General Chemistry - Molecular Geometry

Molecular Geometry     Biological Activity and Molecular Shape o There is no way to superimpose these two molecules. This is due to the chiral carbon o These molecules are enantiomers  Different biological functions o Bond angle – angle (in degrees) defined by lines joining the centers of two atoms to the center of a third atom to which they are covalently bonded To determine geometry o Draw Lewis structure and determine the steric number (SN) o Determine optimal spatial arrangement of electron pairs (bonding + nonbonding) to minimize repulsion Determine the molecular geometry of SO2 o VSEPR gives electron geometry o Molecular electron-pair geometry Molecular Geometry: SN = 4 o Bond angles decrease as number of lone pairs increases  Determine the molecular geometry of SF4 (SN=5)  Bond Polarity o Requirements for polar molecule:  Polar bonds (covalent bond between atoms w ∆EN)  Non-uniform distribution of polar bonds Predicting Polarity o Lewis Structure o VSEPR o Determine dipoles based on EN. Non-uniform distribution of electrons Magnets o A molecule needs to have an unpaired spin (electron) that can align with a magnetic field VSEPR (Valence Shell Electron Pair Repulsion) o Gives the predicted structure based only on electron pair repulsion o Does not explicitly explain how atoms share electrons in a covalent bond H2 – the Simplest Bonding Situation o How does this diagram relate to bonding?  1. Atoms need to be close, but not too close  2. Optimal distance gives the most stable energy values       3. The 1s orbitals on H2 overlap at the “best” distance Valence Bond Theory o Describes a covalent bond as the overlap of half-filled atomic orbitals (each  containing a sing electron) that form a bonding electron pair hared between two atoms Orbitals o Each p-orbital aligns with an axis in Cartesian coordinates o Combinations of orbitals form bonds, and relate to Lewis Structures o σ -bonds – electrons concentrated along bonding axis   π-bonds – electrons concentrated on opposite sides of bonding axis (nodal plane) The orbital wave functions mix to form hybrid orbitals Orbitals Mix – Valence Bond Theory  SN = 4 (H2O, NH3, CH4) sp3  sp2 Hybridization o pi (π) bond – a covalent bond in which electron density is greatest around-not along- the bonding axis Linear: sp Hybridization C2H2 o     sp orbitals = 25% s and 75% p Hybrid Orbitals Based on Steric Number Delocalization of Electron o Delocalization – spreading of electrons in alternating single and double bonds over three or more atoms in a molecule  This alternating pattern allows for overlap of the p orbitals and results in a “delocalization” of the electrons in the p orbitals, so that they can move about among all the C atoms in the molecules  This delocalization and planar configuration contributes to the observed biological activity of these molecules  Unpaired Electrons in O2 o A Gouy balance compares the mass of a sample in the presence of a magnetic field with the mass with the electromagnet turned off to determine the number of unpaired electrons in a sample ψ  i=c 1 i χ 1+c 2i χ 2 +…+ c mi χ m=∑ c μi χ μ μ Wave function (molecular orbital) wave function (atomic orbital)   Bone order = ½(# of bonding e- – # of antibonding e-) General Rule: stronger (more stable) than π bonds  VSEPR – structure/geometry depends on how electrons repel o Bonding not explicitly detailed Valence Bond – localized bonds from half-filled overlapping orbitals o Hybridization -> structure Molecular Orbital – delocalized bonds in molecular orbitals (combinations of atomic orbitals) o MO Diagrams -> bond order -> electron configuration  