21.4: Group 13: The Bor family (2023)

Production and general properties of group 13 elements

As reducing agents, the Group 13 elements are less potent than the alkali metals and alkaline earth metals. However, its compounds with oxygen are thermodynamically stable, and large amounts of energy are required to isolate even the two most accessible elements, boron and aluminum, from their oxide minerals.

Although boron is relatively rare (about 10,000 times rarer than aluminium), concentrated deposits of borax [Na₂B₄o₅(OH)₄8h₂O] occur in ancient lake beds (Figure \(\PageIndex{1}\)) and were used in antiquity for the production of glass and glass ceramics. Boron is produced on a large scale by reacting borax with acid to give boric acid [B(OH)].₃], which is then converted to the oxide (B₂o₃). Reduction of the oxide with magnesium or sodium produces amorphous boron that is only 95% pure:

\[\mathrm{Na_2B_4O_5(OH)_4\cdot8H_2O(s)}\xrightarrow{\textrm{Saure}}\mathrm{B(OH)_3(s)}\xrightarrow{\Delta}\mathrm{B_2O_3(s)} } \label{Eq1}\]

\[\mathrm{B_2O_3(s)}+\mathrm{3Mg(s)}\xrightarrow{\Delta}\mathrm{2B(s)}+\mathrm{3MgO(s)}\label{Eq2}\]

However, pure crystalline boron is extremely difficult to obtain due to its high melting point (2300°C) and the highly corrosive nature of liquid boron. It is usually done by reducing pure BCl.₃with hydrogen gas at high temperatures or by thermal decomposition of boron hydrides such as diborane (B₂H₆):

\[\mathrm{BCl_3(g)}+\frac{3}{2}\mathrm{H_2(g)}\rightarrow\mathrm{B(s)}+\mathrm{3HCl(g)} \label{Gl .3 }\]

\[B_2H_{6(g)} \rightarrow 2B_{(s)} + 3H_{2(g)} \label{Eq4}\]

The reaction shown in the equation \(\ref{Eq3}\) is used to produce boron fibers, which are stiff and light. As such, they are used as structural reinforcing materials in objects as diverse as the American space shuttle and lightweight bicycle frames used in races such as the Tour de France. Boron is also an important component of many heat-resistant borosilicate glass and ceramics, such as B. Pyrex, which is used for ovenware and laboratory glassware.

Unlike boron, the presence of aluminum minerals such as bauxite, a hydrated form of Al₂o₃, they are abundant. With an electrical conductivity approximately twice that of copper by weight, aluminum is used in more than 90% of overhead power lines in the United States. However, because aluminum and oxygen compounds are stable, extracting metallic aluminum from bauxite is an expensive process. Aluminum is extracted from oxide minerals by treatment with a strong base, which generates the soluble hydroxide complex [Al(OH)].₄]⁻. Neutralization of the resulting solution with CO gas₂leads to precipitation of Al(OH)₃:

\[2[Al(OH)_4]^−_{(aq)} + CO_{2(g)} \rightarrow 2Al(OH)_{3(s)} + CO^{2−}_{3( aq)} + H_2O_{(l)} \label{Eq5}\]

Thermal Dehydration of Al(OH)₃produces aluminum₂o₃, and metallic aluminum is obtained by the electrolytic reduction of Al₂o₃Use ofProceso de Hall-Heroult. Of the Group 13 items, only aluminum is widely used: every Boeing 777 aircraft, for example, consists of about 50% aluminum.

Abbildung \(\PageIndex{2}\): Quelle: Thomas D. Kelly und Grecia R. Matos, „Historical Statistics of Mineral and Material Raw Materials in the United States“, US Geological Survey Data Series 140 , 2010, abgerufen am 20. Jul 2011, pubs.usgs.gov/ds/2005/140/.

The other members of Group 13 are quite rare: gallium is about 5,000 times rarer than aluminium, and indium and thallium are even rarer. As a result, these metals are mostly byproducts of processing other metals. However, gallium's extremely low melting point (29.6 °C) makes it easy to separate from aluminum. Due to its low melting point and high boiling point, gallium is used as a liquid in thermometers, which have a temperature range of nearly 2200°C. Indium and thallium, the heavier Group 13 elements, are found as trace impurities in lead and zinc sulfide ores. Indium is used as a brittle seal for high-vacuum cryogenic equipment, and its alloys are used as low-melting solders on electronic circuit boards. Thallium, on the other hand, is so toxic that the metal and its compounds are rarely used. Both indium and thallium oxides are released in combustion dust when sulfide minerals are converted to metal oxides and SO.₂. Until relatively recently, these and other toxic elements were allowed to become airborne, creating large “dead zones” upwind of a melt. The dust from the flies is now trapped and serves as a relatively rich source of elements like In and Tl (as well as Ge, Cd, Te, and As).

The table \(\PageIndex{1}\) summarizes some important properties of the group 13 elements. Note the vast differences between boron and aluminum in size, ionization energy, electronegativity, and standard reduction potential, according to with the observation that boron behaves chemically like a nonmetal and aluminum like a metal. All elements in group 13 have ns²public notary¹valence electron configurations, and all tend to lose their three valence electrons to form compounds in the +3 oxidation state. The heavier elements in the group can also form compounds in the +1 oxidation state, as a result of the formal loss of the single np valence electron. Because Group 13 elements generally contain only six valence electrons in their neutral compounds, these compounds are all moderately strong Lewis acids.

Neutral compounds of Group 13 elements are electron deficient, so they are generally moderately strong Lewis acids.

See Also

Group 13 Elements: The Boro Family - GeeksforGeeksBoron family: Group 13 elements, properties and characteristicsElements of Group 13: The Boro Family (Properties, Notes)

In contrast to Groups 1 and 2, Group 13 elements do not show consistent trends in ionization energies, electron affinities, and reduction potentials, while electronegativities actually increase from aluminum to thallium. Some of these anomalies, especially for the Ga, In, Tl series, can be explained by the increase in the effective nuclear charge (Z_effect), which results from the poor shielding of the nuclear charge by the filled (n − 1)d¹⁰y (n − 2)f¹⁴lower shells. Consequently, although the actual nuclear charge increases by 32 as we go from indium to thallium, the shielding of the filled 5d and 4f subshells is so poor that Z_effectincreases significantly from indium to thallium. Thus the first ionization energy of thallium is actually greater than that of indium.

The anomalies in the periodic trends between Ga, In and Tl can be explained by the increase in the effective nuclear charge due to poor shielding.

Boron reactions and compounds

Elemental boron is a semimetal that is remarkably unreactive; in contrast, the other elements of Group 13 exhibit metallic properties and reactivity. Therefore, we consider the reactions and compounds of boron separately from those of other elements in the group. All Group 13 elements have fewer valence electrons than valence orbitals, generally resulting in delocalized metallic bonding. However, with its high ionization energy, low electron affinity, low electronegativity, and small size, boron does not form a metal lattice with delocalized valence electrons. Instead, boron forms unique, intricate structures containing multicentric bonds in which a pair of electrons hold three or more atoms together.

Elemental boron forms multicentric bonds, while the other Group 13 elements have metallic bonds.

The basic component of elemental boron is not the single boron atom, as would be the case for a metal, but B₁₂icosahedron. Since these icosahedrons do not pack very well, the structure of solid boron contains voids, which leads to its low density (Figure \(\PageIndex{3}\)). Elemental boron can be made to react with many non-metals to give binary compounds that have a variety of uses. For example, boron carbide plates (B₄C) It can stop a 30 caliber armor-piercing bullet, but weighs 10-30% less than normal armor. Other important compounds of boron with nonmetals are boron nitride (BN), which is formed by heating boron with excess nitrogen (Equation \(\ref{Eq22.6}\)); boron oxide (B₂o₃), which is formed when boron is heated with excess oxygen (equation \(\ref{Eq.22.7}\)); and boron trihalides (BX₃), which are formed by heating boron with excess halogen (equation \(\ref{Eq.22.8}\)).

\[\mathrm{2B(s)}+\mathrm{N_2(g)}\xrightarrow{\Delta}\mathrm{2BN(s)}\label{Gl22.6}\]

\[\mathrm{4B(s)} + \mathrm{3O_2(g)}\xrightarrow{\Delta}\mathrm{2B_2O_3(s)}\label{Gl22.7}\]

\[\mathrm{2B(s)} +\mathrm{3X_2(g)}\xrightarrow{\Delta}\mathrm{2BX_3(g)}\label{Gl22.8}\]

As is typical of elements that fall close to the dividing line between metals and nonmetals, many boron compounds are amphoteric, soluble in acids or bases.

Boron nitride is similar to elemental carbon in many ways. With eight electrons, the B-N unit is isoelectronic with the C-C unit, and B and N have the same size and average electronegativity as C. The most stable form of BN resembles graphite and contains six-membered B.₃norte₃Rings arranged in layers. At high temperature and pressure, hexagonal BN transforms into a cubic, diamond-like structure, one of the hardest substances known. boron oxide (B₂o₃) contains layers of trigonal plane BO₃Groups (similar to BX₃) in which oxygen atoms bond two boron atoms. Dissolves many metallic and non-metallic oxides, including SiO.₂to obtain a wide range of commercially important borosilicate glasses. A small amount of CoO gives the characteristic deep blue color of "cobalt blue" glass.

At high temperatures, boron also reacts with virtually all metals to form metallic borides, which contain regular three-dimensional lattices or clusters of boron atoms. The structures of two metal borides: ScB₁₂and cabin₆—are shown in the figure \(\PageIndex{4}\). Because metal-rich borides like ZrB₂and TiB₂Hard and resistant to corrosion even at high temperatures, they are used in applications such as turbine blades and rocket nozzles.

Boron hydrides were only discovered in the early 20th century when the German chemist Alfred Stock undertook a systematic study of binary boron and hydrogen compounds, although binary hydrides of carbon, nitrogen, oxygen, and fluorine have been known since the 18th century. Between 1912 and 1936, Stock oversaw the preparation of a series of boron and hydrogen compounds with unprecedented structures that could not be explained by simple bond theories. All of these compounds contain multicentric bonds. The simplest example is diborane (B₂H₆), which contains two bridging hydrogen atoms (part (a) in Figure \(\PageIndex{5}\). An extraordinary variety of boro-hydrogen polyhedral groups is now known; one example is B₁₂H₁₂²⁻Ion having a polyhedral structure similar to icosahedron B₁₂A unit of elemental boron where each boron atom is bonded to a single hydrogen atom.

A related class of polyhedral groups, the carboranes, contain CH and BH units; here is an example. Replacing the hydrogen atoms attached to carbon with organic groups results in substances with novel properties, some of which are currently being investigated for use as liquid crystals and in cancer chemotherapy.

The enthalpy of combustion of diborane (B₂H₆) is −2165 kJ/mol, one of the highest known values:

\[B_2H_{6(g)} + 3O_{2(g)} \rightarrow B_2O_{3(s)} + 3H_2O(l)\;\;\; ΔH_{Kamm} = −2165\; kJ/mol \label{Gl. 22.9}\]

Consequently, in the 1950s and 1960s, the US military explored the use of borohydrides as rocket propellants. These efforts were ultimately abandoned because borohydrides are unstable, expensive, and toxic, and most importantly, B₂o₃it proved to be very abrasive to rocket nozzles. However, the reactions performed during this investigation indicated that borohydrides exhibit unusual reactivity.

Because boron and hydrogen have nearly identical electronegativities, borohydride reactions are governed by subtle differences in the electron density distribution in a given compound. In general, two distinct types of reactions are observed: electron-rich species such as BH₄⁻ion are reducing agents, while electron-deficient species such as B₂H₆act as an oxidizing agent.

Reactions and compounds of the heavier Group 13 elements

The four heaviest Group 13 elements (Al, Ga, In, and Tl) react rapidly with the halogens to form compounds with a 1:3 stoichiometry:

\[ 2M_{(s)} + 3X_{2(s,l,g)} \rightarrow 2MX_{3(s)} \text{ oder } M_2X_6 \label{Eq10}\]

An exception is the reaction of Tl with iodine: although the product has the stoichiometry TlI₃, is not thallium(III) iodide, but a thallium(I) compound, Tl⁺Sal de triioduro (I₃⁻). This compound arises because iodine is not a strong enough oxidizing agent to oxidize thallium to the +3 oxidation state.

Of the halides, only fluorides show typical behavior of ionic compounds: they have high melting points (>950°C) and low solubility in nonpolar solvents. On the other hand, aluminum, gallium and indium trichorides, tribromides and triiodides as well as TlCl₃and TlBr₃, have a more covalent character and form halogen-bridged dimers (part (b) in Figure \(\PageIndex{4}\)). Although the structure of these dimers is that of diborane (B₂H₆) the bond can be described as an electron pair bond rather than a delocalized electron deficiency bond in diborane. Bridging halides are poor electron pair donors, therefore Group 13 trihalides are strong Lewis acids that readily react with Lewis bases such as amines to form a Lewis acid-base adduct:

\[Al_2Cl_{6(soln)} + 2(CH_3)_3N_{(soln)} \rightarrow 2(CH_3)_3N:AlCl_{3(soln)} \label{Eq11}\]

In water, Group 13 metal halides hydrolyze to metal hydroxide (\[M(OH)_3\)):

\[MX_{3(s)} + 3H_2O_{(l)} \rightarrow M(OH)_{3(s)} + 3HX_{(aq)} \label{Eq12}\]

In a related reaction, Al₂(SO₄)₃used to clarify drinking water by precipitating hydrated Al(OH)₃, which captures particles. The heavier metal halides (In and Tl) are less reactive with water due to their lower charge-radius ratio. Instead of forming hydroxides, they dissolve and form the ions of the hydrated metal complex: [M(H₂Ö)₆]³⁺.

Of the Group 13 halides, only the fluorides behave as typical ionic compounds.

Like boron (Equation \(\ref{Eq22.7}\)), all of the heavier Group 13 elements react with excess oxygen at elevated temperatures to form the trivalent oxide (M₂o₃), although Tl₂o₃is unstable:

\[\mathrm{4M(s)}+\mathrm{3O_2(g)}\xrightarrow{\Delta }\mathrm{2M_2O_3(s)} \label{Eq13}\]

Aluminum oxide (Al₂o₃), also known as aluminum oxide, is a hard, chemically inert, high-melting insulator used as a ceramic and as an abrasive in sandpaper and toothpaste. Replacement of a small amount of Al³⁺Ions in crystalline alumina with Cr³⁺Ions form the Ruby gem while Al replaces himself³⁺with a mixture of faith²⁺, fe³⁺, and you⁴⁺produces blue sapphires. The gallium oxide compound MgGa₂o₄it emits the bright green light familiar to anyone who has ever operated a xerographic copier. All oxides dissolve in dilute acid, but Al₂o₃and Ga₂o₃they are amphoteric, consistent with their position along the diagonal line of the periodic table, and they also dissolve in a concentrated aqueous base to form solutions containing M(OH).₄⁻ions

Group 13 trihalides are strong Lewis acids that react with Lewis bases to form a Lewis acid-base adduct.

Aluminum, gallium, and indium also react with the other Group 16 elements (chalcogens) to form chalcogenides with M stoichiometry.₂Y₃. However, since Tl(III) is too strong an oxidant to form a stable compound with electron-rich anions like S²⁻, Se²⁻, and you²⁻, Thallium forms only the thallium(I) chalcogenides with the stoichiometry Tl₂Y. Only aluminum, like boron, reacts directly with N.₂(at very high temperatures) to AlN, used in transistors and microwave ovens as a non-toxic heat sink due to its thermal stability; GaN and InN can be made using other methods. All metals, again with the exception of Tl, also react with the heavier Group 15 (pnicogenic) elements to form so-called III-V compounds like GaAs. These are semiconductors whose electronic properties, such as their bandgap, differ from those achievable with pure or doped Group 14 elements. For example, gallium arsenide doped with nitrogen and phosphorus (GaAs_1-x-yPAG_Xnorte_j) is used in the displays of calculators and digital clocks.

All group 13 oxides dissolve in dilute acid, but Al₂o₃and Ga₂o₃they are amphoteric

Unlike boron, the heavier Group 13 elements do not react directly with hydrogen. Only aluminum and gallium hydrides are known, but they must be prepared indirectly; AlH₃is an insoluble polymeric solid that decomposes rapidly in water, while GaH₃it is unstable at room temperature.

Complexes of group 13 elements

Boron has a relatively limited tendency to form complexes, but aluminum, gallium, indium, and to some extent thallium form many complexes. Some of the simplest are the hydrated metal ions [M(H₂Ö)₆³⁺], which are relatively strong Brønsted-Lowry acids that can donate a proton to form M(H₂Ö)₅(OH)²⁺Ion:

\[[M(H_2O)_6]^{3+}_{(aq)} \rightarrow M(H_2O)_5(OH)^{2+}_{(aq)} + H^+_{(aq)} } \label{Eq14}\]

Group 13 metal ions also form stable complexes with species containing two or more negatively charged groups, such as the oxalate ion. The stability of such complexes increases as the number of coordination groups provided by the ligand increases.

Summary

Compounds of Group 13 elements with oxygen are thermodynamically stable. Many of the anomalous properties of the Group 13 elements can be explained by the increase in Z_effectgo down in the group. Isolation of group 13 elements requires a high expenditure of energy since compounds of group 13 elements with oxygen are thermodynamically stable. Boron behaves chemically like a nonmetal, while its heavier congeners display metallic behavior. Many of the observed inconsistencies in the properties of Group 13 elements can be explained by the increase in Z_effectthis is due to poor shielding of the nuclear charge by the full (n − 1)d¹⁰y (n − 2)f¹⁴lower shells. Rather than forming a metallic lattice with delocalized valence electrons, boron forms unique aggregates containing multicentric bonds, including metallic borides, in which boron bonds with other boron atoms to form three-dimensional lattices, or clusters with regular geometric structures. . All neutral compounds of the Group 13 elements are electron deficient and behave as Lewis acids. Trivalent halides of the heavier elements form halogen-bridged dimers that contain electron-pair bonding rather than the delocalized electron-deficient bonding characteristic of diborane. Its oxides dissolve in dilute acid, although aluminum and gallium oxides are amphoteric. None of the elements in Group 13 react directly with hydrogen, and the stability of hydrides produced by other routes decreases as we move up the group. Unlike boron, the heavier Group 13 elements form a large number of complexes in the +3 oxidation state.