Book Notes Chapter 3

Book Notes Chapter 3 - Chapter 3 Chemical Bonds...

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Chapter 3 – Chemical Bonds Empirical-based on rules suggested by experimental observation rather than on fundamental principles. Bond angles-angles between bonds. Valence-shell electron-pair repulsion model-extend Lewis structure theory to account for molecular shapes by adding rules that account for bond angles. Electron arrangement-“most distant” locations of regions of electron concentration. Molecular shape considers only positions of atoms, not any lone pairs on central atoms. Multiple bonds are no different from single bonds when determining molecular shape. According to the VESPR model, regions of high electron concentration take up positions that maximize their separations; electron pairs in a multiple bond are treated as a single unit. The shape of the molecule is then identified from the relative locations of its atoms. VESPR formula- AX n E m ; A=central atom; X=attached atom; E=lone pairs on the central atom. Axial lone pair-lies on the axis of the molecule. Equatorial lone pair lies on the molecule’s equator. In a molecule that has lone pairs or a single nonbonding electron on the central atom, the valence electrons contribute to the electron arrangement about the central atom but only bonded atoms are considered in the identification of the shape. Lone pairs distort the shape of a molecule so as to reduce lone pair— bonding pair repulsions. Polar molecule-molecule with a nonzero dipole moment. Nonpolar molecule-molecule that has no electric dipole moment. Homonuclear diatomic molecules-molecules containing two atoms of only one element, which are always nonpolar (O 2 , F 2 , etc.). A diatomic molecule is polar if its bond is polar. A polyatomic molecule is polar if it has polar bonds arranged in space in such a way that the dipole moments associated with the bonds do not cancel. Valence-bond theory-description of covalent bonding to be devised in terms of atomic orbitals; quantum mechanical description of the distribution of electrons in bonds that goes beyond Lewis’s theory and VESPR models by providing a way of calculating the numerical values of bond angles and bond lengths. Sigma bond-cylindrically symmetrical, with no nodal planes containing the internuclear axis. Overlap of orbitals-merging of two atomic orbitals. All single covalent bonds are sigma bonds. Pi bond- a bond in which the two electrons lie in two lobes, one on each side of the internuclear axis (a pi bond has a single nodal plane containing the internuclear axis).
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A single bond is a σ-bond. A double bond is one σ-bond and one π-bond. A triple bond is one σ-bond plus two π-bonds.
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Book Notes Chapter 3 - Chapter 3 Chemical Bonds...

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