**4. Major physicochemical phenomena used to mitigate the generation of AMD**

The AMDs characteristics vary from site to site, however, for simplicity and without loss of generality the pH usually ranges between acid (close to pH 1 to 3) and slightly acid (around pH 5) which after wastewater treatment needs to reach pH values between 6.5 and 9 to be adequately discarded [60]. The acidity stabilises a concentration of certain heavy metals in aqueous solutions such as copper, zinc, chromium among others and metalloids like arsenic. Several physicochemical phenomena used to remove ions from an aqueous phase are succinctly reviewed in this section from both thermodynamic and kinetic classic modelling standpoints.

### **4.1 Neutralisation: coupled or not with chemical precipitation**

Perhaps this is the most straightforward strategy to treat AMD wastewaters coming from mine sites. The idea is to remove from the aqueous phase dissolved species that might be toxic to human beings and the environment by adding salts with high saturation or solubility product constant (*Kps*) values but very small ones for the ions to be withdrawn. Considering that the AMD is acid, one of the priorities is to bring the pH of the water close to neutral and that is usually accomplished by dosing with an alkali reagent. The release of hydroxyl ions (*OH*� ð Þ *aq* ) to the aqueous solution enhances the hydrolysis of the metal ions promoting the generation of neutral species which will eventually precipitate. The most used reagent at an industrial scale is the quicklime which in presence of water is transformed into slaked lime [61]. In **Figure 9** a usual AMD containing a high concentration of acid (*H*3*O*<sup>þ</sup> ð Þ *aq* ), heavy metals (*M<sup>n</sup>*<sup>þ</sup> ð Þ *aq* ) and sulphate (*SO*<sup>2</sup>� <sup>4</sup>ð Þ *aq* ) is presented. When adding quicklime, ultimately dissolved calcium ions would react with sulphate forming anhydrite-like structures while hydroxyl ions would form metal hydroxides and simultaneous adjust the pH to neutral values. The precipitation of sparingly soluble species would follow two sequential steps namely nucleation and crystal growth being the former, in many cases, more energy consuming than the latter. That is the reason why seeds and/or surface roughness of different materials would be desirable to improve the precipitation thermodynamic spontaneity and kinetics [62].

### *4.1.1 Thermodynamics*

An initial concentration of metal is *CMe*,0 and of a counter ion (anion, which can be sulphate) is *Cci*,0. The flow rate of liquid enriched in these two ions is *Q*. For

**Figure 9.** *Sketch of the role of different ions formed during the dosing of quicklime to acidic aqueous solutions.*

computation purposes, the dose of leachate of a metal hydroxide (for simplicity, calcium hydroxide or slaked lime) of concentration *Cx* will be considered at a flow rate of *q*. Usually, as *Cx* is relatively high, the following flowrate relationship holds *q* ≪ *Q*. The reagent with higher solubility compared to that of the salt to be precipitated, undertake an ionisation reaction such as that presented in Eq. (13).

$$\text{M(OH)}\_{n'(s)} \longleftrightarrow \text{M}^{n'+}\_{(aq)} + n'\text{OH}^-\_{(aq)}\tag{13}$$

The mathematical condition required for the precipitation to occur is given by Eq. (13).

$$\frac{C\_x Q\_r}{Q} C\_{ci,0} > Kps^{(1)}\tag{14}$$

Then, the efficiency of the treatment can be computed using Eq. (14).

$$\rho = 100 \left\{ \frac{\mathcal{C}\_{\rm x} Q\_r}{2 \mathcal{C}\_{ci,0} Q} + \frac{1}{2} \left[ \sqrt{\left( \frac{\mathcal{C}\_{\rm x} Q\_r}{\mathcal{C}\_{ci,0} Q} - 1 \right)^2 - \frac{K p s^{(1)}}{\mathcal{C}\_{ci,0}^2}} - 1 \right] \right\} \tag{15}$$

The physical units of the different parameters in these equations need to be consistent. Plus, bear in mind that the efficiency of Eq. (15) is overestimated as only the major species were considered in this computation. Simultaneously, the precipitation of metals can be computed from Eq. (16). For the metal precipitation, Eq. (16) needs to be solved.

$$Kps^{(2)} = (\mathcal{C}\_{Me,0} - \mathcal{y})(n'\mathcal{C}\_x - n\mathcal{y})^n \tag{16}$$

Eq. (16) does admit an analytical solution only in very specific conditions, so it is better to solve it numerically. Although it looks like this strategy is quite promising as *Fundamentals and Practical Aspects of Acid Mine Drainage Treatment: An Overview from Mine… DOI: http://dx.doi.org/10.5772/intechopen.104507*

it would be able to remove sulphate, heavy metals and even neutralise the acidic conditions of the AMD, it is somehow misleading for at least four reasons [63, 64]:


During mine closure conditions this strategy is not directly recommended since the reagent dosing control is difficult to implement without personnel in place. However, the fundamentals behind this mechanism still hold. In the scenario of mine closure, the main idea would be to incorporate some sparingly soluble minerals with alkaline behaviour such as silicates, or others.

### *4.1.2 Kinetics*

There are several studies focused on determining the dissolution rate of solids, especially that of quicklime in water. The first stage of dissolution is usually modelled by Eq. (17) [65].

$$\frac{da\_i}{dt} = k(a\_{i, \text{sat}} - a\_i) \tag{17}$$

where *ai* is the activity of the ion "*i*" as a function of time, *ai*,*sat* is the activity of the ion "*i*" in saturation conditions, *t* represents the time, and *k* is a specific rate constant obtained per unit of volume that has been defined as a function of the diffusion coefficient of the ion being transferred from the solid surface to the aqueous solution bulk (*D*), the thickness of the diffusion layer (*δ*), and the specific surface *<sup>S</sup> V* as presented in Eq. (18) [66].

$$k = \left(\frac{\mathcal{S}}{V}\right)\frac{D}{\delta} \tag{18}$$

The classic shrinking core model with reaction control can also be used (Eq. 18) [67–69].

$$\mathbf{1} - (\mathbf{1} - a)^{\natural \downarrow} = k\_{ct} \mathbf{t} \tag{19}$$

where *α* represents the conversion of the reaction which corresponds to the volume fraction of solid that has been dissolved, and *kcc* is the kinetic rate constant when the process is governed by chemical control.

### **4.2 Adsorption: Chemisorption**

This mechanism aims at removing pollutants from an aqueous phase by fixing them onto a surface of a solid which is stable when immersed in the wastewater. The adsorption mechanism is one of the preferred reactions for wastewater treatment not only because low-cost adsorbents consisting of by-products or wastes from other industries may be used, but also because it may reach high removing efficiencies of dissolved molecules with final concentrations of a few parts per billion [69]. **Figure 10** presents several aspects to consider when picking up this mechanism. Different reactions between the adsorbent and the aqueous solution lead to the partial dissolution of the adsorbent affecting the local pH near its surface and its stability of the adsorbent suspension whenever forming small particles may occur [70].

There are several drawbacks behind the implementation of adsorption-based technologies. This technology requires optimising the contact between the solid phase and the aqueous phase containing the species to be adsorbed. Usually, piling up of adsorbent material in a column disposed of vertically or horizontally is preferred [71] but maintaining the permeability of the porous medium with time could become a challenge. It is also desirable to implement technologies using chemisorption rather than physisorption. Chemisorption has many advantages such as its specificity exemplified in **Figure 10** by the single and double binding shown for the metal and sulphate. That is, different adsorption sites would be used by different types of adsorbates reducing the competition for adsorption sites. Plus, the relatively high binding energy

**Figure 10.** *Diagram of adsorption processes used in wastewater treatment.*

*Fundamentals and Practical Aspects of Acid Mine Drainage Treatment: An Overview from Mine… DOI: http://dx.doi.org/10.5772/intechopen.104507*

associated with the adsorption process turns it quite irreversible from a kinetic standpoint which reduces the chances of pollutants desorption. Nevertheless, the main disadvantage would be that the eventual saturation of the adsorbent may be reached needing to move forward to a desorption stage to regenerate the adsorbent [72].

### *4.2.1 Thermodynamics*

The thermodynamics of the adsorption process is explained in terms of the adsorption isotherm [73]. The adsorption isotherm is usually plotted in a graph where the y-axis represents the maximum quantity of adsorbed species per unit of the dry mass of adsorbent (also known as specific adsorption) while the x-axis presents the concentration of the species in equilibrium with the specific adsorption measured. All the data is obtained experimentally at a constant temperature and solids percent. There are many mathematical models that can be used to describe the process having each of the conditions and assumptions that as much as possible must represent the specifics of the process under study. The most used adsorption isotherms cited by researchers are the Langmuir and Freundlich isotherms as Eqs. (20) and (21) [74].

$$q\_{i,eq,L} = \frac{\mathcal{S}\_m K\_L a\_{i,eq}}{\mathbf{1} + K\_L a\_{i,eq}} \tag{20}$$

where *qi*,*eq*,*<sup>L</sup>* corresponds to the volume (or mol) of adsorbate *i* at the surface of the adsorbent in equilibrium conditions, *ai*,*eq* is the equilibrium real concentration of the adsorbate in *molL*�<sup>1</sup> , *KL* is the Langmuir isotherm constant commonly associated with the binding energy, *Sm* es the amount of adsorbate required to form a monolayer.

$$q\_{i,eq,F} = K\_F a\_{i,eq}{}^{\vee\_n} \tag{21}$$

where *KF* is the Freundlich empirical constant usually associated with the sorption capacity, and *n* is the sorption intensity.

### *4.2.2 Kinetics*

One general case to model adsorption kinetics is presented in Eq. (23) [75]

$$r = \frac{(kinetic\ factors)(process\ potential)}{(adsortion\ factor)}\tag{22}$$

wherein the numerator there is a description of the classic law of mass action and in the denominator, the inhibition of the adsorption rate procured by the blockage of surface sites of other species in the system is incorporated. For example, the mathematical model for adsorption kinetics of one species labelled with the underscore "*i*" can be described as in Eq. (24) in the case of chemical reaction control.

$$r = k\theta\_i = k \frac{K[a\_i]}{1 + K[a\_i] + \sum\_{j \neq i} K\_j[a\_j]} \tag{23}$$

where *k* is the specific kinetic constant, *θ<sup>i</sup>* is the fractional occupancy of adsorption sites by the main species "*i*", *K* and *ai* are the Langmuir adsorption constant referred

**Figure 11.** *Redox reactions in wastewater treatment.*

to the main ion and its activity in the aqueous phase while *K <sup>j</sup>* and *a <sup>j</sup>* are their equivalent but for other ions competing for adsorption sites.
