Boiling point = - 92.6 degrees Celsius Spontaneously flammable in air Instantly hydrolyzed by water forming hydrogen gas and boric acid -
-
University of Bristol Homepage:
University of Bristol:School of Chemistry
My E-mail Address:
-
BRIEF BACKGROUND
Boron forms an extensive range of molecular hydrides, from the gaseous diborane ( a chain consisting of only two Borons) to the solid decaborane (a chain cosisting of ten Borons). The first boranes were prepared between 1912 and 1936 using the vacuum line technique developed by Alfred Stock specifically for the preperation of the boranes. This was significant as boranes are very reactive materials and therefore were hard to produce experimentally. Stocks' original method is still used to prepare hexaborane, but has been superseded for the preperation of all other boranes. As an example of the reactivity of the boranes, and thus the reasons for the great dificulty in preparing them, here is a list of the properties of diborane:
INSIDE THE BORANES
The uniqueness of the boranes stems from the structures which they form, which can only be described as unlike any other hydride family. The structural anomaly which gives the boranes their individuality is that there are not sufficient electrons to allow the formation of conventional two-electron bonds between all adjacent pairs of atoms. This causes an electron deficiency which is rationalized by using the concept of the multicenter bonding.
CONCEPT OF MULTICENTER BONDING.
In the simplest borane - diborane - there are eight adjacent atoms and only six pairs of electrons available for bonding. If there is to be the normal distribution of two-center two-electron (2c-2e) bonding, then eight bonds are required. A resonance structure could be used to account for this, with two canonical forms. The conclusions from this would have to be that there is one electron pair distributed over two B-H bonds, resulting in a bond order of one half. This would effectively give a B-H-B bridge. The rest of the B-H bonds in the molecule would be normal 2c-2e bonds. Alternatively, the problem could be looked at using the Molecular Orbital(MO) theory. This approach to the problem only encompasses the bridging system, and treats the terminal B-H groups separately as localized elctron-pair bonds. Therefore, an H-B-H group, which has equivalent sp(3) hybridized orbitals, has two ordinary B-H bonds, with a single hybridized orbital available. If two such groups of H-B-H were brought together, with the free single hybridized orbitals in each molecule facing each other, the remaining hydrogens would assume a bridging position, and an orbital would form which would bridge over the hydrogens. Below is a graphical representation of this.
The resulting molecule formed is diborane in this case, with a formula of B(2)H(6). Both of the B--H--B orbitals are three-centered and have no nodes. This means that the orbitals are capable of bonding together three atoms. There are four electrons over the two three-centerd orbitals (one from each of the borons and one from each of the hydrogens). Therefore, there are is one electron pair for each three-centered orbital. The type of bond resulting is called a 3 center,2 electron bond, or 3c-2e fro short. These types of bonds have approximately one half of the strength of a normal 2c-2e bond, as each electron pair is spread over three atoms instead of two. This is equivalent to a bond order of one half obtained in the resonance treatment.
A DEEPER LOOK INTO THE 3c-2e BOND.
If we consider the sp(3) orbital on each boron atom and the 1(s) orbital on each bridging Hydrogen atom, we see that in a B-H-B bond, there are three orbitals cotributing to the bond. Therefore there are three Molecular Orbitals (MO's) which can be drawn up to represent the possibiities. The first MO is a bonding orbital, and is the one which has been discussed. The second MO is an antibonding orbital, which has a node between each adjacent pair of atoms (the hydrogen assumes the bridging position, but has opposite charge to the borons). The third is a non-bonding orbital which is the signs of the two sp(3) orbitals out of phase . This results in no net overlap with the 1s orbital of the hydrogen, and thus no bond.
This little section is devoted to the person that brought you this little look into the life of Boron - me. I am Kieran Parkinson and I am student. You are probably wondering why a student is spending his time making a website, when he should be working hard for his degree. Well, I will fill you in some things about me before you jump to any conclusions. I am studying chemistry(MSci) at the University of Bristol and I am in my first year. I had never used the internet or e-mail or anything like that before I came to Bristol. I suppose that doesn't really explain why I am sitting here, investing my time and energy, making this web page. Although chemistry is my honours course, I have to do two other subsiduary subjects, which are maths and computer methods, and as part of the computer course, I have had to spend a large part of the last 10 weeks creating this page and everything in it. I have researched into various aspects of boron chemistry and the further I dive into it, the bigger pool of information gets. I had to learn how to write in HTML as well, and search the internet for various sites which were connected to the subject of my page. If that wasn't enough, I am going to get examined on this in a written paper at the end of May as part of the first year credit points system. Don't get me wrong though, as it has been fun. Well , thats enough about me, so if you want to write to me, here is my e-mail address, and the University of Bristol address. I hope you enjoyed the page, and if you didn't you have to remember that I am a beginner, and I have had to fit this in around my intense studying program. Goodbye.
http://www.bris.ac.uk
http://www.bris.ac.uk/depts/chemistry/
kp6881@bris.ac.uk