Theories of Covalent Bonding

Theories of Covalent Bonding

Theories of Covalent Bonding 11.1 VSEPR Theory 11.2 Valence Bond Theory 11.3 Molecular Orbital (MO)Theory

A - central atom X -surrounding atom E -nonbonding valence electron-group integers VSEPR - Valence Shell Electron Pair Repulsion Theory Each group of valence electrons around a central atom is located as far away as possible from the others in order to maximize repulsions. These repulsions maximize the space that each object attached to the central atom occupies. The result is five electron-group arrangements of minimum energy seen in a large majority of molecules and polyatomic ions. The electron-groups are defining the object arrangement,but the molecular shape is defined by the relative positions of the atomic nuclei. Because valence electrons can be bonding or nonbonding, the same electron-group arrangement can give rise to different molecular shapes. AXmEn

Valence Bond Theory • Lewis structures and VSEPR do not explain why a bond forms. • How do we account for shape in terms of quantum mechanics? • What are the orbitals that are involved in bonding? • We use Valence Bond Theory: Chapter 9

The Central Themes of VB Theory Themes A set of overlapping orbitals has a maximum of two electrons that must have opposite spins. • The greater the orbital overlap, the stronger (more stable) the bond. • Types of overlap: • 1. head-on overlap • -formed by sigma bonds • Sigma bonds • -formed by pi bonds There are two electrons of opposite spin in the orbital overlap

Valence Bond Theory Orbital overlap and spin pairing in Hydrogen, H2 Chapter 9

Valence Bond Theory • As two nuclei approach each other their atomic orbitals overlap. • As the amount of overlap increases, the energy of the interaction decreases. • At some distance the minimum energy is reached. • The minimum energy corresponds to the bonding distance (or bond length). • At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron). • As the two atoms get closer, their nuclei begin to repel and the energy increases. Chapter 9

Valence Bond Theory • Could an overlap occur between diatomic molecules with different orbitals? • Example: HCl • H: 1s1 • Cl: 1s2 2s2 2p6 3s2 3p5 • 1s2 2s2 2p6 3s2 3p2 3p2 3p1 • By definition, covalent bonding forms by electron pairing (sharing) and since one 3p orbital has an unpaired electron, therefore: Chapter 9

Valence Bond Theory Hydrogen chloride, HCl Fluorine, F2 • Note: • The electrons in the overlap region are of opposite spin. Chapter 9

Hybrid Orbitals • Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding. • Key points: • The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. • The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. Chapter 9

Hybrid Orbitals • Hybridization is determined by the electron domain geometry • Types of Hybrid Orbitals: • sp • sp2 • sp3 • sp3d • sp3d2 Chapter 9

Hybrid Orbitals • sp Hybrid Orbitals • Consider the BeCl2 molecule (experimentally known to exist): • Be has a 1s22s2 electron configuration. • Cl has a [Ne] 3s2 3p5 electronic configuration. • Be has no unpaired electron available for bonding. • We conclude that the atomic orbitals are not adequate to describe orbitals in molecules. • We know that the Cl-Be-Cl bond angle is 180 (VSEPR theory). • We also know that one electron from Be is shared with each one of the unpaired electrons from Cl. Chapter 9

Hybrid Orbitals • sp Hybrid Orbitals • How could BeCl2 occur then? • How does a covalent bond form between Be and l atoms? • How could a linear geometry be possible? Chapter 9

Hybrid Orbitals • sp Hybrid Orbitals • We could promote an electron from the 2s orbital on Be to the 2p orbital to get the two unpaired electrons for bonding. Chapter 9

Hybrid Orbitals • We have solved the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital. • The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital. Chapter 9

Hybrid Orbitals • How about the geometry? • How does an sp orbital look like? Chapter 9

The sp hybrid orbitals in gaseous BeCl2. atomic orbitals hybrid orbitals orbital box diagrams

Hybrid Orbitals • Since the central atom, Be has two equal hybrideized orbitals at 1800 from each other and that each sp hybrid orbital has an unpaired electron, therefore, BeCl2 with a linear geometry could be formed. Chapter 9

Hybrid Orbitals • sp2 Hybrid Orbitals • Important: when we mix n atomic orbitals we must get n hybrid orbitals. • sp2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.) • The large lobes of sp2 hybrids lie in a trigonal plane. • All molecules with trigonal planar electron pair geometries have sp2 orbitals on the central atom. Chapter 9

Hybrid Orbitals • sp3 Hybrid Orbitals • sp3 Hybrid orbitals are formed from one s and three p orbitals. Therefore, there are four large lobes. • Each lobe points towards the vertex of a tetrahedron. • The angle between the large lobes is 109.5. • All molecules with tetrahedral electron pair geometries are sp3 hybridized. Chapter 9

sp2 and sp3 Hybrid Orbitals

Theories of Covalent Bonding