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Hydrogen | Group 1 - Alkali Metals | Group 2 - Alkaline Earth Metals | Group 12
It was a long trek, but the hikers boron.
Ah well, whatever; let's boron till we strike gold...
Boron in the form of borax has been known since antiquity. The element itself, a black or brown powder second only to diamond in hardness, was isolated by Sir Humphry Davy in 1807. Boron is unique amongst the elements of group 13 in that it is a non-metal. The element has two naturally occurring isotopes, 10B and 11B, the ratios of which vary depending upon the mineral in which it is found and also on where it is mined but the proportion is 10B tends to be approximately 20%. The minerals in which boron is principally found are ulexite (NaCa[B5O6(OH)6]·5 H2O), borax (Na2 [B4O5(OH)4]·8 H2O), colemanite (Ca2 [B3O4(OH)3]2·2 H2O)and kernite (Na2 [B4O5(OH)4]·2 H2O). Large deposits of borax are found in the Mojave desert in California and there are other large mineral deposits in Turkey.
Preparation and Physical Properties
The minerals can be treated to form the oxide, B2O3, from which elemental boron can be extracted by reaction with magnesium
B2O3 + 3 Mg → 2 B + 3 MgO
2 BCl3 + 3 Zn → 3 ZnCl2 + 2 B
2 BX3 + 3 H2→ 6 HX + 2 B
There are a few different allotropes (physical forms of the pure element) which are either chestnut-coloured amorphous powders or black crystalline solids. All the allotropes consist of B12 units where the boron atoms form an icosohedron but differ their packing. They all have very high melting point (around 2300 °C) and have low density and electrical conductivity. Crystalline allotropes of boron are very hard, fairly inert and only react with hot solutions of oxidising agents. Amorphous boron is much more reactive however. For example, it reacts with ammonia to produce a white waxy compounds with the overall formula BN. This compounds has a graphite like structure with layers of sheets that are made up of interlocking hexagonal arrangements of alternating B and N atoms.
The chemistry of boron is extremely diverse and complex. The element has extremely high ionisation energies (first three ionisation energies 801, 2427 and 3659 kJ mol-1) which can not be offset by the energies given out due to enthalpies for lattice formation or hydration. Because of this there are no compounds containing discrete B3+ ions. The element has a similar electronegativity to those of carbon and hydrogen and so boron compounds are highly covalent. It is also said to be electron deficient having only three valence electrons. It forms coordinatively unsaturated trigonal planar compounds of the form BX3 which can become tetravalent and therefore saturated by acting as a Lewis acid, accepting electron from a donating molecule to form what is called an adduct. Boron also copes with it's electron deficiency by forming compounds containing cage or basket like polyhedral arrangements of boron atoms in which the available electrons are shared out.
Uses of Boron and its Compounds
Boron itself has limited uses, but amorphous boron is used in pyrotechnic flares to provide a highly visible green colour. Boron filament has properties of high-strength, light weight and compressibilty, making it suitable for advanced aerospace structures and sports goods.
Boron is sometimes included as a component of alloy steels to increase their hardenability.
Boron is also a key component in iron, neodymium, boron (FeNdB) magnets, which are low-cost, high-performance devices. Indeed, they are among the most powerful permanent magnets known, and are therefore used in 'maglev' trains. They also find application in automobiles which, nowadays contain about 30 magnets, and in electric motors in the home. (The average house contains about 70 electric motors).
Boron compounds, however, have many uses. Borax (di-Sodium tetraborate, Na2B4O7.10H20) has been known in Europe since the Middle Ages and probably even earlier in the east. Up until the 19th Century it was imported to Britain from Tibet under the name of 'tincal'. It occurs naturally in salt lakes, but is now extracted mainly from boron-containing minerals whose chief source is the California desert. It is used as a water softener in washing powders and as a flux in welding. Borosilicate glasses (Pyrex) are made by dissolving silica, SiO2, in molten boron oxide, B2O3. When heated, they expand far less than most other glasses and, as a result, can withstand heat changes.So, from 1892 onwards they were used for making laboratory glassware and thermometer tubes. Borosilicate glasses also make better lenses than other glasses and so are used for a variety of optical instruments including telescope lenses. Coloured glasses are made by dissolving transition metal oxides into the mix.
Boric acid (H3BO3) was found to have antiseptic properties by Joseph Lister in 1875 and is useful for cuts, abrasions, slight eye infections, nose and throat sprays and insect bites. Boric acid is also a highly effective fungicide and can therefore used to treat athletes foot and, in the form of pessaries, to control thrush, especially when other interventions have failed1. These two infections are caused by similar organisms known as Candida, and so are forms of candidiasis. Both boric acid and borax were formerly used as food preservatives (E284 and E285 respectively), boric acid being particularly used for margarine, butter and bacon, but this was forbidden in the United Kingdom in 1925. In North America, it is used as an insecticide for the control of cockroaches, silverfish, ants, fleas, and other insects.
Borax and boric acid are flame retardants and are used in cotton mattresses and chipboard.
Sodium perborate (NaBO3) is the bleaching agent used in most European powdered detergent formulations as it is less aggressive than sodium hypochlorite (bleach) and therefore less damaging to fabrics and their dyes.
Metal boride compounds, such as chromium, molybdenum and tungsten, and TiB2, are important materials finding applications as gas turbine blades and rocket motor nozzles due to their highly inert nature and ability to withstand high temperatures. Borides of iron, manganese and nickel are used as coatings to certain devices to increase their wear-resistance and corrosion-resistance properties. Magnesium diboride is a superconductor with a comparatively high transition temperature of 39K.
Because of the uniquely high capacity of the 10B isotope to readily absorb neutrons, metal borides and boron carbides find uses in the nuclear industry as nuclear shielding materials and in reactor control rods. Boron hydrides are sometimes used as high energy fuels, in liquid form for jet aircraft and in solid form as rocket fuels. Boron nitride is used in the thermal tiles on the Space Shuttle. A form of boron nitride has lubricating properties similar to graphite and, in a special form called borazon (see below), to replace industrial diamonds, being harder than diamond and having a much higher melting point.
Biological Role of Boron
Boron is an essential plant micronutrient, and plants grown in an adequate supply of boron are better able to resist disease-causing organisms. The main functions of boron relate to cell wall strength and development, cell division, fruit and seed development, sugar transport, and hormone development.
Chemistry of Boron
Oxygen Containing Compounds of Boron
There are many examples of oxygen containing boron compounds and most of the minerals in which boron is found are borates. These often consist of interconnected trigonal planar BO3 and tetrahedral BO4 units that form rings and chains of alternating B and O atoms, some examples are shown in figure 1. These arrays are often anionic and many of the minerals also incorporate water molecules into their structure as they crystallise.
Figure 1: Some examples of borate anions present in boron containing minerals.
An important oxygen compound of boron is boric acid, B(OH)3. It is acidic when dissolved in water as it forms B(OH)4- and liberates a proton.
B(OH)3 + H2O → B(OH)4- + H+
The anion B(OH)4- occurs in many borate type minerals. Boric acid reacts with alcohols in the presence of sulphuric acid to give alkyl orthoborate compounds, B(OR)3, which can be reacted with alkali metal hydrides to give borohydride anions of the form [BH(OR)3]- which are useful reducing agents. It will also form stable compounds by reacting with dialcohols, eg (HO)H2C-CH2(OH), to give cyclic species, eg see figure 2.
Figure 2: The cyclic product from the reaction of boric acid with ethane diol.
The best known are the trihalides, BX3. They are all strong Lewis acids acting as electron pair acceptors. They are quite volatile with BF3 and BCl3 being gases at room temperature. They are all water-sensitive and undergo hydrolysis to give boric acid and HX though BF3 is only partially hydrolysed giving a mixture of boric acid and the tetrafluoroborate anion, BF4-. For example
BCl3 + 3 H2O → B(OH)3 + 3 HCl
4 BF3 + 6 H2O → 3 H3O+ + B(OH)3 + 3 BF4-
Lewis bases, such as amines (R3N), phosphines (R3P), ethers (R-O-R) or anions, react with the boron trihalides by donating lone pairs of electrons forming larger associated molecules called adducts. In these reactions, the boron goes from being trigonal planar and hence sp2 hybridised to being tetrahedral and therefore sp3.
BX3 + :N(CH3)3→ X3B-N(CH3)3
Since the most electronegative halide is fluorine we would expect BF3 to form the most stable adducts as the boron would more readily accept a lone pair of electrons, however this is not the case. It is the heavier boron halides such as BBr3 that form the more favoured Lewis acid/base adducts. This reverse order in the expected reactivity comes down to the nature of the bonding in the BX3 reactant. There is a partial double bond formed between the B and X atoms. When X is Cl or Br, this partial double bonding involves the 3p and 4p orbitals of the halogens which are larger and differ in energy to the 2p orbitals of the boron atom. Since fluorine is in the same row of the periodic table as boron, the 2p orbitals it uses are of comparable size and energy to those on boron and so there is much better overlap than there is for Cl and Br giving a greater double bond character. This somewhat stabilises the unsaturated BF3 as a lot more energy is required to overcome this partial double bonding than in BBr3.
If the Lewis base contains an acidic proton then instead of adduct formation we can get a substitution reaction that results in the elimination of HX and can sequentially lead to complete solvolysis, for example
BCl3 + C2H5OH → Cl2B-OC2H5 + HCl →→ B(OC2H5)3
Another feature of the reactivity of the boron trihalides is halide exchange. If one halide of boron, BX3 is mixed with a another halide of boron, BY3 then we form mixed halide boron compounds, eg BX2Y. The reaction is thought to proceed by the formation of transient dimers where a halide one on boron atom forms a bond to bridge to the boron of another molecule. With the new B-X bond formed, the old B-X bond is broken and so the halide atom is transferred though these transient species have not been detected.
There are also many examples of subhalide compounds of boron where the ratio of halide atom to boron atoms is less than 3. We can get simple molecules such the unstable monohalides BF and BCl which can be generated in the gas phase, and the diboron tetrahalides, B2X4 in which there is a B-B bond. We can also get compounds of the form BnXn such as B9Cl9 which are cluster molecules in which the boron atoms are bonded to each other forming a complex polyhedron.
Nitrogen Compounds of Boron
Compounds containing a B-N bond (boron nitrides) are interesting chemical species. Since N has one more electron and B has one less electron than carbon, a B-N fragment has the same number of electrons as a C-C fragment (they are said to be isoelectronic) and may be thought of as being analogous. As we have seen before, the reaction of ammonia, NH3, with elemental boron at high temperatures leads to the formation of compound with the same structure as graphite but with alternating B and N atoms. In fact, boron nitride possesses three polymorphic forms, one analogous to graphite, another to diamond and another to the fullerenes. The diamond-like allotrope, although one of the hardest materials known, is softer than diamond.
Another allotrope of boron nitride, borazon, is harder than diamond. This was first produced Robert H Wentorf of the General Electric Company in 1957, by heating equal quantities of boron and nitrogen at temperatures in excess of 1800 °C at 7 GPa pressure.
The simplest boron-nitrogen compounds, however, are the amine boranes. These are Lewis acid/base adducts of amines and boranes (see later) but can also be formed from the reaction of ammonium salts with a tetrahydroborate, BH4, containing compound
[RNH3]Cl + NaBH4→ LiCl + H2 + RH2N-BH3
These compounds with a single B-N bond are somewhat analogous to alkanes containing C-C bonds. The strength of the bond depends on the nature of the groups already on N and B. An amine borane that has a weak bond that is essentially a donor-acceptor interaction would have the bond drawn as an arrow, ie N→B. If the electrons from nitrogen are completely shared then we will have a stronger full bond. Since the nitrogen atom has fully donated it's lone pair and forms four bonds we draw it as having a positive charge whilst the boron atom which accepts the electrons is drawn with a negative charge, ie N+-B-.
Just as the reaction of boron trihalides with alcohols results in elimination of HX, amine boranes can also eliminate HX to give aminoboranes. The bonding interaction between the B and N atoms can be imagined to be somewhere between being a single B-N bond and a B-=N+ double bond. These compounds can therefore be thought of as being analogous to alkenes containing C=C double bonds, eg
(CH3)2NH + BCl3→ (CH3)2HN→BCl3→ (CH3)2N+=B-Cl2 + HCl
One of the main reactions of aminoboranes is substitution of the groups on boron. If the product aminoborane in the example above is treated with a Grignard reagent, RMgX (see organometallic compounds of the alkaline earth metals), the Cl atoms will be replaced by organic groups labelled R.
(CH3)2N+=B-Cl2 + 2 RMgCl → (CH3)2N+=B-R2 + 2 MgCl2
Aminoboranes can also condense to form cyclic dimeric compounds, see figure 3. If this condensation also occurs with elimination of HX then we can get formation of six-membered cyclic molecules called borazines.
Figure 3: Structures of cyclic aminoborane dimers, borazine and it's carbon based structural analogue, benzene.
Borazine is structurally similar and is isoelectronic with it's carbon based analogue, benzene. The two have similar physical properties, however borazine is much more reactive than the fairly inert benzene. This is due to the polarity in the B-N bonds of the ring as there are significant positive and negative charges on the nitrogen and boron atoms respectively.
Boron Hydrides - Boranes
Boron has an extensive chemistry with hydrogen forming a multitude of compounds called boranes. These have varying ratios of B to H atoms and form polyhedral structures with complex bonding between the boron atoms and can also incorporate atoms of ther elements (for pictures of typical boranes follow this link. In naming boranes, we first give the name with the prefix for the number of boron atoms and then give the number of hydrogen atoms in brackets, for example B4H10tetraborane(10). The hydrogen atoms in these compounds can be terminal, bonding to just one boron atom, or can occupy bridging positions where it is bound to two B atoms.
The bonding in boranes can't be described conventionally as being wholly in terms of two electron between two atoms. The polyhedral cluster formation in boranes is due to their electron deficiency, there are not enough valence electrons to go round to form such two centre two electron (2c-2e) bonds between all the atoms. Instead many atoms come together and share their electrons engaging in so called multicentre bonding, ie one chemical bond can be formed between a number of atoms using two electrons. In the polyhedral boranes (and the previously mentioned polyhedral monohalides), several multicentre bonds hold the structures together.
The simplest borane is diborane(6), B2H6. This molecules exist as a dimer of two BH3 units in which two H atoms bridge the two B atoms which are tetrahedral (see figure 4). Each B-H-B fragment forms a three centre two electron (3c-2e) bond whilst the terminal hydrogen atoms form conventional 2c-2e bonds.
Figure 4: Structure of diborane(6)
Diborane(6) is a gas that is spontaneously flammable on contact with air and instantly hydrolyses on addition of water to boric acid and hydrogen gas. It can be formed by reacting sodium borohydride, NaBH4 with either BF3 or by reduction with iodine. On industrial scales it can be made by reducing BF3 with sodium hydride.
3 NaBH4 + 4 BF3→ 2 B2H6 + 3 NaBF4
2 NaBH4 + I2→ B2H6 + 2 NaI + H2
2 BF3 + 6 NaH → B2H6 + 6 NaF
Diborane(6) is a useful reagent for making organoboron compounds which are employed in organic synthesis and can be used to form higher boranes by controlled heating and loss of hydrogen. The heavier boranes are mostly liquids and have generally have greater stability to air and water with increasing molecular weight but their thermal stability is variable, see table 1.
Table 1: Chemical and physical properties of some boranes.
|Reaction with air, thermal stability and reaction with water at 25 °C|
|Diborane(6), B2H6||-165||-93||Spontaneously flammable, thermally stable, hydrolyses instantly|
|Tetraborane(10), B4H10||-120||18||Not flammable if pure, decomposes at 25 °C, hydrolyses over 24 h|
|Pentaborane(9), B5H9||-47||48||Spontaneously flammable, thermally stable at 25 °C but decomposes on heating, hydrolyses with heating|
|Penatborane(11), B5H11||-123||63||Spontaneously flammable, rapid thermal decomposition at 25 °C, fast hydrolysis|
|Hexaborane(10), B6H10||-62||108||Spontaneously flammable, slow decomposition at 25 °C, hydrolyses with heating|
|Hexaborane(12), B6H12||-82||80-90||Stable for a few hours at 25 °C, hydrolyses to give hydrogen, boric acid and B4H10|
|Octaborane(12), B8H12||-20||-||Decomposition above -20 °C|
|Decaborane(14), B10H14||100||213||Non flammable, stable up to 150 °C, slow hydrolysis|
There are also examples of anionic boron hydrides, the simplest being borohydrides containing the BH4- ion. The most common example is sodium borohydride, NaBH4 which is an important reagent in both organic and inorganic chemistry as a source of H- and as a reducing agent. There also a range of polyhedral anionic boranes which have the general formula BnHn2-. Their structure and bonding follows the same rules as for the neutral boranes. There are also a range of neutral compounds called carboranes in which two BH- fragment are replaced with CH or C-alkyl group. These are called carboranes and have the general formula Bn-2C2Hn. These can be produced by reacting a borane with a thioether (R-S-R) to displace H2, followed by an alkyne, for example
B10H14 + 2 R2S → B10H12(R2S)2 + H2
B10H12(R2S)2 + RC≡CR → B10H10(CR)2 + 2 R2S + H2
The chemistry of boranes, borane anions and carboranes is truly vast and we are only able to have the briefest glimpse here. There is an enormous range of compounds of these types that have been produced and there are even more interesting examples in which fragments of transition metal complexes can be introduced into the structure in place of BH units. There are also examples of anionic carboranes which can act as ligands for transition metal ions.