Introduction

Cyanide plays an important role in many industrial activities, such as electroplating, metal cleaning, and extraction of gold from its ore (Castilla-Acevedo et al. 2020; Moussavi et al. 2018; Zhang et al. 2015; Kuyucak and Akcil 2013; Korte et al. 2000). Cyanide has been frequently detected in soils and aquifers due to spills or leakage during cyanide production and utilization, or to improper disposal of cyanide-containing solid residues (e.g., gas manufacturing by-products and blast furnace sludge) (Wei et al. 2018; Sut-Lohmann et al. 2014; Zhang et al. 2010; Rennert et al. 2007; Rennert and Mansfeldt 2002a; Theis et al. 1994). Anthropogenic cyanide in polluted soils reportedly exists mainly as iron–cyanide complexes (i.e., ferrocyanide [Fe(CN)6]4− and ferricyanide [Fe(CN)6]3−), with the content of free cyanide (HCN and CN−) accounting for only a small percentage (Wei et al. 2020; Sut-Lohmann and Raab 2017; Mansfeldt and Höhener 2016; Manar et al. 2011; Johnson et al. 2008). These iron–cyanide complexes, although normally regarded as thermodynamically stable and nontoxic compounds (Shifrin et al. 1996), present a potential health risk due to their tendency to release toxic free cyanide when exposed to sunlight (Rader et al. 1993; Meeussen et al. 1992). In addition to the potential risk posed by dissociation of these iron–cyanide complexes, they also show direct ecotoxicity to some aquatic species, e.g., high chronic toxicity to fleas (Manar et al. 2011). Therefore, more attention should be paid to the geochemical behavior of these metal–cyanide complexes in the soils of cyanide-polluted sites.

In contaminated site soils, the mobility of pollutants is usually controlled by many processes including sorption, complexation, precipitation, and co-precipitation. Of all the processes involved, adsorption normally plays the most significant role, by reducing aqueous concentrations in soil systems. This is particularly true for the ionic pollutants. Gibbsite (Al(OH)3) and goethite (FeOOH) are ubiquitous soil minerals (Cornell and Schwertmann 2003; Violante and Pigna 2002), and these two minerals are normally regarded as the most active soil components controlling the mobility of many cationic (e.g., Pb2+, Cd2+) and anionic pollutants (e.g., PO43−, AsO43−, CrO42−) in soil systems (Arroyave et al. 2018; Xiong et al. 2018; Antelo et al. 2005). Considering their widespread presence and high affinity for many anionic pollutants, sorption of these anions with metal oxides has attracted great attention in the fields of geochemistry and environmental sciences.

The affinity of metal oxides and ionic pollutants is a result of either outer-sphere or inner-sphere complexation. Outer-sphere complexation occurs in such a manner that the ions are physically bound to the mineral surface via electrostatic attractions; ions fixed in this manner are susceptible to environmental conditions and are thus more mobile. Inner-sphere complexation, however, is formed via chemical binding to the mineral surface, becoming part of the surface. In most cases, complete separation between the two mechanisms is difficult, as both outer- and inner-sphere complexation often exist. A study discriminating these two mechanisms is important to understand the mobility and potential risks of many pollutants. From this point of view, it is of great significance to understand the factors (e.g., soil components, pH, ionic strength, etc.) that affect the geochemical behavior of these metal–cyanide complexes.

With regard to the iron–cyanide complexes considered in this study, experiments focusing on individual Al and Fe oxides have been conducted to evaluate their ability to retain these complexes by sorption and/or precipitation (Rennert et al. 2005; Bushey and Dzombak 2004; Rennert and Mansfeldt 2002b, 2001). However, most studies have been concerned only with sorption behavior in the context of one specific soil component, and no systematic comparison has been made among different minerals. Moreover, the contribution of these metal oxides to the fixation of the cyanide complexes in real soil samples has not been systematically studied. Thus, the effect of these metal oxides on Fe–CN retention remains unclear.

The aim of this study was to evaluate the relative importance of iron and aluminum oxides in the retention of iron–cyanide complexes in soils. For this purpose, ferrocyanide sorption behaviors on synthesized metal oxides (goethite and gibbsite), and on five different soil samples, were examined via batch adsorption experiments conducted under various cyanide loading, pH, and ionic strength conditions. During the adsorption experiments, the aqueous cyanide concentration was monitored to evaluate the amount of adsorbed cyanide, and to compare the maximum sorption capacity between the selected metal oxides and soil samples, thus providing an estimation of the relative importance of these metal oxides in ferrocyanide immobilization.

Materials and methods

Preparation of goethite and gibbsite

The goethite used throughout this work was synthesized following a method used by previous researchers (Hiemstra et al. 1989). Briefly, a NaOH solution (5 M) was gradually added (dropwise) to a Fe(NO3)3·9H2O solution (0.5 M) under continuous stirring until reaching pH 12. The suspension was stabilized at 60 °C for 90 h and then cooled and centrifuged to remove the liquid phase. The obtained solid phase was subsequently dialyzed with ultrapure water for about 2 weeks until the solution conductivity was < 10 μS/cm. Finally, the suspension was freeze-dried to obtain a goethite powder. Gibbsite was purchased from Sinopharm Chemical Reagent Co., Ltd. The chemicals and reagents utilized in the present study were of AR grade.

Soil sampling and characterization

To investigate the adsorption of ferrocyanide by soils, five soil samples with different physiochemical properties were collected (SPB, JC, FJ, ZY, and BS). The five soils were sampled from four different provinces, namely Chongqing (SPB and FJ), Shanxi (JC), Guizhou (ZY), and Jilin (BS). The SPB, JC, FJ, and ZY soils were sampled from the topsoil (approximately 0–20 cm in depth), and the BS soil, which contained a large amount of soil organic matter (SOM), was purchased from Shenzhibei Agricultural Science and Technology Co., Ltd. All soils were dried at ambient temperature, passed through a 2-mm-mesh sieve, and then gently ground for homogenization.

The properties of the five soils tested are shown in Table 1. The soils were neutral or moderately alkaline except the BS soil, which was acidic. The SOM of the five soils ranged from 0.93 to 22.32%. The cation exchange capacity (CEC) was in the range of 10.10–58.54 cmol/kg. The Al and Fe extracted with acid ammonium oxalate (OX) and dithionite–citrate–bicarbonate (DCB) were quantified by ICP-AES. The oxalate method extracts the amorphous and poorly crystalline fraction of the natural (hydr)oxide fraction (referred to as FeOX and AlOX) (Roden and Zachara 1996). The DCB method extracts the FeOX, AlOX, and crystalline Fe and Al oxide phases (referred to as FeDCB and AlDCB) (Roden and Zachara 1996; Mehra and Jackson 1958). FeDCB, FeOX, AlDCB, and AlOX in the five soils were determined to be 8.55–29.03 g/kg, 0.95–15.08 g/kg, 1.12–3.11 g/kg, and 0.29–1.82 g/kg, respectively. The fractions of sand, silt, and clay in the soils, as shown in Table 1, were measured using the hydrometer method (NY/T1121.3-2006) and sorted following the soil textural classification method suggested by the International Society of Soil Science (Kaddah 1974).

Table 1 Physical and chemical properties of the five soils in this study
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Adsorption experiments

Adsorption of ferrocyanide onto the synthesized metal oxides was performed in batch experiments using varying initial concentrations, ionic strengths, and pH. For each experiment, a solution containing the desired concentration of ferrocyanide was prepared at the pre-determined ionic strength and pH. The ferrocyanide solution was then transferred to a brown HDPE bottle containing a goethite (or gibbsite) suspension of the same ionic strength and pH. The volume of the reaction system was set to 25 mL, and the concentrations of goethite and gibbsite were maintained at 5 g/L and 25 g/L, respectively, for all adsorption experiments. The pH after the addition of adsorbent was adjusted using HNO3 or NaOH solutions. The brown HDPE bottle was placed in a horizontal oscillator at 25 °C and a vibration speed of 180 r/min for 24 h. Then, the solution was filtered through a 0.22 μm membrane filter, and the supernatant was used to determine the cyanide concentration. With this procedure, three sets of adsorption curves were obtained. The first set evaluated the adsorption at electrolyte concentrations of 0.001, 0.01, 0.05, and 0.1 M, with KNO3 as the electrolyte. The second set was conducted at pH ranging from 3 to 13, and the third set was performed at initial ferrocyanide concentrations of 2 × 10−6 to 4 × 10−3 M. All experiments were repeated twice, and the average values were used.

To evaluate the ferrocyanide sorption capacity of the five soils, batch adsorption experiments were conducted at a solid/liquid ratio of 1:10 (10 g soil to 100 mL water). No pH adjustment or ionic strength control was performed throughout the experiments with the soil reaction systems. The other operational conditions (e.g., temperature, vibration speed, filtration method, etc.) were the same as in previously described experiments with metal oxides.

Analytical methods

The mineralogical characteristics of the two metal oxides (goethite and gibbsite) were determined by X-ray diffraction (XRD) with a PANalytical X’Pert powder diffractometer using Cu Kα radiation at a voltage of 40 kV and current of 40 mA. Minerals were identified using the Inorganic Crystal Structure Database (ICSD). The morphological properties of the two sorbents were investigated by scanning electron microscopy (SEM) with the Zeiss ∑IGMA HD microscope. The FTIR spectra of the samples were determined using the Bruker Vertex 70 spectrophotometer. The specific surface area of the adsorbent was obtained by the BET method using the Quantachrome Quadrasorb 2 MP instrument, and the charging characteristics of the two metal oxides were determined by plotting a zeta potential–pH curve using a zeta potentiometer (Zetasizer Nano ZS90; Malvern). The cyanide concentration in the solution was analyzed following a spectrometric method (HJ484-2009). Briefly, the cyanide-bearing solution was acid-distilled to dissociate the cyanide complexes, and the released free cyanide was absorbed by an alkaline solution. The scrubber solution was then analyzed by a colorimetric method.

Results and discussion

Characterization of goethite and gibbsite

Figure 1a and b presents the morphology of the synthesized goethite sample. Figure 1a shows irregular agglomeration of goethite crystals at a scale of 1 μm. The enlarged SEM image in Fig. 1b shows that the synthesized goethite has a needle-like or rod-like shape, with a crystal size of 200–300 nm, corresponding well with the goethite synthesized in other studies (Xu et al. 2019; Liu et al. 2012). The electron microscope analysis shown in Fig. 1c shows that the purchased gibbsite minerals are randomly distributed, with a thicker hexagonal plate form about 250–400 nm in diameter, which is typical for the gibbsite, as described by other researchers (Hong et al. 2016; Sweegers et al. 2001).

Fig. 1
figure 1

Electron microscopic images of goethite (a–b) and gibbsite (c–d)

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The XRD patterns of goethite and gibbsite are shown in Fig. S1. The XRD data of goethite indicated a well-crystallized goethite mineral. No diffraction peaks of other oxide or hydroxide minerals were found in this study. The XRD pattern of the purchased gibbsite also indicated that it was a well-crystallized mineral.

Generally speaking, the specific surface area (SSA) of a sorbent is quite important, as a higher SSA normally results in greater adsorption. For goethite, the SSA is known to be related to the drop rate of the alkali solution and the aging time during synthesis (Wu et al. 2012). Within a given time range, a longer aging period normally results in a smaller mineral size and larger SSA of goethite (Liu et al. 2012). When the SSA is less than 60 m2/g, the goethite crystal changes from nano-size to micro-size. The SSA of the synthesized goethite in this study was 92.76 ± 2.19 m2/g, similar to the results reported in other studies using the Hiemstra method (Krumina et al. 2016; Xiong et al. 2015). The SSA for the commercially purchased gibbsite was 178.92 ± 4.35 m2/g, higher than that of gibbsite samples in previous studies (Xiong et al. 2018; Wu et al. 2000).

The isoelectric point (IEP) values, calculated via linear interpolation using the pH–zeta potential curves (see Fig. S2), were 9.1 and 8.7 for goethite and gibbsite, respectively. The IEP values for goethite and gibbsite in this study were consistent with previously reported values. Previously reported IEP values ranged from 6.4 to 9.5 for goethite (Szewczuk-Karpisz et al. 2018a; Madigan et al. 2009; Antelo et al. 2005) and 7.6 to 10.5 for gibbsite (Szewczuk-Karpisz et al. 2018b; Essington and Stewart 2016; Ladeira and Ciminelli 2004).

Effect of ionic strength on ferrocyanide adsorption

The influence of ionic strength on reactions at aqueous interfaces can provide insight into the nature of the chemistry involved. Varying the ionic strength had similar effects on the adsorption of ferrocyanide by goethite and gibbsite (Fig. 2). Generally, ferrocyanide adsorption was marginally affected by ionic strength, except at pH 4.13. At this low pH level, the adsorption ratio decreased from 95 to 90% for goethite and from 15 to 9.5% for gibbsite with increasing electrolyte concentrations from 0.001 to 0.1 M. This phenomenon may be partially explained as follows. First, at pH 4.13, a significant proportion of ferrocyanide presents as [HFe(CN)6]3−, which is expected to have a lower affinity to the positively charged goethite surface; second, under such an acidic condition, a proportion of ferrocyanide is assumed to have been transformed to ferricyanide. Previous researchers reported that ferrocyanide and ferricyanide could easily be inter-transformed under appropriate conditions (Huang et al. 2007). The formation of ferricyanide was supported by the appearance of a light yellow color in the adsorption system at pH 4.13. Additional research conducted by our group found that ferricyanide was more susceptible to ionic strength than ferrocyanide. Under other pH scenarios, with increased ionic strength, the ferrocyanide adsorption fraction on goethite and gibbsite varied in the ranges of 84–83% and 5.5–5.2% at pH 5.00, 79–78% and 3.9–3.8% at pH 6.30, and 62–62% and 2.7–2.6% at pH 7.3, respectively.

Fig. 2
figure 2

Effects of ionic strength on adsorption by goethite a and gibbsite b (0.1 mM ferrocyanide, 5 g/L goethite, and 25 g/L gibbsite)

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The effects of the electrolyte concentration on adsorption can be explained in terms of the complexes formed between the adsorbate and the adsorbent. Ions of the supporting electrolyte compete for active sites on the mineral surface with the target adsorbates that adsorb via outer-sphere complexation. Consequently, an elevated electrolyte concentration normally results in a decrease in adsorption. For example, CrO42−, SO42−, and SeO42− are believed to form outer-sphere complexes with γ-Al2O3, because their adsorption behavior depends strongly on ionic strength (Wu et al. 2000). In contrast, adsorbates that adsorb via inner-sphere complexation are intrinsically chemically bound with the active sites on the mineral surface and are therefore less influenced by changes in the electrolyte concentration. Antelo et al. (2005) evaluated the effects of two different electrolyte concentrations (0.01 and 0.1 M KNO3) on phosphate and arsenate sorption onto goethite. They observed no obvious salt effect on arsenate adsorption at all pH conditions studied (4.5, 7.0, 9.0, and 10.0). However, the behavior was somewhat different for phosphate: no salt effect was observed under acidic conditions, whereas under neutral and alkaline conditions, an increase in adsorption with increasing electrolyte concentration could be seen. Similar results were found by other researchers (Xu et al. 2019). This is due to the fact that an increased electrolyte concentration decreases electric potential at the particle–liquid interface, consequently reducing the electrostatic repulsion of the ions toward the charged particle surface and thus promoting sorption. However, no such enhancement was observed in the pH ranges studied here, for goethite or gibbsite. The present study found that ionic strength had little influence on ferrocyanide sorption. This result at least showed that inner-sphere complexation played a dominant role when ferrocyanide adsorbed onto goethite or gibbsite.

Effect of pH on ferrocyanide adsorption

Figure 3a and b presents data for the pH-dependent fraction of sorbed ferrocyanide (%), with initial ferrocyanide concentrations ranging from 0.01 to 0.15 mM for the goethite system and from 0.005 to 0.1 mM for the gibbsite system. For the two metal oxides, ferrocyanide adsorption declined with increasing pH, which was identical to the general pH-dependent sorption behavior of anions (e.g., PO43−, AsO43−, etc.) when performing adsorption onto oxide minerals (Xu et al. 2019; Arroyave et al. 2018; Wu et al. 2000). The goethite and gibbsite systems differed, as ferrocyanide adsorption on goethite was higher across the pH range and was still significant at pH over the IEP. When the ferrocyanide concentration was 0.01 mM, the sorbed fraction by goethite remained above 99% until the pH reached 10.5. Above pH 10.5, the adsorption ratio declined rapidly, dropping to 3.5% at pH 11. When ferrocyanide was 0.05 mM, the adsorption ratio was 87% at pH 10.5 and quickly dropped to 3% at pH 11; at 0.1 mM, the adsorption ratio was 72% at pH 9.5 and then dramatically decreased to 1% at pH 11; and when ferrocyanide was 0.15 mM, a similar but slightly lower sorption ratio was observed compared to that at 0.1 mM. From these results, we concluded that at lower adsorbate concentrations, the position of the adsorption edge at a given solid concentration is independent of the adsorbate concentration, as observed in the present study for ferrocyanide concentrations up to 0.05 mM. At higher adsorbate concentrations, the pH edge was shifted to lower pH points, which may result, for example, from saturation effects of surface binding sites (e.g., ≡FeOH for goethite). For gibbsite, a quite different sorption trend and much lower sorption ratio were observed. When the pH approached its IEP (8.7 for gibbsite), the sorption ratios were below 10% under all four scenarios.

Fig. 3
figure 3

pH-dependent ferrocyanide sorption onto goethite a and gibbsite b in 0.05 M NaNO3. Different symbols represent different ferrocyanide concentrations: 0.01–0.15 mM for the goethite system and 0.005–0.1 mM for the gibbsite system

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The affinity between adsorbent and adsorbate is highly dependent on the characteristics of the adsorbent (e.g., SSA and charging characteristics), as well as on the chemical speciation of the adsorbate. The pH-dependent ferrocyanide sorption onto goethite and gibbsite may have been caused by pH-dependent variation in adsorbent and adsorbate characteristics. All pKa values for ferrocyanic acid were reportedly lower than 4.5 (e.g., pK3 = 2.2 and pK4 = 4.2), implying that the pH had little influence on the speciation of ferrocyanide in most experiments in this study. In other words, most ferrocyanide ions present as [Fe(CN)6]4−, except under experimental scenarios with pH below 4.5. The significant influence of pH on ferrocyanide sorption was thus mainly ascribed to the charging characteristics of the sorbent. For amphoteric minerals such as goethite and gibbsite, higher pH leads to a more negatively charged mineral surface, resulting in stronger repulsion between the adsorbent and anionic adsorbate, whereas a lower pH results in more positive charge on the mineral surface, thus making the adsorbent more attractive to the anions. A striking finding was that the ferrocyanide sorption was still observed at pH values higher than the IEP of goethite, which further indicated that mechanisms other than electrostatic interactions, such as specific adsorption, were involved.

FTIR analysis

The FTIR spectra of goethite before and after sorption are shown in Fig. 4. The FTIR analysis of the gibbsite system provided no valuable information due to the low sorption capacity of gibbsite even at low pH; therefore, the data are not shown here. The FTIR spectrum of goethite before sorption further indicated a pure mineral without contamination. The spectra of the goethite minerals after ferrocyanide sorption were not a simple overlaying of the spectra of goethite (before sorption) and [Fe(CN)6]4−, as the characteristic spectrum of [Fe(CN)6]4− at 2044 cm−1 split into two peaks (at 2044 and 2026 cm−1) after sorption. This is an indication that ferrocyanide was not solely adsorbed as outer-sphere complexes, since a new bond was detected spectroscopically. Two bands, at 2026 cm−1 and 2043 cm−1, were found for the four samples adsorbed at different pH conditions. These bands, although varying in strength, did not change position or shape. The increasing band strength with decreased pH should be related to the increased sorption amount of ferrocyanide. Therefore, the band at 2026 cm−1 can be ascribed to the chemical binding of ferrocyanide to the goethite surface, e.g., formation of a monodentate or bidentate inner-sphere complex, while the spectrum at 2044 cm−1 represents sorbed ferrocyanide via an electrostatic attraction with the goethite surface.

Fig. 4
figure 4

FTIR spectra of ferrocyanide adsorbed on goethite at different pH values

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The interaction of ferrocyanide and metal oxides may include two-dimensional adsorption (accumulation of ferrocyanide at the oxide–solution interface) and three-dimensional precipitations (formation of an insoluble product). However, no precipitation of cyanide-containing products (e.g., Berlin blue, with a characteristic band at 2092 cm−1) as a result of the reaction of ferrocyanide with metal oxides was found in the present study.

Sorption capacities of ferrocyanide: metal oxides versus soil samples

The results presented above provide some essential information for understanding the mechanisms involved when ferrocyanide reacts with goethite and gibbsite. However, the relative importance of these metal oxides in terms of the retention of ferrocyanide in real soil samples remains unclear. Therefore, we performed additional isothermal adsorption experiments with varying cyanide concentrations and then fitted a Langmuir model to obtain the maximum sorption capacities of both synthesized oxides and soil samples. By combining the content of FeDCB, FeOX, AlDCB, and AlOX in the soils used in this study, we can roughly estimate the contribution of these metal oxides to the immobilization of ferrocyanide in natural soils.

Figure 5a and c shows the sorption isotherms of ferrocyanide on goethite and gibbsite. The sorbed fraction of ferrocyanide on these two minerals is also plotted in Fig. 5b and d for better differentiation of their sorption behavior at different pH values. The adsorption results were fitted to the Langmuir model and Freundlich model (Eqs. (1 and 2)), and the results are summarized in Table 2.

$$q_{e} = \frac{{Q_{\max } K_{L} C_{e} }}{{\left( {1 + K_{L} C_{e} } \right)}},$$
(1)
$$q_{e} = K_{F} C_{e}^{\frac{1}{n}}$$
(2)

where qe is the amount of sorbed ferrocyanide (mmol/kg); Ce is the concentration of ferrocyanide in the liquid phase after equilibrium (mmol/L); Qmax and KL are Langmuir parameters related to the monolayer sorption capacity (mmol/kg) and sorption affinity coefficient (L/mmol), respectively (Wei et al. 2014; Yan et al. 2010; Wisawapipatet al. 2009); n (dimensionless) and KF (mol1−(1/n)L1/n/kg) are empirical Freundlich parameters representing the sorption intensity and capacity, respectively (Wei et al. 2014).

Fig. 5
figure 5

Adsorption isotherms of ferrocyanide onto goethite (a–b) and gibbsite (c–d) in 0.01 M NaNO3

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Table 2 Langmuir and Freundlich fitting parameters of ferrocyanide adsorption onto pure metal oxides and soil samples
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The purpose of this set of experiments was to determine the maximum sorption amount of ferrocyanide on these two minerals. The Qmax values calculated from the Langmuir model revealed a great disparity between goethite and gibbsite under comparable pH conditions, e.g., the Qmax for goethite at pH 5.1 was 28.9 mmol/kg, which is 35 times higher than that for gibbsite. These results seem incompatible with the SSA of these two minerals, which are 92 m2/g for goethite and 178 m2/g for gibbsite. A study conducted by Bushey and Dzombak (2004) also noted the low sorptive capacity of ferrocyanide on gibbsite. However, it should be noted that the gibbsite used in that study had a low SSA (0.1 m2/g), which is quite different from the gibbsite used here. The Freundlich adsorption capacity parameters (KF) of the goethite sample were all higher than those of gibbsite at comparable pH conditions, which is indicative of the stronger affinity of goethite for ferrocyanide adsorption. Considering the results above, we concluded that the contribution of gibbsite to ferrocyanide retention is negligible in soil systems despite its surface characteristics. Therefore, only goethite was considered when evaluating the contribution of these metal oxides to the retention of ferrocyanide by soils.

The five soils used in this study were characterized by different physiochemical properties (Table 1). The isothermal adsorption curves of ferrocyanide by the five soils are presented in Fig. 6. The Langmuir and Freundlich fitting parameters are listed in Table 2. The adsorption data revealed that the five soils showed quite different sorption capacities of ferrocyanide, ranging from 11.90 to 41.73 mmol/kg. However, no correlation was found between the sorption capacity and total content of active iron oxides (FeDCB). To evaluate the contribution of goethite to the retention of ferrocyanide in real soils, we assumed that all active iron oxides present were goethite. Based on this assumption, the contribution of goethite (R%) was roughly estimated using Eq. (3).

$$R\% = \frac{{89 \times Q_{\max - G} Fe_{{{\text{DCB}}}} }}{{10 \times 56 \times Q_{\max - S} }} \% ,$$
(3)

where Qmax−G and Qmax−S are the adsorption capacities of goethite and soil samples derived from the Langmuir fitting results (mmol/kg), respectively; FeDCB is the amount of total active iron oxides in soils; and 89 and 56 in the equation are the molecular weights of Fe and FeOOH, respectively.

Fig. 6
figure 6

Adsorption isotherms of ferrocyanide onto different soil samples

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The calculated results are shown in Table 3. It is obvious that the contribution of goethite was no more than 9% for all soils investigated, revealing the low significance of goethite with respect to the retention of ferrocyanide in natural soils. These results suggest the presence of other mechanisms controlling the mobility of Fe–CN in soils (e.g., precipitation). The sorption capacity was highest for BS soil, which had the lowest pH, and lowest for the JC soil, which had the highest pH. This was consistent with trends for the active metal oxides in soils (e.g., goethite and gibbsite) on the sorption of ferrocyanide, whereby a decreased pH normally leads to increased sorption of ferrocyanide. Furthermore, acidic soil may favor the precipitation of cyanide-bearing compounds (e.g., Berlin blue) (Sut-Lohmann et al. 2015), which further promotes the partitioning of cyanide into the soil solid phase. Additionally, the high content of SOM in the BS may also be a non-negligible factor affecting its sorption behavior. However, further research is required to verify this hypothesis.

Table 3 Estimated contribution of goethite to the retention of ferrocyanide in soils
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We demonstrated the strong affinity of goethite for ferrocyanide in Sect. “Effect of ionic strength on ferrocyanide adsorption”; however, the sorption capacity of goethite with ferrocyanide was quite low when compared to other equivalent anions (e.g., PO43−), which led directly to the low R% values (Table 3). Phosphate is an anion commonly found in soils. Moreover, the association of phosphate with many metal oxides (e.g., goethite, gibbsite, ferrihydrite, etc.) has been extensively studied (Tiberg et al. 2020; Arroyave et al. 2018; Wei et al. 2014). Comparison of the sorptive behavior of phosphate and ferrocyanide should provide more insight into the interaction mechanisms of ferrocyanide with metal oxides. Table 4 summarizes the reported Qmax for phosphate on goethite and gibbsite samples. The Qmax of phosphate on goethite ranged from 70 to 336 mmol/kg with a pH of 4–10 and electrolyte concentration of 0.01 M KCl/KNO3, which is dozens of times higher than that for ferrocyanide. Moreover, gibbsite has also been shown to be an effective sorbent for phosphate, with Qmax ranging from 225 to 389 mmol/kg, which is approximately 100 times that of ferrocyanide. We cannot yet provide a definitive explanation of this discrepancy. When we make such comparisons, we should always keep in mind that the Qmax value stands for an apparent sorption maximum under specific conditions (e.g., ionic strength and pH) and is not necessarily identical to the total number of active groups on the mineral surface. One plausible explanation for the dissimilar sorptive behaviors of phosphate and ferrocyanide might be the steric match (or geometric match) between the active site and sorbed species.

Table 4 Comparison of the sorptive capacities of ferrocyanide and phosphate on goethite and gibbsite
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Tiberg et al. (2020) evaluated the sorption behavior of arsenate and phosphate on poorly crystalline iron (Fh) and aluminum hydroxide (Alhox), and found a higher affinity of phosphate on Alhox and a higher affinity of arsenate on Fh. They ascribed the specific affinities of arsenate/Fh and phosphate/Alhox to the steric match. Arsenate and phosphate are believed to exist mainly as bidentate complexes on Alhox and Fh via ligand exchange for –OH2 and –OH on mineral surfaces. The distance between two oxygen atoms in the phosphate and arsenate tetrahedra was compared with the distance between two neighboring oxygen atoms on the mineral surfaces. The authors found that the edge length of an Fe octahedron (in Fh) was approximately 2.6–3.0 Å, fitting better with the arsenate tetrahedron (~ 2.7 Å), whereas the Al octahedron, with a smaller edge length (approximately 2.5–2.9 Å), fit better with the phosphate octahedron (~ 2.5 Å). Ferrocyanide has an octahedron geometry, and the distance between two adjacent N atoms in the octahedron was calculated to be 4.55 Å based on its geometry (Prampolini et al. 2014), which is much larger than that of phosphate (~ 2.5 Å). This may be the major factor inhibiting its combination with goethite and gibbsite.

Conclusion

To evaluate ferrocyanide migration in natural systems, the sorption performance and nature of sorption onto goethite and gibbsite at different pH levels and electrolyte concentrations were evaluated. Changes in the ionic strength had no obvious effect on ferrocyanide sorption by goethite and gibbsite. Ferrocyanide adsorption onto the two minerals declined with increasing pH values, which was consistent with the sorption patterns of other high-valence anions (e.g., PO43−, AsO43−, etc.). For ferrocyanide, the sorption capacity of goethite was much greater than that of gibbsite. One possible explanation for this is that the geometric size of ferrocyanide is better matched to that of goethite than gibbsite. An FTIR spectrum analysis proved the existence of chemical bonding of ferrocyanide with the goethite surface. However, the nature of sorbed ferrocyanide species on goethite and gibbsite over a broad range of environmental conditions still needs further exploration. Despite the strong affinity of ferrocyanide for goethite, the sorption capacity Qmax is low compared to that of other anions. Based on a rough estimation using laboratory experiment-derived parameters (including Qmax and FeDCB), we consequently concluded that the contribution of these metals oxides to the retention of ferrocyanide in natural soils is relatively small.