a brown amorphous powder. The aluminium and chromium salts were insoluble in water and various concentrations of sulphuric acid, but on heating in acid solution they apparently decomposed to give violet-red solutions. After standing for a short time manganese dioxide was precipitated. Potassium manganic sulphate dissolved readily in 3N and higher concentrations of sulphuric acid, but soon deposited manganese dioxide; the precipitation occurred less readily as the acid concentration increased. Phosphoric acid was found to cnhancc the stability, and when I ml was added for each IOO ml of GN or CJV sulphuric acid, manganese dioxide precipitation was dc)laycd for one or two days. After this period the manganese dioxide continued to precipitate and the strength of the solution fell accordingly. For example, a solution originally o.ogr4N fell to 0.0889~V on the twelfth day after preparation, and to o.o748N on the fifty-eighth day. Various attempts were made to improve the stability. The amount of phosphoric acid was increased, and pcrchloric and hydrofluoric acids were tried in its place. None of these mcasurcs effected any improvement; pcrchloric acid was ineffective in promoting stability, whilst hydrofluoric acid behaved similarly to phosphoric acid. The salt was never sufficiently pure to be used as a primary standard. When the theoretical amount necessary to give a o.xN solution was weighed out, the initial strength of the solution was generally about o.oggN. The impurity was probably sulphuric acid (in which medium the salt was prepared), for it is difficult to wash out the last traces. The new reagent was used for the titration of ferrous salts (in the prescncc of chlorides), oxalates, nitrites and hydrogen peroxide. 1n every case the reaction proccedcd more smoothly than that with permanganate. Titrations of ferrous salts containing amounts of hydrochloric acid up to 21V in strength gave end-points corresponding to theoretical, although when the hydrochloric acid exceeded o.5N it was necessary to add a drop or two of phosphoric acid to decolorise the ferric chloride formed as otherwise the end point was masked. With pcrmanganate under the same conditions the titration figures obtained were II per cent above potassium manganic sulphate theoretical in 2N hydrochloric acid. Although stabilised by phosphoric acid was an improvement on manganic sulphate, it was not found possible to improve its stability; and since the preparation of the salt is a long and tedious process, further work on the reagent was abandoned. The stabilisation

oj trivnhrt

ntangunese

by formation

of a com#dcx

Previous attempts to produce a stable reagent containing trivalent manganese having failed, attention was turned to the possibility of forming stable complex compounds. The most promising appeared to be that formed with the pyrophosphate ion. KOLTHOPF AND WATTER~ have shown that divalent manganese can Rcfevences p. 332.

R, BELCHOR,

324

T.

S.

WEST

VOL.

6 (1952)

be oxidiscd to the trivalent state in pyrophosphatc solution, and the solution is stable over long periods, However, the concentrations they used were much lower than would be practicable ‘for an oxidimetric reagent, for their purpose was to determine manganese polarographically by reduction of trivalent manganese at the dropping electrode. KOLTHOFFAND WATTHRS consider that a tridihydrogen pyrophosphatomanganiate is formed, having the formula

PW~J&?WP.

In the earlier parts of this work, mnnganic solutions were prepared

by oxidation. permanganate and by lead convenient and was used for

Two methods were used, oxidation by potassium dioxiclc. The latter oxidant was found to be more all further preparations. The effect of varying the 1~yrol~~~osl~I~ate concentration

was studied, but it was found that below a concentration of 0,4Jl, the oxidation of the manganese was incomplete, and the solution lacked stability. Solutions having clifferent p11 values were prcparcd and the greatest stability was found to be within the 1x1 range 4-G. fZclow prr 4 the solutions were unstable, whilst at values greater than ~7x1 6 the colour of the reagent became brownish and it deposited large quantities of potassium sulphate during the course of titration and also during storage. The reagent containing pyrophosphatc did not possess as great a tinctorial power as the reagents described previously, ancl it was necessary to use an indicator. The titrc of the reagent against standard ferrous sulphate remained ~lncll~n~~cl for six weeks, and there was no apparent photo-decomposition. This method of preparation by osidation of a manganous salt was considered troublcsomc howcvcr, and attention was turned to the possibility of using a stable manganic salt which could bc easily prepared. The only suitable salt fulfilling these requirements appeared to be manganic phosphate. This salt is insoluble in strong acicls, it is scarcely attacked by strong alkali, and also it is insoRuble in water. It can, however, be brought into solution by adding a strong soltrtion of pyrophosphate, and bringing to the boil for a few seconds. Varying quantities of acid may then be added to bring the reagent to the desired 1x1 value, and the precipitate of potassium sulphate filtered off after cooling, followed by dilution to the required value. Fluoride and orthophosphnte ions were also tried as cotnplcxing reagents but were found to be inferior to ~~yropl~osphate. Ethylencc~i~~mine tetra-acetic acid was also considered, but was not tried since the work of other investigator@ indicates that it does not form a sufficiently stable complex. TITRATIONS

USXNG hIAh’Gr\NIC PYROPHOSPI-IATE

Il”ilval&t Solutions

of ferrous

from manganic

iron were titrated

phosphate,

~f+%?~Ci% $% 332.

o/ ferrom

using barium

iron

satisfactorily with the reagent prepared diphenylamine sulphonate as indicator.

VOL.

6

(Iqj2)

TRIVALENT

AS AN 0XII)XMETRXC

BUNGANISE

The reaction was found to be stoichiometric Results are recorded in Tables I and II. TABLE TXTRAPION

ml of 0.1

N

OF

Fe+’

0.1

N

using o.ogN

ml of 0.1

N Mn+3 30.00

25.00 20.00

25.00 20.00

15*.00 10.00

OF

Fe+*

soln.

10.02 5.00

TABLE

N

solutions.

15.00

5.00

ml of 0.05

o.rN

SULPHATE

30.00

TITRATION

and

‘325

I FXSRROUS

soln.

REAGENT

0.05

N

11 FERROUS

SUtPllATlS

ml of 0.05

soln.

N Mn+3

30.00

30.00

25.00 20.00 X5.00 IO.00 5.00

25.00 20.00 15.00 10.00 5.00

sob.

In the prescncc of hydrochloric acid the accuracy was unaffected. Even at a concentration of 5.N hydrochloric acid no difference in titre could be observed. The results, compared with titres using potassium pcrmanganatc are given in Table III. TABLE TITRATION

WC1 concentration o.oN

OF

0.1

N

ml 0fo.r

FERROUS

0.5iv 0.75N o.gN

x5.00 15.00 X5.00

I.2N

r5.00

1.5N

15.00 15.00 X5.00 15.00

2N

References

p.

332.

SULPHATE

N Fc+*soIn. xg.00 15.00

0.2rv

III IN

mlofo.:

IfYDROCllLORXC

NMn+a

15.00 15.00 IS.00 x5.00 X5.00 X5.02 x5.00 15.00 X5.00 rjj.00

ACID

ml of 0.1 N MnO,rjj.00

15.30 15.60 15.65 15.70 X5.72 15.80 -

32G

R.

BDLCHER,

T.

S.

WEST

VOL.

6 (x952)

Titration of vanadiwn In the presence of phosphoric acid, solutions of vanadium, reduced to the tetravalent state by snlphur dioxidecould not be titrated directly with manganic pyrophosphatc. By adding a measured excess of ferrous sulphatc to the vanadatc solution acidified with phosphoric acid, the vanadium could be determined by back titrating the cxccss of ferrous iron. Results arc recorded in Table IV. TABLE TITRATION -

ml of 0.x

N

IV

OF

vanaclnte

V~NADATI:

ml of

0.1

.N

30.00

30.00

25.00 20.00 f 5.00

25.02 20.00

Mn+3

15.00 9.98 5.00

10.00 5.00

Titration of nitrile Solutions of nitrite wcrc satisfactorily titrated with manganic pyrophosphatc, but as in permanganate titrations of nitrite, it was found prcfcrable to reverse the procedure and titrate measured amounts of the reagent with the nitrite solution. The reaction procccdcd equally as well in cold as in hot solution. Results arc recorded in Table V. TABLE TITRATION

ml of 0.1

N

nitrite

OF

V NITRITE

ml

of 0.1

30.00 25.00

N

Mn+=

30.02 25.00 20.00 15.00 10.00

20.00 X5.00 10.00

5.00

5.00

Titva lion of oxalate The initial reaction between manganic pyrophosphate and an acid solution of osalate at an elcvatcd temperature was rapid, but near the equivalence point the reaction slowed down considerably. The reagent was insufficiently coloured to act as its own indicator, and the indicators examined were irreversibly attacked by the reagent at the temperature employed. Although titrations were effected without indicators, the procedure w‘as not considered to be satisfactory. References

p.

332.

VOL.

6

(1952)

TRIVALENT

MANGANESE

AS AN

Titration

OXIDIMETRIC

REAGENT

327

of peroxide

The initial reaction between manganic pyrophosphatc and an acidified solution peroxide was rapid, but near the equivalence point the rate of reaction fell off considerably. Acidification with phosphoric acid markedly slowed down the reaction and, as this suggested that the phosphate ion was interfering, an excess of ferric iron was added to the peroxide solution bcforc titration, to complex the pyrophosphate content of the reagent. This had a favourablc effect on the reaction, and it was found that the end point could easily bc located using the ferrous ortl~ophcnanthroline complex as indicator. ‘I’hc end-points obtained by this method did not correspond to the cquivalencc points however*, and it was concluded that the reagent was not suitable for the dctcrmination of pcrosidc.

. of hydrogen

Titratiou

of arsenilc

Solutions of arscnitc could not be titrated directly with manganic pyrophosphate, but the reaction could bc carried out iodometrically by allowing a solution of the manganic pyrophosphatc to liberate iodine from potassium iodide, and titrating the liberated iodine directly with the arsenitc in a bicarbonate buffered medium. The direct titration was effected in the prcscncc of osmic acid as a catalyst. The catalysed reaction took place smoothly, and the end-point was sharp. Iicsult s are recorded in Table VI. In subsequent work, pure arsenious oxide was used to standardisc manganic pyrophosphatc solutions. TABLE TITRATION

nrl of

0.1

N

OF

arsenitc

VI ARSJ3NITE

ml of 0.1

N Mn+3

30.00

30.00

25.00 20.00 15.00

25.00 20.00

14.98 IO.00 5.00

10.00 5.00 Redox

fiotcntial

of reagent

The above experiments indicated that although manganic pyrophosphate solution is more stable than manganic sulphate, it is a much weaker oxidant. The result of stabilising the manganic ion against decomposition has been to stabilise it to some extent against reduction. The decrease in oxidative power was confirmed by determining the redox potentials of manganic pyrophosphate and sulphate reagents. The corresponding potentials were found to be 1.22 and ca. 1.4 volt at 25’ C. The latter figure can only be regarded as approximate due l

The redox

potentials

References

potential of are destroyed.

p. 332.

the

indicator

is too

high.

Indicators

having

suitable

328

R. BELCHER,

T. S. WEST

VOL.

6 (x952)

.

to the instability of the reagent, but these values confirm that the oxidative 1 power of the manganic ion has been reduced by formation of a complex. In conclusion, although this investigation has revealed that it is possible to stabilise the manganic ion by forming a complex in order to use it as an oxidimetric reagent, it is doubtful whether the reagent will have any practical application, as its stability is inferior to that of potassium dichromate or ceric sulphate, which can also be used for the determination of ferrous iron in the presence of hydrochloric acid.

PREPARATION

EXPERIMENTAL OF MANGANIC

PYROPHOSPHATE

30 g of syrupy phosphoric acid and IO g of concentrated nitric acid were added to a solution of 34.2 g of manganous chloride in 50 ml of water. When this solution was concentrated nearIy to dryness a finely divided grey powder was obtained. This was filtered off after cooling and adding 50 ml of water. It was thoroughly washed with water and finally with acetone and ether, dried at IXO’ C, and stored in a reagent bottle. The reagent was prepared from manganic pyrophosphate as follows: Approximately Q g of manganic phosphate were weighed out and dissolved in 200 ml of 2nl pyrophosphate. The solution was bought to the boil slowly with constant stirring, and boiled for half a minute. During this time the colour of the solution turned to a deep reddish brown. The necessary amount of 50 per cent. sulphuric acid was added (approx. 40 ml) to bring the PH to 4 and the solution cooled thoroughly under the tap. It was allowed to stand for a few minutes to ensure precipitation of potassium sulphate, filtered and diluted to I: litre. The normality of the reagent prepared thus was approximately 0.05~. Xitration Solutions

of ferrous

iron

required

o.siW ferrous ammonium sulphate in o.xN o.xAf manganic pyrophosphate reagent 50 per cent. phosphoric acid (v/v) 0.2 per

cent. aqueous

barium

diphenylamine

sulphuric acid

sulphonate.

Measured amounts of ferrous ammonium sulphate were treated with 15-20 ml of 50 per cent. phosphoric acid, and a few drops of barium diphenylamine sulphonateindicatoradded. Thesolution wasdilutedto ~oomlandtitrateddirectly with manganic pyrophosphate from a burette, till the appearance of the violet oxidation colour indicated the end-point. Themanganicpyrophosphatesolution was standardised against pure arsenious &ferenccs

$, 332*

VOL.

(X952)

6

TRIVALENT

MANGANESE

AS AN OXIDIMETRIC

REAGENT

329

oxide as described below, and the ferrous ammonium sulphate standardised against potassium dichromate. Titrations in the presence of hydrochloric acid were carried out as above, except that the necessary amount of concentrated hydrochloric acid was added before diluting the solution to IOO ml. The corresponding permanganate titrations were carried out with potassium perrnanganate standardised against pure arsenious oxide, other experimental conditions being identical.

Titvation Solutions

of vanadale

vequired

Approx. o.rN ferrous ammonium sulphate in o.rN sulphuric sodium metavanadatc o. xN manganic pyrophosphate 50 per cent. phosphoric acid 0.2 per cent. aqueous barium diphenylamine sulphonatc.

acid

o.rN

Procedwe Measured amounts of sodium metavanadate were treated with x5/20 ml of 50 per cent. phosphoric acid, and a measured excess of ferrous ammonium sulphate added, A few drops of indicator were added and the solution diluted to IOO ml and back titrated with manganic pyrophosphate. The titre of the ferrous solution was evaluated against the manganic pyrophosphatc as described previously, the difference in titre being equivalent to the amount of vanadium. The metavanadate solution was standard&d by reducing with sulphur dioxide and titrating with standard permanganate. Titration Sohtiom

of nitvite

reqarircd sodium nitrite manganic pyrophosphatc 50 per cent. phosphoric acid.

O.IN o.rN

Procedwe Measured amounts of standard manganic pyrophosphate were placed in a conical flask and acidified with 50 per cent. phosphoric acid. The nitrite solution was added from a burette whilst shaking the flask thoroughly and the rate of addition was reduced in the neighborhood of the end-point. When the titration was carried out carefully the end-point was quite sharp. The nitrite solution was standardised against standard pcrmanganate.

Referertces p. 332.

330

R. BELCHER,

VOL.

T. S. WEST

6 (x952)

Tilvalion of oxalate Solulions required sulphuric acid o.riV manganic pyrophosphate A.R. sodium oxalate. 2N

Pvocedtive

Varying amounts of sodium oxalate were weighed out accurately and dissolved in approximately zoo ml of 2N sulphuric acid in a conical beaker. The solution pyrophosphate. was warmed to 50” C and titrated with standard manganic Indicators such as barium diphenylamine sulphonate were irreversibly destroyed. Titration of

ficroxide

required

Solz4lions

O.IN hydrogen peroxide o.xN manganic pyrophosphate 5N sulphuric acid ferrous orlhophenanthroline indicator A.R. ferric ammonium sulphate. Procedtcvc

Varying amounts of peroxide were delivered into a conical flask, the solution diluted to IOO ml with 5N sulphuric acid, and an excess of solid ferric ammonium sulphate was added. 3-4 drops of ferrous orthophenanthroline indicator were added, and the solution titrated with manganic pyrophosphate. The end-point had to be approached carefully, but was sharp to within one drop of reagent. The peroxide solution was standardised against potassium permanganatc. Tilralion of arscnile o .IN arsenious oxide O.IN manganic pyrophosphate 1.0 per cent. osmic acid catalyst 50 per cent. phosphoric acid 0.2 per cent. aqueous barium diphenylamine

sulphonate

indicator.

Procedure Varying amounts of arsenious oside were delivered into a conical flask; 3 drops of I per cent. osmic acid were added and 20-30 ml of 50 per cent. phosReferences p. 332.

VOL.

6

(1952)

TRIVALENT

MANGANESE

AS AN OXIDIMETRIC

REAGENT

33=

phoric acid. The volume was made up to IOO ml with distilled waler and 4-5 drops of indicator added, followed by direct titration with manganic pyrophosphate. The end-point was sharp to within one drop of reagent. Deternrimtion

of redox

$otential

py7ophosphate a. Manganic Solutions required o. I N ferrous sulphatc O.IN manganic pyrophosphate 50 per cent. phosphoric acid. ProcecEzcre 25 ml portions of mnnganic pyrophosphate solution, acidified with phosphoric acid, were accurately titrated with ferrous ammonium sulphate using barium diphenylamine sulphonate as indicator. After having established the titre, a measured amount of manganic pyrpphosphate was placed in a beaker together with half the equivalent amount of ferrous sulphate plus a few ml of phosphoric acid*. The redox potential of the solution was then measured against a saturated calomel electrode using a bright platinum indicator electrode on a Marconi mains pH meter. The value of the saturated calomcl electrode was checked against a normal hydrogen electrode immediately before use. All solutions used were maintained at 25O in an accurately-set thermostat tank.

b. Mamgavtic sdfihatc Solutions reqllived o.ogN manganic sulphate o.ogN ferrous sulphate 50 per cent. phosphoric acid. Procediivc solution was prepared according to the method of sulphate The solution was filtered and used immediately. The standard titre against ferrous sulphate and the oxidation potential were measured as described previously. Owing to the instability of the reagent, the value for E, thus obtained can only be regarded as approximate. Manganic

UBBELOHDE~.

ACKNOWLEDGEMiNT

of us (T.S.W.) is indebted to Monsanto this investigation to be carried out. One

* SuHicient to bring Rejcvemes p. 332.

the

acidity

to

I

N.

Chemicals, Ltd., for a grant enabling

332

R. BELCHER,

T.

S.

WEST

VOL.

6 (1952)

SUMMARY A study of the use of trivalent manganese as an oxidimetric reagent has been made. The manganic ion was stabilised with pyrophosphate. The stability of this reagent over a period of six weeks was estabhshed. The reagent may be used in conjunction with barium diphenylaminc sulphonate for the titration of ferrous iron even in the resence of 5N hydrochloric acid. It may also be used for the titration of vana CF lum, nitrite, oxalate, peroxide and arsemte, but the roceclures for oxalate and peroxide arc unsatisfactory. The redox otential of t g e reagent was 1.22 volt at 25O C, which was lower than that foun s for manganic sulphate (cu. I .4 volt). Les auteurs ont dtudid lcs possibilites cl’utilisation du manganese trivalent comme reactif d’oxydo-reduction. Lc cation Mn+a est stabilis6 par l’anion pyrophosphorique. La stabilite de ce rdactif a etd contrOlCc clurant six semaines. Ce reactif p_eut dtre utilise (en prBsence de diphhnylamincsulfonate de baryum) pour Fti;;;;atlon du fer(II), m&me dans l’acidc.chlorhydriquc 5 N. 11 pcut aussi &tre our lc dosage clu vanadium, des nitrites, cles oxalates, des peroxydes ct des ars 8.mtes. Pour les oxalates et les pcroxydcs, la mdthode n’est pas satisfaisantc. Le potentiel rddox clu rdactif est 1.22 volt h 25” C, il cst plus bas clue celui du sulfate manganique (env. x.4 volt). . ZUSAMMENBASSUNG Die Anwcndung von dreiwertigem Mangan als oxyclimctrisches Reagens wurclc untersucht. Das Mangani-Ion wurde mit Pyrophosphat stabilisiert. Die Stabilitat dieses Reagens wurde wahrend sechs Wochen kontrolliert. Dieses Reagens kann in Gegcnwart von diphen lsulfosaurem Barium zur Titration von Ferro-Ionen sogar in Gegenwart von 5 Kr Salzs%me verwendet werclen. Es kann such zur volumetrischen Bestimmung von Vanadium, Nitrit, Oxalat, Peroxyd und Arsenit verwendet werden; ftir die Oxalate und Pcroxyde ist die Mcthode jedoch nicht befriecligend. Das Redox-Potential cles Rcagens betragt x.22 Volt bei 25” C, was niedrigcr ist als das Potential von Manganisulfat (cu. 1.4 Volt). REFERENCES l s a 4 8 o

A. R. A. R. 1. R.

R. J. P. UBDELOXD~, J. Chetn. SOL, (1935) 1605. BGLCH~R AND J. ‘TOWN~ND, unpublished work. ETARD, Compt. rend., 86 (1878) 1400. J, MEYER AND H. BEST, Z. a,rovg. Clrejn., 22 (I 899) 22. M. KOLTIIOFF AND J. L. WATTERS, Ind. 6vtg. C/rem., Axnl. PRIBL, Colt. Caeck. CIJeni. Co?tam., 14 (1949) 323. Received

Ed.,

November

15 (1943) goth,

8.

1951