Lecture 9. 1. January 27, 2010 – Lecture 9. 10.6 Valence Bond Theory: Orbital
Overlap as a. Chemical Bond. 10.7 Valence Bond Theory: Hybridization of
Atomic.
Chem 1011
Dr. L. Dawe
Chem 1011
Dr. L. Dawe
Winter 2010
January 27, 2010 – Lecture 9 10.6 Valence Bond Theory: Orbital Overlap as a Chemical Bond 10.7 Valence Bond Theory: Hybridization of Atomic Orbitals – sp3 Hybridization – sp2 Hybridization and Double Bonds – sp Hybridization and Triple Bonds – sp3d and sp3d2 Hybridization – Writing Hybridization and Bond Schemes
Chem 1011
Dr. L. Dawe
Winter 2010
Lecture 9
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What A Bonding Theory Should Do
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Lewis theory is simple and structures can be determined rapidly. –
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It does not account for odd-electron species, resonance structures or the magnetic and spectral properties of molecules.
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•
VSEPR theory allows shape predictions
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Neither yield quantitative information about bond
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lengths or energies
Chem 1011
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Dr. L. Dawe
Winter 2010
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Introduction to Valence Bond Theory •
Valence-bond method: treats a covalent bond in terms of the overlap of pure or hybridized orbitals. Electron probability (or electron charge density) is concentrated in the area of overlap. –
This theory tells us what a covalent bond is and correlates molecular shapes to the interactions of atomic orbitals.
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The basic principle of valence bond theory is that a covalent bond forms when half filled orbitals on two different atoms (atomic orbitals) overlap. Example: H2
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Introduction to Valence Bond Theory
Lecture 9
___________________________________ http://cwx.prenhall.com/petrucci/medialib/medi a_portfolio/text_images/038_H2BondForm.MO V ___________________________________ ___________________________________ ___________________________________
Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Introduction to Valence Bond Theory
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Chem 1011
Dr. L. Dawe
Winter 2010
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Introduction to Valence Bond Theory •
Localized electron model: according to valence bond theory, core electrons and lone-pair electrons retain the same orbital locations as in the separated atoms. Charge density of the bonding electrons is concentrated in regions of orbital overlap. Example: Bonding in H2S.
•
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–
Note: (+) and (-) signs denote phase signs, not charges!
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Introduction to Valence Bond Theory
Lecture 9
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Using the Valence-Bond Method to Describe a Molecular Structure.
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Describe the phosphine molecule, PH3, by the valence-bond method..
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Step 1: Draw valence shell orbital diagrams for the separate atoms.
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Introduction to Valence Bond Theory
___________________________________ Step 2: Sketch the orbitals of the central atom (P) that are involved in the overlap.
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Recall:
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Introduction to Valence Bond Theory
___________________________________ Step 3: Complete the structure by bringing together the bonded atoms and representing the orbital overlap. Step 4: Describe the structure.
PH3 is trigonal-pyramidal.
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Three H atoms lie in the same plane. The P is situated at the top of the pyramid. The three H-P-H bond angles are 90o.
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(Experimentally, they are measured to be between 93 and 94o, vs VSEPR that predicts slightly less than 109.5o.)
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Introduction to Valence Bond Theory •
Note that this simple approach does not explain bonding in methane!
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals •
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Hybrid orbital properties – – –
The number of hybrid orbitals equals the total number of atomic orbitals that are combined. Hybridization rationalizes experimentally determined shape, it is not an actual physical phenomenon. Atomic orbital energy is conserved upon hybridization. •
Example: For tetrahedral C, the p orbitals each move down ¼ of the energy difference between the s and p orbitals, while the s orbitals move up by ¾.
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Hybridization of Atomic Orbitals
___________________________________ sp3 hybrid orbital: these are four orbitals formed by the hybridization of one s and three p orbitals. The angle between any two of the orbitals is the tetrahedral angle, 109.5o. –
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Think “four electron groups four hybrid orbitals”
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Hybridization of Atomic Orbitals
___________________________________ Bonding and structure of tetrahedral methane (CH4) – an sp3 hybridized molecule. Note that we are only considering the hybridization of the central atom’s orbitals! In the treatment in Chem 1011, we assume the terminal atoms are unhybridized.
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Hybridization of Atomic Orbitals
___________________________________ Bonding and structure of trigonal pyramidal methane (NH3) – an sp3 hybridized molecule.
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Note that hybrid orbitals can accommodate lone pair electrons as well as bonding pairs.
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals sp2 hybrid orbital: these are the three orbitals formed by the hybridization of one s and two p orbitals. The angle between any two of the orbitals is 120o. – –
This hybridization scheme is common to most boron containing compounds. Think “three electron groups three hybrid orbitals”
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
___________________________________ ___________________________________ Bonding and structure of trigonal planar BF3 – an sp2 hybridized molecule.
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals sp hybrid orbital: these are the two orbitals formed by the hybridization of one s and one p orbital. The angle between the two orbitals is 180o. – –
This hybridization scheme is common to most beryllium containing compounds. Think “two electron groups two hybrid orbitals”
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Hybridization of Atomic Orbitals
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals
___________________________________ ___________________________________ ___________________________________ Bonding and structure of linear BeCl2 – an sp hybridized molecule.
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
Lecture 9
___________________________________ http://cwx.prenhall.com/petrucci/medialib/medi a_portfolio/text_images/056_Hybridization.MO V ___________________________________ ___________________________________ ___________________________________
Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals sp3d hybrid orbital: these are five orbitals formed by the hybridization of one s, three p, and one d orbital. The five orbitals are directed to the corners of a trigonal bipyramid. – –
Example: Hybridization of phosphorus, P. Think “five electron groups five hybrid orbitals”
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Hybridization of Atomic Orbitals
___________________________________ Bonding and structure of trigonal bipyramidal PCl5 – an sp3d hybridized molecule. This hybridization scheme also accounts for the shapes of seesaw, tshaped and some linear molecules.
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals sp3d2 hybrid orbital: these are the six orbitals formed by the hybridization of one s, three p and two d orbitals. The six orbitals are directed to the corners of a regular octahedron. –
Think “six electron groups six hybrid orbitals”
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybridization of Atomic Orbitals Bonding and structure of octahdral SF6 – an sp3d2 hybridized molecule. This hybridization scheme also accounts for the shapes of square pyramidal and square planar molecules.
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Hybridization of Atomic Orbitals
___________________________________ Note: Five electron groups Five sp3d orbitals
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Note: Six electron groups Six sp3d2 orbitals
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Hybridization of Atomic Orbitals
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
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Hybrid Orbitals and VSEPR Theory
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VSEPR method uses empirical data to give an approximate molecular geometry, whereas the valence bond method relates to the orbitals used in bonding based on a given geometry.
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We can choose the likely hybridization scheme for a central atom in a structure in the valence-bond method by:
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1. 2. 3.
writing a plausible Lewis structure for the species of interest using VSEPR theory to predict the probable electron-group geometry of the central atom. selecting the hybridization scheme corresponding to the electron-group geometry.
Chem 1011
Dr. L. Dawe
Winter 2010
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Hybrid Orbitals and VSEPR Theory Problem: Predict the shape of the following molecules and a hybridization scheme consistent with this prediction (a) SiF4 (b) XeF4
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Problem: Describe the molecular geometry and propose a plausible hybridization scheme for the central atom in the ion: (a) Cl2F+ (b) BrF4+
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Chem 1011
Dr. L. Dawe
Lecture 9
Problem: Predict the shape of the following molecules and a hybridization scheme consistent with this prediction (a) SiF4 (b) XeF4
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Chem 1011
Dr. L. Dawe
Lecture 9
Problem: Describe the molecular geometry and propose a plausible hybridization scheme for the central atom in the ion: (a) Cl2F+ (b) BrF4+
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Chem 1011
Chem 1011
Dr. L. Dawe
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds sigma (σ) bonds: results from the end-to-end overlap of simple or hybridized atomic orbitals along the straight line joining the nuclei of the bonded atoms. pi (π) bonds: results from the side-to-side overlap of p orbitals, producing a high electron charge density above and below the line joining the bonded atoms.
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
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Chem 1011
Dr. L. Dawe
Winter 2010
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Multiple Covalent Bonds
___________________________________ C2H6 - Molecule should be tetrahedral (sp3 hybridized) about each C atom.
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σ bonding; end-to-end overlap of an hybridized orbital from each carbon
sp3
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
Lecture 9
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C2H6 - Molecule should be tetrahedral hybridized) about each C atom.
(sp3
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C2H4 - Molecule should be trigonal planar (sp2 hybridized) about each C atom.
Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
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Multiple Covalent Bonds
___________________________________ ___________________________________ Molecular shape is determined by the orbitals forming the σ-bonds (σframework). Rotation about the double bond is severely restricted, and the double bond is rigid.
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Twisting one of the –CH2 groups out of plane would reduce the amount of porbital overlap and weaken the π bond.
Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Multiple Covalent Bonds
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
Lecture 9
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C2H6 - Molecule should be tetrahedral hybridized) about each C atom.
(sp3
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C2H4 - Molecule should be trigonal planar (sp2 hybridized) about each C atom.
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C2H2 - Molecule should be linear (sp hybridized) about each C atom.
Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Multiple Covalent Bonds
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Multiple Covalent Bonds
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Chem 1011 Chem 1011
Dr. L. Dawe Dr. L. Dawe
Winter 2010
Multiple Covalent Bonds
Lecture 9
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Chem 1011
Dr. L. Dawe
Winter 2010
___________________________________
Multiple Covalent Bonds
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Problem: Describe the types of bonds and orbitals present for:
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(a) HCN (b) CO2
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Chem 1011
Dr. L. Dawe
Winter 2010
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Looking Ahead: January 29, 2010 – Lecture 10 11.2 Solids, Liquids, and Gases: A Molecular Comparison – Changes between phases 11.3 Intermolecular Forces: The Forces That Hold Condensed Phases Together – Dispersion Force – Dipole-Dipole Force – Hydrogen Bonding – Ion Dipole Force
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Chem 1011
Dr. L. Dawe
Lecture 9
Problem: Describe the types of bonds and orbitals present for: (a) HCN (b) CO2
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