21.4: Group 13: The Bor family (2023)

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    learning goals
    • Understand trends in the properties and reactivity of Group 13 elements.

    Group 13 is the first group to straddle the line between metals and nonmetals, so its chemistry is more diverse than Groups 1 and 2, which only include metallic elements. With the exception of the lightest element (boron), the Group 13 elements are all relatively electropositive; That is, they tend to donate electrons during chemical reactions rather than gain them. Although Group 13 contains aluminium, the most abundant metal on earth, none of these elements were known until the early 19th century because they are never found in a free state in nature. Elemental boron and aluminum first prepared by reduction of B2o3y AlCl3with potassium, it could not be prepared until potassium was isolated and shown to be a strong reducing agent. Indium (In) and thallium (Tl) were discovered using spectroscopic techniques in the 1860s, long before methods for their isolation became available. Indium, named for its indigo (deep purple-blue) emission line, was first observed in the spectrum of zinc minerals, while thallium (from the Greek thallos, meaning "a green plant shoot and young") was named for its bright green emission line. . Gallium (Ga; Mendeleev's eka-aluminium) was discovered in 1875 by French chemist Paul Émile Lecoq de Boisbaudran in a systematic search for Mendeleev's "missing" element in Group 13.

    Group 13 elements never occur naturally in the free state.

    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.

    21.4: Group 13: The Bor family (1)

    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:

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


    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.

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    21.4: Group 13: The Bor family (2)

    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.

    Table \(\PageIndex{1}\): Selected properties of group 13 elements
    Property bor Aluminum* Gallium indio waist
    *This is the name used in the United States; the rest of the world puts an extra i on it and calls it aluminum.
    The configuration shown does not contain filled d and f subshells.
    The values ​​given apply to six-coordinate ions in the most common oxidation state, with the exception of Al.3+, for which the value of the four-coordinate ion is given. The b3+Ion is not a known species; the given radius is an estimate of four coordinates.
    §X is Cl, Br or I. Reaction with F2gives trifluorides (MF3) for all elements of group 13.
    atomic symbol B Alabama Georgia In Tl
    atomic number 5 13 31 49 81
    Atomic mass (one) 10.81 26.98 69.72 114.82 204.38
    valence electron configuration 2s22p1 3 seconds23p1 4s24p1 5 seconds217 o'clock1 6s218:001
    Melting Point/Boiling Point (°C) 2075/4000 660/2519 29.7/2204 156.6/2072 304/1473
    Density (g/cm3) a 25 °C 2.34 2.70 5.91 7.31 11.8
    Atomic radius (pm) 87 118 136 156 156
    first ionization energy (kJ/mol) 801 578 579 558 589
    most common oxidation state +3 +3 +3 +3 +1
    Ionenradio (pm) −25 54 62 80 162
    Electron affinity (kJ/mol) −27 −42 −40 −39 −37
    electronegativity 2.0 1.6 1.8 1.8 1.8
    Standard reduction potential (E°, V) −0,87 −1,66 −0,55 −0,34 +0.741 of M3+(ac)
    Reaction product with O2 B2o3 Alabama2o3 Georgia2o3 In2o3 Tl2o
    type of rust Sauer amphoteric amphoteric amphoteric Basic
    Reaction product with N2 BN AlN GaN Posada none
    Reaction product with X2§ BX3 Alabama2X6 Georgia2X6 In2X6 TLX

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

    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 (Zeffect), 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 Zeffectincreases 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.

    21.4: Group 13: The Bor family (3)

    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 \(\ref{Eq.22.8}\)).


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

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

    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.

    21.4: Group 13: The Bor family (4)

    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.

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    21.4: Group 13: The Bor family (5)

    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.

    21.4: Group 13: The Bor family (6)

    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.

    21.4: Group 13: The Bor family (7)

    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 BH4ion are reducing agents, while electron-deficient species such as B2H6act as an oxidizing agent.

    Example \(\PageIndex{1}\)

    For each reaction, explain why the given products are formed.

    1. B2H6(g) + 3O2(g) → B2o3(s) + 3H2O(l)
    2. BC3(l) + 3H2O(l) → B(OH)3(aqueous) + 3HCl (aqueous)
    3. \(\mathrm{2BI_3(s)}+\mathrm{3H_2(g)}\xrightarrow{\Delta}\frac{1}{6}\mathrm{B_{12}(s)}+\mathrm{6HI( GRAMO)}\)

    Given:balanced chemical equations

    Asked:why the specified products are formed


    Classify the type of reaction. Explain why reaction products are formed using periodic trends in atomic properties, thermodynamics, and kinetics.


    1. Molecular oxygen is an oxidizing agent. If the other reactant is a potential reducing agent, we expect a redox reaction to occur. Although B.2H6contains boron in its highest oxidation state (+3), it also contains hydrogen in the -1 oxidation state (the hydride ion). Since the hydride is a strong reducing agent, a redox reaction is likely to occur. We hope that Hwith oxide and H+y-oh2se reduce a O2−, but what are the actual products? A reasonable guess is B2o3yh2Oh, both connections stable.
    2. Neither BCl3water is also not a strong oxidizing or reducing agent, so a redox reaction is unlikely to occur; a hydrolysis reaction is more likely. Nonmetallic halides are acids and react with water to form a solution of the hydrohalic acid and a nonmetallic oxide or hydroxide. In this case, the most likely boron-containing product is boric acid [B(OH)3].
    3. We normally expect a boron trihalide to behave like a Lewis acid. In this case, however, the other reactant is elemental hydrogen, which usually acts as a reducing agent. Iodine atoms in BI3they are in the lowest accessible oxidation state (-1) and boron is in the +3 oxidation state. Consequently, we can write a redox reaction in which hydrogen is oxidized and boron is reduced. Since boron compounds are rare in lower oxidation states, we expect boron to be reduced to elemental boron. Therefore, the other product of the reaction must be HI.
    Exercise \(\PageIndex{1}\)

    Predict the products of the reactions and write a balanced chemical equation for each reaction.

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    1. \(\mathrm{B_2H_6(g)}+\mathrm{H_2O(l)}\xrightarrow{\Delta }\)
    2. \(\mathrm{BBr_3(l)}+\mathrm{O_2(g)}\flecha derecha\)
    3. \(\math{B_2O_3(s)}+\math{Ca(s)}\xrightarrow{\Delta}\)


    1. \(\math{B_2H_6(g)}+\math{H_2O(l)}\xrightarrow{\Delta}\math{2B(OH)_3(s)}+\math{6H_2(g)}\)
    2. \(\math{BBr_3(l)}+\math{O_2(g)}\rightarrow\mathrm{sin reaction}\)
    3. \(\mathrm{6B_2O_3(s)}+18\mathrm{Ca(s)}\xrightarrow{\Delta}\mathrm{B_{12}(s)}+\mathrm{18CaO(s)}\)

    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).4ions

    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 (GaAs1-x-yPAGXnortej) 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.

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    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:

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

    21.4: Group 13: The Bor family (8)

    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.

    Example \(\PageIndex{2}\)

    For each reaction, explain why the given products are formed.

    1. \(\mathrm{2Al(s)} + \mathrm{Fe_2O_3(s)}\xrightarrow{\Delta }\mathrm{2Fe(l)} + \mathrm{Al_2O_3(s)}\)
    2. \(\mathrm{2Ga(s)} + \mathrm{6H_2O(l)}+ \mathrm{2OH^-(aq)}\xrightarrow{\Delta}\mathrm{3H_2(g)} + \mathrm{2Ga( OH)^-_4(aq)}\)
    3. \(\mathrm{In_2Cl_6(s)}\xrightarrow{\mathrm{H_2O(l)}}\mathrm{2In^{3+}(aq)}+\mathrm{6Cl^-(aq)}\)

    Given:balanced chemical equations

    Asked:why the specified products are formed


    Classify the type of reaction. Explain why reaction products are formed using periodic trends in atomic properties, thermodynamics, and kinetics.


    1. Aluminum is an active metal and a strong reducing agent, and Fe2o3contains Fe(III), a potential oxidizing agent. Therefore, a redox reaction is likely to occur producing metallic Fe and Al.2o3. Since Al is a main group element above the transition element Fe, it should be a more active metal than Fe. Therefore, the reaction should proceed to the right. In fact, this is the reaction of the thermite, which is so violent that it produces molten Fe and can be used for soldering.
    2. Gallium sits just below aluminum on the periodic table and is amphoteric, so it dissolves in acid or base, producing hydrogen gas. Since gallium is similar to aluminum in many of its properties, we predict that gallium will dissolve in the strong base.
    3. The metallic character of the group 13 elements increases with increasing atomic number. Therefore, indium trichloride should behave like a typical metal halide, dissolving in water to form the hydrated cation.
    Exercise \(\PageIndex{2}\)

    Predict the products of the reactions and write a balanced chemical equation for each reaction.

    1. LiH(s) + Al2Kl6(sol)→
    2. Alabama2o3(s) + OH(ac)→
    3. Al(s) + N2(g) \(\xrightarrow{\Delta}\)
    4. Georgia2Kl6(soluble) + Cl(sol)→


    1. 8LiH(s) + Al2Kl6(solution)→2LiAlH4(solution) + 6LiCl(s)
    2. Alabama2o3(s) + 2OH(aqueous) + 3H2O(l) → 2Al(OH)4(ac)
    3. 2Al(s) + N2(g) \(\xrightarrow{\Delta}\) 2AlN(s)
    4. Georgia2Kl6(soluble) + 2Cl(solution) → 2GaCl4(soluble)


    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 Zeffectgo 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 Zeffectthis 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.

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