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Quantitative inorganic chromatography
JOURNAL OF CHROMATOGRAPHY
196
QUANTITATIVE
INdRGAkC
CHROMATOGRAPtiY
PAl$T VI. APPLICATION TO KINETIC STUDIES: THE OXIDATION OF HYDRAZINE BY FERRIC IRdti IN ACID AQUEOUS SOLUTIONS I?. H. POLLARD Depavlment
of Physical
and Inorganic (Received
AND
G. NICKLESS
Chemistry,
The University,
November
znd, 1959)
Brktol
(Great Bvilain)
The use of ion-exchange and/or paper chromatography to elucidate the mechanism of inorganic reactions has not, up to the present, been extensively used, though there are clearly fields where its application might do much to establish the correctness of postulates made from the classical physicochemical approach, or to throw light on hitherto unsuspected intermediates or products which would not be revealed by classical analytical procedures. A number of reactions have been studied to show how chromatography can help in this field, and this paper represents the first of such investigations. When hydrazine is oxidised in aqueous solution, the principal nitrogenous products are ammonia, nitrogen and hydrazoic acidl. The proportions of these products vary considerably with the oxidising agent and the conditions of the reaction. In a paper sumrnarising much of the early work on the oxidation of hydrazine, KIRIC AND BROWNER pointed out a number of common features, and two main classes of oxidising agent were distinguished : (i) those capable of producing hydrazoic acid, ammonia and nitrogen, (ii) those incapable of producing hydrazoic acid. Ferric. iron in acid solution oxidises hydrazine and belongs to the second class, and the overall equation may be written: NsHs+ + e + NH4+ + &N2 + H+ The stoichiometry’ of the reaction (i.e. the change in ‘equivalent concentration of oxidising agent divided by the corresponding change in molecular concentration of hydrazine) is approximately equal. to 1.3. CUY~ proposed that the reaction ‘proceeded via the intermediate N,H,, the hydrazyl radical, which he termed an odd molecule1 ; it contains an uneven number of electrons. KIRIC AND BROWNER suggested that the following hypothetical mechanism was adequate to account for the products obtained with class (ii) oxidising agents.’
N2E_I4+% N4Ho
NaH3 + &N4Ho (HaNo NE-IvNH* NHz) zNzHs e Jy4H0 + zNH3 + Nz
‘(’ cquiv’)+
HgN-
N:N*
NH2
+
N2H4
+
N2
J. Clwomalog.,
4 (1960) lg6ko5’
QCANTITATIVE INORGANIC CHROMATOGRAPHY. VI.
I97
Two independent kinetic investigations 495 of, the slow reaction between iron(II1) and hydrazine in ,acid’ solution have been .carried, out ,to elucidate the mechanism, of the -reaction, but slightly different conclusions were. reached. CAHN AND, .P,OWELLQ conclude that the reaction ,. ‘, competes
zNzH3
with the reaction
-+ N&b
+
N&b
2N2H3 + NdHo
whilst HIGGINSON AND WRIGHTS believe the disproportionation reaction minor importance. The reaction mechanisms proposed by the two,groups of workers are: ., Fe(II1) Fe(I1)
3; N2H4 A-+
Fe(I1)
2N2H2 kr followed
by
A-+ z(1
by
'N&b
Fe(II1) followed
+ N&h
c---1)*
N4Ho
(2)
N4Ho + 2NH2 + N2 2Na&
followed
(1)
NaH3
j-
+ N2H3,k-l-+ Fe(II1)
to be of
by
N2H4 cquiv.)
N2H2
N2H2
(3)”
--f Na
rnpid
-j- N2H2 A+
+
Fe(I1)
+
N2H2
(4)
2(x ‘eq’uiv.).
rnpkl -+Na
.
HIGGINSON’ AND WRIGHTS measured the rate of production of ferrous iron calorimetrically with o-phenanthroline and, the ferric iron remaining spectrophotometrically as the sulphate complex at 303 m,u. CAHN AND POWELL* determined the overall stoichiometry of the reaction only; by employing a large excess of the oxidising agent, the ferrous iron produced was determined by titration with’potassium permanganate. The disagreement between the two mechanisms is because HIG~INSON AND . WRIGHTS state that the’ reaction (- I) must be taken into account. If this is so, ._ ., when hydrazyl radicals or ferrous ions are removed ‘from the reaction scheme, then reaction (- 1) may be neglected. CAHN AND POWELL* showed that cupric ‘copper increases the stoichiometry of the oxidation of hydrazine by ferric ‘iron to 3.9 and ascribe this to the reactions N2&+
cu(II),+
Cu(I): + Fe(II1)’
cu(I)
+ Cu(II)
since cupric copper dots not oxidise hydrazine tant, the rate of disappearance ‘of hydrazine
+NaH2
(54
+’ F&II)
(56)
alone. Thus, if reaction (- r) is impor-* : will be, considerably, increased in the
-,
;
* HIGGI~ON AND WRIGHT'S,~IX~ELII~~~ ""CA~-&'A~~D PdLvk~i's %ch$,nism b&y.
O”lY. ,,’
,,
‘\ .‘,
.
J; Clcromatog., 4, (19.60) Ig6-2,os
198
.-F. H. POLLARD, G.‘NICI
presence of Cu(II), since reaction. will decrease the stationary concentration of hydrazyl radicals and hence the extent of reaction (- I). Whilst the, present investigation was in progress, ROSSEINSKY~ reported that the rate of disappearance of hydrazine was increased by the presence of Cu( II), and thus independent evidence for the importance of reaction (- I) has been obtained. EXPERIMENTAL All water was redistilled and stored in Jena glass aspirators. 99 y0 hydrazine hydrate (B.D.H.) was standardised by titration with potassium iodate solution using the normal Andrew’s conditions. Stock hydrazine, solutions, of varying molarities were prepared by dilution of the hydrazine hydrate with either I M hydrochloric acid or 0.5 M sulphuric acid. Iron(I1) chloride and iron(I1) ‘sulphate solutions were prepared by dissolving weighed amounts of pure iron wire in excess of hydrochloric or sulphuric acids respectively. Iron(II1) chloride or iron(II1) sulphate solutions were prepared by oxidation of the corresponding iron( II) solutions with 100: volume hydrogen peroxide, the excess peroxide was destroyed by gentle heating for some hours. Iron(I1) or iron(II1) cbncentrations were determined in these stock solutions by titration ,with potassium dichromate in dilute sulphuric acid or gravimetrically as the oxide respectively. The iron solutions were outgassed with nitrogen and sealed until used. All other stock solutions ,were prepared from AnalaR reagents where possible, and the concentrations determined by appropriate methods.. In the kinetic experiments in. chloride solutions, a constant initial hydrogen ion concentration was essential, and it was necessary to know the second dissociation constant of hydrazine, ” ,’ CNiib+l CH+l K2 = .’ ; [N2H02f] since the initial hydrazine concentration was varied from 0.242-0.024 M in the investigation.’ The value of I<, was assumed to be 6 & 2 moles 1-l (see HIGGINSON +ND ‘WRIGYT~). The initial hydrogen ion concentration in sulphate solutions was calculated by assuming that’ sulphuric acid was ,,completely dissociated. This is not true, but the variation of. [C, for sulphuric acid with ionic strength of the solution is such as to render the calculation of the exact hydrogen ion concentration very difficult’. The’ reactions‘ were carri&l ‘out in a thermostat at either 25” & ,o.IO, or 50: & 0.25~. A ‘mixture of all reagents except the appropriate hydrazine stock solution was outgassed with nitrogen and’allowed to come to e+ilibrium in the thermostat. The re,action was started by addition ‘of the hydrazine. solution. Samples were withdrawn for ‘ammonia and,iron analyses within a few seconds of mixing. As the reaction proceeded; Samples were removed-for analysis when convenient, i.e. for reactions at 25” /every 2 hours, and&very .hour when’ at ‘50”. O, : ” ; .’ ’ Ammonia was estimated in the following manner. In one of the samples,’ hydrazine was destroyed by addition of a small excess &acid ‘potassium iodatb, and., the’ iodine l
J. CkYo?naiog.,4 (1cy50)
196-205
QUANTITATIVE INORGANIC~CHROMATOGRAPHY. VI.
r’99
produced was destroyed by sodium sulphite. Excess sulphite, was destroyed,by- beiling. The,, solution.was transferred to a,MARHHAM, still8 (with ground~glass joints) and,made alkaline with 19 N caustic soda.solution,. The ammpnia was distilled. off into, z Oh w/v boric. acid solution and titrated with 0.~1 M hydrochloric’ acid #using a mixed methyl red7methylene blue indicatprfl. . :. 1 . I “. I, : ,. , Iron(U) and iron(II1) were estimated using the anion-exchange and calorimetric procedure described by POLLARD, MCOMIE, NICKLESS AND HANSON~O. The paper chromatographic separationll was also employed but iron(II1) trailed backwards towards ‘the starting line, making quantitative r’csults impossible. The prob’able reason for this is that hydrazine accompanies the iron(II1) up, the’.paper.and iron(I1) is formed by continued reaction between the two ionic species. As’irdn(I1) has a low Xp value in the solvent used, it is left behind by the solvent, thus forming a long streak on the paper. In the anion-exchange procedure, however, hydra&e-(as N,&+) .,‘:, : ., is unadsorbed and accompanies the iron( .. The iron(I1) results were checked by measurement Cf the optical density at 510, rnp of the tris-o-phenanthroline complex in citrate buffer?. Iron(II1) results were checked by the measurement of change in optical d,ensity at 620 rnp of the Ferron complex at, pH 2.6 .13”Iron(II), and iron(II1) may be determined to an’ accuracy of & 2 o/o,and ammonia to within & 3 o/o, Plots of the iron(I1) and ammonia concentrations against time were made from these results and the gradients found at times given in the Tables below: ,”
‘R(NHi)’=
d [N&l
aJnd
dt
R(Fc)
=
d FWWI dl
were estimated to be within ‘A 5 % .’ ,* The rate of’disappeara,nce of ,hydrazine was observed by measuring the change in optical density at 455 m,6 of the azine formed between hydrazine and p-dimethylaminobenzaldehydel4~ is in dilute hydrochloric acid solution This method ,was considered more accurate than t,he .procedure described by ROSSEINSKY~, since iron( iron(II1) or copper did not interfere. RESULTS
AND
DISCUSSION
In agreement with earlier .mechanisms, hydrazine and the derived radicals are depicted as possessing zero charge. .Almbst certainly this is not the case, but there is no evidence necessitating that the radicals should react in a particular form. ‘The formulae used are merely to represent the oxidation state of the nitrogen atoms concerned and not, the detailed structure of the radical. Fig. I shows a plot of L loglo[WIWi
I :.
,,
,,,
.I_
.
.’
CNd%lo
[WWt
, [Fe(III)lo; .’
against time (in ,hours) for a reaction mixture composed and: [N,H,] = 0,094 M in chloride solutions at 257.
of [l?e(~I’II)], ‘= 0.100 M,
J.‘Ckromatog.,
,4.
(1960)
Ig6-205
F. H. POLLARD,
200 Egtxt
G. .NICKLESS
of ftwYOzcS~ iron
When increasing amounts of ferrous chloride are added to such a reaction mixture as-above, cdeviations from bimole~ularity~ are ‘im’mediately apparent. As the cdncentration of,ferrols iron is increased, so the deviktion increases, viz. the reaction becomes is > 3, the reaction almost ceases slower, indeed if the ratio [l?e(II)]Of[Fe(III)], ,’ (see Fig. I).
Fig; I; The effect of time and added salts upon the initial reaction rate in chloride solutions at
Fig. 2. The effect: of time upon the initial reaction rate in sulphate solutions at ~5~.
29.
When copper(I1) chloride is added to the reaction solutions, then the deviation from bimolecularity is less than when copper is absent ‘(see Fig. ‘I). As the copper concentration is increased, the deviation at any given time becomes less than when copper(I1) is absent. This effect is probably due to co@per(II) removing hydrazyl rkdicals from the reaction scheme (see eqn’gaj faster than does ir~~(III~. .: ,’ ReacZio~~vdocity The variation in velocity constant of’reaction k- X) may be neglected, is given in Table I. TABLE Time
VeIo&y
Ii No Fe (II)
0.005
(I)
with time, assuming that reaction
I k;‘fh
M I?c(lf)
I moles-* -Ii-*) l
0.050
M Re(IZ)
0.100
1cfFe(IZ)
4 6
0.197 0.197
0.180
0.12Q
-
0.196
0.137
0.1x3’
0.075’
8~
o-173 o-159 o*r47
0.129 O.?s+Z
0.109 0.10s 0.102
-_ 0.068
o-057 .-
o-043 -
IO 12
24 36 48
0.105 0.083
‘0.073.
.-.
0.1~7 0.086 ” 1,‘.
J. Chroialog.,
4 (IgGo) Ig6-zo5
QUANTITATIVE1NORGANiCCHROMATOGRAPHV. VI.
,20x
HIGGINSON AND WRIGHTS quote a value of 12, = 0.06.1. moles -1 h-l for low ferrous iron concentrations at 25” ; unfortunately this has very little meaning unless the ferrous iron concentration is also quoted; The results for sulphate-containing solutions are very similar and a plot of l
log10 CWI~Wt l?W<
CNd510
[Fe(III)lo
against time is given in Fig. 2. Retardation by ferrous iron is much greater than in chloride solutions. ‘The velocity of reaction (I), kl varies from 1.36 1 moles-l h-l with no iron(I1) present to 0.16 l*moles- l* h-l after 48 h (i.e. 20 o/o reaction). l
Qzcantitative investigatiout
of the reaction
l
mechanism
At 25O the ammonia analyses are insufficiently accurate, over the range of stoichiometries investigated, to permit an unambiguous kinetic interpretation. Consequently, the temperature was raised to 50” to enable the reaction to proceed at a faster rate. From $he proposed reaction mechanism, and with Cu(I1) absent R(NHs) = R(Fe)
= 81 [Fe(III)] [NaH4]--R-l
h2 [N2Hs12
(6)
[Fe(II)l [N&Is] i- AsCNdG12 -I- 1~4 CFe(III)I
m2HSl
(7)
‘,
where R(NH3) = d [NHJdt; and R(Fe) = d [Fe(II)]/di. Making the normal stationary state assumption that o = d [NaHa]/dt = Izl [Fe(III)] [NzH~] - k-1 [Fe(II)] [NzHs] k’CNzHs12- k4 CNaHd CFe(III)l Subtract (8) from (7) ’ R(Fe)
=
?z2
/ZZ[NzH3]2
[NsHs]2 + aks [NBHs]~+ 2k&Fe(IIIJ] [NsHe]
(8)
(9)
Divide (9) by (6) A!(
Eliminating
Fe)/R(NHs) = r. +
zh/h
+
2k4 [Fe(IIi)]/ka
(10)
[Nib]
[N,,H,] by means of (6) R(Fe)lR(NHs)
L= r + 2ks/& + 2k4 [Fe(IH)1/2/ksR(NH3)
(11)
Thus a plot of R(Fe)/R(NH,) against [Fe(III)]/dm) should be a straight line of intercept I’ + 2ks/k,. The intercept is, nearly unity, and was evaluated by rewriting (II) in the form R(Fe)
CWWI
dR(NH3) ‘,
4 plot of the left-h:a,nd side ,of eqn. (12) against dF(Nl&)/[Fe(III)]
is shown in Fig; 3 for exl%5ments at 500 in chloride solutions, ,the data for &hich,are su&narised ,.: ‘, ,. in’ Table II.‘. All’ rates’ ‘&em measured‘ by &.&hg the &i,n&nt ti ‘he’ &r&$&&n I, ‘I( /.\ vB1cszcs time curves at t =’ 0.’ ‘, :
J. Chvomalog.,:4
(1960)
196-205
I?.H. POLLARD,
202
G. NlCI
,TABLISII CHLORIDE,SOLUTIONS
AT
50”
[l?e(II)J, = 0 Ex~cvimenl
No.
[I;% (III)
~/R(NH,)
/
moles d-’
[Fc(ZZZ)
I
0.100
0.094'
z
0.100
o-047
0.0194 0.0125
0.0187 0.012G
0.019
0.0047
0.0048
4 :
0.098
0.050 d.020 0.050
0.098 0.049'
0.0116 o.oogG to.0063
Ry Fe) J
[Fe(ZIT.]
I.412
I.379 0.122 0.691 2.150 3.606 .11.583
0.0115 0.0052 0.0051
d R(NH.)
1.100 0.690 2.165 3.632 1';606
-.
Fig. 3 gives a gradient
of I + z&J&
=
1.025=t=o.oz,whence 0.010.
3.04 - .oP ‘iirr b t =?7 &l.O-
1
XI
I OO Fig. 3. The
test of equation (12) in chloride solutions at 5o",
The results for sulphate solutions summarised in Table III.
Fig. 4, The
TARLE
rnoles.l-’
moles 4-’
, I 2“' 3 4
Whence
: 0.100 0.100,
0.0268
0.050 0.020
0.0134 0.0268 0.0268
o
R(NZ3.J
moles4-~~h-~
..
from experiments
50°
R(Fe>
CNsHt /
’
III
[Fe(II)], = [Fe(ZZZ) /
20
test of equation (12) in sulphatc solutions at 50".
at 50” are shown in Fig.,4
SULPHATESOLUTIONSAT
Expevitneni No.
I
d&
~WNW
R (I7e) --[Fe(ZZI)JZ/R(NHIj
tnoles~F~h-~
[FtS(ZZZ~l
0.0062 0.0088, ,
0.0052
0.721
1.250
0.0048 0.0020
0.0027 0.0030 0.0013
0.519 1.096 I.970
11083 1.65d 2.560
,’ ,:
a.; ’
Its@
=-'o.oI~
f
0.015
at 50~.
J. Ckromatog.,
4 (1960) x96-205
QUANTITATWEINORGANICCHRGMATOGRAPWY. VI.
203
The linear plots obtained,- provide good evidence !hat the reaction by HIGGINSONAND %kW-rT6 is valid.
scheme proposed
Efect
of high come&atiom
of fewws
Goi
At high concentrations of ferrous iron, the rate of disappearance of hydrazyl radicals by reactions (z), (3) and (4) will be negligible as compared to back reaction (- I). Then as an approximation, if K = k;jkz
Frqm (13) and (6)
CNd-hl = K [Fe( III)] [N2H~]/[I?o(II)]
(13)
R(NH3) = kd3? CFe(III)-js[NsH4]s/[Fq( II) Js
(14)
from (13) and (IO) ~(Fe)/R(NH3) = r + z&/&z +
2h4
tN&W
CF+I)1/=2
(15)
Thus for equation (14) a plot of R(NHa) jersey [Fe(III)J2 [N2H,J2/[Fc(II)]2 will. be a straight line passing through the origin. Similarly for equation (15) a plot of R(Fe)/R(NH,) against [Fe(II)]/[N2H4] .should be ‘a straight line, of intercept ” I + zk,/k,, ‘whence 12,/12, may,be deduced. Plots for equations (14) and (IS) are shown in. Figs. 5 and 6 respectively, from experiments at ‘50” in Chloride solutions, which are summarised in Table IV.
4.010.0 3.0 -
8.0 -
Fig. 5. The test of equation (14) ‘in chloride solutions at 5o”; where reaction .(- I) S> reactions (2, 3, or4).
Fig.
6. The test of equation (15) in chloride solutions at 50~.
From Fig.. 6, lz,/lz2 = 0.025 f o.org which is in excellent agreement with the value of k3/k2 calculated using equation (x2)! It was thought that the linearity of ,these plots, i.e. ,Figs. 5 and 6, gave. additional
evidence for the validity of the proposed reaction mechanism. E&pt ,of cq5rz’c co@er Stoickiometry. It, has already been’ shown ‘ (Fig. I), that .the’ presence of cppper(I1) tends to keep the deviations from bimolecularity smaller ,than if copper(I1). is absent. J, ClWomalog.,4 (x960) rgG-205
F. R. POLLARD,
204
G. NICKLESS
To account for this, it has :been postulated* that copper(I1) removes the hydrazyl radicals through equation @a), faster than does iron(II1) by reaction (4). Hence’, the amount of 4-electron oxidation must increase i.e. the stoichiometry when copper(I1) is absent must be less than when copper(I1) is present. &LIZ
IV
CHLORIDESOLUTIONS
Bx#&mctll
I
2 3
: 87
9, IO II 12 (13
[FC=(IIZ) ] nroles4-1
0.0500 0.100 0.0800 0.0800~ 0.0400 0.200 0.100 0.100
0.050 0.0800 0.200 0.200 0.200
[I?e(III)J’[NpHt]a
fFe(ZI) I nwlcs .I-’
ix7
0.0480 0.0966 0,112s 0.1128
0.1128 o.og4.0 0.0980 0.0980 0.0113
AT
”
0.1594 o.og40 0,0752 0.0752
. -1 J
0.0973 0.0438 0.2425 0.1212 0.2425 0.0485 0.1488 0.1212 0.1950 0.1950 0.0243 0.0243 9.0486
[Pe(IZ) J* molcs’*l-’
R(NH,) moles +~*lr-l
0.0103 0.00258 0.02958 0.00739 0.00739 0.0106 0.0220 0.0153 o.org1 0.0109
SO0
0.00301 0.00083 0.0100 0.0020
'
0.00265 0.00417 0.01670
0.0023 0.00334 o.oo6go 0.00505 0.00625 0.00350 0.000850 o.oor4o 0.00543
R (Fe) nroles4-l~Jr~
R(l;e)
[Fe(Iqi]
R (NH,)
CN,lf,J
0.00425
0.00225 0.01426 0.003 70 0.00340 o.oogog 0.0110 0.00863 0.0102
0.00600 d.00365 0.00508 0.0124
1.41 2.70 I.43 I .84 I .36 2.72 1.60 1.71 1.64 1.71 4.29 3.63 2.26
0.492 I.970 0.465. 0.93 f o-465 1.940 ~~644 0.755 0.750 0.771 3.86 3.09 * :54
Reaction mixtures containing ferric chloride and hydrazine Exj5erimental. dihydrochloride in hydrochloric acid were prepared, the ferric chloride being in large excess. The mixtures were heated for 4 h at go”, and the resultant ferrous iron determined by titration with potassium dichromate in the normal manner. IThis procedure was repeated with copper(I1) chloride also present ; again the ferrous iron concentration was determined by titration against potassium dichromate. Results Staichziometry.
(I)
Copper(I1)
absent.
1.25; 1.29;
1.28. (2) Copper(I1)
present
3.94,
3.90, 3-91. Rate of disappearance of Itydrazine. If the mechanism favoured by HIGGINSON AND WR~G& is correct, the presence of copper(I1) should increase d [N,H,]/dt by reducing the extent ,of the back reaction (- I). d, [N2H,J/dt has been measured. in the ‘absence of copper( and with different amounts of copper(I1) present. ‘The results are summarised in Table V. The resultsprovide clear evidence for’ the large catalytic effect ,ofs copper( and hence the back reaction (- I) must be appreciable;- since in ‘similar experiments with iron(II1) absent; copper(U) did not oxidise hydrazine.’ : 1 -’ Initial rates are similar in all the’ experiments, since in the early stages of. the reaction, the. back reaction (- I), is negligible. As the reaction proceeds, and iron accumulates,, tee ‘difference in the ,rates in the presence ,and &s&c& of’ co&r(II) iricreases;‘:which- is in : accordance’ with ‘ROSSEINSI
QUANTITATIVE
INORGANIC
CIXROMATOGRAPWY.
TABLE \~ALUES OF d [N,H,] /dt (moles 1.00 M, initial [N,H,]
WC11 =
l
l- ’ ’ h-1) AT VARIOUS MOLAR CONCENTRATIONS OF COPPlZR(II) = 0.047 M, initial [l?e(II.I)] = 0.100 M. Temperature = 50~ ldt (moks.l-’
.h-1)
~o~,[N~H,J(ymks~E-~)
0
0.025 0.050 o-075
205
V
d [N,HJ
I 2
VI.
4.00
3.75
0.015
0.013 0.013 0.013 0.014
0.015 0.015 0.015
3.00
3.50
0.0084 0.0084 0.0087 ,o.oogo
. 3.50
0.002g 0;0038 0.0043 0.0054
0.0047 0.005’5 0.0065 0.0070
3.00
0.0020
0.0028 0.0034 0.0040
r-75
0.0012
0.0024 0.0030 0.0035
SUMMARY
The mechanism proposed by HIGGINSON AND WRIGHT for the slow reaction between fe,r+c irqn an&hydra+ne in acid aque,ous solutions hasbeen confirmed, using an anion, exchange prdqedure for ‘the separation and quantitative estimation ,of ferrous and ferric iron. ,, REFERENCES
‘.
1 W. C. E.’ HIGGINS&,’ Recent Aspects of, the Inorganic Chemistvy of Nilvogen, Publ., No,. ,1o. (1957) 95. ” 8 R. E. KIRK AND A. W. BROWNB, J. Am. Chem. Soc., 59 (1928) 337. : 8 E. J; CVY, J. Am.‘Chem. ‘SOL, 46 (1924) 1810.
Chem.‘Soc., 76 (1954) 2568. Chem. Sac., (1955) ,1551. R. ROSSIEINSKY, J.‘Chem. SOIL; (1957) 4685. C. BRAY AND H. A. LIEBI-IAFSKY, J. Am. Chem. SOL, 57 (1935) 51. MARKHAM, Biockem. J., 36 (1942) ,790. J.’ CONWAY, Micro-di@sion Methods and ,VoZumetvic EWOYS, 2nd Ed., Crosby
4 J. W.
CAI-IN, AND
R. E.
6 W. C. E. HIGGINSON 6 D.
7 W. 8 R.
9 E:
Chem. So& Spec.
AND
POWELL, J. Am. P. WRIGHT, J.
Sons, London, 1947. ’ i0 F.. H. POLLARD, J. F. w. iI F.’ H.
POLLARD, W. BRANDT
MCOMIE,
J; F. W. McOMIE. AND G. F.,SMITH,
G. &(~ICKLESSAND P. HANSON, A.
J. BANISTER
AND
Lockwood
‘&
J. Chvomatog., 4 (1960) 108. Analyst, 82 (1937) 780.
G. NICKLESS,
12 W. Anal. Chem,, 21 (1949) 1313. 1s H. W. SWANK AND M. G. MELLON, Ind. Ercg. Chem., Anal. Ed., g (1937) 406. 14 G. W. WhTT AND J. D. CRISP, Anal. Chem., 24 (1952) 2006. lb P: R. WOOD, Anal. Chem., 25 (1953) 1879.
J. Chvomatog., 4 (1960) 196-205
196
QUANTITATIVE
INdRGAkC
CHROMATOGRAPtiY
PAl$T VI. APPLICATION TO KINETIC STUDIES: THE OXIDATION OF HYDRAZINE BY FERRIC IRdti IN ACID AQUEOUS SOLUTIONS I?. H. POLLARD Depavlment
of Physical
and Inorganic (Received
AND
G. NICKLESS
Chemistry,
The University,
November
znd, 1959)
Brktol
(Great Bvilain)
The use of ion-exchange and/or paper chromatography to elucidate the mechanism of inorganic reactions has not, up to the present, been extensively used, though there are clearly fields where its application might do much to establish the correctness of postulates made from the classical physicochemical approach, or to throw light on hitherto unsuspected intermediates or products which would not be revealed by classical analytical procedures. A number of reactions have been studied to show how chromatography can help in this field, and this paper represents the first of such investigations. When hydrazine is oxidised in aqueous solution, the principal nitrogenous products are ammonia, nitrogen and hydrazoic acidl. The proportions of these products vary considerably with the oxidising agent and the conditions of the reaction. In a paper sumrnarising much of the early work on the oxidation of hydrazine, KIRIC AND BROWNER pointed out a number of common features, and two main classes of oxidising agent were distinguished : (i) those capable of producing hydrazoic acid, ammonia and nitrogen, (ii) those incapable of producing hydrazoic acid. Ferric. iron in acid solution oxidises hydrazine and belongs to the second class, and the overall equation may be written: NsHs+ + e + NH4+ + &N2 + H+ The stoichiometry’ of the reaction (i.e. the change in ‘equivalent concentration of oxidising agent divided by the corresponding change in molecular concentration of hydrazine) is approximately equal. to 1.3. CUY~ proposed that the reaction ‘proceeded via the intermediate N,H,, the hydrazyl radical, which he termed an odd molecule1 ; it contains an uneven number of electrons. KIRIC AND BROWNER suggested that the following hypothetical mechanism was adequate to account for the products obtained with class (ii) oxidising agents.’
N2E_I4+% N4Ho
NaH3 + &N4Ho (HaNo NE-IvNH* NHz) zNzHs e Jy4H0 + zNH3 + Nz
‘(’ cquiv’)+
HgN-
N:N*
NH2
+
N2H4
+
N2
J. Clwomalog.,
4 (1960) lg6ko5’
QCANTITATIVE INORGANIC CHROMATOGRAPHY. VI.
I97
Two independent kinetic investigations 495 of, the slow reaction between iron(II1) and hydrazine in ,acid’ solution have been .carried, out ,to elucidate the mechanism, of the -reaction, but slightly different conclusions were. reached. CAHN AND, .P,OWELLQ conclude that the reaction ,. ‘, competes
zNzH3
with the reaction
-+ N&b
+
N&b
2N2H3 + NdHo
whilst HIGGINSON AND WRIGHTS believe the disproportionation reaction minor importance. The reaction mechanisms proposed by the two,groups of workers are: ., Fe(II1) Fe(I1)
3; N2H4 A-+
Fe(I1)
2N2H2 kr followed
by
A-+ z(1
by
'N&b
Fe(II1) followed
+ N&h
c---1)*
N4Ho
(2)
N4Ho + 2NH2 + N2 2Na&
followed
(1)
NaH3
j-
+ N2H3,k-l-+ Fe(II1)
to be of
by
N2H4 cquiv.)
N2H2
N2H2
(3)”
--f Na
rnpid
-j- N2H2 A+
+
Fe(I1)
+
N2H2
(4)
2(x ‘eq’uiv.).
rnpkl -+Na
.
HIGGINSON’ AND WRIGHTS measured the rate of production of ferrous iron calorimetrically with o-phenanthroline and, the ferric iron remaining spectrophotometrically as the sulphate complex at 303 m,u. CAHN AND POWELL* determined the overall stoichiometry of the reaction only; by employing a large excess of the oxidising agent, the ferrous iron produced was determined by titration with’potassium permanganate. The disagreement between the two mechanisms is because HIG~INSON AND . WRIGHTS state that the’ reaction (- I) must be taken into account. If this is so, ._ ., when hydrazyl radicals or ferrous ions are removed ‘from the reaction scheme, then reaction (- 1) may be neglected. CAHN AND POWELL* showed that cupric ‘copper increases the stoichiometry of the oxidation of hydrazine by ferric ‘iron to 3.9 and ascribe this to the reactions N2&+
cu(II),+
Cu(I): + Fe(II1)’
cu(I)
+ Cu(II)
since cupric copper dots not oxidise hydrazine tant, the rate of disappearance ‘of hydrazine
+NaH2
(54
+’ F&II)
(56)
alone. Thus, if reaction (- r) is impor-* : will be, considerably, increased in the
-,
;
* HIGGI~ON AND WRIGHT'S,~IX~ELII~~~ ""CA~-&'A~~D PdLvk~i's %ch$,nism b&y.
O”lY. ,,’
,,
‘\ .‘,
.
J; Clcromatog., 4, (19.60) Ig6-2,os
198
.-F. H. POLLARD, G.‘NICI
presence of Cu(II), since reaction. will decrease the stationary concentration of hydrazyl radicals and hence the extent of reaction (- I). Whilst the, present investigation was in progress, ROSSEINSKY~ reported that the rate of disappearance of hydrazine was increased by the presence of Cu( II), and thus independent evidence for the importance of reaction (- I) has been obtained. EXPERIMENTAL All water was redistilled and stored in Jena glass aspirators. 99 y0 hydrazine hydrate (B.D.H.) was standardised by titration with potassium iodate solution using the normal Andrew’s conditions. Stock hydrazine, solutions, of varying molarities were prepared by dilution of the hydrazine hydrate with either I M hydrochloric acid or 0.5 M sulphuric acid. Iron(I1) chloride and iron(I1) ‘sulphate solutions were prepared by dissolving weighed amounts of pure iron wire in excess of hydrochloric or sulphuric acids respectively. Iron(II1) chloride or iron(II1) sulphate solutions were prepared by oxidation of the corresponding iron( II) solutions with 100: volume hydrogen peroxide, the excess peroxide was destroyed by gentle heating for some hours. Iron(I1) or iron(II1) cbncentrations were determined in these stock solutions by titration ,with potassium dichromate in dilute sulphuric acid or gravimetrically as the oxide respectively. The iron solutions were outgassed with nitrogen and sealed until used. All other stock solutions ,were prepared from AnalaR reagents where possible, and the concentrations determined by appropriate methods.. In the kinetic experiments in. chloride solutions, a constant initial hydrogen ion concentration was essential, and it was necessary to know the second dissociation constant of hydrazine, ” ,’ CNiib+l CH+l K2 = .’ ; [N2H02f] since the initial hydrazine concentration was varied from 0.242-0.024 M in the investigation.’ The value of I<, was assumed to be 6 & 2 moles 1-l (see HIGGINSON +ND ‘WRIGYT~). The initial hydrogen ion concentration in sulphate solutions was calculated by assuming that’ sulphuric acid was ,,completely dissociated. This is not true, but the variation of. [C, for sulphuric acid with ionic strength of the solution is such as to render the calculation of the exact hydrogen ion concentration very difficult’. The’ reactions‘ were carri&l ‘out in a thermostat at either 25” & ,o.IO, or 50: & 0.25~. A ‘mixture of all reagents except the appropriate hydrazine stock solution was outgassed with nitrogen and’allowed to come to e+ilibrium in the thermostat. The re,action was started by addition ‘of the hydrazine. solution. Samples were withdrawn for ‘ammonia and,iron analyses within a few seconds of mixing. As the reaction proceeded; Samples were removed-for analysis when convenient, i.e. for reactions at 25” /every 2 hours, and&very .hour when’ at ‘50”. O, : ” ; .’ ’ Ammonia was estimated in the following manner. In one of the samples,’ hydrazine was destroyed by addition of a small excess &acid ‘potassium iodatb, and., the’ iodine l
J. CkYo?naiog.,4 (1cy50)
196-205
QUANTITATIVE INORGANIC~CHROMATOGRAPHY. VI.
r’99
produced was destroyed by sodium sulphite. Excess sulphite, was destroyed,by- beiling. The,, solution.was transferred to a,MARHHAM, still8 (with ground~glass joints) and,made alkaline with 19 N caustic soda.solution,. The ammpnia was distilled. off into, z Oh w/v boric. acid solution and titrated with 0.~1 M hydrochloric’ acid #using a mixed methyl red7methylene blue indicatprfl. . :. 1 . I “. I, : ,. , Iron(U) and iron(II1) were estimated using the anion-exchange and calorimetric procedure described by POLLARD, MCOMIE, NICKLESS AND HANSON~O. The paper chromatographic separationll was also employed but iron(II1) trailed backwards towards ‘the starting line, making quantitative r’csults impossible. The prob’able reason for this is that hydrazine accompanies the iron(II1) up, the’.paper.and iron(I1) is formed by continued reaction between the two ionic species. As’irdn(I1) has a low Xp value in the solvent used, it is left behind by the solvent, thus forming a long streak on the paper. In the anion-exchange procedure, however, hydra&e-(as N,&+) .,‘:, : ., is unadsorbed and accompanies the iron( .. The iron(I1) results were checked by measurement Cf the optical density at 510, rnp of the tris-o-phenanthroline complex in citrate buffer?. Iron(II1) results were checked by the measurement of change in optical d,ensity at 620 rnp of the Ferron complex at, pH 2.6 .13”Iron(II), and iron(II1) may be determined to an’ accuracy of & 2 o/o,and ammonia to within & 3 o/o, Plots of the iron(I1) and ammonia concentrations against time were made from these results and the gradients found at times given in the Tables below: ,”
‘R(NHi)’=
d [N&l
aJnd
dt
R(Fc)
=
d FWWI dl
were estimated to be within ‘A 5 % .’ ,* The rate of’disappeara,nce of ,hydrazine was observed by measuring the change in optical density at 455 m,6 of the azine formed between hydrazine and p-dimethylaminobenzaldehydel4~ is in dilute hydrochloric acid solution This method ,was considered more accurate than t,he .procedure described by ROSSEINSKY~, since iron( iron(II1) or copper did not interfere. RESULTS
AND
DISCUSSION
In agreement with earlier .mechanisms, hydrazine and the derived radicals are depicted as possessing zero charge. .Almbst certainly this is not the case, but there is no evidence necessitating that the radicals should react in a particular form. ‘The formulae used are merely to represent the oxidation state of the nitrogen atoms concerned and not, the detailed structure of the radical. Fig. I shows a plot of L loglo[WIWi
I :.
,,
,,,
.I_
.
.’
CNd%lo
[WWt
, [Fe(III)lo; .’
against time (in ,hours) for a reaction mixture composed and: [N,H,] = 0,094 M in chloride solutions at 257.
of [l?e(~I’II)], ‘= 0.100 M,
J.‘Ckromatog.,
,4.
(1960)
Ig6-205
F. H. POLLARD,
200 Egtxt
G. .NICKLESS
of ftwYOzcS~ iron
When increasing amounts of ferrous chloride are added to such a reaction mixture as-above, cdeviations from bimole~ularity~ are ‘im’mediately apparent. As the cdncentration of,ferrols iron is increased, so the deviktion increases, viz. the reaction becomes is > 3, the reaction almost ceases slower, indeed if the ratio [l?e(II)]Of[Fe(III)], ,’ (see Fig. I).
Fig; I; The effect of time and added salts upon the initial reaction rate in chloride solutions at
Fig. 2. The effect: of time upon the initial reaction rate in sulphate solutions at ~5~.
29.
When copper(I1) chloride is added to the reaction solutions, then the deviation from bimolecularity is less than when copper is absent ‘(see Fig. ‘I). As the copper concentration is increased, the deviation at any given time becomes less than when copper(I1) is absent. This effect is probably due to co@per(II) removing hydrazyl rkdicals from the reaction scheme (see eqn’gaj faster than does ir~~(III~. .: ,’ ReacZio~~vdocity The variation in velocity constant of’reaction k- X) may be neglected, is given in Table I. TABLE Time
VeIo&y
Ii No Fe (II)
0.005
(I)
with time, assuming that reaction
I k;‘fh
M I?c(lf)
I moles-* -Ii-*) l
0.050
M Re(IZ)
0.100
1cfFe(IZ)
4 6
0.197 0.197
0.180
0.12Q
-
0.196
0.137
0.1x3’
0.075’
8~
o-173 o-159 o*r47
0.129 O.?s+Z
0.109 0.10s 0.102
-_ 0.068
o-057 .-
o-043 -
IO 12
24 36 48
0.105 0.083
‘0.073.
.-.
0.1~7 0.086 ” 1,‘.
J. Chroialog.,
4 (IgGo) Ig6-zo5
QUANTITATIVE1NORGANiCCHROMATOGRAPHV. VI.
,20x
HIGGINSON AND WRIGHTS quote a value of 12, = 0.06.1. moles -1 h-l for low ferrous iron concentrations at 25” ; unfortunately this has very little meaning unless the ferrous iron concentration is also quoted; The results for sulphate-containing solutions are very similar and a plot of l
log10 CWI~Wt l?W<
CNd510
[Fe(III)lo
against time is given in Fig. 2. Retardation by ferrous iron is much greater than in chloride solutions. ‘The velocity of reaction (I), kl varies from 1.36 1 moles-l h-l with no iron(I1) present to 0.16 l*moles- l* h-l after 48 h (i.e. 20 o/o reaction). l
Qzcantitative investigatiout
of the reaction
l
mechanism
At 25O the ammonia analyses are insufficiently accurate, over the range of stoichiometries investigated, to permit an unambiguous kinetic interpretation. Consequently, the temperature was raised to 50” to enable the reaction to proceed at a faster rate. From $he proposed reaction mechanism, and with Cu(I1) absent R(NHs) = R(Fe)
= 81 [Fe(III)] [NaH4]--R-l
h2 [N2Hs12
(6)
[Fe(II)l [N&Is] i- AsCNdG12 -I- 1~4 CFe(III)I
m2HSl
(7)
‘,
where R(NH3) = d [NHJdt; and R(Fe) = d [Fe(II)]/di. Making the normal stationary state assumption that o = d [NaHa]/dt = Izl [Fe(III)] [NzH~] - k-1 [Fe(II)] [NzHs] k’CNzHs12- k4 CNaHd CFe(III)l Subtract (8) from (7) ’ R(Fe)
=
?z2
/ZZ[NzH3]2
[NsHs]2 + aks [NBHs]~+ 2k&Fe(IIIJ] [NsHe]
(8)
(9)
Divide (9) by (6) A!(
Eliminating
Fe)/R(NHs) = r. +
zh/h
+
2k4 [Fe(IIi)]/ka
(10)
[Nib]
[N,,H,] by means of (6) R(Fe)lR(NHs)
L= r + 2ks/& + 2k4 [Fe(IH)1/2/ksR(NH3)
(11)
Thus a plot of R(Fe)/R(NH,) against [Fe(III)]/dm) should be a straight line of intercept I’ + 2ks/k,. The intercept is, nearly unity, and was evaluated by rewriting (II) in the form R(Fe)
CWWI
dR(NH3) ‘,
4 plot of the left-h:a,nd side ,of eqn. (12) against dF(Nl&)/[Fe(III)]
is shown in Fig; 3 for exl%5ments at 500 in chloride solutions, ,the data for &hich,are su&narised ,.: ‘, ,. in’ Table II.‘. All’ rates’ ‘&em measured‘ by &.&hg the &i,n&nt ti ‘he’ &r&$&&n I, ‘I( /.\ vB1cszcs time curves at t =’ 0.’ ‘, :
J. Chvomalog.,:4
(1960)
196-205
I?.H. POLLARD,
202
G. NlCI
,TABLISII CHLORIDE,SOLUTIONS
AT
50”
[l?e(II)J, = 0 Ex~cvimenl
No.
[I;% (III)
~/R(NH,)
/
moles d-’
[Fc(ZZZ)
I
0.100
0.094'
z
0.100
o-047
0.0194 0.0125
0.0187 0.012G
0.019
0.0047
0.0048
4 :
0.098
0.050 d.020 0.050
0.098 0.049'
0.0116 o.oogG to.0063
Ry Fe) J
[Fe(ZIT.]
I.412
I.379 0.122 0.691 2.150 3.606 .11.583
0.0115 0.0052 0.0051
d R(NH.)
1.100 0.690 2.165 3.632 1';606
-.
Fig. 3 gives a gradient
of I + z&J&
=
1.025=t=o.oz,whence 0.010.
3.04 - .oP ‘iirr b t =?7 &l.O-
1
XI
I OO Fig. 3. The
test of equation (12) in chloride solutions at 5o",
The results for sulphate solutions summarised in Table III.
Fig. 4, The
TARLE
rnoles.l-’
moles 4-’
, I 2“' 3 4
Whence
: 0.100 0.100,
0.0268
0.050 0.020
0.0134 0.0268 0.0268
o
R(NZ3.J
moles4-~~h-~
..
from experiments
50°
R(Fe>
CNsHt /
’
III
[Fe(II)], = [Fe(ZZZ) /
20
test of equation (12) in sulphatc solutions at 50".
at 50” are shown in Fig.,4
SULPHATESOLUTIONSAT
Expevitneni No.
I
d&
~WNW
R (I7e) --[Fe(ZZI)JZ/R(NHIj
tnoles~F~h-~
[FtS(ZZZ~l
0.0062 0.0088, ,
0.0052
0.721
1.250
0.0048 0.0020
0.0027 0.0030 0.0013
0.519 1.096 I.970
11083 1.65d 2.560
,’ ,:
a.; ’
Its@
=-'o.oI~
f
0.015
at 50~.
J. Ckromatog.,
4 (1960) x96-205
QUANTITATWEINORGANICCHRGMATOGRAPWY. VI.
203
The linear plots obtained,- provide good evidence !hat the reaction by HIGGINSONAND %kW-rT6 is valid.
scheme proposed
Efect
of high come&atiom
of fewws
Goi
At high concentrations of ferrous iron, the rate of disappearance of hydrazyl radicals by reactions (z), (3) and (4) will be negligible as compared to back reaction (- I). Then as an approximation, if K = k;jkz
Frqm (13) and (6)
CNd-hl = K [Fe( III)] [N2H~]/[I?o(II)]
(13)
R(NH3) = kd3? CFe(III)-js[NsH4]s/[Fq( II) Js
(14)
from (13) and (IO) ~(Fe)/R(NH3) = r + z&/&z +
2h4
tN&W
CF+I)1/=2
(15)
Thus for equation (14) a plot of R(NHa) jersey [Fe(III)J2 [N2H,J2/[Fc(II)]2 will. be a straight line passing through the origin. Similarly for equation (15) a plot of R(Fe)/R(NH,) against [Fe(II)]/[N2H4] .should be ‘a straight line, of intercept ” I + zk,/k,, ‘whence 12,/12, may,be deduced. Plots for equations (14) and (IS) are shown in. Figs. 5 and 6 respectively, from experiments at ‘50” in Chloride solutions, which are summarised in Table IV.
4.010.0 3.0 -
8.0 -
Fig. 5. The test of equation (14) ‘in chloride solutions at 5o”; where reaction .(- I) S> reactions (2, 3, or4).
Fig.
6. The test of equation (15) in chloride solutions at 50~.
From Fig.. 6, lz,/lz2 = 0.025 f o.org which is in excellent agreement with the value of k3/k2 calculated using equation (x2)! It was thought that the linearity of ,these plots, i.e. ,Figs. 5 and 6, gave. additional
evidence for the validity of the proposed reaction mechanism. E&pt ,of cq5rz’c co@er Stoickiometry. It, has already been’ shown ‘ (Fig. I), that .the’ presence of cppper(I1) tends to keep the deviations from bimolecularity smaller ,than if copper(I1). is absent. J, ClWomalog.,4 (x960) rgG-205
F. R. POLLARD,
204
G. NICKLESS
To account for this, it has :been postulated* that copper(I1) removes the hydrazyl radicals through equation @a), faster than does iron(II1) by reaction (4). Hence’, the amount of 4-electron oxidation must increase i.e. the stoichiometry when copper(I1) is absent must be less than when copper(I1) is present. &LIZ
IV
CHLORIDESOLUTIONS
Bx#&mctll
I
2 3
: 87
9, IO II 12 (13
[FC=(IIZ) ] nroles4-1
0.0500 0.100 0.0800 0.0800~ 0.0400 0.200 0.100 0.100
0.050 0.0800 0.200 0.200 0.200
[I?e(III)J’[NpHt]a
fFe(ZI) I nwlcs .I-’
ix7
0.0480 0.0966 0,112s 0.1128
0.1128 o.og4.0 0.0980 0.0980 0.0113
AT
”
0.1594 o.og40 0,0752 0.0752
. -1 J
0.0973 0.0438 0.2425 0.1212 0.2425 0.0485 0.1488 0.1212 0.1950 0.1950 0.0243 0.0243 9.0486
[Pe(IZ) J* molcs’*l-’
R(NH,) moles +~*lr-l
0.0103 0.00258 0.02958 0.00739 0.00739 0.0106 0.0220 0.0153 o.org1 0.0109
SO0
0.00301 0.00083 0.0100 0.0020
'
0.00265 0.00417 0.01670
0.0023 0.00334 o.oo6go 0.00505 0.00625 0.00350 0.000850 o.oor4o 0.00543
R (Fe) nroles4-l~Jr~
R(l;e)
[Fe(Iqi]
R (NH,)
CN,lf,J
0.00425
0.00225 0.01426 0.003 70 0.00340 o.oogog 0.0110 0.00863 0.0102
0.00600 d.00365 0.00508 0.0124
1.41 2.70 I.43 I .84 I .36 2.72 1.60 1.71 1.64 1.71 4.29 3.63 2.26
0.492 I.970 0.465. 0.93 f o-465 1.940 ~~644 0.755 0.750 0.771 3.86 3.09 * :54
Reaction mixtures containing ferric chloride and hydrazine Exj5erimental. dihydrochloride in hydrochloric acid were prepared, the ferric chloride being in large excess. The mixtures were heated for 4 h at go”, and the resultant ferrous iron determined by titration with potassium dichromate in the normal manner. IThis procedure was repeated with copper(I1) chloride also present ; again the ferrous iron concentration was determined by titration against potassium dichromate. Results Staichziometry.
(I)
Copper(I1)
absent.
1.25; 1.29;
1.28. (2) Copper(I1)
present
3.94,
3.90, 3-91. Rate of disappearance of Itydrazine. If the mechanism favoured by HIGGINSON AND WR~G& is correct, the presence of copper(I1) should increase d [N,H,]/dt by reducing the extent ,of the back reaction (- I). d, [N2H,J/dt has been measured. in the ‘absence of copper( and with different amounts of copper(I1) present. ‘The results are summarised in Table V. The resultsprovide clear evidence for’ the large catalytic effect ,ofs copper( and hence the back reaction (- I) must be appreciable;- since in ‘similar experiments with iron(II1) absent; copper(U) did not oxidise hydrazine.’ : 1 -’ Initial rates are similar in all the’ experiments, since in the early stages of. the reaction, the. back reaction (- I), is negligible. As the reaction proceeds, and iron accumulates,, tee ‘difference in the ,rates in the presence ,and &s&c& of’ co&r(II) iricreases;‘:which- is in : accordance’ with ‘ROSSEINSI
QUANTITATIVE
INORGANIC
CIXROMATOGRAPWY.
TABLE \~ALUES OF d [N,H,] /dt (moles 1.00 M, initial [N,H,]
WC11 =
l
l- ’ ’ h-1) AT VARIOUS MOLAR CONCENTRATIONS OF COPPlZR(II) = 0.047 M, initial [l?e(II.I)] = 0.100 M. Temperature = 50~ ldt (moks.l-’
.h-1)
~o~,[N~H,J(ymks~E-~)
0
0.025 0.050 o-075
205
V
d [N,HJ
I 2
VI.
4.00
3.75
0.015
0.013 0.013 0.013 0.014
0.015 0.015 0.015
3.00
3.50
0.0084 0.0084 0.0087 ,o.oogo
. 3.50
0.002g 0;0038 0.0043 0.0054
0.0047 0.005’5 0.0065 0.0070
3.00
0.0020
0.0028 0.0034 0.0040
r-75
0.0012
0.0024 0.0030 0.0035
SUMMARY
The mechanism proposed by HIGGINSON AND WRIGHT for the slow reaction between fe,r+c irqn an&hydra+ne in acid aque,ous solutions hasbeen confirmed, using an anion, exchange prdqedure for ‘the separation and quantitative estimation ,of ferrous and ferric iron. ,, REFERENCES
‘.
1 W. C. E.’ HIGGINS&,’ Recent Aspects of, the Inorganic Chemistvy of Nilvogen, Publ., No,. ,1o. (1957) 95. ” 8 R. E. KIRK AND A. W. BROWNB, J. Am. Chem. Soc., 59 (1928) 337. : 8 E. J; CVY, J. Am.‘Chem. ‘SOL, 46 (1924) 1810.
Chem.‘Soc., 76 (1954) 2568. Chem. Sac., (1955) ,1551. R. ROSSIEINSKY, J.‘Chem. SOIL; (1957) 4685. C. BRAY AND H. A. LIEBI-IAFSKY, J. Am. Chem. SOL, 57 (1935) 51. MARKHAM, Biockem. J., 36 (1942) ,790. J.’ CONWAY, Micro-di@sion Methods and ,VoZumetvic EWOYS, 2nd Ed., Crosby
4 J. W.
CAI-IN, AND
R. E.
6 W. C. E. HIGGINSON 6 D.
7 W. 8 R.
9 E:
Chem. So& Spec.
AND
POWELL, J. Am. P. WRIGHT, J.
Sons, London, 1947. ’ i0 F.. H. POLLARD, J. F. w. iI F.’ H.
POLLARD, W. BRANDT
MCOMIE,
J; F. W. McOMIE. AND G. F.,SMITH,
G. &(~ICKLESSAND P. HANSON, A.
J. BANISTER
AND
Lockwood
‘&
J. Chvomatog., 4 (1960) 108. Analyst, 82 (1937) 780.
G. NICKLESS,
12 W. Anal. Chem,, 21 (1949) 1313. 1s H. W. SWANK AND M. G. MELLON, Ind. Ercg. Chem., Anal. Ed., g (1937) 406. 14 G. W. WhTT AND J. D. CRISP, Anal. Chem., 24 (1952) 2006. lb P: R. WOOD, Anal. Chem., 25 (1953) 1879.
J. Chvomatog., 4 (1960) 196-205