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Hybridization. Hybridization. In Lewis theory, a covalent bond is a shared pair of electrons. This suggests that there is a region of electron density between two atoms as a result of the electron pair being shared.
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Hybridization • In Lewis theory, a covalent bond is a shared pair of electrons. • This suggests that there is a region of electron density between two atoms as a result of the electron pair being shared. • Atomic orbital theory suggests that electrons exist in atomic orbitals that are regions of space surrounding the nucleus of an atom.
Hybridization • Each atomic orbital can hold no more than two electrons. • How can we use atomic orbitals to explain bonding and to account for the shapes of molecules? • This leads to a model of chemical bonding called valence-bond theory.
Overlapping Orbitals • In valence-bond theory, a region of electron density between atoms is explained in terms of the overlap of two atomic orbitals. • Two orbitals from different atoms, each containing an unpaired electron, can merge in the region of space between the two atoms. • The overlapping of two orbitals creates a bonding orbital between the two atoms.
Overlapping Orbitals • When two p orbitals overlap in the same axis, or an s orbital and a p orbital or even two s orbitals overlap, the bond formed is called a σ (sigma) bond.
Overlapping Orbitals • This overlapping orbitals explains the difference in bond length between two single bonds, such as an F-F bond and a Cl-Cl bond. • The two orbitals overlapping in the case of the F-F bond are both 2p orbitals, whereas in the case of the Cl-Cl bond, the overlapping orbitals are both 3p orbitals.
Overlapping Orbitals • The 3p orbitals are larger than the 2p orbitals, so when two 3p orbitals overlap, they do so at a greater distance form the nucleus than do two 2p orbitals. • This causes the bond length in the Cl2 to be greater than in the F2.
Overlapping Orbitals • Recall that the three p orbitals are at right angles to each other. • When two px orbitals are overlapping in a σ bond, the other two p orbitals, py and pz are not overlapping.
Overlapping Orbitals • When two atoms come close enough together, the py and/or the pz orbitals are able to overlap in a sideways fashion. • This sideways overlapping of two orbitals is called a π (pi) bond. • In a π bond the overlapping is not as great as in a σ bond, so a π bond is not as strong as a σ bond.
Overlapping Orbitals • π bonds occur most commonly when a multiple bond (a double or a triple bond) is formed. • A double bond is formed by one σ bond and one π bond. • Since a π bond is not as strong as a σ bond, a double bond is less than twice as strong as a single covalent bond between the same two elements.
Overlapping Orbitals • A triple bond is formed by one σ bond and two π bonds and is less than three times as strong as a single bond.
Hybrid Orbitals • Bond lengths can tell us much about the equality of covalent bonds between atoms. • When the bond lengths of the four C-H bonds in methane, CH4, are measured, they are all found to be equal. • This suggests that they are all equal in nature.
Hybrid Orbitals • The electronic configuration of carbon is 1s22s22p2 and that of hydrogen is 1s1. • An orbital that overlaps with another must have only one electron in it, so that the new overlapping (bonding) orbital has no more than 2 electrons in it. • The valence configuration of carbon causes problems.
Hybrid Orbitals • There are only 2 orbitals with just one electron in them. • The 2s orbital is full and the third 2p orbital has no electrons in it at all. • If the four bonds that carbon is commonly observed to make are to exist, there must be four unpaired electrons in the valence shell. • This is made possible if the atomic orbitals in the valence shell are assumed to mix together to make hybrid orbitals.
Hybrid Orbitals • The shape of a hybrid orbital is different from that of the two orbitals from which it is formed, but the total number of orbitals in the valence shell remain the same.
Hybrid Orbitals • Note that there is not energy change involved in forming hybrid orbitals; the is, the total energy is the same before and after hybridization has occurred. • The process of mixing atomic orbitals as atoms approach each other is called hybridization.
sp3 Hybridization • The case of carbon’s four equal bonds can be explained by hybridization. • The four orbitals in the valence shell of carbon, 2s, 2px, 2py and 2pz, are mixed together. • Whenever we mix a certain number of atomic orbitals, the same number of hybrid orbitals is formed. • Hence, we get four sp3 hybrid orbitals.
sp3Hybridization • The shape of an sp3 hybrid orbital is not the same as either the s orbital or the p orbital. • Like p orbitals it has two lobes, but, unlike p orbitals, one lobe is much bigger than the other. • When put together, the four lobes of the sp3 orbitals take up a tetrahedral shape in space. • This shape allows the electron pairs to be as fair apart as possible.
sp3 Hybridization • The H-C-H bond angle is 109.50. • In this manner the hybridization of atomic orbitals agrees with the requirements of VSEPR theory. • Each orbital is equal in size, hence allowing the four bonds to be equal in length and strength when they overlap with eh orbital of another atom.
Methane • A molecule of methane is formed when each of the four sp3 hybrid orbitals of carbon overlap in σ bonds with the s orbitals of four hydrogen atoms.
sp2 Hybridization • Boron is an element that only forms three bonds. • It has three electrons in its valence shell and will form three covalent bonds with elements such as hydrogen and fluorine. • These three equal covalent bonds can be explained by the existence of hybrid orbitals. • The electron configuration of boron is 1s22s22p1. • If the 2s orbital and the two 2p orbitals are mixed, the hybrid orbitals will be called sp2 hybrid orbitals.
sp2 Hybridization • The sp2 hybrid orbital will be arranged around the central atom in such a way as to minimize electrostatic repulsion. • This results in a trigonal planar arrangement.
Ethene • Ethene, C2H4, has a C-C double bond between the two carbon atoms and there is a trigonal planar arrangement of atoms around each carbon. • Although carbon has four orbitals in the valence shell, only three of them (2s, 2px and 2py) are involved in the hybridization. • The fourth orbital, the 2pz orbital, exists at right angles to the plane of the hybrid orbitals.
Ethene • The sp2 hybridized orbitals of the carbon atom in ethene form σ bonds with hydrogen 1s orbitals and with the sp2 hybridized orbital of the other carbon atom. • The atoms are sufficiently close that the 2pz orbitals of the two carbon atoms can overlap sideways forming a π bond. • The combination of the σ bond between the hybridized orbitals and the π bond between the two 2pz orbitals creates a double bond between the two carbon atoms.
sp Hybridization • Beryllium has only two electrons in its valence shell. • Its electron configuration is 1s22s2. • It makes two equal-sized bonds with elements such as hydrogen and the halogens. • In order for two equal bonds to be made, the 2s and one 2p orbital must be mixed to form two sp hybrid orbitals. • This results in a linear arrangement.
sp Hybridization • In beryllium hydride, BeH2, the sp hybrid orbitals of beryllium form σ bonds with the 1s orbitals of hydrogen. • The H-Be-H bond angle is 180o. • The shape of these molecules is linear.
Hybridization • Generally speaking, if there are four negative charge centres around the central atom, the hybridization will be sp3; if there are three negative charge centres, the hybridization will be sp2; and if there are two negative charge centres, the hybridization will be sp.
Sample Problems • Both σ and π bonds involve the overlapping of orbitals. Draw a diagram to show the difference between the overlapping that constitutes a σ bond and the overlapping to make a π bond.
Sample Problems • Copy and complete the following table.
