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CHEM 230 (5)
Lecture

Class Notes 2 Summer 2012.pdf

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Department
Chemistry
Course
CHEM 230
Professor
Jennifer Warriner
Semester
Summer

Description
Covalent Bonding and Molecular Structure 1. Introduction (a) Bonding in Inorganic Compounds • Electronegativitydifference • Bonding types: (i) Covalent (ii) Ionic (iii) Metallic (iv) Molecular (v) Hydrogen bonds • Bonding is usually mixed and competitive (b) Covalent Bonding • How do atoms combine to form molecules? • Molecular bonding and structures ⇒ (i) Evolution of Covalent Bonding Theory Valence Bond [VB] Theory ⇒ • Lewis (1916): H , O , N ⇒ 2 2 2 • Heitler & London (1927):Application of quantum mechanics 19 • Gillespie: Valence Shell Electron -Pair Repulsion • Pauling (1930’s): Hybridization of atomic orbitals Molecular Orbital [MO] Theory • Mulliken & Hund (1932): (ii) Formation of Covalent Bonds • Electron pair linking two nuclei ⇒ From Pauli Exclusion Principle • Potential Energy graph E(r) of electron interaction • When r ↓↓ ⇒ nucleus repulsive force ↑↑ • At equilibrium: r → r (o 0.74 Å in H ) 2 • Formation of molecules ⇒ - • For two e sf parallel spins⇒ repulsion 20 (iii) Bonding Parameters  Bond Energy E b • Energy needed to break a bond: • Eb= ∆H diss -∆H fstandard state, 1 mol) • E ↑ ⇒ stronger bonding ⇒ b Bond lengths L bnd Covalent Radii: • Interatomic (nucleus) distance L = r + r b A B • rA, B ⇒ covalent atomic radius; • Smaller L b High attractive energy⇒ N-N: L b 146 pm N=N: L = 125 pm b N≡N: L =b110 pm  Bond angle • Angle between the bonds ⇒ • H 2: CO :2  Bond Polarity • Polarity ⇒ Non-symmetric distribution of charges ⇒ 21 • Origin: Difference in electronegativity ∆χ = • Higher χ → (excess) –ve charges; Dipole moment: µ = q.e.d ⇒  Molecular Polarity • Diatomic molecules: Polar bond ⇒ polar molecules • Polyatomic molecules: • CO ,2C-O: • H 2, H-O: polar bonds at 104.5° • Polar molecules (SO , N2 ,…) 3 soluble in polar solvents • Homonuclear diatomic molecules: H , O , 2 , 2l ,2.. 2 • Mol. dipole moment determination ⇒ 22 (iv) Characteristics of Covalent Bonding • Formation of electron pair(s) with antiparallel spins • Saturation of bonding: • Directionality: Direction of maximum overlap of orbitals To minimize energy of the system⇒ 2. Modern Valence Bond Theory- • esremain localized between pairs of atoms or on a particular atom • Bonding pairs: • Lone pairs: (a) Lewis Structures • Atoms achieve a noble gas e-configuration • Only valence e are involved s • Lewis Diagrams: - Proper connection and arrangement of atoms - Distribution of valencee s • Rules for writing Lewis structures: - Sum the total valence e fros all atoms - Use a pair of e - s -Arrange the remaining e ⇒ s- - Central atom: O & H are often peripheralatoms 23 (i) Saturated systems: ⇒ Single bond diagrams: CH , NH , H O, HF, ….. 4 3 2 (ii) Unsaturated systems: Valencee not suffscient for single bond • Formation of multiple bonds to satisfy the octet rule - • NO , 3 (iii) Exceptions to the Octet rule: Electron deficient molecules: nd • 2 row atoms Be & B: • High reactivity with molecules having lone pair e s- - • Valence shell expansion: > 8 e aroundsan atom • 3 row and heavier atoms: PCl . SbF , BrF 5 6 3 • Extra electrons to empty d-orbitals • Third row and heavier elements 24 • Resonance structures: Molecular Structure: how are atoms arranged in a molecule? • How are the molecular bonds arranged in 3-D? • What are the bond angles? ⇒Molecular structures⇒ character and usefulness 1. Valence Shell Electron Repulsion (VSEPR) Theory 2. Orbital Hybrid Approach (b) Valence Shell Electron Pair Repulsion (VSEPR) Theory (i) Electron Pair Repulsion - • Electron pairs: Concentration of e density ⇒ Repulsions between the e pairs ⇒ (ii) Stability of Molecular Structures • To minimize the electron-pair repulsions ⇒ • Both bonding and non-bonding pairs ⇒ ⇒ 25 (iii) Applications of VSEPR • Space occupancy [x + y] of the atomA in a moleculeAB E x y • Single bonds: [x + y] = No. of electron pairs • Multiple bonds? ⇒ • Geometry of the e-pairs ⇒ • Name of the structures⇒ - Linear Geometry: AB ; ABE 2 - • BeCl ;2CN - Trigonal Planar Geometry: AB ; AB E; 3BE 2 2 • BF , [NO ] , [CO ] : 2- 3 3 3 - • AB E2 [NO ] :2Trigonal geometry ⇒ • Angle ∠ONO = 115° < 120° ? • Lone pair e attached to a single nucleus⇒ s ⇒ Require more room nearA than bonding pairs ⇒ More repulsive to the bonding pairs ⇒ Bonding pairs squeeze together: 26 + - • Comparison between: [NO ] , N2 , [NO 2 2 - Tetragonal Geometry:AB ; AB E; AB E 4 3 2 2 • AB : CF , Tetrahedral geometry= 4 4 • AB E3 NH : T3trahedral geometry with one lone pair • Actual structure⇒ • Angle ∠HNH = 107° ⇒ • •AB E2 2
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