1.4 chemical bonding_3 - UNIVERSITY OF ALBERTA INTRODUCTORY...

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1 UNIVERSITY OF ALBERTA INTRODUCTORY UNIVERSITY CHEMISTRY I (CHEM 101) Theories of Covalent Bonding I Introduction The VSEPR is only good for predicting molecular shapes It does not provide the roles of the atomic orbitals in bond formation. It does not provide explanations for bond energy, magnetic and spectral properties in covalent compounds. There are two theories for covalent bond formation; 1. Valence Bond (VB) Theory which focuses on overlapping of atomic orbitals 2. Molecular Orbital (MO) Theory which focuses on orbital energy levels. It assigns electrons to a series of molecular orbitals which belong to the whole molecule. The two theories compliment each other for a full picture of covalent bond formation II Valence Bond Theory Introduced in the 1930s, the VB Theory provides a mostly qualitative approach to covalent bonding. There are three main themes for the VB Theory: 1. The sharing of electrons is only possible when the two atomic orbitals overlap H H H H + 1s 1s 1s 1s F F 2p 2p + F F 2p 2p F F F F F H 1s 2p H F + 1s 2p F HF Please note that the orbitals must be of the same phase for overlapping. The more effective is the overlapping, the stronger is the bond
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2 2. For covalent bond formation, the electron must be of opposite spin . 3. For effective overlapping, the atomic orbitals hybridize. Consider the formation of methane (CH 4 ) Electron configuration of carbon is: 1s 2 2s 2 2px 1 2py 1 The orbital box diagram of the valence electrons is: C 2s 2p Only two valence electrons are available to form bonds with hydrogen which is inconsistent with the molecular formula of CH 4 . However, one of the 2s electrons can be excited to an empty 2p orbital and thus four electrons are now available. The process of atomic orbital mixing is called hybridization and the new atomic orbitals are called hybrid orbitals . In this case with carbon, it forms four equivalent hybrid orbitals (sp 3 orbitals) sp 3 signifies one s orbital and three p orbitals are combined Fig 11.4, p414 Four sp 3 orbitals give rise to a tetrahedron (109.5 o ) (consistent with four electron groups from VSEPR) Four single bonds ( σ bonds): sp 3 (C) overlaps 1s (H)
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3 Hybridization The number of hybrid orbitals is equal to the number of participating atomic orbitals The type of the hybrid orbital is dependent on the type of participating atomic orbitals. For example: sp, sp 2 , sp 3 etc Formation of ammonia (NH 3 ) One of the sp 3 hybridized orbital contains a lone pair of electron. Fig. 11.5, p415 Other sp 3 hybridized compounds: NH 4 + , H 2 O sp 2 hybridization Formation of BF 3 B 2s 2p Excitation B 2s 2p Hybridization 2p sp 2 B sp 2 sp 2 This is sp 2 hybridization with an empty p orbital that does not hybridize Structure of BF 3 + Hybridization + Three single bonds ( σ bonds): sp 3 (N) overlaps 1s (H) One lone-pair (N)
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4 B Empty 2p orbital Empty 2p orbital F F F B Empty 2p orbital Empty 2p orbital F F F Three σ bonds: sp 2 (B) overlaps with 2p (F)
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1.4 chemical bonding_3 - UNIVERSITY OF ALBERTA INTRODUCTORY...

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