NOTES
AND
COMMENT
ON THE MEASUREMENT OF FERROUS IRON IN NATURAL WATERS?
INTRODUCTION
Measurements of iron in natural waters
are uncertain because of the lability
of
the two states of oxidation
of iron,
because ferrous iron ( Fe2+) is soluble and
ferric iron ( Fe3+) is not, and because of
the probable (but not understood)
involvement of the iron with organic matter.
Nevertheless, it is generally believed that
in surface waters freed of plankton the
predominant form of iron is as Fe3+ hydroxide, either as a precipitate or as a
peptized sol ( Hutchinson
1957; Shapiro
1964 ) .
The above belief has been challenged
by Gjessing ( 1964)) who suggested that
the iron in colored surface waters near
Oslo, Norway, is mostly Fe2+ in organic
complexes.
Because Gjessing’s results
were so unexpected and extreme (in one
case 92% of the iron was Fe2+), a critical
examination has been made of his method
to determine possible systematic errors
in it.
METHOD
OF
LEE
AND
iron from organic matter or from mixed
ferrous-ferric
oxides.” Gjessing modified
this step by increasing the boiling time to
15 min. He states (p. 273), “For the estimation of the ferrous and ferric iron in the
water types examined, it was found necessary to pretreat the sample with a high
concentration of hydrochloric
acid, or to
increase the boiling time with the acid,
to insure a complete release of the iron
from the complexes.”
EXAMINATION
OF THE
OF
EFFECTS
BOILING
In considering possible sources of error,
attention was turned first to the boiling
step. Boiling a sample of colored water
0.6
0.5
STUMM
0.4
Gjessing used the bathophenanthroline
technique ( Lee and Stumm 1960)) which
can be summarized as follows :
L
g
1. Collect sample and add HCl ( 2 ml/liter)
to stabilize Fe’+ during storage.
2. Add 1 ml coned HCl per 25 ml and boil
for approximately
5 min.
3. Add sodium acetate solution to pH 4.
4. Add bathophenanthroline.
5. Extract
color with hexanol,
dilute, and
measure absorbance.
0.3
E”
I
z
1 0.2
0.1
The second step, that of boiling the
sample with acid, is recommended by Lee
and Stumm (p. 1572), “. . . to free the
0.0
0
1 Contribution
No. 24 from the Limnological
Research Center, Univ. of Minnesota.
The work
was supported
by USPHS Research Grant WPMrs. Margaret
Weeks
performed
the
00771.
analyses with exemplary care.
4
Time
8
12
16
- minutes
FIG. 1. Ferric and ferrous iron concentrations
as a function of boiling time of an acidified sample of lake water (Cedar Bog Lake, Minnesota).
293
294
NOTES
0
5
IO
AND
COMMENT
15
Time
20
-
25
30
35
40
minutes
FIG. 2. Reduction
of ferric iron to ferrous iron as a function of boiling time of an acidified
tion of yellow acids and ferric chloride.
Total iron was determined by addition.
with acid is an uncertain procedure. As
the color is due to organic substances
(Shapiro 1957), under the conditions of
acidity
(favoring
solution of Fe3+ and
stabilization
of Fez+), high temperature
( increased reaction speed ) , and boiling
(elimination of dissolved oxygen from the
sample ) , Fe3+ could be reduced to Fe2+,
and the reduction would be more extensive the longer the boiling time. This was
demonstrated
by analyzing
samples of
water from Cedar Bog Lake, Minnesota
(Fig. 1). The samples were boiled with
acid ( 1 ml HC1/25 ml) for various lengths
of time and divided into two subsamples.
The Fe’+ concentration
was determined
by the Lee and Stumm method and the
Fe3+ by the addition of potassium thiocyanate. Careful
attention
was paid to
blank and inherent color corrections, The
Fe2+ iron concentration increased and the
Fe3+ decreased to almost the same degree.
In other words, the Fe2+ was apparently
coming not from Fe2+ complexes but from
reduction.
The data in Fig. 2 strengthen this con-
solu-
elusion. In this case, a solution of 30 mg/
liter of yellow organic acids previously
extracted from Linsley Pond, Connecticut
(Shapiro 1957), was used to simulate a
lake water sample. Fe3+ chloride was
added to give a Fe3+ concentration of 1.6
mg/liter.
The solution was acidified as
before ( 1 ml HC1/25 ml) and subsamples
were boiled for varying lengths of time
after which they were cooled in ice water,
divided, and analyzed for Fe3+ and Fe2+.
There is a remarkable inverse correspondence between the two curves and a constant sum of the Fe3+ and Fe2+. There is
no possibility that the Fe2+ originated from
Fe2+ complexes, because the organic matter (tested separately)
contained
only
enough total iron to increase the concentration by 0.018 mg/liter,
whereas the
actual increase in Fe2+ was 1.02 mg/liter.
EXPERIMENTS
WITH
DILUTED
SAMPLES
The results described thus far involved
rather highly colored samples. Cedar Bog
Lake has a color ( at pH 8) of 135, and
the solution of yellow acids also had a
NOTES
AND
Time
COMMENT
-
minutes
FIG. 3,. Reduction
of ferric iron to ferrous iron as a function of boiling
tion of yellow acids and ferric chloride.
The five curves refer to different
w/w ratio of acids : iron is indicated on each curve.
high color. As both of these are more
highly
colored than Gjessing’s samples
( 19-59 Pt-Co units),
experiments were
carried out in samples having less color.
A sample of Cedar Bog Lake water (collected under the ice) was serially diluted
with distilled water to provide concentrations of 100, 50, 25, 12.5, and 6.25%. The
samples were acidified, boiled for exactly
5 min to allow only partial reduction, and
analyzed for Fe3+ and Fe2+. The results
showed that the proportion
of iron reduced was substantially the same in each
sample.
EFFECT
OF
ORGANIC
MATTER
: IRON
RATIO
In the experiments described, the ratio
of organic matter : iron remained constant
at all dilutions. The curves in Fig. 3 illustrate the results of experiments designed
to test the effect of different ratios. They
show the rate of reduction of the Fe3+ as
a function of boiling time in acidified
solutions containing ratios of organic mat-
295
time of an acidified soluexperiments
in which the
ter : iron ranging from 0 : 1 (acidified distilled water) to 457 : 1 (w/w ) . The solutions were prepared
by adding
Fe3+
chloride solution to solutions containing
different concentrations of yellow acids in
distilled water. To ensure comparability,
the concentration
of iron in all experiments was between 0.153 and 0.168 mg/
liter, and the ratio was changed by changing only the yellow acid concentration.
The results show clearly that increasing
the ratio of organic matter : iron increases
the rate of reduction of Fe3+ and that,
except at very low ratios, the error involved is large. The 15 min boiling time
used by Gjessing, while more effective in
reduction than the 5 min time used by
Lee and Stumm, is not three times so
because of the change in slope of the
curves. Thus, even 5 min boiling reduces
a significant proportion of the Fe3+, and
in extreme cases it may reduce all of the
iron.
The ratios tested are not unnaturally
296
TABLE
ratios
NOTES
1.
Colors
and derived
AND
COMMENT
organic mutter : iron
(Data
1964)
of several waters near Oslo, Norway.
from
Gjessing
Organic
matter : iron
(w/w)
Water
Aurevann,
Brook 1
Brook 2
Brook 18
Brook 19
Lake
Sandungen, Lake
Halsj@en, Lake
Maridalsvannet,
Lake
30
59
29
35
38
27
26
19
185
227
151
159
119
375
79
150
6 80
.L
.;
70
f 60
0
.-'0 50
t
$ 40
E
z
k
30
20
IO
high. Using the water color and iron data
of Gjessing ( 1964) and the color : weight
relationship in Fig. 11 of Shapiro 1957, it
is possible to arrive at an approximation
of the organic matter : iron ratio in the
waters studied by Gjessing ( Table 1).
All are high enough to result in significant
reduction of Fe 3+ during a 15 min period
of boiling.
In fact, if these ratios are
plotted against the per cent of iron found
by Gjessing to be in the Fe2+ state, there
is a direct relationship indicating a greater
per cent reduction at the higher ratios.
This relationship also serves as a logical explanation for Gjessing’s finding that
there is a marked decrease of Fe2+ during
storage in polyethylene
bottles. I have
frequently noted that such bottles of colored water form a precipitate that clings
very tightly to the sides and bottom of the
bottle. Thus, if the precipitate is largely
organic matter, the ratio organic matter :
iron will decrease, or, if it is made up of
iron and organic matter, the iron concentration in the water will decrease. In both
cases, samples analyzed
by Gjessing’s
method will appear to have less Fe2+ the
longer they have been stored. It is also
evident from his comparison of bottles
stored at 4C with those stored at room
temperature that the “decrease” of Fe2+ is
due not to a chemical phenomenon such
as oxidation, in which case the 010 would
be much higher (see Stumm and Lee
1961)) but to a physical phenomenon in
which temperature plays only a small role.
0
I2
3
4
5
6
7
8
9
IO
PH
FIG. 4. Reduction
of
iron by boiling a solution
and ferric iron for 15 min
The open circles are from
ALTERATION
OF
ferric
iron to ferrous
containing yellow acids
at different
pH values.
a second experiment.
BOILING
CONDITIONS
An attempt was made to find conditions
under which boiling-a
useful procedure
for dissolving compounds-could
be done
safely. One approach that was tried involved boiling at somewhat higher “pH,
followed by lowering the pH for addition
of reagents. Fig. 4 shows the results that
were obtained with a solution containing
a mixture of yellow acids and Fe3+ at an
iron concentration of 0.155 mg/liter and
an organic matter : iron ratio of 77 : 1. The
samples were boiled for exactly 15 min
during which the pH, adjusted with KOH
and HCl, changed only slightly.
It is
apparent that boiling samples having a
pH value of less than 6 introduces considerable error, particularly at the lower pH
values. It also seems that boiling at high
pH introduces error through reduction.
Thus, the only safe ;pH range seems to be
near neutrality.
Because most complexes
or chelates tend to dissociate either under
acid or alkaline conditions,
depending
upon their nature, boiling at neutrality is
a compromise that is unlikely to be very
useful.
Hem (1960) found a similar pH dependence in the reduction of Fe3+ by tannic
NOTES
AND
297
COMMENT
1.6 1
.*..*a.
1.0
. . . . . . . . . . . . . . . . . . . ..-.
‘-J. . . . . . . . . . . . . . .TOTAL
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . -. . . . . . . . . . . . . . . 0
-
FERRIC
0.8 -
l
0.6 0.4 / 0
0.2 0.0
tf
-0
/-
0 0 0
’
0
RR
0 /'
0I
FERROUS
---------
____--
---e-o
/H
Ferrous
at
pH
I
I
I
20
40
60
I
I
I
80
Time
I
too
-
I
I
120
I
140
I
7
,A
160
hours
FIG. 5. Reduction
of ferric iron to ferrous iron during storage of acidified
samples of lake water.
The single point labeled pH 7 shows the ferrous iron concentration
of a sample that had not been
acidified.
-
acid solutions. However, Hem’s time scale
was considerably longer, because his experiments,
done at room temperature,
required several weeks.
The only way that boiling is acceptable
is when Fe3+ and Fe2+ are measured before
and after boiling. Fe2+ in the boiled sample, in excess of that amount of Fe”+
which has been reduced, can then legitimately be assumed to have come from a
Fe”+ complex.
DETERMINATIONS
WITHOUT
BOILING
The question arises as to whether boiling is necessary. Considering the solubilities of Fe3+ phosphate, Fe3+ hydroxide,
Fe2+ sulfide, and Fe2+ hydroxide and the
pH values normally required for dissociation of reversible complexes (for example, Fe3+ EDTA dissociates at “pH 2)) it
appears that lowering the “pH to 0.5 to
0.2 for a few minutes is sufficient to release virtually all iron from a sample that
has been Millipores
(0.45 ,u) filtered. For
example, a sample boiled with acid (2 ml
coned HC1/50 ml) for 15 min and treated
with hydroxylamine
had the same amount
of bathophenanthroline-reactive
iron pres-
ent ( absorbancy = 0.755) as a similar
sample that was not boiled but was
merely acidified and treated with hydroxylamine after 15 min (absorbancy = 0.750).
Because this sample was collected under
the ice of Cedar Bog Lake where anoxic
conditions prevail, and was highly colored, it could be expected to have contained iron in several of the above-mentioned forms. Nonetheless, boiling failed
to liberate any more iron than was liberated in the unboiled sample. Samples
containing plankton must be boiled, but
their wet oxidation with perchloric acid is
preferable.
STORAGE
OF
SAMPLES
The fact that boiling samples with acid
results in conversion of Fe3+ to Fe2+ has
significance
to storage and subsequent
analysis of lake water samples. A sample
of Cedar Bog Lake water was acidified
(1 ml coned HCl/liter)
and stored in the
dark at room temperature
(25C).
Subsamples were removed from time to time
and their Fe2+ and Fe3+ concentrations
measured. The results (Fig. 5) illustrate
the changes. It is usually assumed that
298
NOTES
AND
adding acid stabilizes the Fe2+, but the
Fe2+ is seen to increase through reduction
of Fe3+.
ORDER
OF
ADDITION
COMMENT
what greater absorbance with Fe3+, but
this can be compensated for as suggested
by Lee and Stumm.
OF REAGENTS
CONCLUSIONS
Early in this work, it was observed that
the order of reagent addition during the
1, lo-phenanthroline
procedure for Fe2+ is
extremely important.
It had been previously determined that the most color was
developed in the shortest time when the
phenanthroline
reagent was added to the
acidified sample (pH 2 or less) and the
pH subsequently raised, either by adding
sodium acetate or NaOH (using Congo
Red paper as an indicator ) . This point
was later made by Davis, Osborne, and
Nash ( 1958), but the differences they observed were much smaller.
Similarly,
Sandell ( 1959 ) , while not making the
point specifically,
describes addition
of
phenanthroline
followed by sodium acetate. It is therefore surprising that Lee
and Stumm suggest adding the bathophenanthroline after the solution has been
brought to “pH 4 with sodium acetate.
Gjessing adopted this procedure, adjusting the pH to 4-5 with ammonia after
first adding sodium acetate. A demonstration of both orders of addition on
identical samples of lake water, and on
distilled water, yielded the following results :
Blank
absorbancy
1. Lee and Stumm, Gjessing:
Water + acid + acetate +
bathophenanthroline
0.009
2. Davis et al., Sandell, Shapiro:
Water + acid + bathophenanthroline + acetate
0.013
Sample
absorbancy
0.036
0.059
A similar demonstration on distilled water
containing a higher concentration of Fe2+
g ave, for Order 1, 0.465;
( 0.5 mg/liter)
for Order 2, 0.634. Therefore, Order 2
was used throughout the work reported in
this paper. The increased sensitivity outweighs the higher blank value. This order
of reagent addition also results in a some-
The recently published results of Gjessing (1964) purporting to show iron mostly
in the Fe2+ state in Norwegian waters appear to be in error because of his analytical technique, in that acidified samples
were boiled; with such treatment, even
small amounts of organic matter will rapidly reduce Fe3+ to Fe2+. The rate of reduction is rapid enough to affect even
acidified samples standing at room temperature. The method of Lee and Stumm
may be made significantly
more sensitive
by changing the order of addition of reagents.
JOSEPH
SHAPIRO
Limnological
Research Center,
University of Minnesota,
55455.
Minneapolis
REFERENCES
N. F., C. E. OSBORNE, AND H. A. NASH.
1958. pH adjustment in calorimetric
iron determinations.
Anal. Chem., 30: 2035.
GJESSING, E. T.
1964. Ferrous iron in water.
Limnol. Oceanog., 9: 272-274.
Hm,
J. D. 1960. Complexes
of ferrous iron
with tannic acid.
US. Geol. Surv., Water
Supply Papers. 1459-D, p. 75-94.
HUTCHINSON,
G. E. 1957. A treatise on limnology, v. I. Wiley, New York, N.Y. 1015 p.
1960. DeterminLEE, G. F., AND W. SIVMM.
ation of ferrous iron in the presence of ferric
J. Am. Wairon with bathophenanthroline.
ter Works Assoc., 52: 1567-1574.
SANDELL,
E. B. 1959. Calorimetric
determination of traces of metals.
Interscience,
New
York, N.Y. 1032 p.
Chemical
and biological
SHAPIRO,
J.
1957.
studies on the yellow organic acids of lake
water.
Limnol. Oceanog., 2 : 161-179.
-.
1964. Effect
of yellow
organic acids
on iron and other metals in water.
J. Am.
Water Works Assoc., 56: 1062-1082.
STUMM,
W., AND G. F. LEE.
1961. Oxygenation of ferrous iron.
Ind. Eng. Chem., 53:
143-146.
DAVIS,