These electrons will be used in the formation of the bonds. Register now! The delocalization of the electrons means that there aren't alternating double and single bonds. The extra energy released when these electrons are used for bonding more than compensates for the initial input. compare the reactivity of a typical alkene with that of benzene. Here, carbon is the central atom. Unless otherwise noted, LibreTexts content is licensed by CC BY-NC-SA 3.0. It is a regular hexagon because all the bonds are identical. The three sp2 hybrid orbitals arrange themselves as far apart as possible - which is at 120° to each other in a plane. The carbon atom is now said to be in an, The difference in benzene is that each carbon atom is joined to two other similar carbon atoms instead of just one. This diagram shows one of the molecular orbitals containing two of the delocalized electrons, which may be found anywhere within the two "doughnuts". Further, the carbon atom lacks the required number of unpaired electrons to form the bonds. Notice that the p electron on each carbon atom is overlapping with those on both sides of it. In the diagram, the sigma bonds have been shown as simple lines to make the diagram less confusing. The nitrogen has a lone pair of electrons perpendicular to the ring. The energetic stability of benzene: This is accounted for by the delocalization. Note that the figure showing the molecular orbitals of benzene has two bonding (π2 and π3) and two anti-bonding (π* and π5*) orbital pairs at the same energy levels. We will look at the details below. . Register now! In practice, 1,3-cyclohexadiene is slightly more stable than expected, by about 2 kcal, presumably due to conjugation of the double bonds. Legal. However, the major constraint is the angle $\ce{C^6-C^1-C^2}$, which is compressed to a mere $111°$. Benzene (\(C_6H_6\)) is a planar molecule containing a ring of six carbon atoms, each with a hydrogen atom attached. This is accounted for by the delocalization. The carbon atoms in the benzene ring are arranged in a trigonal planar geometry. Only a part of the ring is shown because the diagram gets extremely cluttered if you try to draw any more. The six carbon atoms form a perfectly regular hexagon. Each carbon atom now looks like the diagram above. In this, 1 s orbital and two p orbitals are hybridized and form three sp2 hybridized orbitals. The six delocalized electrons go into three molecular orbitals - two in each. Evidence for the enhanced thermodynamic stability of benzene was obtained from measurements of the heat released when double bonds in a six-carbon ring are hydrogenated (hydrogen is added catalytically) to give cyclohexane as a common product. This extensive sideways overlap produces a system of pi bonds which are spread out over the whole carbon ring. What happens next is the promotion of one 2s2 electron pair to the empty 2pz orbital. The six-membered ring in benzene is a perfect hexagon (all carbon-carbon bonds have an identical length of 1.40 Å). The hybridization is sp 2 type. For more information contact us at info@libretexts.org or check out our status page at https://status.libretexts.org. For more information contact us at info@libretexts.org or check out our status page at https://status.libretexts.org. The delocalization of the electrons means that there aren't alternating double and single bonds. The next diagram shows the sigma bonds formed, but for the moment leaves the p orbitals alone. The delocalization of the electrons means that there aren't alternating double and single bonds. The LibreTexts libraries are Powered by MindTouch® and are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. Notice that the p electron on each carbon atom is overlapping with those on both sides of it. Because the electrons are no longer held between just two carbon atoms, but are spread over the whole ring, the electrons are said to be delocalized. You may wish to review Sections 1.5 and 14.1 before you begin to study this section. Orbitals with the same energy are described as degenerate orbitals. Benzene, however, is an extraordinary 36 kcal/mole more stable than expected. As a general principle, the more you can spread electrons around - in other words, the more they are delocalized - the more stable the molecule becomes. The delocalization of the p-orbital carbons on the sp2 hybridized carbons is what gives the aromatic qualities of benzene. This is easily explained. Here, two structurally and energetically equivalent electronic structures for a stable compound are written, but no single structure provides an accurate or even an adequate representation of the true molecule. Problems with the chemistry. The conceptual contradiction presented by a high degree of unsaturation (low H:C ratio) and high chemical stability for benzene and related compounds remained an unsolved puzzle for many years. Legal. The shape of benzene: Benzene is a planar regular hexagon, with bond angles of 120°. 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It is a regular hexagon because all the bonds are identical. Free LibreFest conference on November 4-6! It is this completely filled set of bonding orbitals, or closed shell, that gives the benzene ring its thermodynamic and chemical stability, just as a filled valence shell octet confers stability on the inert gases. In chemistry, the Z-matrix is a way to represent a system built of atoms.A Z-matrix is also known as an internal coordinate representation.It provides a description of each atom in a molecule in terms of its atomic number, bond length, bond angle, and dihedral angle, the so-called internal coordinates, although it is not always the case that a Z-matrix will give information regarding bonding since the matrix itself … Among the many distinctive features of benzene, its aromaticity is the major contributor to why it is so unreactive. Because each carbon is only joining to three other atoms, when the carbon atoms hybridize their outer orbitals before forming bonds, they only need to hybridise three of the orbitals rather than all four. state the length of the carbon-carbon bonds in benzene, and compare this length with those of bonds found in other hydrocarbons.

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