If we look only at the outer electrons as "electrons-in-boxes":. There are 3 unpaired electrons that can be used to form bonds with 3 chlorine atoms. The four 3-level orbitals hybridise to produce 4 equivalent sp 3 hybrids just like in carbon - except that one of these hybrid orbitals contains a lone pair of electrons.
Each of the 3 chlorines then forms a covalent bond by merging the atomic orbital containing its unpaired electron with one of the phosphorus's unpaired electrons to make 3 molecular orbitals.
You might wonder whether all this is worth the bother! Probably not! It is worth it with PCl 5 , though. You will remember that the dots-and-crosses picture of PCl 5 looks awkward because the phosphorus doesn't end up with a noble gas structure. This diagram also shows only the outer electrons. In this case, a more modern view makes things look better by abandoning any pretense of worrying about noble gas structures. If the phosphorus is going to form PCl 5 it has first to generate 5 unpaired electrons.
It does this by promoting one of the electrons in the 3s orbital to the next available higher energy orbital. Which higher energy orbital?
It uses one of the 3d orbitals. You might have expected it to use the 4s orbital because this is the orbital that fills before the 3d when atoms are being built from scratch.
Not so! Apart from when you are building the atoms in the first place, the 3d always counts as the lower energy orbital. The 3-level electrons now rearrange hybridise themselves to give 5 hybrid orbitals, all of equal energy. They would be called sp 3 d hybrids because that's what they are made from. The electrons in each of these orbitals would then share space with electrons from five chlorines to make five new molecular orbitals - and hence five covalent bonds.
Why does phosphorus form these extra two bonds? It puts in an amount of energy to promote an electron, which is more than paid back when the new bonds form. Put simply, it is energetically profitable for the phosphorus to form the extra bonds. The advantage of thinking of it in this way is that it completely ignores the question of whether you've got a noble gas structure, and so you don't worry about it.
Nitrogen is in the same Group of the Periodic Table as phosphorus, and you might expect it to form a similar range of compounds. In fact, it doesn't. For example, the compound NCl 3 exists, but there is no such thing as NCl 5. The reason that NCl 5 doesn't exist is that in order to form five bonds, the nitrogen would have to promote one of its 2s electrons. The problem is that there aren't any 2d orbitals to promote an electron into - and the energy gap to the next level the 3s is far too great.
Triple-bonding is not the only way the researchers got boron to mimic its superstar neighbour, carbon, though. They also coaxed boron atoms into forming a chain. Previous attempts to do this with boron failed and resulted in messy clusters. Next, Braunschweig and his team hope to free this boron chain from its scaffold and increase the chain length to form the boron equivalent of polyethylene, a common plastic.
A world of previously forbidden chemistry beckons. Trending Latest Video Free. X-ray crystal structures confirmed that the compound possesses true triple bond character. As expected, the distance between the boron atoms was shorter in the triple-bonded structure than in the compound with a double bond, matching closely with predicted figures.
The molecule is also linear, just as the equivalent triple-bonded carbon compound would be. Boron-containing compounds are already used in commercial production of organic light-emitting diodes, for example. Braunschweig, H. Science , — Zhou, M. Wang, Y. Download references. You can also search for this author in PubMed Google Scholar. Inorganic chemistry: Two-armed silicon May Cleaner, greener fireworks Apr Tough competition Jul And whereas a quadruple bond should be stronger than a triple bond between the same atoms, calculations suggested that the connection in C2 was actually weaker than the triple bond in ethyne C 2 H 2.
A few years ago the researchers made a compound containing a triple bond between bismuth and boron, and wondered if they could swap bismuth for a transition metal. They started their search with rhodium, largely because it only has one naturally occurring isotope, which makes it easier to work with experimentally and computationally. So they fired a laser at a disk containing powdered rhodium and boron to produce a stream of vaporized compounds, and separated them in a mass spectrometer.
Then they studied selected compounds by photoelectron spectroscopy, which measures the kinetic energies of electrons ejected from ionized samples and serves as a fingerprint of the structure and bonding in the molecules. This process allowed the researchers to identify two curious compounds: a boronyl-coordinated rhodium boride, RhB BO - , and RhB itself.
The team also used theoretical calculations to study how the orbitals of each atom in RhB overlapped to form molecular orbitals. This showed that the atoms were connected by two sigma bonds and two pi bonds.
Taken together, Wang says, all their evidence suggests that RhB is the first diatomic molecule containing a quadruple bond to a boron atom.
RhB is not a new molecule, though. It was made more than a decade ago Mol. This was because researchers assumed that one of the sigma-bonding molecular orbitals was a non-bonding lone pair on the boron atom, Wang says. Morse at the University of Utah, who has previously worked with RhB. That delay meant they were unable to claim the first ever quadruple bond to boron—another team reported quadruple bonding in BFe CO 3 - late last year Nat.
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