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Bonding and Hybrid Orbitals MOLECULAR ORBITALS: Molecular orbitals can be thought of as the region of space that two electrons that are shared between atoms inhabit. Typically these orbitals are a combination of the atomic orbitals that provided the electrons that are now shared between the atoms. Bonding occurs when atomic orbitals overlap. Orbitals that overlap imply that a pair of electrons, each of which was originally associated with only a single atom, now is likely within a region of space between two atoms. When one end of an orbital overlaps a second end of an orbital, and both orbitals are centered on a common axis, a sigma bond (σ bond) is formed. All single bonds occurring between atoms can be described as a sigma bond. Double bonds require more than this first, single sigma bond. Obviously it requires a second bond. This occurs when a p atomic orbital from each of the two atoms participating in the bond “overlap” forming weaker bonds on either side of the central sigma bond. This overlap of p-orbitals (and remember, these p orbitals will be at right angles to the initial sigma bond), will form what is called a pi bond (π bond) is formed. The double bond thus consists of one sigma and one pi bond. Triple bonds are just like double bonds except two pi bonds are formed between each atom in addition to the central sigma bond. Each atom participating in the triple bond must have two half-filled p-orbitals available to participate in the bonding. VSEPR THEORY (Valence Shell Electron Pair Repulsion) This theory is helpful when attempting to predict molecular structure. In a nut shell – pairs of electrons wish to separate themselves as far apart from each other as possible. This is a direct consequence of the fact that all electrons share the same electrical charge (negative). This results in an electrostatic force that repels all pairs of electrons away from one another. As a result, the pairs of electrons around any given atom will, in most cases, form one of the following structures for most covalent molecules. Linear (2 electron groups): The atoms in the molecule are in a straight line. This can be either need to be three atoms to get a bond angle) or because the three atoms are lined up in a straight line (corresponding to a 180 degree bond angle). Trigonal planar (3 electron groups): A central atom connected to three additional atoms. Angles between any two bonds are 120 degrees apart, and all four atoms will lie within a single plane. Boron commonly forms this arrangement (as the central atom). When it does, it becomes an exception to the octet rule, only needing six electrons to be happy. Tetrahedral (4 electron groups): Tetrahedral molecules look like pyramids with four faces. Each point on the pyramid corresponds to an atom that's attached to the central atom. Bond angles are 109.5 degrees. Pyramidal (4 electron groups): It's like a tetrahedral molecule, except flatter. It looks kind of like a squished pyramid because one of the atoms in the pyramid is replaced with a lone pair. Bond angles are 107.5 degrees (it's less than tetrahedral molecules because the lone pair shoves the other atoms closer to each other). Bent (4 electron groups): They look, well, bent. Bond angles can be either 118 degrees for molecules with one lone pair or 104.5 degrees for molecules with two lone pairs. Trigonal bipyramidal (5 electron groups): It looks like the hood ornament of a Mercedes automobile, or like a peace sign with that bottom-most line gone. The bond angles are 120 degrees. Octahedral (6 electron groups): It will look like a square pyramid (with four sides) with a o pointed top as well as a pointed bottom. Bond angles will typically be 90 . HYBRIDIZATION In order to have the orbitals necessary to complete the bonding you might predict when using dot structures, you need to be prepared to visualize orbitals that are combinations of the existing atomic orbitals of the valence electrons prior to bonding. Once again, in a nutshell, during hybridization, this model suggests that an atom’s single s-orbital and 3 p-orbitals combine, mix, and form new orbitals which then will bond via sigma bonding with orbitals from other atoms. There are five possible kinds of hybridizations (listed below). The hybridization that results depends upon how many “things” are bonding to the atom whose orbitals are undergoing hybridization. Often, carbon (as shown in the table below in the first three columns) is the atom undergoing hybridization. 3 3 2 Phosphorous is known to have sp d hybridization at times, and sulfur is known to experience sp d hybridization at times. Possible Hybrids: Summary 3 2 3 3 2 Type of Hybrid sp sp sp sp d sp d Atomic orbitals s, p, p, p s, p, p s, p s, p, p, p, d s, p, p, p, d, d used Number of hybrid orbitals 4 3 2 5 6 formed Number of 5 (but not C 6 (but not C) atoms bonded to 4 3 2 (P) (S) the C Number of 4 3 2 5 6 sigma bonds Number of left 0 1 2 0 0 over p orbitals Number of pi 0 1 2 0 0 bonds | \ =C= Bonding pattern -C- C= or | / -C≡ Note that the number of atoms bonded to the carbon atom and the number of sigma bonds are equal to the number of hybrid orbitals. Also, the number of pi bonds formed is equal to the number of p orbitals that were not hybridized. In each case, the total number of bonds (sigma and pi) and the total number of orbitals (hybrid and p) is four. These are some relationships to keep in mind as you sort through the bonding arrangements and hybridization that occurs in various compounds.
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