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# New approach for Iron (III) stabilization in water, for its potential quantification by UV-Visible spectroscopy

Baia Pierre Clément, Robert Mauricot, David Neumeyer, Marc Verelst

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Baia Pierre Clément, Robert Mauricot, David Neumeyer, Marc Verelst. New approach for Iron (III) stabilization in water, for its potential quantification by UV-Visible spectroscopy. 1st thematic school "Materials for environmental depollution", Feb 2025, Natitingou, Benin. hal-05392965

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# New approach for Iron (III) stabilization in water, for its potential quantification by UV-Visible spectroscopy

Pierre-Clément Baia

CEMES-CNRS

Toulouse, France

pierreclement.baia@cemes.fr

Robert Mauricot

CEMES-CNRS

Toulouse, France

robert.mauricot@cemes.fr

**Main Corresponding Author:** Marc Verelst

**Abstract**—(English):

Iron is naturally present in high quantity in water. It is a peculiar element, as it is both essential for nearly all organisms and, at the same time, a known pollutant among the “Heavy Metal” category. Iron in water tends to be in its oxidated form, Fe<sup>III</sup>. It is a tricky element, very labile and versatile, which precipitate easily. To prevent this and be able to quantify ferric ions in solution, the common approach consists of reducing ferric ions (Fe<sup>3+</sup>) into ferrous ions (Fe<sup>2+</sup>) and to use a colored complex. However, this method is incompatible to study Fe<sup>III</sup> ions as unbound adsorbates and their interaction with our depolluting material. Iron evolution may involve different pathways depending on environmental conditions. Among them, Fe<sup>III</sup> ions can polymerize in a succession of discreet transient steps until visible precipitation. In this work, several techniques were used simultaneously (UV-Visible spectroscopy, pH-measurements, Dynamic Light Scattering...) to corroborate our results with similar studies using more advanced means (X-Ray Adsorption Spectroscopy, Mossbauer ...). This allowed us to illustrate the observed phenomenon via a step-by-step demonstration and provide a clearer understanding of the key-factors leading to Fe<sup>III</sup> polymerization. The goal was to understand how to avoid this polymerization/precipitation and be able to study Fe<sup>III</sup> and its interaction with any adsorbent via UV-Visible spectrophotometry. Through an empirical approach, a simple yet efficient way to stabilize trivalent iron in water was proposed: This new method only requires purging the dissolved gases before adding the iron salt. There, we boiled the water under a nitrogen atmosphere. This simple precaution stabilizes the iron III solutions for at least 24 hours. We hope that this work also contributes to a better understanding of the early steps of Fe<sup>III</sup> polymerization since this phenomenon also concerns other trivalent pollutants like Cr<sup>III</sup>.

**Résumé** – (Français):

Le fer est naturellement présent en large quantités dans l'eau. C'est un élément particulier car il est simultanément, essentiel à la quasi-totalité des organismes, et un polluant connu parmi la catégorie des « métaux lourds ». Le fer est naturellement présent, en milieu aqueux, sous forme oxydée Fe<sup>III</sup>. C'est un élément délicat, très labile et versatile, qui précipite facilement. Pour prévenir cela et pouvoir quantifier les ions ferriques, l'approche usuelle consiste à réduire les ions ferriques (Fe<sup>III</sup>) en ions ferreux (Fe<sup>II</sup>) et d'utiliser un complexant coloré. Cependant, cette méthode est incompatible pour étudier les ions Fe<sup>III</sup> comme adsorbats non-liés et leurs interactions avec le matériau dépolluant. L'évolution du fer peut impliquer différents cheminements selon les conditions environnementales. Parmi elles, le Fer III peut polymériser en une succession d'étapes transitoires discrètes jusqu'à ce que la précipitation soit visible. Dans cette étude, plusieurs techniques ont été utilisées simultanément (spectroscopie UV-Visible, mesures pH-métriques, Diffusion

David Neumeyer

CEMES-CNRS

Toulouse, France

david.neumeyer@cemes.fr

Marc Verelst

CEMES-CNRS

Toulouse, France

marc.verelst@cemes.fr

Dynamique de la Lumière...) afin de corroborer les résultats à des études similaires employant des moyens plus avancés (Spectroscopie d'absorption des rayons X, spectroscopie Mössbauer...). Cela nous a permis d'illustrer le phénomène observé à travers une démonstration étape par étape et de fournir une interprétation plus claire des facteurs-clés conduisant à la polymérisation du fer III. L'objectif étant de de comprendre comme éviter cette polymérisation/précipitation afin de pouvoir étudier le fer III et son interaction avec quelconque adsorbant via la spectrophotométrie UV-Visible. A travers une approche empirique, une solution simple mais efficiente pour stabiliser le fer III dans l'eau a été proposé. Cette nouvelle méthode ne nécessite que de purger les gaz dissous dans le solvant, avant l'addition du sel métallique. Ici, nous avons porté l'eau à ébullition sous atmosphère d'azote. Cette simple précaution stabilise les solutions de fer III pour au moins 24 heures. Nous espérons que cette avancée facilite également l'interprétation des étapes initiales de la polymérisation du fer III, un phénomène commun à d'autres polluants métalliques trivalents comme le chrome III.

**Keywords** : UV-Vis. Spectroscopy, Iron Hydrolysis, Ferrihydrite

## 1. Introduction

Iron is a queer element in water, as it is a well-known essential nutriment needing a daily intake<sup>1-5</sup> and simultaneously a common pollutant, able to easily degrade the apparent quality of water<sup>6</sup>. Iron is among the most recurrent pollutants known as “heavy metals” or “trace metals”<sup>7-10</sup>. A big friction for the detection and quantification of traces of metallic pollutants in water is the availability of the technics. Indeed, despite many solutions<sup>11</sup> for trace analysis, such as Induced Couple Plasma spectroscopies, most of them are too expensive to maintain and are reserved to dedicated laboratories. UV-Visible spectroscopy was selected as a potential alternative, more affordable, for regular detection and quantification of Fe<sup>III</sup> ions. This approach has satisfying results with divalent heavy metals. UV-Vis allows us to indirectly evaluate the capacity of any adsorbent, by measuring the ion concentration before and after contact with the sorbent (figure 1).

<figure><img src="image_2.png" /><figcaption>Figure 1: Illustration of UV-spectroscopy to evaluate adsorbent capacity using Beer's Law.</figcaption></figure>

To ensure adsorption to an equilibrium, trial solutions usually rest for several hours. Thus, adsorbates like Fe<sup>III</sup> need to remain stable in this timeframe. Most studies for iron quantification by UV-Vis, tend to use complexants<sup>12-14</sup>. However, this is not adequate for
adsorption studies, since it would mean studying the complexed ions and not the free species available in natural conditions. Fewer studies¹⁵,¹⁶ were made on Fe<sup>III</sup> ions in water by UV-visible spectroscopy without complexing agent. These studies introduce iron hydrolysis, Fe(OH)<sub>y</sub>(<sup>3-y</sup>)<sup>+</sup> with (y≤4), and the formation of a dimer Fe<sub>2</sub>(OH)<sub>5</sub><sup>4+</sup> as detectable species but remained elusive to the specific behavior of iron that could explain the common use of complexants. Through an attempt to establish a linear range of concentration, the formation of a precipitate in an iron solution was observed after several hours. This led us to do a study by Dynamic Light Scattering (figure 2) of a solution for 10 hours after its preparation. Our result showed a very early formation of colloids not visible to the naked eye, in a pH remaining below 3. This is problematic since iron in water evolve even without the adsorbent. Furthermore, the apparition of particles meant that the observed concentration of free adsorbates diverges from the initial quantity of iron introduced in the solvent.

<figure><img src="image_3.png" /><figcaption>**Figure 2:** Evolution of the “particles” after dissolution of 1,5.10<sup>-3</sup> M Fe<sup>III</sup> nitrate in distilled water by Dynamic Light Scattering</figcaption></figure>

As illustrated above (figure 2a) an evolution appears, akin to nucleation-growth, from our solution with an evolution of “particles” reaching 2 micrometers. This is problematic since the quantity of iron used to form the particles is unknown. On a linear scale, the evolution seemed like a classical scheme of growth with a drastic jump around 2 hours after the dissolution of iron salt in the solvent where it reaches its critical size. This result was coherent with another observation : the coloration of iron (III) in water is unusual from other metallic ions at low concentration. Usually, the coloration is immediate since the species absorb parts of the visible spectrum. For iron, an orange coloration was visible from the start in higher concentrations. The large absorption peaks are in UV but border the visible range, thus it was coherent when the solution became colorless at low concentrations. However, the orange coloration reappears after several hours, at 1.5mM. Thus, a former interpretation was that the coloration is not from the initial iron species in solution but from the “particles” growing in the solution akin to silver<sup>17</sup> or gold<sup>18</sup> nanoparticles . The identification of these particles is not trivial either for neophytes, as it happens at pH lower than 3 when the prime suspect, Fe(OH)<sub>3</sub>, is only stable<sup>16</sup> for a pH between 6 and 10.

The same data along a logarithmic scale (figure 2b) indicated the coexistence of several populations of particles (1nm, 10nm and 2µm), only few minutes after dissolution. These populations remain

stable in the observed timeframe and contradict the idea of visible coloration past a certain size from a unique new species. To perform iron quantification, it needed to be stable, thus the priority was to identify this transformation to prevent it.

Fe<sup>III</sup> is unstable in water, relevant terms such as its polymerisation<sup>15</sup> and its succeeding form, a ferric gel known as ferrihydrite<sup>16</sup> were briefly mentioned without being explicit. Iron may change its oxidation state and/or complex with other recurrent anions in water to form other precipitates<sup>19,20</sup>. This imply a myriad of non-exclusive possibilities<sup>21,22</sup> from Fe<sup>III</sup> ions in water, where ferrihydrite formation remains generally an hypothesis among others. Due to the specific conditions (low pH and initial concentration), our focus was centered on ferrihydrite and its precursors formation. Even if the precipitate does not form in a timeframe of few hours, they evolve greatly through polymerization, rendering the monomeric species' concentration impossible to follow. To study Fe<sup>III</sup> ion concentration, the imperative is to prevent the formation of the dimer and thus polymerization. Despite numerous studies<sup>23-33</sup> that greatly contributed to chart it, Fe<sup>III</sup> ion is too labile to follow the polymerization in its early stages<sup>34</sup>. Iron is very electrophilic, due to its small ionic radius and half-filled valence orbital. This makes it very instable in water, iron pulls the electrons towards it from its water ligands, which in turn weakens the O-H bonds and generate the hydrolyzed forms. To stabilize further, the hydrolyzed monomers must bond with each other. Polymerisation<sup>21,30</sup> can be resumed to a stochastic succession of reversible bonding, olation and oxolation, between hydrolyzed iron species to form higher weight polycations. These successive bonding, with a continuous release of protons, allows it to stabilize through a diffusion of the charge along a bigger structure<sup>32</sup>. This implies lower repulsion among polycations, generating clusters and eventually aggregation until precipitation of the ferric gel. Unfortunately, the early phases (from dimers to clusters) remain unproven, due to the impossibility to stabilize iron without strong complexants<sup>28,31</sup>, implying the impossibility to halt polymerization once it occurs. Thus, to prevent polymerization we had to research how to disrupt the “starting reactions”, the oxolation/olation bonding.

In this work, a water treatment was explored, by purging the dissolved gases through boiling the solvent under nitrogen atmosphere before salt addition. This solution was approached due to repeatability problems. Indeed, the solvent used is distilled water stocked in non-hermetic bottles. This leads to variable proportions of trapped CO<sub>2</sub> in the solvent which alter directly the pH of the solvent and of the initial solution. This implies various proportions of the iron species, and thus various kinetics of evolution.

## 2. Simulation or experimental part

The, ACS reagent grade, iron nitrate nonahydrate (Fe(NO<sub>3</sub>)<sub>3</sub>,9H<sub>2</sub>O) from Merck was dissolved in distilled water. Solutions were prepared by quick salt addition into the solvent in a volumetric flask. For degassed solutions, the solvent was purged through boiling and cooling under nitrogen atmosphere, then sealed as soon as the salt were added. Each solution was studied at [Fe<sup>III</sup>]<sub>Total</sub>= 1,5.10<sup>-3</sup> M.

All solutions were studied simultaneously by UV-Visible spectrophotometry and pH-measurements with a maximum delay of five minutes between the dissolution and each first measurement. The UV-vis absorption spectra were obtained by a Cary 5000 UV-VIS-NIR spectrophotometer from Agilent with the proprietary software Cary WinUV. For long datasets of 10 hours, we programmed a measurement every five minutes in a wavelength range between 250-450 nm. For punctual dataset, we extended the range to 250-800 nm. For each measurement, we used the distilled water as blank solution. pH-measurements were acquired by a Titroline Easy apparatus, from Fisher Scientific, coupled to an in-house software, to collect the pH in solution every 30 seconds for 10
hours. Dynamic Light Scattering studies were realized via a zeta-
sizer from Malvern with their proprietary software.

3. Results and discussion

Fe<sup>III</sup> ion in water adopt an octahedral geometry with six aqueous ligands Fe(H<sub>2</sub>O)<sub>6</sub>^{3+} and is very acidic. It pulls toward it the electrons from the water ligands, which in turn weaken the bond O-H, leading to the releasing of protons : Fe(H<sub>2</sub>O)<sub>6</sub>^{3+} ↔ Fe(H<sub>2</sub>O)<sub>5</sub>(OH)<sub>2</sub>^{2+} + H<sup>+</sup>. This reaction is known as its hydrolysis<sup>34</sup>, and can happen successively, Fe(H<sub>2</sub>O)<sub>6-n</sub>(OH)<sub>n</sub>(<sup>3-n</sup>)<sup>+</sup> (with 0 ≤ n ≤ 4), depending of the pH of the solution. For simplification, water molecules will be omitted later.

<figure><img src="image_4.png" /><figcaption>Figure 3 : UV-Vis absorption spectra of Fe<sup>III</sup> nitrate 1,5,10<sup>-3</sup> M dissolved in ambient (left) and degassed (right) distilled water at different timestamps.</figcaption></figure>

On Figure 3, we show the behaviors of Fe<sup>III</sup> ions in water with and without degassing. Figure 3.a, represents the evolution of aqua-iron species, without degassing, for 10 hours through three selected measures : the first one, just after preparation (in blue); the second one (in yellow) one hour later, and the last one (in red), ten hours after the preparation. The final deviation of absorbance observed at 297nm is also reported. In contrast, figure 3.b, represents the evolution of the Fe<sup>III</sup> ions species in degassed distilled water.

Figure 3.a illustrates the usual behavior observed for iron in literature. The pH-measurement indicates that the pH starts at 2,9 and then decreases significantly around three hours after the preparation. Based on this measured pH and initial Fe<sup>III</sup> concentrations, Fe<sup>III</sup> coexist under three species Fe<sup>3+</sup>, Fe(OH)<sup>2+</sup>, Fe(OH)<sub>2</sub><sup>+</sup> with molar proportions<sup>16,34</sup> around 15%, 70% and 15% and molar absorptivities<sup>15</sup> of 4640, 2030 and 1800 M<sup>-1</sup>.cm<sup>-1</sup> respectively. Fe<sup>3+</sup> is not directly detectable since its absorption peak maximum at 240 nm is out of the studied range. However, the bottom of this peak, near 250 nm, can be perceived. For Fe(OH)<sup>2+</sup> and Fe(OH)<sub>2</sub><sup>+</sup>, the determination is less clear since they both share a same adsorption peak<sup>15</sup> at 297 nm. It is assumed that Fe(OH)<sub>2</sub><sup>+</sup> is negligible due to its low molar absorptivity and its low molar proportion at a pH inferior to 3. Thus, the blue curve indicates mainly Fe(OH)<sup>2+</sup> species at 297 nm. The iron dimer Fe<sub>2</sub>(OH)<sub>2</sub><sup>4+</sup> with an absorption peak at 335 nm is not yet observable on this first measurement. It is expected since dimer formation requires the bonding of two hydrolyzed iron monomers which usually requires higher concentration<sup>28,31</sup>. After 1 hour (yellow curve), the iron

species already evolve significantly. The shoulder around 250nm is
higher, meaning a higher proportion of Fe³⁺, despite a pH stable
around 2,9. As introduced, monomers are able to bond among
themselves to generate a dimer³²,³⁴, the latter is then able to bond
with monomers form higher polymers. Higher polymers also induce
slower dynamics³⁵ and more potential collisions between fast
monomers and slower polymers. Thus, the following polymers,
since there are just transient states, are more difficult to identify. The
diffuse elevation of the absorbance between 335 and 420nm are
interpreted as these transient states, the polycations clusters. These
growing clusters are less soluble³⁵, deviating light and generating a
turbidity in the solution, elevating homogenously the absorption on
the whole visible spectrum. Through polymerization, polycations
progressively stabilize their charge by releasing protons³², which are
then available to rehydrate hydrolyzed monomers. This explains the
higher proportion of Fe³⁺ despite a pH still constant after one hour.
The formation of these polycations also explains the visible
coloration of the solution. Iron monomers all absorb in the
ultraviolet spectrum. When the solution shifts to yellow-orange, it
indicates that polymerization is underway. Ten hours after the
preparation (red curve), the impact of turbidity becomes significant
as it eclipses the absorption peaks of the monomers. The absorbance
from polycations, between 335 and 430 nm is clearly higher than
between 297 and 335 nm or between 430 and 450 nm. Thus, it is
impossible to follow the quantity of iron (by its initial monomers)
by UV-visible in these conditions.

By contrast, figure 3.b shows the same study but now with degassed
distilled water. In this case, the graph now contains all the measures
from 5 minutes to 10 hours after the preparation. With this minute
precaution, **the Fe**<sup>III</sup> **solution is now stable for 10 hours, with a
maximum deviation of less than 2%.** Moreover, punctual UV-
Visible spectra showed that it remained stable 24 hours after
preparation (but only for a few days). The origin of this partial
inhibition of iron evolution has yet to be determined, but it is
estimated that the purged gases played an important role.

4. Conclusion

Iron ability to polymerize is a known phenomenon since 1938³⁶ and
has been strongly followed in the last 50 years³⁷, however the critical
conditions to prevent its occurrence are still elusive. To summarize,
trivalent iron ions are very unstable in water and tend to polymerize
and later precipitate in a very short timeframe, even in an acid
environment. This phenomenon is quite insidious since it happens
through fast and discreet steps (hexa-aqua-ferric ions → hydrolyzed
monomers → dimers → polycations clusters → ferrihydrite). This
prevents the quantification of Fe<sup>III</sup> species in water by UV-Visible
spectroscopy without resorting to complexing agents. As shown
here, it is possible to outline this complex phenomenon with just
UV-Visible spectroscopy and pH measurements. This result will
allow faster comprehension of iron's atypical behavior and is the
first step toward iron quantification without using complexing
agents.

This work also allowed to highlight a critical information: Fe<sup>III</sup> ions are capable to precipitate in acidic conditions to form a ferric gel known as ferrihydrite. Such information is essential for past and future adsorption studies<sup>36-44</sup> as it may invalid results if iron is considered trapped on the depolluting materials when it may just precipitate around it. This also implies that iron may cover the adsorption sites without occupying them, thus preventing more concerning elements from being captured. On the contrary, ferrihydrite is also a natural adsorbent able to trap other metals<sup>40</sup>. Thus, in multi-metallic studies, sorbent capacity towards other pollutants would be overestimated or underestimated in presence of iron. Consequently, we recommend not neglecting the influence that Fe<sup>III</sup> could play in water decontamination processes, because even if the latter does not constitute a very worrying toxicant, it is sometimes present in very large quantities in natural waters and is likely to distort many results.
CREDIT AUTHORSHIP CONTRIBUTION STATEMENT

This study was done by P-C. Baia with D. Neumeyer. Then reviewed and revised by R. Mauricot and M. Verelst before its final submission.

DECLARATION OF COMPETING INTEREST

The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this study.

ACKNOWLEDGMENT

The authors would like to thank our sponsor, the CNRS (National Center for Scientific Research), for financing this study and all the committee members of this event.

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