With nitrogen, however, there are five rather than four valence electrons to account for, meaning that three of the four hybrid orbitals are half-filled and available for bonding, while the fourth is fully occupied by a (non-bonding) pair of electrons. Just like the carbon atom in methane, the central nitrogen in ammonia is sp 3-hybridized. The sp 3 bonding picture is also used to described the bonding in amines, including ammonia, the simplest amine. In another module we will learn more about the implications of rotational freedom in sigma bonds, when we discuss the ‘conformation’ of organic molecules. This means, in the case of ethane molecule, that the two methyl (CH 3) groups can be pictured as two wheels on a hub, each one able to rotate freely with respect to the other. All of these are sigma bonds.īecause they are formed from the end-on-end overlap of two orbitals, sigma bonds are free to rotate. The carbon-carbon bond, with a bond length of 1.54 Å, is formed by overlap of one sp 3 orbital from each of the carbons, while the six carbon-hydrogen bonds are formed from overlaps between the remaining sp 3 orbitals on the two carbons and the 1 s orbitals of hydrogen atoms. Both carbons are sp 3-hybridized, meaning that both have four bonds arranged with tetrahedral geometry. In the ethane molecule, the bonding picture according to valence orbital theory is very similar to that of methane. (It will be much easier to do this if you make a model.) In the images below, the exact same methane molecule is rotated and flipped in various positions. Imagine that you could distinguish between the four hydrogens in a methane molecule, and labeled them H a through H d. Interactive 3D model of the bonding in methane This system takes a little bit of getting used to, but with practice your eye will learn to immediately ‘see’ the third dimension being depicted. Normal lines imply bonds that lie in the plane of the page. A dashed wedge represents a bond that is meant to be pictured pointing into, or behind, the plane of the page. In this convention, a solid wedge simply represents a bond that is meant to be pictured emerging from the plane of the page. To do this on a two-dimensional page, though, we need to introduce a new drawing convention: the solid / dashed wedge system. While previously we drew a Lewis structure of methane in two dimensions using lines to denote each covalent bond, we can now draw a more accurate structure in three dimensions, showing the tetrahedral bonding geometry. The length of the carbon-hydrogen bonds in methane is 1.09 Å (1.09 x 10 -10 m). This is simply a restatement of the Valence Shell Electron Pair Repulsion (VSEPR) theory that you learned in General Chemistry: electron pairs (in orbitals) will arrange themselves in such a way as to remain as far apart as possible, due to negative-negative electrostatic repulsion.Įach C-H bond in methane, then, can be described as an overlap between a half-filled 1 s orbital in a hydrogen atom and the larger lobe of one of the four half-filled sp 3 hybrid orbitals in the central carbon. This geometric arrangement makes perfect sense if you consider that it is precisely this angle that allows the four orbitals (and the electrons in them) to be as far apart from each other as possible. The larger lobes of the sp 3 hybrids are directed towards the four corners of a tetrahedron, meaning that the angle between any two orbitals is 109.5 o. Unlike the p orbitals, however, the two lobes are of very different size. The sp 3 hybrid orbitals, like the p orbitals of which they are partially composed, are oblong in shape, and have two lobes of opposite sign. In the new electron configuration, each of the four valence electrons on the carbon occupies a single sp 3 orbital. In this picture, the four valence orbitals of the carbon (one 2 s and three 2 p orbitals) combine mathematically (remember: orbitals are described by equations) to form four equivalent hybrid orbitals, which are named sp 3 orbitals because they are formed from mixing one s and three p orbitals. In order to explain this observation, valence bond theory relies on a concept called orbital hybridization. How does the carbon form four bonds if it has only two half-filled p orbitals available for bonding? A hint comes from the experimental observation that the four C-H bonds in methane are arranged with tetrahedral geometry about the central carbon, and that each bond has the same length and strength. Recall the valence electron configuration of the central carbon: Now let’s turn to methane, the simplest organic molecule. Hybrid orbitals: sp 3 hybridization and tetrahedral bonding
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