6.14: Group 13 Halides

Trihalides, MX₃

As shown in Table \(\PageIndex{1}\) all the combinations of Group 13 element (M) and halogen (X) exist for the trihalides (MX₃), except thallium(III) iodide. It should be noted that while there is a compound with the general formula TlI₃, it is actually a thallium(I) compound of I₃^-.

Table \(\PageIndex{1}\): Selected physical properties of the Group 13 trihalides, MX₃.

Element	Mp (°C)	Bp (°C)
BF3	-126.8	-100.3
BCl3	-107.3	12.6
BBr3	-46.3	91.3
BI3	49.9	210
AlF3	1291	-
AlCl3	192.4 (anhydrous), 0.0 (hexahydrate)	120 (hexahydrate)
AlBr3	97.8	265
AlI3	189.4 (anhydrous) 185 dec. (hexahydrate)	300 subl.
GaF3	800	1000
GaCl3	77.9	201
GaBr3	121.5	278.8
GaI3	212	345
InF3	1172	-
InCl3	586	800
InBr3	220	-
InI3	210 subl.	-
TlF3	300 dec.	-
TlCl3	40 dec.	-
TlBr3	40 dec.	-

The trihalides of boron are all monomers with a coordination number of 3 (Table \(\PageIndex{2}\)), as evidence from their low melting points (Table \(\PageIndex{1}\)). In contrast, the fluorides and chlorides of the heavier Group 13 elements (except GaCl₃) are generally ionic or have a high ionic character, with a coordination number of 6 (Table \(\PageIndex{2}\), Figure \(\PageIndex{1}\) and Figure \(\PageIndex{2}\)). The bromides and iodides (except InBr₃) are generally dimeric with a coordination number of 4 (Table \(\PageIndex{2}\)) and have molecular structures involving halide bridging ligands (Figure \(\PageIndex{3}\) and Table \(\PageIndex{3}\)). AlCl₃ is unusual in that in the solid state it has an ionic structure, but it is readily sublimed, and in the vapor phase (and liquid phase) it has a dimeric structure (Figure \(\PageIndex{3}\)).

Table \(\PageIndex{2}\): The Group 13 element coordination numbers for the trihalides, MX₃.

Element	Fluoride	Chloride	Bromide	Iodide
B	3	3	3	3
Al	6	6 (4)	4	4
Ga	6	4	4	4
In	6	6	6	4
Tl	6	6	4	-

Figure \(\PageIndex{1}\): Solid state structure of MF₃ (M = Al, Ga, In). Grey spheres represent the metal, while pale blue spheres represent fluorine.

Figure \(\PageIndex{2}\): Solid state structure of MCl₃ (M = Al, In, Tl) and MBr₃ (M = In). Grey spheres represent the metal, while green spheres represent the halogen.

Figure \(\PageIndex{3}\): Structure of GaCl₃, MBr₃ (M = Al, Ga, Tl), and MI₃ (M = Al, Ga, In). Grey spheres represent the metal, while green spheres represent the halogen.

Table \(\PageIndex{3}\): Selected bond lengths and angles for dimeric M₂X₆ compounds. ^aX_t and X_b indicate terminal and bridging halides, respectfully.

Compound	M-Xt (Å)^a	M-Xb (Å)^a	Xt-M-Xt (°)^a	Xb-M-Xb (°)^a	M-X-M (°)
Al2Br6	2.21	2.33	115	93	87
In2I6	2.64	2.84	125.1	93.7	86.3

Synthesis

Boron trifluoride (BF₃) is manufactured commercially by the reaction of boron oxides with hydrogen fluoride, (6.14.1). 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, (6.14.2).

\[ \text{B}_2\text{O}_3 \text{ + 6 HF} \rightarrow \text{2 BF}_3 \text{ + 3 H}_2\text{O}\]

\[ \text{PhN}_2\text{[BF}_4\text{]} \rightarrow \text{PhF + BF}_3 \text{ + N}_2 \]

Boron trichloride is also made from boron oxide, but in the presence of carbon, (6.14.3).

\[ \text{B}_2\text{O}_3 \text{ + 3 C + 3 Cl}_2 \rightarrow \text{2 BCl}_3 \text{ + 3 CO}\]

Many of the trihalides are readily prepared by either the direct reaction of the metal with the appropriate halogen, (6.14.4) - (6.14.6), or the acid, (6.14.7) and (6.14.8). Thallium tribromide can be prepared in CH₃CN by treating a solution of the monobromide with bromine gas, (6.14.9).

\[ \text{2 Al + 3 Cl}_2 \rightarrow \text{2 AlCl}_3\]

\[ \text{2 Al + 3 Br}_2 \rightarrow \text{2 AlBr}_3\]

\[ \text{2 Al + 3 I}_2 \rightarrow \text{2 AlCl}_3\]

\[ \text{2 Al + 6 HCl} \rightarrow \text{2 AlCl}_3 \text{ + 3 H}_2\]

\[ \text{2 Al + 6 HBr} \rightarrow \text{2 AlBr}_3 \text{ + 3 H}_2\]

\[ \text{TlBr + Br}_2 \rightarrow \text{TlBr}_2 \]

Reactivity

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

• Lewis acid-base complex formation.
• Hydrolysis.

There are, however, a number of reactions involving halide exchange reactions. Aluminum tribromide reacts with carbon tetrachloride at 100 °C to form carbon tetrabromide, (6.14.10), and with phosgene yields carbonyl bromide and aluminum chlorobromide, (6.14.11).

\[ \text{4 AlBr}_3 \text{ + 3 CCl}_4 \rightarrow \text{4 AlCl}_3 \text{ + 3 CBr}_4 \]

\[ \text{AlBr}_3 \text{ + COCl}_2 \rightarrow \text{COBr}_2 \text{ + AlCl}_2\text{Br}\]

Group 13 halides are used as synthons for their organometallic derivatives, (6.14.12) and (6.14.13).

\[ \text{MX}_3 \text{ + 3 RMgX} \rightarrow \text{MR}_3 \text{ + MgX}_2\]

\[ \text{2 MR}_3 \text{ + MX}_3 \rightarrow \text{3 MXR}_2\]

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, (6.14.12). 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.

\[ \text{MX}_3 \text{ + L} \rightleftharpoons \text{X}_3\text{M-L}\]

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

Figure \(\PageIndex{4}\): 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 \(\PageIndex{5}\): 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 \(\PageIndex{3}\)) 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, (6.14.13). In the case of the fluoride, the HF formed reacts with BF₃ to form fluoroboric acid, (6.14.14). However, there is also a minor equilibrium (2-3%) resulting in the formation of the BF₃ complex of OH^- and H₃O⁺, (6.14.15).

\[ \text{BX}_3\text{ + H}_2\text{O} \rightarrow \text{B(OH)}_3\text{ + 3 HX}\]

\[ \text{HF + BF}_3 \rightarrow \text{HBF}_4\]

\[ \text{BF}_4 \text{ + H}_3\text{O} \rightleftharpoons \text{F}_3\text{B-OH}_2 \rightleftharpoons \text{[BF}_3\text{OH][BF}_3\text{(OH}_3\text{)]}\]

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 \(\PageIndex{6}\)). 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 \(\PageIndex{6}\): The structure of 3-methyl-2-hexenoic acid.

Boron trihalides: a special case

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:

\[ \text{BF}_3 \text{ < BCl}_3 \text{ < BBr}_3\text{(strongest Lewis acid)}\]

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 \(\PageIndex{7}\)). 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, (6.4.12).

Figure \(\PageIndex{7}\): 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 \(\PageIndex{7}\)a). In contrast, the overlap for the large (diffuse) p-orbitals on the chlorine is poorer (Figure \(\PageIndex{7}\)b). As a result, the π-donation of F is greater than that of Cl. Interestingly, as may be seen from Table \(\PageIndex{4}\), 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 \(\PageIndex{4}\)).

Table \(\PageIndex{4}\): The B-X bond distances in the boron trihalides, BX₃, as compared to the sum of the covalent radii. ^aCovalent radius of B = 0.84(3) Å.

Compound	B-X (Å)	X covalent radius (Å)	Sum of covalent radii (Å)^a	Δ (Å)
BF3	1.313	0.57(3)	1.41	0.097
BCl3	1.75	1.02(4)	1.86	0.11
BBr3	1.898	1.20(3)	2.04	0.142
BI3	2.125	1.39(3)	2.23	0.105

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

• Directional relationship of the orbitals.
• Relative symmetry of the orbitals.
• Relative energy of the orbitals.
• Extent of orbital overlap

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: (6.14.16). 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.

Anionic halides

The trihalides all form Lewis acid-base complexes with halide anions, (6.14.17), and as such salts of BF₄^-, AlCl₄^-, GaCl₄^-, and InCl₄^- are common.

\[ \text{MX}_3 \text{ + X}^- \rightarrow \text{MX}_4^-\]

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

\[ \text{2 GaCl}_4^- \rightleftharpoons \text{Ga}_2\text{Cl}_7^- \text{ + Cl}^-\]

Figure \(\PageIndex{8}\): 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 \(\PageIndex{8}\)).

Low valent halides

Oxidation state +2

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, (6.14.19). An alternative route is by the co-condensation of copper as a reducing agent with BCl₃, (6.14.20).

\[ \text{2 BCl}_3\text{ + 2 Hg} \rightarrow \text{B}_2\text{Cl}_4 \text{ + Hg}_2\text{Cl}_2\]

\[ \text{2 BCl}_3\text{ + 2 Cu} \rightarrow \text{B}_2\text{Cl}_4 \text{ + 2 CuCl}\]

B₂F₄ has a planar structure (Figure \(\PageIndex{9}\)) with D2h symmetry, while B₂Cl₄ has the same basic structure it has a staggered geometry (Figure \(\PageIndex{10}\)). 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 \(\PageIndex{9}\): The planar structure of B₂F₄.

Figure \(\PageIndex{10}\): The staggered structure of B₂Cl₄.

Oxidation state +1

Boron forms a number of halides with cluster structures, B_nCl_n where n = 4 (Figure \(\PageIndex{11}\)), 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 \(\PageIndex{11}\): Structure of B₄Cl₄.

Intermediate halides

The dihalides (MX₂) of gallium, indium, and thallium do not actually contain the metal in the +2 oxidation state. Instead they are actually mixed valence compound, i.e., M⁺[MX₄]^-. The dihalides of gallium are unstable in the presence of water disproportionating to gallium metal and gallium(III) entities. They are soluble in aromatic solvents, where arene complexes have been isolated and the arene is η⁶-coordinated to the Ga⁺ ion. InBr₂ and InI₂ are greenish and yellow crystalline solids, respectively, which are formulated In(I)[In(III)X₄]. TlCl₂ and TlBr₂ both are of similar formulations.

Ga₂X₃ (X = Br, I) and In₂Br₃ are formulated M(I)₂[M(II)₂X₆]. Both anions contain a M-M bond where the metal has a formal oxidation state of +2. The Ga₂Br₆^2- anion is eclipsed like the In2Br6^2- anion, whereas the Ga₂I₆^2- anion is isostructural with Si₂Cl₆ with a staggered conformation. In₂Cl₃ is colorless and is formulated In(I)₃[In(III)Cl₆].

Ga₃Cl₇ contains the Ga₂Cl₇^- ion, which has a structure similar to the dichromate, Cr₂O₇^2-, ion with two tetrahedrally coordinated gallium atoms sharing a corner (Figure). The compound can be formulated as gallium(I) heptachlorodigallate(III), Ga(I)[Ga(III)₂Cl₇].

Figure \(\PageIndex{12}\): The crystal structure of Ga₃Cl₇, better described as gallium(I) heptachlorodigallate(III), Ga(I)[Ga(III)₂Cl₇]. Grey spheres represent the gallium, while green spheres represent the chlorine.

In₄Br₇ is light sensitive (like TlCl and TlBr) decaying to InBr₂ and In metal. It is a mixed salt containing the InBr₄^- and InBr₆^3- anions balanced by In⁺ cations. It is formulated In(I)₅[In(III)Br₄]₂[In(III)Br₆]. In₅Br₇ is a pale yellow solid formulated as In(I)₃[In(II)₂Br₆]Br. The In(II)₂Br₆^2- anion has an eclipsed ethane like structure with an In-In bond length of 2.70 Å. In₅Cl₉ is formulated In(I)₃[In(III)₂Cl₉], with the In₂Cl₉^2- anion having two 6 coordinate indium atoms with 3 bridging chlorine atoms, face sharing bioctahedra. Finally, In₇Cl₉ and In₇Br₉ have a structure formulated as InX₆[In(III)X₆]X₃.

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