Do alkaline earth metals form cations or anions?

Post-column Reaction Detection of Cations

While alkali metal and alkaline earth cations are detected routinely using conductivity detection, transition metals and lanthanoids are most commonly detected using a post-column reaction (PCR). In most cases, this involves a post-column addition of a colour-forming ligand, generally a metallochromic dye. Typical examples of such dyes are 4-(2-pyridylazo)resorcinol (PAR), which is used for the detection of transition metals, and 2,7-bis(2-arsonophenylazo)-1,8-dihydroxynaphthalene-3,6-disulfonic acid (Arsenazo III), which is used for the detection of lanthanoids. Both dyes react rapidly with a wide range of metal ions to form strongly absorbing complexes that facilitate sensitive detection without the need for complicated reactors or mixing devices. Figure 12 shows a typical chromatogram obtained using Arsenazo I as the PCR reagent.

1.1.2.2 Coordination Geometry and Number

The directional preferences for coordination to the alkali metal and alkaline earth cations is obviously related to the number of substituents coordinated to the cation. As yet there is little predictability of the coordination number among these cations. For example, the first member of this series, the Li+ cation, is the best characterized with well over 500 X-ray crystal structures containing this ion. Coordination numbers to Li+ ranging from two through seven and all values in between can be found. The Li+ cation is also found symmetrically π-complexed to the faces of aryl anions and to conjugated linear anions (see ref.11). At present enough evidence exists to deduce only that the coordination number to the alkali metal and alkaline earth cations, and consequently the coordination geometry about these cations, is governed primarily by steric factors. Unfortunately the predictability of any individual unknown structure is relatively low.

In general the metal cation to substituent distances are found spanning a range of values. A working criterion for coordination to the metal cations is that the M

Table 1. Carbanion–Metal Bond Lengths

BondMeanS.D.Minimum (Å)Maximum (Å)NobsC

This table includes all examples listed in the CSD (version 4.20, 1990) located by a fragment search (i.e. CONNSER) for C

Related structural aspects of metal ion coordination geometry are covered in some recent publications and are worthy of note. The directional preferences of ether oxygen atoms towards alkali and alkaline earth cations are reported by Chakrabarti and Dunitz.37 The conclusion of this work is that the larger cations show an apparent preference to approach the ether oxygen along a tetrahedral lone pair direction, whereas Li+ cations tend to be found along the C

9.3.4.5.3 Dissolution inhibitors

As is true for growth, the capacity of divalent alkaline earth cations to inhibit calcite dissolution rate is also complex (see also Morse and Arvidson, 2002), site specific, and not independent of dissolved carbon concentrations.

Manganese. Mn(II) has been shown to interact strongly with the calcite, effectively halting the dissolution process at concentrations of ≲2 μM (Arvidson et al., 2003; Lea et al., 2001). In addition, at 2 μM [Mn2 +], Lea et al. (2003) observed epitaxial nucleation with the AFM of a MnxCa1 − xCO3 solid solution on the calcite cleavage surface parallel to the 22¯1direction. These oriented overgrowths could be seen forming simultaneously with dissolution of the host calcite. Vinson et al. (2007) also observed a zero rate at Mn(II) concentrations of ~ 2 μM, but noted that the effectiveness of Mn(II) as an inhibitor was critically dependent on dissolved carbon concentration. In fact, in solutions where carbonate ion is submicromolar, Mn(II) additions of 2 μM yield an ‘increase’ in overall dissolution rate. The increased destabilization of the surface in the presence of the divalent metal ion where dissolved carbonate ion is lacking is also observed for magnesium (see below).

Strontium. The effect of strontium on calcite dissolution is quite different from either manganese or magnesium (discussed below). Vinson and Luttge (2005) observed that the additions of strontium up to 250 μM to carbonated solution (pH ~ 8.7) led to passivation of the calcite surface, with reductions in the extent of etch-pit nucleation and step velocity, and increases in terrace width; 4¯41−steps showed greater relative reduction compared to 4¯41+steps. Nucleation of a strontian phase within etch pits was also described (see also Astilleros et al., 2003). However, regardless of the changes associated with etch-pit populations induced by strontium additions, the ‘overall’ retreat of the surface measured by VSI showed almost negligible reduction. This result underscores the role of rough (kink-rich) monolayer steps, typically originating at etch-pit margins, as rate defining. In the larger view, it emphasizes the importance of understanding how the surface interacts as a whole during dissolution: rate observations restricted to etch pits alone may not express the dissolution rate.

Magnesium. Not surprisingly, the interactions of magnesium ion with the dissolving calcite surface are also quite complex. Many of the descriptions of its inhibitory role during growth are largely thermodynamic (e.g., incorporation leading to enhanced solubility, Berner, 1975; Davis et al., 2000; Wasylenki et al., 2005b), but the inhibition of dissolution cannot derive from the same mechanism. In combined AFM and VSI work, Arvidson et al. (2006) observed significant inhibition of calcite dissolution by magnesium, although, as with manganese inhibition of Vinson and Luttge (2005), the relationship between etch-pit generation, morphology, and overall dissolution rate is complex. In carbonated solutions, inhibition is observed at [Mg2 +] of 50 μM; at higher concentrations distinct pinning of 4¯41+steps is observed, with essentially zero step speeds at 0.8 mM; by contrast, 4¯41−step speeds are little changed. This condition produced unique etch-pit morphologies, with curved step obtuse edges produced by acquisition of excess, pinned 4¯41+/+and 481¯+/+kink sites. These features were not observed in AFM work by Xu and Higgins (2011) (see below).

Two additional observations are of interest here. First, in the experiments described above, addition of magnesium at concentrations of 0.8 mM, although clearly reducing (obtuse) step speeds, seems to ‘activate’ the surface, bringing about a rapid increase in the rate of etch-pit nucleation. The effect was rapid and reproducible, with removal of Mg burden bringing about an immediate restoration of smoother surfaces, suggesting an impermanent residence on the surface sensitive to dissolved concentrations. Second, the addition of magnesium at the same concentration (0.8 mM) to carbon-free solutions had little inhibitory effect, similar to observations of Alkattan et al. (2002) in the acid regime.

The effect of dissolved magnesium on calcite dissolution was explored with the AFM at higher concentrations than those above (total Mg > 1 M, with negligible carbonate alkalinity) by Ruiz-Agudo et al. (2009). Their AFM results show that increasing Mg concentrations, while producing some reductions in step velocity, are able to drive enhanced overall dissolution, due to a higher nucleation rate of new pits. The effect was most pronounced in the case of MgSO4 addition, suggesting a possible role for sulfate in the dissolution mechanism, at least in the absence of carbonate ion.

The inhibition described above under highly undersaturated becomes still more complex as equilibrium is approached. In 50 ᵒC AFM experiments over the range of 0.05 ≤ Ω ≤ 1.2, Xu and Higgins (2011) observed no inhibition for [Mg2 +] concentrations of 0.100 mM, but abrupt arrest of ‘both’ acute and obtuse steps at [Mg2 +] = 1.00 mM and Ω ≥ 0.2. The authors attributed this behavior to a mechanistic ‘switch’ defined by a critical ΔG. This critical point, above which (i.e., 0 > ΔG > ΔGcrit) the residence time of Mg attachment at step edges exceeds the reciprocal of either the nucleation rate of new kinks or the detachment rate of extant kinks, could in fact define two mechanistic regions. This relationship between ΔG and rate is similar to that identified for silicates (e.g., albite, see Figure 2 in Arvidson and Luttge, 2010; Lüttge, 2006), where, under conditions sufficiently close to equilibrium, there is insufficient chemical potential to open hollow cores at screw dislocations; in this case, Al3 + may occupy a role analogous to Mg2 +.

8.3.2.4 Scaling (Lenntech)

Scaling means the deposition of particles of salts such as calcium carbonate, calcium sulfate, and sulfates of other alkaline earth cations. Most natural waters contain relatively high concentrations of calcium, sulfate, and bicarbonate ions. When a large amount of water is recovered by RO, the concentration of gypsum (calcium sulfate dehydrate) and calcite (calcium carbonate) in the retentate becomes more than their solubility limit in water leading to precipitation and crystallization on the membrane surface. The surface blockage causes flux decline as well as narrowing the feed water passage in the RO module. There are three ways of preventing scaling. Acidification, ion exchange, and addition of antiscalant. Carbonate ions are destroyed by adding acid. Thus, acidification is effective to reduce calcium carbonate but not very effective to remove calcium sulfate. By ion exchange, the divalent cations are exchanged to monovalent sodium ion, thus preventing the precipitation of calcium carbonate and sulfate, etc. Antiscalant interferes the salt precipitation in the following three different ways:

Threshold inhibition: the supersaturation of salt in water is maintained.

Crystal modification: the antiscalant distorts the shape of salt crystals.

Dispersion: antiscalants are adsorbed to the surface of colloidal particles to impart high anionic charges. Thus, crystals are kept separated.

Antiscalants are mostly proprietary mixtures of various molecular weight polycarboxylates and polyacrylates.

9.30.9.2.1.(i) Template syntheses

Few template syntheses seem efficacious for crown thioethers (although the potential of these methods has probably not been fully explored). Alkali and alkaline earth cations, successful templates for polyether macrocycles, have no utility here as they do not bind the α,ω-dithiolate precursors. Occasionally, transition-metal fragments are used as templates, but in most cases the potential starting dithiolate complexes are bridged polynuclear species that tend to be neither soluble nor particularly nucleophilic 〈52JCS146, 78IC1296, 86JCS(D)1747, 91IC3700, 91ZN(B)209〉. In the most noteworthy exceptions, 9S3 has been synthesized efficiently around a Mo(CO)3 template 〈84AG(E)807, 85JOM(289)57〉 and a template method for some S4 crowns has been described 〈70JA1935〉. Reaction of [(dpttd)Fe(CO)] (dpttdH2 = 2,3,11,12-dibenzo-1,4,7,10,13-pentathiatridecane) with 1,5-dibromo-3-thiapentane (Equation (3)) gives dibenzo-18S6 in 55–75% yield 〈86AG(E)1107〉; preparation of the tetra(t-butyl) analogue proceeds in 62% yield 〈90ICA231〉. 2,3,8,9-Dibenzo-15S5 can be prepared either from [(dpttd)Fe(CO)] and dibromoethane or (better) from [(dttd)Fe(CO)2] and 1,5-dibromo-3-thiapentane (dttdH2 = 2,3,8,9-dibenzo-1,4,7,10-tetrathiadecane).

(3)

1 INTRODUCTION

The alkaline‐earth fluoride crystals show space group Oh5, where the alkaline‐earth ions are eightfold coordinated. Trivalent rare‐earth ions (RE3+) readily substitute for the divalent alkaline‐earth cations and charge compensation is required. In CaF2:Sm3+ and SrF2:Sm3+ crystals, the well‐known C4v(F‐) center is predominant [1]. This center is composed of a Sm3+‐F‐ pair with the charge‐compensating fluorine ion located in the nearest‐neighbor position along the [001] direction from the Sm3+ ion [1–4]. After oxidization, these samples have trigonal C3v(O2‐) centers, consisting of a Sm3+‐O2‐ pair with the charge‐compensating oxygen ion located in the nearest‐neighbor position along the [111] direction from the Sm3+ ion [1,2].

The effect of co‐doping CaF2:RE3+ and SrF2:RE3+ crystals with LiF, NaF or KF was reported during the 1960ʼs by groups in the former Soviet Union, see for example [5], and more recently by Jones, Reeves and co‐workers [6,7]. From a combination of infrared absorption and laser selective excitation spectroscopy, the later workers illustrated profound changes to the defect distribution; namely that the population of regular C4v(F‐) centers is reduced/eliminated, cubic centers with remote charge compensation are significantly enhanced, and new orthorhombic symmetry centers with monovalent alkali ions located in the Ca2+ or Sr2+ site in the [110] direction from the RE3+ ion are created.

This paper presents electron paramagnetic resonance (EPR) spectroscopy upon Sm3+ centers in CaF2:Sm3+ and SrF2:Sm3+ crystals co‐doped with Li+ or Na+ in order to better understand the microscopic structure, since the anisotropy of the g tensors yields indirect information on the possible ligand positions. From this, we propose models of the Sm3+‐Li+ /Na+ centers.

Do alkaline earth metals like to form cations?

What type of ions do alkaline earth metals form?

Can alkali metals form an anion?

Do alkali metals make cations?