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13.1: Introduction to Group 13 Elements

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    34120
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    Introduction

    • Sources:
      • Ulexite: NaCa[B5O6(OH)6].5H2O
      • Borax: Na2[B4O5(OH)4].8H2O
      • Colmanite: Ca2[B3O4(OH)3].2H2O
      • Kenite: Na2[B4O5(OH)4].2H2O
      See the figure for the borate anions in te above.
    • There are no ionic B3+ compounds, c.f. Mg2+ and Li+.
    • The compounds which are coordinatively unsaturated (e.g.BCl3) are very strong Lewis bases.
    • Tetrahedral adducts and anions are common, for example:
      F3B:O(C2H5)2 BF4- B(C6H5)4-
    • The boron hydrides (boranes) are a class unto themselves together with the carboranes and all the anions. They frequently feature closed and open polyhedra based on fragments of an icosahedron, but not always.
    • Boron resembles silicon more than aluminum in some ways:
      • B2O3 and B(OH)3 are acidic rather like SiO2 and Si(OH)4 whereas the Al compounds are weakly amphoteric.
      • The borates have some features in common with the silicates.
      • The halide compounds of B and Si are readily hydrolysed (except BF3) whereas the halide compounds of Al are only partly hydrolysed.
      • The B and silicon hydrides are volatile molecular compounds, which inflame in air, while AlH3 is an involatile solid.

    Manufacture and Properties of Boron

    Boron is quite difficult to isolate, because it is refractory and reactive at high temperature and so it is difficult to contain:

    1. B2O3 + 3Mg alt 2B + MgO (98% - Wash with NaOH, HCl and HF)
    2. 2BCl3 + 3Zn alt 3ZnCl2 + 2B (900 oC)
    3. 2BX3 + 3H2 alt 6HX + 2B (Tantalum catalyst)

    Boron is rather inert in most forms which contain icosahedral cages. It is attacked by hot oxidizing acids.

    Amorphous boron is more reactive, if white hot, and is attacked by NH3 to form boron nitride which is isomorphous with graphite.

    Oxygen Compounds of Boron

    See Figure 12-1.

    The "anhydrous" borates involve the ions: BO33-, B3O63-, (BO2)nn-, and larger aggregates.

    The hydrated ones feature: BO3 units which are planar and BO4 unitsd which are tetrahedral and formally have a -ve charge on the boron. The charge on the ion is equal to th number of these latter units. The structures without BO4 units hydrate readily.

    Boric Acid

    See Figure 12-2 for a summary.

    • Note that B2O3 + SiO2 is "pyrex".
    • Note that B(OH)3, normally a weak acid, can be considerably strengthened to the point where it can be titrated with NaOH by "chelation" by an organic compound with neighbouring OH groups e.g. glycerol: alt
    • In general, equilibria involve the processes:
      B-O-H + H-O-B alt B-O-B + H2O
      B-O-H + O-H- alt B-O- + H2O

    Halides of Boron

    Trihalides

    BF3 is the most important and is used on an industrial scale. It is a gas boiling at -101 oC.

    It is a very strong Lewis acid:

    BF3 + :F- alt BF4-

    BF3 + :L alt F3B:L

    Unlike the others, BF3 is only partly hydrolysed:

    4BF3 + 6H2O alt 3H3O+ + 3BF4- + B(OH)3
    BF4- + H2O alt BF3(OH) + HF
    (BCl3 + 3H2O alt B(OH)3 + 3HCl)

    In synthetic organic chemistry it is used as follows:

    The conversion of ethers or alcohols with acids to esters e.g:

    H+ + RCOOH alt [RCOOH2]+
    [RCOOH2]+ + BF3 alt [RCO}+ + F3B:OH2
    [RCO]+ + R'OH alt RCOOR' + H+

    Friedel-Crafts alkylations and acylations:

    RX + BF3 alt R+ + BF3X-
    R+ + PhH alt PhR + H+
    H+ + BF3X- alt BF3 + HX

    Fluoroboric acid: "HBF4" is sold as a 40% solution in water. It is a strong acid and source of BF4- ions useful for crystallizations where a coordinating anion is to be avoided.

    Reactions of the Trihalides of Boron.

    1. Adduct formation - the main thing to remember that the order of acid strength is counter to naive expectations because of p-bonding effects whnich are strongest for B-F bonds and lead to BF3 being least willing to go from sp2 to sp3. hybridized.
    2. Halide exchange reactions:
      BCl3 + BBr3 alt BCl2Br + BBr2Cl

      The exchange, presumably through a bridged intermediate, is very facile, so pure mixed compounds cannot be obtained.

    3. Elimination of halide - covers the various solvolyses in addition to hydrolysis:
      BCl3 + 3C2H5OH alt B(OC2H5)3 + 3HCl
      BCl3 + 3NH(C2H5)2 alt B(N(C2H5)2)3 + 3HCl

      This will happen with any solvent with an exchangeable H.

    Subhalides of Boron

    They have a B:X ratio less than 1:3

    • There are halides BF and BCl.
    • There are halides B2X4 known for F, Cl, Br and I. Rotation about the B-B bond is easy. Any multiiple bonding tendency involves the boron to halogen bonds as in BX3. (This is in contrast with the aminoboranes - see below.)
    • BnXn compounds have polyhedral cages of boron each carrying one X group. The largest range of compounds is known for Cl where n = 4, 8, 9, 10 and 11. B4Cl4 is tetrahedral, B8Cl8 is a triangular dodecahedron, and B9Cl9 is a rectangularly tricapped trigonal prism. Look them up!

    The Hydrides of Boron - the Boranes

    Table 12-1 lists the hydrides up to B10H14

    Figure 12-4 shows some of their structures as perspective drawings. Note that the lines are intended to clarify the shape of the molecule and do not necessarily represent 2e- - 2-centre bonds.

    Note also the nomenclature - the prefix gives the number of boron atoms and the number in parentheses the number of hydrogen atoms, e.g. pentaborane(9) is B5H9.

    Synthesis

    Diborane (b.p. -92.6 oC can be made by several methods:

    3NaBH4 + 4BF3 alt 2B2H6 + 3NaBF4
    2NaBH4 + I2 alt B2H6 + 2NaI + H2
    BF3 + 6NaH alt B2H6 + 6NaF

    The last is the main industrial method. The higher boranes are made by thermolysis of diborane under various conditions.

    Structures

    The connectivities in the boranes cannot be explained using 2e- - 2-centre bonding only, that is the molecules are electron deficient. Valence bond theory has been "extended" by designating three types of 2e- - 3-centre bonding in addition to a normal 2e- - 2-centre B-H and B-B bonds:

    alt

    An example of the use of this scheme for B10H14 is shown below:

    altEach hydrogen must have one bond ending at it and each boron must have a total of 4 bonds ending at it. If an atom is in the middle of a three-centre bond, the curved line passing through it counts as only one bond.

    Thus, for example, B6 has one normal bond to a terminal hydrogen and one normal bond to B2, plus it is at the end of two three-centre bonds through the bridging hydrogens for a total of four bonds i.e eight electrons.

    B2 has one normal bond from B6, one normal bond from its terminal hydrogen, is at the end of a "closed" three-centre bond from B1 and B2, and it is in the centre of an "open" three-centre bond from B5 to B7. This is also equivalent to four bonds.

    In some cases, more than one "resonance" (canonical) structure can be formulated to account for the observed molecular shape.

    Reactions of the boranes

    Diborane

    1. With oxygen (explosive):
      B2H6 + 3O2 alt B2O3 + 3H2O
    2. With water:
      B2H6 + 3H2O alt B(OH)3 + 6H2
      or alcohols:
      B2H6 + 3HOR alt B(OR)3 + 6H2
    3. Substitution reactions:
      B2H6 + HCl alt B2H5Cl + H2
      B2H6 + 6Cl2 alt 2BCl3 + 6HCl
    4. Cleavage with Lewis bases:

      Symmetrical:

      B2H6 + N(CH3)3 alt 2H3BN(CH3)3
      Unsymmetrical:
      B2H6 + 2NH3 alt [H2B(NH3)2]+[BH4]-
    5. Reduction:
      2B2H6 + 2Na alt NaBH4 + NaB3H8
      B2H6 + NaBH4 alt NaB3H8 + H2
      5B2H6 + 2NaBH4 alt Na2B12H12

    Pentaborane(9)

    This molecule illustrates two general trends: Attack by bases can remove the somewhat acidic bridging hydrogen:

    B5H9 + NaH alt NaB5H8 + H2

    The pyramidal B5H9 loses one of its four bridging hydrogens and the resulting ion is "fluxional", that is, the location of the missing bridge is not stationary, and all the atoms in the base of the pyramid (four borons, four terminal hydrogens and three bridging hydrogens) appear equivalent on the time scale of nmr experiments which might otherwise have distinguished them.

    Attack by electrophiles can lead to substitution at the apex of the pyramid:

    B5H9 + I2 alt B5H8I(apical) + HI

    Decaborane(14)

    Once again the bridging hydrogens can be removed by base:

    B10H14 + OH- alt B10H13- + H2O

    or converted to terminal hydrogens by reducing agents:

    B10H14 + 2Na alt Na2B10H14

    In this reaction, the product has two bridging hydrogens between B1 and B5 and B7 and B8.

    Other nucleophiles will add at B6 and B9 with loss of two bridging hydrogens. Again, the two that are left bridge between B1 and B5 and B7 and B8:

    B10H14 + 2CH3CN alt 6,9-(CH3CN)2B10H12 + H2

    Electrophiles substitute terminal hydrogens at the bottom of the "basket" in the 1 and 3 or 2 and 4 positions, e.g.:

    B10H14 + I2 alt 2,4-I2B10H12 + 2HI

    There are two reactions that lead to a closed cage:

    B10H14 + 2Et3NBH3 alt [Et3NH]+2[B12H12]2-

    and

    (SEt2)2B10H14 + HCCH alt B10C2H12 + H2 + 2SEt2

    Polyhedral Borane Anions and Carboranes

    Realize that two carbon atoms can replace two B-'s in a closo-borane anion BnHn2-. Derived molecule or molecule ions are the nido structures, which are missing one vertex relative to the closo structure and the more open arachno structures which are missing two vertices. Figure 12-12 shows structures from B4 to B12. see also Figure 12-8.

    Skip the chemistry of these species.

    The Tetrahydroborate Ion (BH4-)

    This is an important reducing agent, source of H-, and reagent to make other less ionic borohydrides.

    NaBH4 is stable in dry air and alkaline aqueous solution. (It will react with water initially but the reaction stops as the concentration of the hydrolysis product, sodium borate, builds up.)

    LiBH4 is similar to NaBH4 but more sensitive to water.

    Al(BH4)3 is liquid which explodes with air or water. It probably has pairs of hydrogen bridges like diborane.

    Zr(BH4)4 is a molecular solid with three bridging hydrogens connecting each boron to the zirconium.

    Boron-Nitrogen Compounds

    The following are really equivalent representations, but the text uses the right-hand one to indicate a weaker bond:

    alt

    Amine Boranes

    This is the class of amine - BH3 Lewis adducts. They contain the unit shown above. The safest synthesis is:

    H3NRCl + LiBH4 alt H2RN:BH3 + LiCl + H2

    Aminoboranes

    These have the structure:

    alt

    The molecules are flat, and rotation about the B—N bond is restricted, therefore the left-hand structure must be is a significant contributor. It is perhaps not correct to represent them as canonical structures since the geometries would be so different. The cleanest synthetic route is, for example:

    (CH3)2NH + BCl3 alt (CH3)2HN:BCl3
    (CH3)2HN:BCl3 alt (CH3)2N:BCl2 + HCl (on heating)
    (CH3)2N:BCl3 + 2RMgBr alt (CH3)2N:BR2 + 2MgClBr

    Borazine

    This six membered ring can be synthesised by several routes:

    alt alt

    Unlike benzene, borazine undergoes addition reactions:

    alt

    Notice where the Hd+ and the Cld- end up: This illustrates how unrealistic the formal charges on the boron and nitrogen atoms really are!

    Like benzene, borazine can form p-complexes with transition metals:

    Aluminium, Gallium, Indium and Thallium

    Aluminium is the most common metallic element in the crust of the earth, but the common minerals, for example, felspars and micas are rather difficult to process. Aluminum is obtained from bauxite, (Al2O3.nH2O) and cryolite, Na3AlF6 by electrolysis. Gallium and Indium occur in traces in bauxite and all three are found in certain sulphite ores of other metals. While aluminium is obviously the most important, gallium is used in gallium arsenide semiconductors. The elements are all much more metallic than boron, but there are a number of borderline covalent compounds. All are trivalent, but for thallium the univalent state (Tl+) becomes the dominant state as covalent bond strenghts diminish down the group. Some thallium III compounds are thermodynamically unstable:

    \[\ce{TlX3 -> TlX + X2}\]

    The MX3 compounds (halides and organometallic) are Lewis bases like boron. The strengths vary in the sequence:\[\ce{B > Al > Ga > (In ~ Tl)}\]

    The trihalides are not monomers like BX3 but are more or less associated e.g. Al2Cl6 and (AlF3)n.

    All give aqua ions [M(H2O)6]3+ in their salts obtained from aqueous solution.

    Occurence, Extraction and Properities of the Elements

    To obtain aluminum, bauxite is dissolved in sodium hydroxide to give sodium aluminate, NaAl(OH)4. The insolubles which include hydrated iron oxide are filtered off, and the pH adjusted (with CO2)to reprecipitate the aluminum as Al(OH)3.3H2O which is then dehydrated to Al2O3 and dissolved in molten cryolite, Na3AlF6 for electrolysis. The other (less reactive) metals can be obtained by electrolysis of aqueous solutions.

    • The metals are all soft and quite reactive.
    • Aluminum is "passivated" by a film of oxide which prevents it reacting with oxygen, water and even dilute nitric acid. (If the surface is amalgamated, reaction with water can occur.)
    • The metals dissolve in non-oxidizing dilute acids. Aluminum and gallium are amphoteric and will dissolve in sodium hydroxide.
    • They react with the halogens and sulphur.
    • Thallium reacts slowly because the Tl+ salts which are formed are often insoluble and coat and passivate the metal surface.

    The Oxides

    • The most important are g-alumina, used as a stationay phase in liquid-solid chromatography and a-alumina, used as a catalyst in petroleum cracking.
    • The gemstone, ruby is Al2O3 contaminated with traces of Cr3+ in place of the Al3+ and sapphire has Fe2+, Fe3+ and Ti4+ replacing some of the Al3+.

    The Halides

    • All the M(III) trihalides are known except the triiodide of thallium. The compound, TlI3 is actually Tl+[I3]-.
    • The coodination numbers are 4 to 6 depending on the relative metal and halogen sizes.
    • The 4-coordinate compounds are molecular compounds and have lower melting-points.
    • They are strong Lewis bases and some can be used as Friedel-Crafts reaction catalysts:
      RCOCl + "AlCl3" alt RCO+ + AlCl4-

      The carbocation goes on to attacks the other organic reagent electrophilicly.

    The Aqua Ions, Oxo salts and Aqueous Chemistry

    The aqua ions all undergo hydrolysis:

    \[\ce{[M(H2O)6]^{3+} -> [M(H2O)5(OH)]^{2+} + H^{+}(aq)}\]

    Element Ka
    Al 1.12x10-5
    Ga 2.5x10-3
    In 2x10-4
    Tl ~7x10-2

    Salts of weak acids cannot exist in solution because the anions would be protonated and the hydroxides would precipitate.

    The "hydroxides" of aluminum and gallium are amphoteric:

    M(OH)3(s) alt M3+ + 3OH-
    M(OH)3(s) alt MO2- + H+ + H2O

    Depending on the conditions, bridging hydroxide is also common:

    2[M(H2O)5(OH)]2+ alt [(H2O)5MOM(H2O)5]4+ + H2O (etc)

    Hydroxides

    The real hydroxides, by extension of the above, are complicated structures involving bridging OH-, terminal H2O and perhaps [M(OH)4]- for some metals.

    Alums

    These are the compounds for which aluminum was originally named. They are double salts of formula MM'(SO4)2.12H2O where M+ is usually an alkali metal ion (not Li+) and M'3+ is Al3+ or another trivalent ion. For example, plain "alum" or "potash alum" is the potassium/aluminum salt and "chrome alum" is the potassium/chromium(III) salt. These compounds are characterized by easily grown octahedral crystals. Each metal ion is 6-coordinated by water.

    Coordination Compounds

    Examples are: [Al(H2O)6]3+, [AlF6]3-, Cl3Al(N(CH3)3)2, [Al(ox)3]3- and Al(8-hydroxyquinolinate)3

    Hydrides

    The metal hydrides are not very stable except "AlH3" which is an air-sensitive polymeric material. The tetrahydroaluminate ion AlH4- is an important reducing agent and hydride source which usually comes as lithium aluminum hydride. The analogous gallium compound exists. The compounds are very sensitive to hydrolysis which is very exothermic and can be explosive.

    The is a series of MH3 Lewis adducts with donor molecules which are generally more stable to, for example, hydrolysis than the parent hydrides.

    Lower Valent Compounds

    This section is mainly about Tl+ which resembles K+ and Ag+ in its chemistry. This section was not covered in depth in lectures. Skip it.

    Summary of the Periodic Trends for the Elements of Group 13

    1. Boron
      1. Forms no simple B3+ cation.
      2. Forms covalent compounds almost exclusively, and polyatomic ions are internally covalently bonded.
      3. Has a maximum covalence of 4 corresponding to an octet.
      4. The trivalent compounds are usually strong Lewis acids.
      5. Its oxide and "hydroxide" are acidic.
      6. Forms many polyatomic borates.
      7. The trihalides are easily hydrolysed.
      8. Forms many hydrides and hydride anions which are polyhedral clusters: the boranes, carboranes and the borane anions. The simplest BH4- is a very important synthetic reagent
    2. Aluminum
      1. Readily forms the Al3+ ion which is usually coordinated.
      2. Much more metallic than boron and forms many ionic compounds.
      3. Forms molecular compounds and ionic lattices with cordination numbers from 4 up to 6 and higher.
      4. Forms oxides which are chemically and thermally fairly inert.
      5. Forms a mainly basic but quite amphoteric hydroxide.
      6. Forms partially hydrolysable halides.
      7. Forms a polymeric hydride and the AlH4- ion. The latter is important.
    3. Gallium, Indium, and Thallium.
      1. Readily form M3+ aquo species and have a rich coordination chemistry.
      2. Form increasingly stable M+ compounds especially thallium. Covalent bonds successively weaken down the group enhancing this trend.
      3. Halides are increasingly aggregated with the increasing size of the metals.
      4. Hydrides and hydride ions are not very important or stable.
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