Give Two Examples of Mineral Acids

Chemical and Physical Properties

James G. Speight , in Reaction Mechanisms in Environmental Engineering, 2018

2.1 Inorganic Acids and Bases

An inorganic acid (also called a mineral acid) is an acid derived from one or more inorganic compounds. All inorganic acids form hydrogen ions and the conjugate base ions when dissolved in water. Commonly used inorganic acids are sulfuric acid (H ₂SO₄), hydrochloric acid (HCl), and nitric acid (HNO₃). Inorganic acids range from superacids (such as perchloric acid, HClO₄) to very weak acids (such as boric acid, H₃BO₃). Inorganic acids tend to be very soluble in water and insoluble in organic solvents.

Inorganic acids are used in many sectors of the chemical industry as feedstocks for the synthesis of other chemicals, both organic and inorganic. Large quantities of these acids—especially sulfuric acid, nitric acid, and hydrochloric acid—are manufactured for commercial use in large plants. Inorganic acids are also used directly for their corrosive properties. For example, a dilute solution of hydrochloric acid is used for removing the deposits from the inside of boilers, with precautions taken to prevent the corrosion of the boiler by the acid. This process is known as descaling.

Although HF can be named hydrogen fluoride, it is given a different name for emphasis that it is an acid—a substance that dissociates into hydrogen ions (H⁺) and anions in water. A quick way to identify acids is to see if there is an H (denoting hydrogen) in front of the molecular formula of the compound. To name acids, the prefix hydro- is placed in front of the nonmetal and modified to end with -ic. The state of acids is aqueous (aq) because acids are found in water. Some common binary acids include:

HF(g) (hydrogen fluoride)   →   HF(aq) (hydrofluoric acid)

HBr(g) (hydrogen bromide)   →   HBr(aq) (hydrobromic acid)

HCl(g) (hydrogen chloride)   →   HCl(aq) (hydrochloric acid)

H₂S(g) (hydrogen sulfide)   →   H₂S(aq) (hydrosulfuric acid)

The term inorganic base represents a large class of inorganic compounds with the ability to react with acids, that is, neutralize acids to form salts. An inorganic base causes an indicator to take on characteristic colors and usually refers to water-soluble hydroxides, such as sodium hydroxide (NaOH), potassium hydroxide (KOH), or ammonium hydroxide (NH₄OH). The term also includes weak bases, such as water-soluble carbonate derivatives (

) or bicarbonate derivatives (

).

Some chemicals can act either as an acid and as a base—an example is water which may either donate a hydrogen ion (to form the hydroxyl ion, OH⁻) or accept a hydrogen ion to form the hydroxonium ion (H₃O⁺), also called the hydronium ion. This property makes water an amphoteric solvent. Briefly, an amphoteric compound is a molecule or ion that can act both as an acid as well as a base, and this property can influence the behavior of chemicals in the environment.

Metal oxides which react with both acids as well as bases to produce salt and water are amphoteric oxides, and include lead oxide (PbO) and zinc oxide (ZnO), among many others such as the oxides of aluminum (Al₂O₃) and copper (CuO). Other examples of amphoteric compounds are oxides and hydroxides of elements that lie on the border between the metallic and nonmetallic elements in the periodic table (Fig. 3.2). For example, aluminum hydroxide [Al(OH)₃] is insoluble at neutral pH (pH   =   7.0) but can accept protons in an acid solution to produce [Al(H₂O)₆]³⁺ or accept a hydroxide ion (OH⁻) in a basic solution to produce [Al(OH)₄₊]⁻ ions. Consequently, aluminum oxide is soluble in acid and in base, but not in neutral water. Other examples of amphoteric oxides are beryllium oxide (BeO), gallium oxide (Ga₂O₃), and antimony oxide (Sb₂O₃). Increasing the oxidation state of a metal increases the acidity of its oxide by withdrawing electron density from the oxygen atoms. For example, antimony pentoxide (Sb₂O₅) is acidic but antimony trioxide (Sb₂O₃) is amphoteric.

Figure 3.2. The periodic table of the elements showing the groups and periods including the Lanthanide elements and the Actinide elements.

One other phenomenon that deserves consideration here is the phenomenon known as solvent leveling, which is an effect that occurs when a strong acid is placed in a solvent such as (but not limited to) water. Because strong acids donate their protons to the solvent, the strongest possible acid that can exist is the conjugate acid of the solvent. In aqueous solution, this is the hydroxonium ion (H₃O⁺). This means that the strength of acids such as hydrochloric acid (HCl) and hydrobromic acid (HBr) cannot be differentiated in water as they both are dissociated 100% to the hydroxonium ion.

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Hydrofluoric Acid

S.E. Gad , D.W. SullivanJr., in Encyclopedia of Toxicology (Third Edition), 2014

Uses

HFA is the inorganic acid of elemental fluorine and commercially is the most important fluorine compound. Its largest use is in the manufacture of fluorocarbons to be used as refrigerants, solvents, and aerosols. HFA is also used in fluoropolymers, aluminum production, stainless steel pickling, uranium processing, glass etching, oil well acidizing, gasoline production, removal of sand and scale from foundry castings, and as a laboratory reagent. Anyone using HFA should understand the safety measures required to protect human health: for example, read the relevant Material Safety Data Sheet (MSDS), call the supplier for additional information if necessary, and confirm that the personal protective equipment (PPE) has been shown to effectively protect against HFA exposure. In addition, the PPE should be checked carefully before each use; for example, a pinhole-sized hole could cause problems because HFA can penetrate deeply into skin and muscle tissue and simply flushing the area with water is not enough.

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FUEL CELLS – PHOSPHORIC ACID FUEL CELLS | Electrolytes

T. Murahashi , in Encyclopedia of Electrochemical Power Sources, 2009

Phosphoric acid is a unique inorganic acid electrolyte, which is generally used in fuel cell applications at around 200  °C in order to obtain higher system efficiency where its concentration is over 100%. A matrix, which is made of SiC, is used to retain the hot phosphoric acid in a cell. Phosphoric acid has advantageous properties as an electrolyte, such as low volatility, good ionic conductivity, stability at relatively high temperatures, carbon dioxide tolerance, and also carbon monoxide tolerance. With all these advantages, there have been several technical electrolyte-related problems with fuel cell stacks. The major issues are volume change, evaporation loss, and electrolyte migration, and these are described in detail. These problems mainly arise due to the fact that phosphoric acid is a liquid electrolyte under fuel cell operation.

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Infrared and Raman Spectroscopies of Clay Minerals

J. Madejová , H. Pálková , in Developments in Clay Science, 2017

13.5.4 Adsorption of Pyridine on Acid-Treated Samples

The interaction of Mt with inorganic acid significantly changes the surface acidity of a mineral. A powerful technique capable to distinguish different types of clay mineral acid sites is MIR spectroscopy. For example, pyridine as a base of moderate strength can be physically and/or chemically adsorbed on the Mt surface. Chemisorbed pyridine can either form H bonds between the pyridine nitrogen atom and OH groups of clay minerals, can be protonated to pyridinium cation at strong Brønsted acid (BA) sites (e.g., from polarized water molecules), or can be coordinately bound through pair of electrons on nitrogen atom to Lewis acid centre (e.g., Al ^{3   +} cations). These various species are easily identified and distinguished by examining the 1700–1400   cm^{−   1} region, where the absorption bands related to ring stretching vibrations (skeletal modes) of pyridine are located.

Pálková et al. (2013) used MIR spectroscopy and pyridine to characterize the acid sites on Na⁺-saturated Cheto Mt surface before (Na⁺-SAz) and after dissolution in 6   M HCl at 80°C for 2, 4, and 8   h (SAz-2   h, SAz-4   h, and SAz-8   h). After pyridine adsorption the spectrum of Na⁺-SAz showed the intense bands at 1441 and 1594   cm^{−   1} corresponding to the H-bonded pyridine, a band near 1487   cm^{−   1} indicating the presence of weak BA sites and the vibration at 1573   cm^{−   1} attributed to the physisorbed pyridine (Yariv, 2002). The modification of the Na⁺-SAz structure in HCl affected the bonds between pyridine and the sample. The spectrum of SAz-2   h showed a decrease of the intensities of the bands related to H-bonded and physi-sorbed pyridine and a new well-developed band at 1541   cm^{−   1} assigned to N⁺-H deformation vibration of pyridinium cation. The pyridine molecules accepted protons from H₃O⁺, present in the sample after HCl treatment, thus confirming the existence of strong BA sites on the mineral surface. The spectrum of SAz-4   h revealed a pronounced decrease of the intensity of the 1541   cm^{−   1} band reflecting the reduction of the amount of the strong BA sites. After 8 hours of HCl treatment, no band was found in the SAz-8   h spectrum due to pyridinium cations. Only bands corresponding to the H-bonded and/or physi-sorbed pyridine trapped on the amorphous silica surface were seen in the 1700–1400   cm^{−   1} region.

The same set of samples was used by Madejová et al. (2015) to examine the potential of NIR spectroscopy to distinguish different pyridine species adsorbed on SAz-1 Mt surface (Fig. 13.8). The NIR spectrum of acid-untreated Na-SAz shows characteristic bands at 7057 and 4515   cm^{−   1} related to the structural OH groups and a combination water band at 5239   cm^{−   1} (Fig. 13.8A,a). The spectra of acid treated samples show a gradual decrease in the intensities of the structural OH bands due to the release of the central atoms from the octahedral sheets and the appearance of the SiOH overtone at 7315   cm^{−   1} confirming the formation of protonated silica (Fig. 13.8A,b–d).

Fig. 13.8. NIR spectra of (A) Na-SAz treated with 6   M HCl at 80°C for (a) 0   h, (b) 2   h, (c) 4   h and (d) 8   h, (B) the same samples as in (A) but after pyridine adsorption, (C) SAz treated with HCl for 2   h with adsorbed pyridine (a) unheated and heated at (b) 110°C, (c) 170°C and (d) 230°C.

The NIR spectra of the samples after pyridine vapors adsorption are presented in Fig. 13.8B. Due to partial replacement of water molecules by pyridine the environment of the structural OH groups of Na⁺-SAz was changed and the 2ν(OH) band is split into two components at 7156   cm^{−   1} and 7059   cm^{−   1} (13.8.B,a). A decreased intensity of the (ν  + δ)H₂O band also confirms the partial replacement of water by pyridine. The first C–H overtone of the aromatic ring and combination bands in the 6200–5800   cm^{−   1} and 4700–4000   cm^{−   1} regions, respectively, correspond to physi-sorbed and/or H-bonded pyridine (Fig. 13.8B,a). A detailed assignment of the individual CH bands/components of pyridine is given in Madejová et al. (2015). After pyridine adsorption on the acid-treated samples, the pronounced modification of the 2ν(SiOH) band is observed. No band at 7315   cm^{−   1} can be identified for SAz-2   h (Fig. 13.8B,b) and only a very weak shoulder near 7305   cm^{−   1} is recognized for SAz-4   h and SAz-8   h (Fig. 13.8B,c and d). The majority of the silanol groups, as weak Brønsted acid sites, form H bonds with pyridine-nitrogen. The 2ν(SiOH) band of H-bonded SiOH groups was shifted to lower wavenumbers and contributed to the broad complex band near 7100   cm^{−   1}. The shape and position of the 2ν(CH) band of pyridine is modified only slightly comparing to the Na⁺-SAz, i.e. acid-untreated sample. In contrast to the MIR region no feature/band due to pyridinium cation can be identified in the NIR spectra presented in Fig. 13.8B.

To examine the strength of the pyridine bonding the samples were heated at 110, 170 and 230°C prior to IR analysis. Fig. 13.8C shows the NIR spectra of heated SAz-2   h sample. The re-appearance of the 7315   cm^{−   1} band assigned to the free Si–OH groups and the increased intensity of the (ν  + δ)H₂O band confirmed that pyridine H-bonded to silanol groups and water molecules was released upon heating. The intensity of the 2ν(CH) overtone decreased due to desorption of H-bonded pyridine and the band was considerably shifted to the higher wavenumbers (Fig. 13.8C,c and d). Pálková et al. (2013) showed that the strong BA sites in SAz-2   h provide protons for the generation of pyridinium cations absorbing at 1539   cm^{−   1}. Moreover, only pyridine molecules adsorbed on these types of acid sites persisted at the heating of the sample at 230°C. Thus the bands at 6096 and 6065   cm^{−   1} in the NIR spectrum of SAz-2   h heated at 230°C were assigned to C–H overtones of the pyridinium cations. Similar C–H positions were observed for SAz Mt directly saturated with pyridinium cations. For SAz-2   h unheated and heated at 110°C the pyridinium bands could not be observed because they were covered with stronger bands of physi-sorbed and/or H-bonded pyridine molecules. The position of the C–H overtone with peaks at 6055 and 6018   cm^{−   1} found for SAz-2   h heated at 170°C indicated the overlapping contribution of both the H-bonded pyridine molecules and the pyridinium cations.

The NIR spectra of heated SAz-4   h (not shown) confirmed the gradual release of physically adsorbed and/or weakly H-bonded pyridine leaving only those molecules that interacted with strong BA sites. Although the intensity of the 2ν(CH) band in the samples heated at 170°C and 230°C was significantly lower than in the corresponding SAz-2   h, the peaks at 6097 and 6066   cm^{−   1} were compatible with the presence of strongly bonded pyridinium cations. The reaction product obtained after 8 hours of decomposition of Na⁺-SAz in HCl contained about 97% of protonated amorphous silica (Madejová et al., 2015). The acidity of the Si–OH groups was not strong enough to provide protons for the creation of a pyridinium cation. After SAz-8   h heating at 230°C the absence of the 2ν(CH) bands confirmed almost complete desorption of pyridine from the sample. The obtained results showed that in addition to MIR, NIR spectroscopy can also distinguish different pyridine species on the Mt surface.

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Phosphorus-based and Intumescent Flame Retardants

S. Hörold , in Polymer Green Flame Retardants, 2014

2.2.3 Melamine polyphosphate

Melamine derivatives are salts with organic or inorganic acids such as boric acid, cyanuric acid, and phosphoric or polyphosphoric acid. These derivatives have higher decomposition temperatures (melamine cyanurate (MC), 310  °C; melamine polyphosphate (MPP), 360   °C) in comparison to melamine (250   °C) alone, are less soluble in water, and can add other features to performance such as nondripping (cyanuric acid) or char formation (boric acid, phosphoric acid). Melamine homologs such as melam, melem, and melon have even higher decomposition temperatures with 400   °C and above, but they have found only experimental use so far.

MC is used in the production of halogen-free flame-retarded unfilled and mineral-filled PAs for electrical and electronic applications, in thermoformable polyurethane foams, and in polypropylene intumescent formulations in conjunction with APP. Melamine phosphate and melamine pyrophosphate are mainly used in intumescent coatings and in glass-filled polyolefins and PAs. Melamine phosphate and melamine pyrophosphate can also be used in paper, textiles, and wood laminates to effect fire-retardant barriers. Melamine borate is mainly used in intumescent coatings, in various polyolefins, in PVC as an antimony-free synergist and smoke suppressant, and in textile backcoatings [4].

Like APPs, melamine phosphates are also substances combining the synergistic effect of nitrogen- with phosphorus-containing components in one salt. Based on increasing thermal stability the melamine phosphates can be ranked as follows: melamine phosphate   <   melamine pyrophosphate   <   melamine polyphosphate [28]. Melamine monophosphate is a salt of melamine and phosphoric acid. Above ∼200   °C melamine phosphate will react with melamine pyrophosphate with release of reaction water, which will result in a heat sink. Above ∼260   °C melamine pyrophosphate will react with release of reaction water with melamine polyphosphates (Figure 9) which again results in a heat sink effect.

FIGURE 9. Structure of melamine polyphosphate (MPP).

Above 350   °C, melamine polyphosphate undergoes endothermic decomposition thus acting as a heat sink and cooling the combustion source. The released phosphoric acid acts to coat and therefore shield the condensed combustible polymer. The phosphoric acid along with the polymer also works to form a char around the fuel source (polymer), thus reducing the amount of oxygen present at the combustion source. The melamine released is also a blowing source, which blows up the char, resulting in an intumescent behavior.

Compared to APP, MPP is more thermally stable and less sensitive to hydrolysis, but its FR effect is lower than that of APP, so it could not replace APP in intumescent coatings and polyolefin FRs. It was originally developed for glass fiber-reinforced (GF) PA 66 and gives a V-0 at 25% dosage, but it fails to improve the flammability characteristics of PA 6 at reasonable levels. Jahromi demonstrated that MPP depolymerizes at above 350   °C, inducing significant cross-linking (as evidenced by increase in viscosity) in PA 66 and leading to a drastic depolymerization of PA 6 (as evidenced by a decrease in viscosity) [29].

MPP is today mostly used in combination with other FRs, such as metal phosphinates (see phosphinate chapter), metal hydroxides, and phosphates [28]. It is characterized by its good thermal stability and a low impact on the glass-transition temperature (Tg). As mentioned, under thermal stress, melamine derivatives decompose endothermically (heat sink) and release inert nitrogen gases (e.g. ammonia) that dilute oxygen and the flammable gases in the flame. Often phosphoric acid is also formed as a decomposition product and promotes the formation of insulating char on the surface of the polymer [30,31].

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Handbook of Clay Science

P. Komadel , J. Madejová , in Developments in Clay Science, 2013

10.1.3 Acid Dissolution of Smectites

Acid treatment of clay minerals with strong inorganic acids resulted in solid products containing unaltered layers and amorphous three-dimensional cross-linked silica, depending on the extent of acid activation. The IR spectra of SWy-1 montmorillonite (SWy-1 Mt) after reaction with 6  M HCl (Fig. 10.1.1) demonstrated the changes of the chemical bonds in the Mt structure. The gradual decrease in the intensities of the OH bending (930–800   cm^{-   1}) and Al–O–Si (524   cm^{-   1}) bands indicated the decomposition of the octahedral sheets. Changes in the tetrahedral sheets were reflected in the position and shape of the Si–O stretching band. In addition to the tetrahedral Si–O band near 1048   cm^{-   1}, the MIR spectra of the acid-treated samples revealed a pronounced absorption near 1100   cm^{-   1} assigned to Si–O vibrations of amorphous silica with a three-dimensional framework formed during acid treatment. The position of the Si–O band at 1103   cm^{-   1} together with the weak inflection near 524   cm^{-   1} observed for the sample dissolved for 30   h reflected a very high yet incomplete dissolution of Mt in 6   M HCl (Madejová et al., 1998). NIR spectrum of SWy-1 Mt presented a broad, complex band at 7083   cm^{-   1} corresponding to the first overtones (2ν _OH) of the structural OH groups and H₂O molecules. The intense band near 5251   cm^{-   1} was attributed to the combination mode ${\left(\nu +\delta \right)}_{{\mathrm{H}}_2\mathrm{O}}$ of the water molecules and the band at 4533   cm^{-   1} to combination modes (ν  + δ)_OH of the structural OH groups (Fig. 10.1.1). Upon acid treatment, the intensities of these bands gradually decreased as a result of the decomposition of the Mt layers. A new band at 7311   cm^{-   1} assigned to 2ν _SiOH confirmed the creation of SiOH groups (Pálková et al., 2003).

Figure 10.1.1. NIR and MIR spectra of SWy-1 Mt treated with 6   M HCl at 95   °C for 0, 8, 12, 18 and 30   h. Q is quartz.

MIR spectra from Madejová et al. (1998).

Early acid-dissolution studies, based on solution analysis, of dioctahedral smectites in HCl by Osthaus (1954, 1956) indicated faster dissolution of octahedral than tetrahedral sheets. Assays of solid reaction products employing advanced spectroscopic techniques provided experimental evidence that acid treatments dissolved central atoms from the tetrahedral and octahedral sheets at similar rates (Luca and MacLachlan, 1992; Tkáč et al., 1994).

Luca and MacLachlan (1992) studied the dissolution in 10% HCl of two nontronites from by Mössbauer spectroscopy. They fitted the spectra either with two octahedral Fe^{3   +} doublets only, or with an additional tetrahedral Fe^{3   +} doublet. Acid treatment appeared to remove octahedral and tetrahedral Fe^{3   +} from the structure at about the same rate. Mössbauer and IR spectroscopies and XRD indicated that the remaining undissolved part was the untreated nontronite. ²⁷Al and ²⁹Si MAS-NMR study on removal of tetrahedral and octahedral Al^{3   +} from Mt by 6   M HCl led to very similar conclusions (Tkáč et al., 1994). The rates of dissolution of tetrahedral and octahedral Al^{3   +} were also comparable. Three different types of structural units were identified in acid-treated samples, including (SiO)₃SiOH units remaining as a result of poor ordering of the framework without the possibility of cross-linking.

The extent of the dissolution reaction depended on both clay mineral type and reaction conditions, such as the acid/clay mineral ratio, acid concentration, time and temperature of the reaction (Komadel, 2003; Sakizci et al., 2011). The composition of the clay mineral layers substantially affected their stability against acid attack; trioctahedral layers dissolved much faster than their dioctahedral counterparts. Higher substitutions of Mg^{2   +} and/or Fe^{3   +} for Al^{3   +} in dioctahedral smectites increased their dissolution rate in acids (Vicente et al., 1994, 1995a; Komadel et al., 1996b; Madejová et al., 1998, 2009b; Steudel et al., 2009a). For 15 dioctahedral smectites, a good correlation of the Mg^{2   +} and Fe^{3   +} contents was obtained with the half-time of dissolution in 6   M HCl at 96   °C (Novák and Číčel, 1978).

Li⁺ dissolved slightly faster than Mg^{2   +} from hectorite layers at low acid concentrations (Komadel et al., 1996b). Thus, protons were preferentially attracted by sites close to Li⁺ (in the octahedral sheet) that were more negative compared to sites adjacent to Mg^{2   +}. This difference disappeared at high acid concentrations when the reaction rates were high. Similarly, octahedrally coordinated Mg^{2   +} cations were preferentially released by HCl in comparison with Fe^{3   +} and Al^{3   +} (Christidis et al., 1997; Gates et al., 2002). The effect of acid anion on dissolution of hectorite is complex and remains uncertain (Komadel et al., 1996b; Van Rompaey et al., 2002). Effects of smectite type, acid concentration and temperature on the half-time of dissolution in 0.2   L HCl/g smectite, acid/clay mineral ratio in closed systems (no substances being added or removed) are summarized in Table 10.1.1. The rate of dissolution of various atoms obtained from chemical analysis of the liquid reaction products indicated the presence of different phases in bentonite. Readily soluble octahedral and tetrahedral constituents and 'insoluble' portions of constituent atoms calculated from the dissolution curves provided information on the distribution of atoms in the sample (Číčel and Komadel, 1994). Readily soluble portions included exchangeable cations and easily soluble admixtures such as goethite (Komadel et al., 1993) and calcite (Komadel et al., 1996b). The most common 'insoluble' phases found in the fine fractions of bentonites were kaolinite, quartz, anatase and volcanic glass. Halloysite was the most decomposed mineral after reaction in sulphuric acid of different concentrations, followed by Mt, pyrophyllite and kaolinite (Kato et al., 1966). The observed low dissolution rate of pyrophyllite compared with Mt was due to (i) low octahedral substitution and (ii) the presence of collapsed non-swelling interlayer spaces in pyrophyllite.

Table 10.1.1. Effects of Smectite Type, Acid Concentration and Temperature on Half-Time of Dissolution in 0.2   L HCl/g Smectite in Closed Systems

Smectite	HCl (M)	T (°C)	t 0.5 (h)
Effect of smectite type
Nontronite	6.0	95	0.16
Mg-rich montmorillonite	6.0	95	6.2
Al-rich montmorillonite	6.0	95	8.0
Effect of acid concentration
Hectorite	0.25	20	4.6
Hectorite	0.50	20	2.6
Hectorite	1.00	20	1.7
Effect of temperature
Fe3   +-beidellite	6.0	50	12.0
Fe3   +-beidellite	6.0	60	6.0

Pentrák et al. (2010) investigated the influence of chemical composition and swelling ability of three dioctahedral clay minerals from the Mt-illite series on their dissolution in 6   M HCl. Mt was completely dissolved within 18   h, while the residues of non-decomposed illite could be distinguished in both samples with prevailing non-swelling interlayer spaces, treated for 36   h. Chemical composition of dioctahedral clay minerals had a greater effect on the dissolution rate than swellability. Illite with higher degrees of substitution of Mg^{2   +} and Fe^{3   +} for Al^{3   +} in the octahedral sheets and of Al^{3   +} for Si^{4   +} in the tetrahedra was more easily soluble in HCl than the illite/smectite with 30% swelling interlayer spaces.

A series of reduced-charge Mt was prepared via Li⁺ fixation at elevated temperatures (the Hofmann–Klemen effect) to explore how the expandability of the interlayer spaces influenced the extent of dissolution. As the negative LC decreased, the content of non-swelling interlayer spaces increased (Komadel et al., 1996a). The dissolution of reduced-charge Mt in HCl indicated that pyrophyllite-like layers surrounded by non-swelling interlayer spaces dissolved more slowly than Mt layers of similar chemical composition located between swelling interlayer spaces (Fig. 10.1.2). This clearly showed that protons attacked the layers from the swollen interlayer spaces also. Non-swelling illite and kaolinite were also more resistant to HCl attack than Mt or vermiculite (Jozefaciuk and Bowanko, 2002).

Figure 10.1.2. Ca^{2   +}-saturated 100% smectite and a reduced-charge smectite with about 10% swelling interlayers. Left: dissolution of Al^{3   +} in 6   M HCl at 95   °C; Right: pyrophyllite-like features in the IR spectra.

From Komadel et al. (1996a).

An in situ observation by AFM showed that the dissolution of hectorite and nontronite in acid solutions occurred inward from the edges; the basal surfaces were unreactive. The hectorite (0   1   0) faces dissolved more slowly than the lath ends. The edges dissolved on all sides and roughened. The (010), (110) and (110) faces on nontronite were stable. Dissolution fronts originating at the edges or defects would quickly became fixed along these faces, after which no more dissolution was observable. All the oxygen atoms on the nontronite stable edge faces were saturated, whereas the connecting oxygen atoms on all hectorite edge faces and nontronite edges were coordinatively unsaturated. This difference in reactivity of these faces suggested that the rate-limiting step of the dissolution process was the breaking of the bonds of connecting oxygen atoms (Bickmore et al., 2001).

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ACIDIFICATION

Bill Freedman , in Environmental Ecology (Second Edition), 1995

Concentration of Soil Solutes.

Soil acidification is caused by the accumulation of soluble inorganic and organic acids, at a faster rate than they can be neutralized. Ionization of these acids produces free H⁺ in the soil solution. Exchangeable H⁺ from cation-exchange surfaces also contributes to acidity. The relative contribution of exchangeable H⁺ is influenced by the total osmotic strength of the soil solution; the more concentrated the extracting solution, the greater the amount of H⁺ that is exchanged.

This fact is important in the measurement of soil pH. If distilled water is used as an extracting solution, the measured pH value is typically 0.5–1.0 pH units higher than that obtained when using a standardized-salt solution as an extractant, usually 0.01 M CaCl₂ (Russell, 1973; Allen et al., 1974).

Similarly, in natural ecosystems the acidity of soil may be influenced by temporal and spatial variations in the total concentration of solutes. In addition, locations with relatively large atmospheric inputs of total ions (e.g., a site near the ocean that is subject to a large rate of sea-salt deposition) may have a larger exchange acidity in soil than more continental sites (Rosenqvist, 1978a,b; Seip and Tollan, 1978).

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Recycling of Li-Ion Batteries for Electric Vehicles

Panpan Xu , ... Zheng Chen , in Reference Module in Earth Systems and Environmental Sciences, 2021

2.2 Hydrometallurgical method

Hydrometallurgical recycling of spent LIBs involves the use of a potent leaching agent like inorganic acids paired with a reducing agent like H ₂O₂ (hydrogen peroxide) depending on the cathode material being recycled. The leaching agent extracts the metal ions from the cathode powder while the reducing agent reduces the oxidation states of particular metals to facilitate facile leaching. The overall hydrometallurgical recycling process is illustrated in Fig. 3A . Before acid leaching, the spent LIBs first undergo mechanical separation, involving discharging, dismantling and separation of the cathode active materials.

Fig. 3. Schematic of hydrometallurgical recycling (A). Regeneration process of mixed spent cathode material after leaching process (B).

Reproduced with permission from Ref. Natarajan S and Aravindan V (2018) Recycling strategies for spent Li-ion battery mixed cathodes. ACS Energy Letters 3: 2101–2103. Copyright 2018 American Chemical Society.

Discharging can be done through physical or chemical routes (He et al., 2017; Nie et al., 2015). In a physical discharging process, spent LIBs are perforated to induce short circuit and release residual charges. The whole process is performed under the cooling protection of liquid nitrogen to dissipate excess heat generated during the short circuit. However, such processes require expensive equipment to handle liquid nitrogen as well as provide sufficient insulation. Chemical discharging on the other hand places the spent LIBs in a conductive solution, such as NaCl (sodium chloride) and Na₂SO₄ (sodium sulfate) to release the remaining energy. For large-scale treatment in the industry, a simple and low-cost chemical discharging is more suitable. The discharged spent LIBs are then subject to crushing followed by separation of cathode particles from plastics, metals as well as anode through density separator and froth flotation. There are numerous methodologies and steps followed for the pretreatment of the batteries prior to leaching but their main aim lies in efficiently separating out the pure cathode active material to enable the leaching process with minimal impurities.

Strong inorganic and mineral acids such as H₂SO₄ (sulfuric acid), HCl (hydrochloric acid) and HNO₃ (nitric acid) are commonly utilized coupling them with powerful reducing agents such as H₂O₂ and NaHSO₃. Initial work in the field focused on the recycling of LCO batteries because of their widespread use in electronics, requiring an urgent solution to their accumulation. It was found that H₂SO₄ coupled with varying concentrations of H₂O₂ can attain very high recovery efficiencies of upwards of 90% for both Co and Li. In these studies, optimal leaching conditions included optimizing concentrations of acid and reducing agent, leaching time, temperature and solid to liquid ratio in order to maximize leaching efficiencies (Sun and Qiu, 2011; Swain et al., 2007).

With increasing use of other battery chemistries like NCM (LiNi_xCo_yMn_zO₂, x   +   y   +   z   =   1) and NCA (LiNi_xCo_yAl_zO₂, x   +   y   +   z   =   1) in EVs, the need to recycle such materials has become of paramount importance. Naturally, recycling studies have shifted their focus towards applying the same H₂SO₄-H₂O₂ system to these chemistries. Results reported show extremely efficient recovery rates and under optimal conditions demonstrating leaching efficiencies of 98%, 99%, 98% and 98% for Co, Li, Ni and Mn respectively (Chen and Ho, 2018).

Despite the high efficiency, a major barrier for hydrometallurgical recycling is the use of these potent, caustic and environmentally hazardous acids which are particularly dangerous to handle at industrial scales. Seeking to make the process more eco-friendly, organic acids became the center of focus to replace the dangerous inorganic acids. Particularly, Li's group have extensively demonstrated the practicality of organic acids for the process, in one of their studies they attempted to recycle LCO waste batteries using CA (citric acid) coupled with H₂O₂ and were able to attain leaching efficiencies greater than 90% for both Co and Li, a performance comparable to that of its inorganic counterparts (Li et al., 2010, 2012). Evidently the use of organic acids appears to be not only a sustainable but also an effective option for the recycling of lithium ion batteries. The only limiting factor for it is the high cost of organic acids which makes it significantly more expensive to use at industrial scales. Another alternative is the usage of alkaline agents for the leaching process, among these a system comprising NH₃ ·H₂O, NH₄HCO₃ and Na₂SO₃ proved to be quite effective reaching leaching efficiencies of 91% for Co and 97% for Li while sole NH₃ ·H₂O showed efficiencies of only up to 80% (Ku et al., 2016; Qi et al., 2020). Although alkaline agents are seemingly an equally inexpensive alternate, they are plagued by the same issue of being highly caustic and at the same time showing comparatively lower leaching efficiencies making them less suitable candidates. To promote eco-friendly processes, bio-leaching was explored which involves the use of living organisms for metal extraction. When Aspergillus Niger-a fungal species was used for the leaching process-moderate efficiencies of 95% for Li, 70% for Mn, 65% for Al and comparatively lower efficiencies of 45% for Co and 38% for Ni were observed (Horeh et al., 2016). Bioleaching, despite being environmentally sustainable is extremely time consuming and uses living matter which is difficult to maintain and use at industrial scales. So, even though the alternatives seem lucrative with respect to their eco-friendly appeal, it is unlikely to replace the existing industrial hydrometallurgical practices with present technologies.

The metal leaching is then followed by solvent extraction and precipitation (or salt precipitation) to separate products. Solvent extraction involves the use of a chemical agents in order to separate out the various metals from the leached solution. Common solvents include Na-Cyanex 272 (dialkyl phosphinic acid extractant used) for the separation of Co and Ni, Na-D2EHPA (a sodium salt of di 2-ethylhexyl phosphoric acid) to separate Co and Mn and DMG (dimethylglyoxime) used often as an analytical reagent, hereto segregate Ni and isolate Li with purities upward of 90% obtained for all the elements (Chen et al., 2015; Kang et al., 2010). In the following step metal ions are precipitated out as metal hydroxides or carbonates and as an alternative solely chemical precipitation can be done using NaOH (sodium hydroxide) to regulate the pH and then at particularly pH with the addition of Na₂CO₃ (sodium carbonate) the carbonate metal salts can be precipitated, using this method greater than 90% of the metals were found to be precipitated from the solution (Fu et al., 2020; Nayl et al., 2017).

The hydrometallurgical recycling method can also work well to leach spent cathodes mixtures sourced from different battery chemistries. As shown in Fig. 3B and 6 mixed spent cathode materials are leached and the metal salts which will serve as precursors are then precipitated; these precursors are then calcined with LiOH to obtain regenerated Mn and Ni rich cathode materials-making the recycling loop fairly direct and simple.

Notably, the hydrometallurgical leaching and precipitation steps are also combined with pyrometallurgy to leach the alloy generated from the high-temperature smelting process. However, the final products for both the recycling technologies are metal hydroxides or carbonates precursors. For the synthesis of fresh cathode materials, addition of a lithium source along with these precursors is required to complete the recycling loop. Hydrometallurgical recycling though showing extremely high efficiencies and currently being the industrial favorite is limited in terms of its long-term environmental sustainability and practicality due to the large number of steps involved. An alternative is needed to reduce the number of steps and materials input required to recycling spent battery materials.

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Biomass and Biofuel Production

A.A. Refaat , in Comprehensive Renewable Energy, 2012

5.13.4.7.3(iii) Organosolv

In this approach an organic or aqueous–organic solvent mixture is used with addition of an inorganic acid catalyst (H ₂SO₄ or HCl), which is used to break the internal lignin and hemicellulose bonds. The hydrolyzed lignin is thus dissolved and recovered in the organophilic phase. However, this acid addition can be avoided for a satisfactory delignification by increasing process temperature (above 185   °C).

Solvents that are used are typically methanol, ethanol, acetone, ethylene glycol, triethylene glycol, and phenol. Due to the volatility of organic solvents, some of these substances are explosive and highly inflammable and thus difficult to handle. So, organosolv pretreatment must be performed under extremely tight and efficient control [233]. Examples of different solvents used for organosolv pretreatment are given in Table 9 .

Table 9. Organosolv pretreatment using different solvents and feedstock

Substrate	Solvent	Results	Reference
Wheat straw	Glycerol	Cellulose recovery 95%; lignin removal >   70%; digestibility 92%	[237]
	Glycerol	More effective than steam explosion	[238]
	Glycerol (crude)	Digestibility >   75% – cost reduction on expense of delignification a	[239]
Sugarcane bagasse	Ethanol–H2SO4	Ethanol yield 92.8%	[241]
Lignocellulosic residues	Acetone–H3PO4	Ethanol yield >   93%	[242]
Hybrid poplar	Ethanol	Digestibility ∼   85%	[234]
Beetle-killed lodgepole pine	Ethanol	Digestibility ∼   97%	[235]
Softwood (Loblolly pine)	Ethanol	Reduced cellulose crystallinity (CrI reduced from 63% to 53%)	[236]
Softwood (Pinus radiata)	Acetone–water	Ethanol yield 99.5%	[240]

a

It was recommended to remove lipophilic compounds from crude glycerol before utilization to overcome decreased delignification.

Compared to other chemical pretreatments the main advantage of organosolv process is the recovery of relatively pure lignin as a byproduct. Most of the hemicellulose and lignin are solubilized, but the cellulose remains as solid [233]. However, solvents need to be separated because they might be inhibitory to enzymatic hydrolysis and fermentative microorganisms [190]. Removal of solvents from the system is necessary using appropriate extraction and separation techniques, for example, evaporation and condensation, and they should be recycled to reduce operational costs.

The high commercial price of solvents is another important factor to consider for industrial applications. For economic reasons, among all possible solvents, the low-molecular weight alcohols with lower boiling points such as ethanol and methanol are favored. Although organosolv pretreatment is more expensive at present than the leading pretreatment processes, it can provide some valuable byproducts. An integral optimization and utilization of byproducts might lead the organosolv pretreatment to be a promising one for biorefining LC feedstock in the future [233].

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Polymers for Advanced Functional Materials

C. Kröhnke , in Polymer Science: A Comprehensive Reference, 2012

8.14.8 Acid Scavengers

Acid scavengers (other technical expression 'antacids') are commonly metal salts of weak organic or inorganic acids. Their corresponding free bases are able to efficiently neutralize acidity. The efficiency of an acid scavenger is determined by the reactivity of the salt achieved in the polar polymer matrix as well as by the acidity of impurities that expel a weaker acid from its salt. Besides traditional salts of fatty acids, for example, Ca-stearate and Zn-stearate, inorganic compounds such as hydrotalcites or still zinc oxide are frequently used. Major reasons to use antacids are the presence of catalyst residues in polymer matrices that can generate free acidity during or after catalyst deactivation by steam stripping or solvent treatment. Antacids neutralize this acidity and prevent a couple of undesired side effects such as corrosion of processing equipment. ⁵¹ Furthermore, particularly the stearate types of acid scavengers can fulfill functioning as slipping agents by reducing shear forces that can be important especially for processing of high-molecular-weight polymers. In addition, sodium and calcium salts of fatty acids such as Ca-stearate influence the crystallization behavior of several technical polymers such as polyolefins, polyethylene terephthalate, and polyamides exhibiting nucleating effects, accelerating crystallization kinetics, and enhancing mechanical properties of the corresponding articles. Besides such main effects, metal stearates also act as lubricants and release agents.

Finally, improvement of the performance of piperidine-based HA(L)S as well as their rectified resistance against degradation by external pollutants, for example, pesticides in greenhouses should be mentioned. ⁵²

Drawbacks of degradation processes of polymers can be restrained by the use of suitable stabilizers as already discussed above. Those stabilizers and stabilizer combinations are usually employed together with co-additives used for the stabilization of a given polymer grade. Base additive packages required for the stabilization of polymers, particularly polyolefins, comprise usually combinations of phenolic antioxidants, phosph(on)ites, and acid scavengers. The cooperative performance of such additive combinations is certainly influenced by a proper choice and concentration of all individual components. ⁵³

Addition of zinc stearate is known to be technically a superior acid scavenger since it helps to avoid early discoloration of a large series polymer formulations, but it holds for being physically irritant in humans. ⁵⁴ Exposure to zinc stearate over a long time may develop extensive fibrosis. Although there is no specific information available on the concentration of the exposure leading to such a condition, it is believed that it is very high. Moreover, aspiration of zinc stearate by infants can be associated with respiratory distress and acute pneumonitis. ⁵⁵

Another group of organic antacids are represented by metal lactates, especially calcium lactate ( Figure 26 ) and calcium stearoyl-2-lactylate. Besides their action principle as acid scavengers, these derivatives are able to form chelate complexes even with traces of residual aluminum and titanium in polymer matrices. Their addition can help to avoid discoloration of polymer formulations, particularly in combination with phenols.

Figure 26. Chemical structure of calcium lactate.

Synthetic hydrotalcites ⁵⁶ as inorganic acid scavengers recently became important co-stabilizer in polymer formulations. They are primarily utilized in replacing heavy metal-based stabilizing components such as lead stearate, lead phosphite, dibasic lead phthalate, or tribasic lead sulfate and the respective cadmium salts.

Mainly two types of hydrotalcites have proven excellent performance in long-term stabilization of polymers, namely, a pure Mg/Al hydrotalcite, empirical formula [Mg₆Al₂(OH)₁₆CO₃·4H₂O], (e.g., the commercial products Sorbacid® 911 and Hycite® 713 available from Süd-Chemie AG) and a zinc-containing derivative, [Zn₂ Mg₄Al₂(OH)₁₆CO₃·4H₂O] (like Sorbacid® 944 of Süd-Chemie AG).

Common to all types are their layered structure (see Figure 27 ) and small particle size (typically 80% <   1   µm) providing excellent dispersibility in the polymer matrix.

Figure 27. Layered structure of hydrotalcites (schematic view).

Hydrotalcites act in scavenging acidic decomposition products of halogenated polymers such as PVC, polychloroprene (CR), chlorosulfonated polyethylene (CSM), chlorinated polyethylene (CPE), epichlorohydrin (ECO), fluoroelastomers (FKM), and halobutyls (bromobutyl rubber (BIIR), chlorobutyl rubber (CIIR)).

In PP and polyethylene, hydrotalcites serve in immobilizing and neutralizing acidic catalyst residues derived from the Ziegler-Natta polymerization process.

Hydrotalcites have been added to flame-retardant systems as an inorganic stabilizer and smoke suppressant, especially with halogenated systems. These additives are also used in electrical applications for its antitracking properties. ⁵⁷

Antacids serve several purposes in additive formulations for polymers such as neutralizing the free acidity of catalyst residues after their deactivation by steam stripping or solvent treatment.

But these compounds also act as internal slipping agents in order to reduce shear forces during extrusion which is important especially for processing of high-molecular-weight polyolefins and production of polyolefin films in general influencing the crystallization behavior of polymers such as polyolefins as well as some engineering plastics such as polyethylene terephthalate and polyamides. Therefore, they exhibit certain nucleating effects, that is, acceleration of crystalline kinetics and enhancement of mechanical properties of finished polymer articles. Another function is the improvement of dispersion of stabilizers in the polymer matrix. In addition, acid scavengers can play an important role for the melt viscosity retention during processing as well as for the long-term stability of the final polymer article.

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Give Two Examples of Mineral Acids

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