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Unit 06: Molecular Structure & Bonding Theories Section B: Molecular Shapes, Polarity, Molecular Orbital Theory Starring: Janice and Laura Formal Charge The charge calculated for that atom based on the Lewis structure of the molecule or ion using the equation... Formal Charge = [original charge of atom or ion] [number of lone pair electrons] + 1/2[number of bonding electrons] Bond Polarity Unless pure covalent bonding occurs, in which atoms share an electron pair equally, the electron pair will be unequally shared and result in a polar covalent bond. But why would the atoms not share? A. Just something they didn’t learn in Kindergarten. B. Because not all atoms hold onto their valence electrons with the same force. C. Just keep clicking the mouse If you said B, you were right! Depending on the atom and the number of valence electrons, the strength at which electrons can be kept or pulled away differs. Does this sound familiar? It should, it is called elecronegativity! But since it isn’t in our section of chapter 6 we will move on. Molecular Polarity In a polar molecule, electron density accumulates toward one side of the molecule depending on the elecronegativity of each atom. The result is that one part of the molecule becomes more negative. If these forces cancel out, then the molecule can be neutral. Whether or not the forces cancel is usually dependant on their shapes. Valence shell electron-pair repulsion (VSEPR) Is a method used to predict the shapes of covalent molecules and poly atomic ions. It works by assuming that electrons in bond pairs and lone pairs will repel each other and try to stay as far apart from each other as possible. Generally not used for transition metals. Valence electron theory The idea that bonds are formed by the overlapping of orbitals Hybridization of Orbitals The first bond on a central atom is a s orbital The subsequent three bonds, or electron lone pairs are p orbitals Any other orbitals are designated as d orbitals Basic Molecular Shapes Orbitals Hybridized as sp When there are two bond pairs a linear shape observed, and the orbitals are in the same plane and located 180° away from each other. For all basic molecular shapes, if the atoms around the central atom are the same, or have the same electronegativity, then the molecule will be non-polar More Basic Molecular Shapes When there are three bond pairs a trigonal planar or triangle shape is observed. The orbitals are still all in the same plane, but are now located 120° away from each other. Orbitals Hybridized as sp2 More Basic Molecular Shapes When there are three bond pairs a tetrahedral shape is observed. Because of repulsion factors, the trigonal planar shape is bent away from the new orbital and between each orbital there now exits 109.5° of separation. Orbitals Hybridized as sp3 More Basic Molecular Shapes When there are five bond pairs a Trigonal Bypyramidal shape is formed with three orbitals existing on the same plane in a triangle shape and a pair of orbitals in a linear formation perpendicular to the triangle. The three orbitals forming the base of the trigonal bypyramidal shape are separated by 120° and the two orbitals existing perpendicular are 90° away from the orbitals of the triangle shape. Orbitals Hybridized as sp3d The Last Basic Molecular Shape Slide When there are six bond pairs an octahedral shape is formed with four orbitals existing on the same plane, separated by 90°, and a pair of orbitals in a linear formation perpendicular to the plane. These orbitals are located 90° away from the plane. Orbitals Hybridized as sp3d2 Now we add electron pairs... The basic shapes we just learned about, deal with only bond pairs. However, many molecules have free electron pairs that also effect repulsion. In fact electrons pairs have a greater repulsion factor than regular bond pairs do. For Molecular shapes that have been altered by electron pairs, there is no chance for cancellation, so all of the following will be polar Basic Shape was Trigonal Planar Two bond pair, one electron pair Basic Shape: Tetrahedral Three bond pairs, one electron pair Two bond pairs, two electron pairs Basic Shape was Trigonal Bypyramidal Four bond pairs, one electron pair Three bond pairs, two electron pairs Two bond pairs, three electron pairs Basic Shape was Octahedron Five bond pairs, one electron pair Four bond pairs, two electron pairs Molecular Orbital Theory An alternative way to view orbitals in molecules using the valence electron hybridized orbitals Specifies that when 1s orbitals of two hydrogen atoms overlap, two molecular orbitals result from the addition and subtraction of other overlapping orbitals. Deals with π bonds and σ bonds and their antitheses π* bonds and σ* bonds. Types of Bonds Must have both π bonds and σ bonds to form a complete σ bonds: molecules – higher energy than π bonds – exist in the x and y direction – exist where there are single bonds in the Lewis structure π bonds: – lower in energy than σ bonds – exist in the z direction – exist where there are double bonds in the Lewis structure Antibonds Atomic orbitals form in pairs One orbital carries the electron dense materials and are the π bonds and σ bonds One orbital does not carry electron materials, and these are the anti π bonds and σ bonds (use a * to denote) In order for orbitals of one atom to form bonds with orbitals of another, the number of antibonds (both π and σ) must not cancel out. First Principal of Molecular Orbital Theory Molecular orbitals are the total valence atomic orbitals of all atoms in the molecule Second Principle The bonding molecular orbital is lower in energy than the parent valence orbital Antibonds are higher in energy Third Principle Electrons are assigned to orbitals of successively higher energy Fourth Principle Atomic orbitals combine to form molecular orbitals most effectively when atomic orbitals are of similar energy Homonuclear diatomic molecules Diagram for second period The End The End The End The End The End The End The End The End