Electrons in a bond are attracted to the nuclei at

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Electrons in a bond are attracted to the nuclei at either end, which tends to confine bonding electrons to a cylindrical region near the bond axis. I Lone pairs only have an atom at one end of the orbital, so they tend to spread sideways more. Consequences: 1. Lone pairs locate themselves where there is the most space. 2. Lone pairs push the bonding electrons away, distorting their geometry away from the ideal geometry for n identical electron groups.
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I To name geometries of molecules with lone pairs, first determine the electronic geometry as one of the basic VSEPR geometries. I Then look at the shape made by the atoms surrounding the central atom and name this shape. I The names of the shapes derived from the basic VSEPR geometries will be given in the following examples: NH 3 , H 2 O, SF 4 , ClF 3 , I - 3 , ClF 5 , XeF 4 , NO - 3
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Larger molecules I We can apply VSEPR theory to each non-terminal atom. I Sometimes, geometric constraints mean that we can’t obtain the “ideal” VSEPR geometry. Examples: C 2 H 4 , CH 3 NH 2 , CH 3 CN, cyclopropane (C 3 H 6 )
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Bond polarity revisited I A dipole is a pair of equal and opposite charges separated by a distance d . I The strength of a dipole is measured by the dipole moment , μ = qd I Polar bonds can be thought of as little dipole vectors. By convention in chemistry, these vectors point toward the negative (more electronegative) end of a bond. This is contrary to the convention in physics.
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Molecular polarity I The overall dipole moment of a molecule is the vector sum of the bond dipoles. I A molecule with a nonzero dipole moment is said to be polar . A molecule with a zero dipole moment is nonpolar . I The positive and negative ends of a molecule are often labeled with the symbols δ +, δ - rather than drawing dipole moment vectors. Examples: HCl, CO 2 , O 3 , BF 3 , CH 4 , PCl 5 , SF 6 , CH 3 Cl, NH 3 , H 2 O, SF 4 , ClF 3 , XeF 4
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