Boron trifluoride (BF₃) is manufactured commercially by the reaction of boron oxides with hydrogen fluoride, Equation 1. The HF is produced in-situ from sulfuric acid and fluorite (CaF₂). On smaller scales, BF₃ is prepared by the thermal decomposition of diazonium salts, Equation 2.

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Boron trichloride is also made from boron oxide, but in the presence of carbon, Equation 3.

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Many of the trihalides are readily prepared by either the direct reaction of the metal with the appropriate halogen, Equation 4 - Equation 6, or the acid, Equation 7 and Equation 8. Thallium tribromide can be prepared in CH₃CN by treating a solution of the monobromide with bromine gas, Equation 9.

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The reaction chemistry of the Group 13 trihalides tends to fall into two categories:

There are, however, a number of reactions involving halide exchange reactions. Aluminum tribromide reacts with carbon tetrachloride at 100 °C to form carbon tetrabromide, Equation 10, and with phosgene yields carbonyl bromide and aluminum chlorobromide, Equation 11.

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Group 13 halides are used as synthons for their organometallic derivatives, Equation 12 and Equation 13.

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Lewis acid-base complexes

All of the trihalides are strong Lewis acids, and as such react with Lewis base compounds to form Lewis acid-base complexes, Equation 14. The extent of the equilibrium is dependant on the Lewis acidity of the trihalide and the basicity of the Lewis base. For example, with BCl₃, oxygen donor ligands result in approximately 50:50 ratio of BCl₃ and BCl₃L, while for nitrogen donor ligands the equilibrium is shifted to the formation of the complex.

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The general structure of the Lewis acid-base complexes is such that the Group 13 element is close to tetrahedral (Figure 4). However, for aluminum and the heavier Group 13 elements, more than one ligand can coordinate (Figure 5) up to a maximum of six.

Figure 4: Molecular structure of a typical Group 13 metal trihalide Lewis acid-base complex: AlBr₃(9-fluorenone). Adapted from C. S. Branch, S. G. Bott, and A. R. Barron, J. Organomet. Chem., 2003, 666, 23.

Figure 5: Molecular structure of AlBr₃(9-fluorenone)₂. Adapted from C. S. Branch, S. G. Bott, and A. R. Barron, J. Organomet. Chem., 2003, 666, 23.

It should be noted that the dimeric form of MX₃ (Figure 3) can be thought of as mutual Lewis acid-base complexes, in which a Lewis basic lone pair of a halide on one MX₃ unit donates to the Lewis acidic metal on another MX₃ unit.

Hydrolysis

Generally the fluorides are insoluble in water while the heavier halides are more soluble. However, BF₃, BCl₃, and BBr₃ all decompose in the presence of water, Equation 15. In the case of the fluoride, the HF formed reacts with BF₃ to form fluoroboric acid, Equation 16. However, there is also a minor equilibrium (2-3%) resulting in the formation of the BF₃ complex of OH^- and H₃O⁺, Equation 17.

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While the boron compounds (and AlBr₃) decompose even in moist air, AlCl₃ reacts more slowly to make aluminum chlorohydrate (ACH) which has the general formula Al_nCl_3n-m(OH)_m. While ACH has been proposed to exist as a number of cluster species, it is actually a range of nanoparticles.

ACH is also known as polyaluminum chloride (PAC). The latter name is often used in water purification, where ACH is preferred over alum derivatives (Al₂(SO₄)₃). The combination of ACH and a high molecular weight quaternized ammonium polymer (e.g., dially dimethyl ammonium chloride (DADMAC)), has been known as an effective combination as a flocculant in water treatment process to remove dissolved organic matter and colloidal particles present in suspension.

Aluminum chlorohydrate (ACH) and aluminum-zirconium compounds, are frequently used as the active ingredient in antiperspirants. The mode of action of most aluminum-based compounds involves the dramatic change in the particle size from nano to micro as a function of pH and electrolyte changes on the skin (as compared to the antiperspirant stick or suspension) and hence forming a gel plug in the duct of the sweat gland. The plugs prevent the gland from excreting liquid and are removed over time by the natural sloughing of the skin. A further mechanism of action involves the absorption of 3-methyl-2-hexenoic acid (Figure 6). Human perspiration is odorless until bacteria ferment it. Bacteria thrive in hot, humid environments such as the underarm. When adult armpits are washed with alkaline pH soaps, the skin loses its acid mantel (pH = 4.5 - 6), raising the skin pH and disrupting the skin barrier. The bacteria thrive in the basic environment, and feed on the sweat from the apocrine glands and on dead skin and hair cells, releasing 3-methyl-2-hexenoic acid, which is the primary cause of body odor. As with all carboxylic acids, 3-methyl-2-hexenoic acid, reacts in a facile manner with the surface of the alumina nanoparticles. Aluminum chloride salts also have a slight astringent effect on the pores; causing them to contract, further preventing sweat from reaching the surface of the skin.

Figure 6: The structure of 3-methyl-2-hexenoic acid.

The three lighter boron trihalides, BX₃ (X = F, Cl, Br) form stable adducts with common Lewis bases. Their relative Lewis acidities can be evaluated in terms of the relative exothermicity of the adduct-forming reaction:

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This trend is opposite to that expected based upon the electronegativity of the halogens. The best explanation of this trend takes into account the extent of π-donation that occurs between the filled lone pair orbital on the halogens and the empty p-orbital on the planar boron (Figure 7). As such, the greater the π-bonding the more stable the planar BX₃ configuration as opposed to the pyramidalization of the BX₃ moiety upon formation of a Lewis acid-base complex, Equation 14.

Figure 7: Schematic representation of π-donation from filled halogen p-orbitals into empty p-orbital in the halides BX₃.

The criteria for evaluating the relative strength of π-bonding are not clear, however, one suggestion is that the F atom is small compared to the Cl atom, and the lone pair electron in p_z of F is readily and easily donated and overlapped to empty p_z orbital of boron (Figure 7a). In contrast, the overlap for the large (diffuse) p-orbitals on the chlorine is poorer (Figure 7b). As a result, the π-donation of F is greater than that of Cl. Interestingly, as may be seen from Table 4, any difference in B-X bond length does not follow the expected trend associated with shortening of the B-X bond with stronger π-bonding. In fact the B-Br distance is the most shortened as compared to that expected from the covalent radii (Table 4).

At the simplest level the requirements for bonding to occur based upon the molecular or atomic orbital are:

In the case of the boron trihalides, the direction (parallel) and symmetry (p-orbitals) are the same, and the only significant difference will be the relative energy of the donor orbitals (i.e., the lone pair on the halogen) and the extent of the overlap. The latter will be dependant on the B-X bond length (the shorter the bond the greater the potential overlap) and the diffusion of the orbitals (the less diffuse the orbitals the better the overlap). Both of these factors will benefit B-F over B-Cl and B-Br. Thus, the extent of potential overlap would follow the order: Equation 18. Despite these considerations, it is still unclear of the exact details of the rationalization of the low Lewis basicity of BF₃ as compared to BCl₃ and BBr₃.

The trihalides all form Lewis acid-base complexes with halide anions, Equation 19, and as such salts of BF₄^-, AlCl₄^-, GaCl₄^-, and InCl₄^- are common.

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In the case of gallium the Ga₂Cl₇^- anion (Figure 8) is formed from the equilibrium:

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Figure 8: Structures of the (a) M₂X₇^- and (b) M₂X₉^3- anions.

As a consequence of its larger size indium forms a wide range of anionic halides addition compounds with trigonal bipyramidal, square pyramidal, and octahedral coordination geometries. For example, salts of InCl₅^2-, InBr₅^2-, InF₆^3-, InCl₆^3- and InBr₆^3- have all been made. The InCl₅^2- ion has been found to be square pyramidal in the salt [NEt₄]₂InCl₅, but is trigonal bipyramidal in the acetonitrile solvate of [Ph₄P]₂InCl₅. The oligomeric anionic halides In₂X₇^- and In₂X₉^3- (X = Cl and Br) contain binuclear anions with tetrahedral and octahedrally coordinated indium atoms, respectively (Figure 8).

Boron forms a series of low oxidation halides containing B-B bonds a formal oxidation state of +2. Passing an electric discharge through BCl₃ using mercury electrodes results in the synthesis of B₂Cl₄, Equation 21. An alternative route is by the co-condensation of copper as a reducing agent with BCl₃, Equation 22.

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B₂F₄ has a planar structure (Figure 9) with D2h symmetry, while B₂Cl₄ has the same basic structure it has a staggered geometry (Figure 10). The energy for bond rotation about the B-B bond is very low (5 kJ/mol) that can be compared to ethane (12.5 kJ/mol). The bromide, B₂Br₄, is also observed to be staggered in the solid state. The staggered conformation is favorable on steric grounds, however, for B₂F₄ the planar geometry is stabilized by the smaller size of the halide, and more importantly the presence of strong delocalized π-bonding.

Figure 9: The planar structure of B₂F₄.

Figure 10: The staggered structure of B₂Cl₄.

Boron forms a number of halides with cluster structures, B_nCl_n where n = 4 (Figure 11), 8, 9, 10, 11, and 12. Each compound is made by the decomposition of B₂Cl₄. For gallium, none of the monohalides are stable at room temperature, but GaCl and GaBr have been produced in the gas form from the reaction of HX and molten gallium. The stability of thallium(I) as compared to thallium(III) results in the monohalides, TlCl, TlBr, and TlI being stable. Each compound is insoluble in water, and photosensitive.

Figure 11: Structure of B₄Cl₄.