SCH4U: Structure and Properties of Matter 2.7 Hybridization Valence Bond Theory A simple approach to describing the location of electrons in bonds is to start with an atom and its valence orbitals and to form covalent bonds in these atoms through the valence orbitals. This method is called the valence bond theory. According to valence bond theory, a covalent bond forms when 2 atomic orbitals, each with an unpaired electron, overlap. When the covalent bond forms, the lowest energy state is obtained when participating electrons are of opposite spin. The simplest example of the valence bond theory is the bonding of 2 hydrogen atoms. The electrons in each of the hydrogen atoms reside in the 1s orbital. When the two 1s orbitals overlap, a covalent bond forms. In this bond, the electrons in the bonded 1s orbital have opposite spins, similar to a filled atomic orbital. HCl also forms a covalent bond. The hydrogen atom contributes 1 electron from its 1s orbital and chlorine contributes 1 electron from its 3 p orbital to form the bond. The p orbital in the chlorine atom has 2 lobes. When the bond forms, 1 lobe of the chlorine atom overlaps with the 1s orbital from the hydrogen atom. SCH4U: Structure and Properties of Matter The valence bond theory can also describe bonding in polyatomic molecules. Ammonia: 1 nitrogen atom and 3 hydrogen atoms arranged in a trigonal pyramidal structure. You would expect… Hybrid Orbitals The chemical bonds in molecules such as methane can be explained by the concept of hybrid orbitals. Hybrid orbitals form by combining orbitals with different shapes. For the carbon atom, you start with 1 s orbital and 3 p orbitals, and finish with four hybrid sp3 orbitals. The process of forming hybrid orbitals is called hybridization. SCH4U: Structure and Properties of Matter Each of the hybrid sp3 orbitals is identical to the other. Since each hybrid orbital contains an unpaired electron, these sp3 orbitals can overlap with the 1s orbital of a hydrogen atom to form the 4 equal covalent bonds in a methane molecule. The sp3 hybrid orbitals have a large lobe and a small lobe pointing in opposite directions. The large lobes of the four sp3 orbitals have the tetrahedral arrangement predicted for the bonding electron pairs in the VSEPR description of the methane molecule. There are many other examples where hybridization is needed to account for the observed properties of chemical bonds. sp Hybrid SCH4U: Structure and Properties of Matter sp2 Hybrid Double and Triple Covalent Bonds sigma s bond: a bond that is formed when the lobes of 2 orbitals directly overlap end to end SCH4U: Structure and Properties of Matter Double and Triple Bonds pi p bond: a bond that is formed when the sides of the lobes of 2 orbitals overlap Ethene SCH4U: Structure and Properties of Matter Ethyne Benzene Worksheet 2.7: Hybridization p. 238, Q 1, 2, 3, 4, 7