DOCSLIB.ORG

Bonding and Molecular Structure: Orbital Hybridization and Molecular

Total Page：16

File Type：pdf, Size：1020Kb

Download full-text PDF Read full-text

Load more

Page III-8-1 / Chapter Eight Lecture Notes Bonding and Advanced Theories of Molecular Chemical Bonding Structure: Orbital Hybridization and Molecular Orbitals Chapter 8 Chemistry 222 Atomic Orbitals Molecules Professor Michael Russell Last update: MAR 8/9/21 MAR Two Theories of Bonding Two Theories of Bonding VALENCE BOND THEORY - Linus MOLECULAR ORBITAL Pauling THEORY - Robert Mulliken (1896-1986) valence electrons are localized between atoms (or are lone valence electrons are pairs). delocalized half-filled atomic orbitals overlap valence electrons are in to form bonds. orbitals (called molecular orbitals) electrons stabilized by 2 nuclei spread over entire molecule. MAR MAR Sigma Bond Formation by Sigma Bond Formation by Orbital Overlap Orbital Overlap Two s Two s orbitals orbitals overlap overlap One s & one p overlap Two p orbitals overlap MAR MAR Page III-8-1 / Chapter Eight Lecture Notes Page III-8-2 / Chapter Eight Lecture Notes Using Valence Bond Theory Bonding in BF3 Bonding in BF3 •• How to account for 3 bonds 120o apart • F • using a spherical s orbital and p orbitals Boron configuration that are 90o apart? B ↑↓ ↑↓ ↑ Pauling said to modify VB approach with •• •• • F F• •• •• • 1s 2s 2p ORBITAL HYBRIDIZATION planar triangle - mix available orbitals to form a new angle = 120o set of orbitals - HYBRID ORBITALS - that will give the maximum overlap in the correct geometry. MAR MAR Bonding in BF3 Bonding in BF3 The three hybrid orbitals are made from 1 s 2s 2p orbital and 2 p orbitals create 3 sp2 hybrids. hybridize orbs. rearrange electrons three sp 2 unused p orbital hybrid orbitals Now we have 3, half-filled HYBRID orbitals that can be used to form planar B-F sigma bonds. MAR MAR Bonding in CH4 Bonding in BF3 ↑↓ ↑ How do we account 2s 2p hydridize orbs. rearrange electrons for 4 C-H sigma ↑ ↑ ↑ o . bonds 109 apart? three sp 2 unused p orbital hybrid orbitals Need to use 4 atomic 109o An orbital from each F overlaps one of the sp2 orbitals - s, px, py, hybrids to form a B-F σ bond. and pz - to form 4 F new hybrid orbitals ↑ ↑ pointing in the F ↑↑ B correct direction. F MAR MAR Page III-8-2 / Chapter Eight Lecture Notes Page III-8-3 / Chapter Eight Lecture Notes Bonding in a Tetrahedron - Bonding in a Tetrahedron - Formation of Hybrid Atomic Formation of Hybrid Atomic Orbitals Orbitals 4 C atom orbitals 4 C atom orbitals hybridize to form hybridize to four equivalent sp3 form four hybrid atomic equivalent sp3 orbitals. hybrid atomic orbitals. MAR MAR Orbital Hybridization Bonding in CH4 Bonds EPG Hybrid REMAINING p orbs? 2 linear sp 2 p 3 trigonal sp2 1 p planar 4 tetrahedral sp3 none 5 trigonal sp3d --- bipyramid 6 octahedral sp3d2 --- see: VSEPR Guide MAR MAR Hybrid Examples Bonding in Glycine 3 sp 2 H O sp •• H N C C •• H H O H sp3 •• start sp3 finish MAR MAR Page III-8-3 / Chapter Eight Lecture Notes Page III-8-4 / Chapter Eight Lecture Notes Multiple Bonds Sigma Bonds in C2H4 Consider ethylene, C2H4 H H H H 2 2 120˚ C C sp 120˚ C C sp H H H H MAR MAR π Bonding in C2H4 π Bonding in C2H4 The unused p orbital on The unused p orbital on each C atom each C atom contains an contains an electron, and this p orbital electron and this p orbital overlaps the p orbital on the neighboring overlaps the p orbital on atom to form the π bond. the neighboring atom to form the π bond. ↑↓ ↑ ↑ ↑ ↑ ↑ ↑ 2s 2p 3 sp 2 p orb. hybrid ↑↓ ↑ ↑ ↑ ↑ ↑ for π ↑ orbitals bond 2s 2p 3 sp 2 p orb. hybrid for π orbitals bond MAR MAR Multiple Bonding in C2H4 σ and π Bonding in C2H2 C2H2 has a triple bond MAR MAR Page III-8-4 / Chapter Eight Lecture Notes Page III-8-5 / Chapter Eight Lecture Notes Consequences of Multiple Consequences of Multiple Bonding Bonding There is restricted rotation around C=C bond. Restricted rotation around C=C bond. trans-but-2-ene MAR MAR n-butane Advantages of MO Theory Molecular Orbital Theory Dioxygen should be electron paired (diamagnetic) by VB • Accounts for paramagnetism, Theory, but dioxygen is color, bonding actually paramagnetic. • Atomic orbitals delocalize into MO Theory accounts for molecular orbitals paramagnetism of O2 • Bonding, Antibonding and Nonbonding orbitals • Quite complicated, need computers; we will only look at diatomics (2 atom systems) from the first and second periods only MAR MAR Four Principles of MO Theory Molecular Orbital Type Principle #1: When two atomic 1s H orbitals combine, * Number of Molecular Orbitals = a bonding (σ) and antibonding (σ ) Number of Atomic Orbitals molecular orbital forms antibonding σ* Two 1s orbitals from two hydrogen atoms create two molecular orbitals in H2 Two 1s orbitals and two 2s orbitals from two lithium atoms create four molecular orbitals in Li2 See Four Principles of MO Handout MAR MAR bonding σ Page III-8-5 / Chapter Eight Lecture Notes Page III-8-6 / Chapter Eight Lecture Notes Four Principles of MO Theory MO Diagram for H2 Principle #2: Bonding MO lower in energy than the Antibonding MO parent orbital Antibonding MO higher in energy Atomic Orbitals than the parent orbital Principle #3: Bonding MO Electrons of molecule assigned to successively higher MOs Two 1s electrons from two H atoms occupy the σ orbital in H2 Use Pauli Exclusion Principle and Hund's Rule MAR when assigning electrons MAR MO Diagram for He2 Bond Order in MO Theory 1 - - Bond Order = /2 (# bonding e - # antibonding e ) Antibonding MO Bond Order > 0, stable molecule Bond Order = 0 or < 0, Atomic Orbitals unstable molecule Bonding MO In H2, Bond Order = 1 Two 1s electrons in σ, /2 (2 - 0) = 1; Two 1s electrons in σ* stable MAR MAR Bond Order in MO Theory Four Principles of MO Theory Principle #4: 1 - - Atomic orbitals combine to give Bond Order = /2 (# bonding e - # antibonding e ) molecular orbitals only when the Bond Order > 0, atomic orbitals are of similar energy stable molecule Bond Order = 0 or < 0, unstable molecule Similar energy = better overlap 1s + 1s = good MO In He2, 1s + 2s = poor MO Bond Order = 2s + 2s = good MO 1 /2 (2 - 2) = 0; 2s + 2p = poor MO unstable 3s + 2s = poor MO ... etc. ... ∴He2 does not exist MAR MAR Page III-8-6 / Chapter Eight Lecture Notes Page III-8-7 / Chapter Eight Lecture Notes p orbitals and π bonds Example: Dilithium, Li2 Three possible p orbitals on each atom - six total Note: no overlap between p MO orbitals 1s and 2s Two p orbitals create 2 σ MO bonds Four remaining p orbitals create 4 MO bonds 1 π Bond Order = /2 (4 - 2) = 1 Stable molecule Four p atomic orbitals create four π molecular orbitals, π = bonding (2) π* = antibonding (2) Would you expect Be2 to exist? Why? ... but there's a catch! MAR MAR p orbitals and π bonds p orbitals and π bonds For O, F and Ne, For B, C and N, σ orbital lower energy than π orbitals π orbitals lower energy than σ orbital π* orbitals lower energy than σ* orbital π* orbitals lower energy than σ* orbital σ*2pz π∗2px π∗2py 2px 2py 2pz 2px 2py 2pz π2px π2py Example: B2 σ2pz σ∗2s Bond Order = 1 2s 2s σ2s σ∗1s 1s 1s σ1s MAR See MO Diagram (B2 - N2) Handout MAR p orbitals and π bonds Paramagnetism For O, F and Ne, σ orbital lower energy than π orbitals Paramagnetism exists when unpaired π* orbitals lower energy than σ* orbital electrons in MO diagram O2 is paramagnetic; Example: O2 unpaired electrons in two π* orbitals Bond Order = 2 See MO Diagram (O2 - Ne2) Handout MAR MAR Page III-8-7 / Chapter Eight Lecture Notes Page III-8-8 / Chapter Eight Lecture Notes Paramagnetism Molecular Orbital Notation Used to abbreviate the MO diagrams Paramagnetism exists when unpaired • Ignore core electrons electrons in MO diagram • Write in order of increasing energy N2 is diamagnetic; all electrons paired For N2: 2 * 2 4 2 [core electrons](σ2s) (σ 2s) (π2p) (σ2p) MAR MAR Sigma and Pi Bonds MO Diagram for Diatomics Determine sigma and pi bonds using: 1 - - # σ bonds = /2 (# σ bonding e - # σ antibonding e ) 1 - - # π bonds = /2 (# π bonding e - # π antibonding e ) and # σ bonds + # π bonds = bond order 2 * 2 4 2 For N2: [core electrons](σ2s) (σ 2s) (π2p) (σ2p) 1 # σ bonds = /2(4 - 2) = 1 σ bond 1 # π bonds = /2(4 - 0) = 2 π bonds Changes in MO diagrams due to bond order = 1/ (8 - 2) = 3 = 1 + 2 bonds MAR 2 σ π MAR s-p mixing and/or electron repulsion Ionic Diatomic Molecules Ionic Diatomic Molecules Predicting Ionic Diatomic MO diagrams simple Predicting Ionic Diatomic MO diagrams simple Use Hund and Pauli Use Hund and Pauli + - Example: O2 Example: O2 Remove electron from Using the O2 diagram on * π 2p orbital the right, where would you place Check bond order, the extra electron? - paramagnetism Is O2 more or less stable than O2? Why? MAR MAR Page III-8-8 / Chapter Eight Lecture Notes Page III-8-9 / Chapter Eight Lecture Notes Application: Vision Application: Band Theory Molecular Orbital Theory helps to In metallic bonding, electrons delocalized describe the process of vision - over metallic lattice - a sea of electrons photochemistry MO energies identical, excellent overlap Helps explain conductivity, malleability, more MAR MAR Important Equations, Constants, and Handouts End of Chapter 8 from this Chapter: • the bond order, bond Valence Bond / Hybridization energy and bond length Theory: types of hybridization relationships still apply to (sp, sp2, etc.), sigma and pi See: both theories bonds MO Diagram for Mo(CO)5=CH2 • Chapter Eight Study Guide • know the advantages and disadvantages of the • Chapter Eight Concept Guide Valence Bond and Molecular Molecular Orbital Theory: • Important Equations (following Orbital theories bonding and antibonding this slide) • see the Geometry and orbitals, sigma bonds and pi Polarity Guide and the two bonds, paramagnetic and • End of Chapter Problems Molecular Orbital Theory diamagnetic, the “NBC” vs.
Recommended publications
  • Direct Preparation of Some Organolithium Compounds from Lithium and RX Compounds Katashi Oita Iowa State College

    Direct Preparation of Some Organolithium Compounds from Lithium and RX Compounds Katashi Oita Iowa State College

    Iowa State University Capstones, Theses and Retrospective Theses and Dissertations Dissertations 1955 Direct preparation of some organolithium compounds from lithium and RX compounds Katashi Oita Iowa State College Follow this and additional works at: https://lib.dr.iastate.edu/rtd Part of the Organic Chemistry Commons Recommended Citation Oita, Katashi, "Direct preparation of some organolithium compounds from lithium and RX compounds " (1955). Retrospective Theses and Dissertations. 14262. https://lib.dr.iastate.edu/rtd/14262 This Dissertation is brought to you for free and open access by the Iowa State University Capstones, Theses and Dissertations at Iowa State University Digital Repository. It has been accepted for inclusion in Retrospective Theses and Dissertations by an authorized administrator of Iowa State University Digital Repository. For more information, please contact [email protected]. INFORMATION TO USERS This manuscript has been reproduced from the microfilm master. UMI films the text directly from the original or copy submitted. Thus, some thesis and dissertation copies are in typewriter face, while others may be from any type of computer printer. The quality of this reproduction is dependent upon the quality of the copy submitted. Broken or indistinct print, colored or poor quality illustrations and photographs, print bleedthrough, substandard margins, and improper alignment can adversely affect reproduction. In the unlikely event that the author did not send UMI a complete manuscript and there are missing pages, these will be noted. Also, if unauthorized copyright material had to be removed, a note will indicate the deletion. Oversize materials (e.g., maps, drawings, charts) are reproduced by sectioning the original, beginning at the upper left-hand comer and continuing from left to right in equal sections with small overiaps.
  • Mineral Processing

    Mineral Processing

    Mineral Processing Foundations of theory and practice of minerallurgy 1st English edition JAN DRZYMALA, C. Eng., Ph.D., D.Sc. Member of the Polish Mineral Processing Society Wroclaw University of Technology 2007 Translation: J. Drzymala, A. Swatek Reviewer: A. Luszczkiewicz Published as supplied by the author ©Copyright by Jan Drzymala, Wroclaw 2007 Computer typesetting: Danuta Szyszka Cover design: Danuta Szyszka Cover photo: Sebastian Bożek Oficyna Wydawnicza Politechniki Wrocławskiej Wybrzeze Wyspianskiego 27 50-370 Wroclaw Any part of this publication can be used in any form by any means provided that the usage is acknowledged by the citation: Drzymala, J., Mineral Processing, Foundations of theory and practice of minerallurgy, Oficyna Wydawnicza PWr., 2007, www.ig.pwr.wroc.pl/minproc ISBN 978-83-7493-362-9 Contents Introduction ....................................................................................................................9 Part I Introduction to mineral processing .....................................................................13 1. From the Big Bang to mineral processing................................................................14 1.1. The formation of matter ...................................................................................14 1.2. Elementary particles.........................................................................................16 1.3. Molecules .........................................................................................................18 1.4. Solids................................................................................................................19
  • Molecular Orbital Theory to Predict Bond Order • to Apply Molecular Orbital Theory to the Diatomic Homonuclear Molecule from the Elements in the Second Period

    Molecular Orbital Theory to Predict Bond Order • to Apply Molecular Orbital Theory to the Diatomic Homonuclear Molecule from the Elements in the Second Period

    Skills to Develop • To use molecular orbital theory to predict bond order • To apply Molecular Orbital Theory to the diatomic homonuclear molecule from the elements in the second period. None of the approaches we have described so far can adequately explain why some compounds are colored and others are not, why some substances with unpaired electrons are stable, and why others are effective semiconductors. These approaches also cannot describe the nature of resonance. Such limitations led to the development of a new approach to bonding in which electrons are not viewed as being localized between the nuclei of bonded atoms but are instead delocalized throughout the entire molecule. Just as with the valence bond theory, the approach we are about to discuss is based on a quantum mechanical model. Previously, we described the electrons in isolated atoms as having certain spatial distributions, called orbitals, each with a particular orbital energy. Just as the positions and energies of electrons in atoms can be described in terms of atomic orbitals (AOs), the positions and energies of electrons in molecules can be described in terms of molecular orbitals (MOs) A particular spatial distribution of electrons in a molecule that is associated with a particular orbital energy.—a spatial distribution of electrons in a molecule that is associated with a particular orbital energy. As the name suggests, molecular orbitals are not localized on a single atom but extend over the entire molecule. Consequently, the molecular orbital approach, called molecular orbital theory is a delocalized approach to bonding. Although the molecular orbital theory is computationally demanding, the principles on which it is based are similar to those we used to determine electron configurations for atoms.
  • Introduction to Molecular Orbital Theory

    Introduction to Molecular Orbital Theory

    Chapter 2: Molecular Structure and Bonding Bonding Theories 1. VSEPR Theory 2. Valence Bond theory (with hybridization) 3. Molecular Orbital Theory ( with molecualr orbitals) To date, we have looked at three different theories of molecular boning. They are the VSEPR Theory (with Lewis Dot Structures), the Valence Bond theory (with hybridization) and Molecular Orbital Theory. A good theory should predict physical and chemical properties of the molecule such as shape, bond energy, bond length, and bond angles.Because arguments based on atomic orbitals focus on the bonds formed between valence electrons on an atom, they are often said to involve a valence-bond theory. The valence-bond model can't adequately explain the fact that some molecules contains two equivalent bonds with a bond order between that of a single bond and a double bond. The best it can do is suggest that these molecules are mixtures, or hybrids, of the two Lewis structures that can be written for these molecules. This problem, and many others, can be overcome by using a more sophisticated model of bonding based on molecular orbitals. Molecular orbital theory is more powerful than valence-bond theory because the orbitals reflect the geometry of the molecule to which they are applied. But this power carries a significant cost in terms of the ease with which the model can be visualized. One model does not describe all the properties of molecular bonds. Each model desribes a set of properties better than the others. The final test for any theory is experimental data. Introduction to Molecular Orbital Theory The Molecular Orbital Theory does a good job of predicting elctronic spectra and paramagnetism, when VSEPR and the V-B Theories don't.
  • 8.3 Bonding Theories >

    8.3 Bonding Theories >

    8.3 Bonding Theories > Chapter 8 Covalent Bonding 8.1 Molecular Compounds 8.2 The Nature of Covalent Bonding 8.3 Bonding Theories 8.4 Polar Bonds and Molecules 1 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 8.3 Bonding Theories > Molecular Orbitals Molecular Orbitals How are atomic and molecular orbitals related? 2 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 8.3 Bonding Theories > Molecular Orbitals • The model you have been using for covalent bonding assumes the orbitals are those of the individual atoms. • There is a quantum mechanical model of bonding, however, that describes the electrons in molecules using orbitals that exist only for groupings of atoms. 3 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 8.3 Bonding Theories > Molecular Orbitals • When two atoms combine, this model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule. 4 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 8.3 Bonding Theories > Molecular Orbitals Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. • A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital. 5 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 8.3 Bonding Theories > Molecular Orbitals Sigma Bonds When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma bond is formed. • Its symbol is the Greek letter sigma (σ).
  • Molecular Orbital Theory

    Molecular Orbital Theory

    Molecular Orbital Theory The Lewis Structure approach provides an extremely simple method for determining the electronic structure of many molecules. It is a bit simplistic, however, and does have trouble predicting structures for a few molecules. Nevertheless, it gives a reasonable structure for many molecules and its simplicity to use makes it a very useful tool for chemists. A more general, but slightly more complicated approach is the Molecular Orbital Theory. This theory builds on the electron wave functions of Quantum Mechanics to describe chemical bonding. To understand MO Theory let's first review constructive and destructive interference of standing waves starting with the full constructive and destructive interference that occurs when standing waves overlap completely. When standing waves only partially overlap we get partial constructive and destructive interference. To see how we use these concepts in Molecular Orbital Theory, let's start with H2, the simplest of all molecules. The 1s orbitals of the H-atom are standing waves of the electron wavefunction. In Molecular Orbital Theory we view the bonding of the two H-atoms as partial constructive interference between standing wavefunctions of the 1s orbitals. The energy of the H2 molecule with the two electrons in the bonding orbital is lower by 435 kJ/mole than the combined energy of the two separate H-atoms. On the other hand, the energy of the H2 molecule with two electrons in the antibonding orbital is higher than two separate H-atoms. To show the relative energies we use diagrams like this: In the H2 molecule, the bonding and anti-bonding orbitals are called sigma orbitals (σ).
  • Paul B. Tomascak Tomáš Magna Ralf Dohmen Advances in Lithium Isotope Geochemistry Advances in Isotope Geochemistry

    Paul B. Tomascak Tomáš Magna Ralf Dohmen Advances in Lithium Isotope Geochemistry Advances in Isotope Geochemistry

    Advances in Isotope Geochemistry Paul B. Tomascak Tomáš Magna Ralf Dohmen Advances in Lithium Isotope Geochemistry Advances in Isotope Geochemistry Series editor Jochen Hoefs, Göttingen, Germany More information about this series at http://www.springer.com/series/8152 Paul B. Tomascak • Tomáš Magna Ralf Dohmen Advances in Lithium Isotope Geochemistry 123 Paul B. Tomascak Ralf Dohmen State University of New York at Oswego Institute of Geology, Mineralogy Oswego, NY and Geophysics USA Ruhr-University Bochum Bochum, Nordrhein-Westfalen Tomáš Magna Germany Czech Geological Survey Prague Czech Republic ISSN 2364-5105 ISSN 2364-5113 (electronic) Advances in Isotope Geochemistry ISBN 978-3-319-01429-6 ISBN 978-3-319-01430-2 (eBook) DOI 10.1007/978-3-319-01430-2 Library of Congress Control Number: 2015953806 Springer Cham Heidelberg New York Dordrecht London © Springer International Publishing Switzerland 2016 This work is subject to copyright. All rights are reserved by the Publisher, whether the whole or part of the material is concerned, specifically the rights of translation, reprinting, reuse of illustrations, recitation, broadcasting, reproduction on microfilms or in any other physical way, and transmission or information storage and retrieval, electronic adaptation, computer software, or by similar or dissimilar methodology now known or hereafter developed. The use of general descriptive names, registered names, trademarks, service marks, etc. in this publication does not imply, even in the absence of a specific statement, that such names are exempt from the relevant protective laws and regulations and therefore free for general use. The publisher, the authors and the editors are safe to assume that the advice and information in this book are believed to be true and accurate at the date of publication.
  • Inorganic Chemistry/Chemical Bonding/MO Diagram 1

    Inorganic Chemistry/Chemical Bonding/MO Diagram 1

    Inorganic Chemistry/Chemical Bonding/MO Diagram 1 Inorganic Chemistry/Chemical Bonding/MO Diagram A molecular orbital diagram or MO diagram for short is a qualitative descriptive tool explaining chemical bonding in molecules in terms of molecular orbital theory in general and the Linear combination of atomic orbitals molecular orbital method (LCAO method) in particular [1] [2] [3] . This tool is very well suited for simple diatomic molecules such as dihydrogen, dioxygen and carbon monoxide but becomes more complex when discussing polynuclear molecules such as methane. It explains why some molecules exist and not others, how strong bonds are, and what electronic transitions take place. Dihydrogen MO diagram The smallest molecule, hydrogen gas exists as dihydrogen (H-H) with a single covalent bond between two hydrogen atoms. As each hydrogen atom has a single 1s atomic orbital for its electron, the bond forms by overlap of these two atomic orbitals. In figure 1 the two atomic orbitals are depicted on the left and on the right. The vertical axis always represents the orbital energies. The atomic orbital energy correlates with electronegativity as a more electronegative atom holds an electron more tightly thus lowering its energy. MO treatment is only valid when the atomic orbitals have comparable energy; when they differ greatly the mode of bonding becomes ionic. Each orbital is singly occupied with the up and down arrows representing an electron. The two AO's can overlap in two ways depending on their phase relationship. The phase of an orbital is a direct consequence of the oscillating, wave-like properties of electrons.
  • Covalent Bonding and Molecular Orbitals

    Covalent Bonding and Molecular Orbitals

    Covalent Bonding and Molecular Orbitals Chemistry 35 Fall 2000 From Atoms to Molecules: The Covalent Bond n So, what happens to e- in atomic orbitals when two atoms approach and form a covalent bond? Mathematically: -let’s look at the formation of a hydrogen molecule: -we start with: 1 e-/each in 1s atomic orbitals -we’ll end up with: 2 e- in molecular obital(s) HOW? Make linear combinations of the 1s orbital wavefunctions: ymol = y1s(A) ± y1s(B) Then, solve via the SWE! 2 1 Hydrogen Wavefunctions wavefunctions probability densities 3 Hydrogen Molecular Orbitals anti-bonding bonding 4 2 Hydrogen MO Formation: Internuclear Separation n SWE solved with nuclei at a specific separation distance . How does the energy of the new MO vary with internuclear separation? movie 5 MO Theory: Homonuclear Diatomic Molecules n Let’s look at the s-bonding properties of some homonuclear diatomic molecules: Bond order = ½(bonding e- - anti-bonding e-) For H2: B.O. = 1 - 0 = 1 (single bond) For He2: B.O. = 1 - 1 = 0 (no bond) 6 3 Configurations and Bond Orders: 1st Period Diatomics Species Config. B.O. Energy Length + 1 H2 (s1s) ½ 255 kJ/mol 1.06 Å 2 H2 (s1s) 1 431 kJ/mol 0.74 Å + 2 * 1 He2 (s1s) (s 1s) ½ 251 kJ/mol 1.08 Å 2 * 2 He2 (s1s) (s 1s) 0 ~0 LARGE 7 Combining p-orbitals: s and p MO’s antibonding end-on overlap bonding antibonding side-on overlap bonding antibonding bonding8 4 2nd Period MO Energies s2p has lowest energy due to better overlap (end-on) of 2pz orbitals p2p orbitals are degenerate and at higher energy than the s2p 9 2nd Period MO Energies: Shift! For Z <8: 2s and 2p orbitals can interact enough to change energies of the resulting s2s and s2p MOs.
  • Structure of Benzene, Orbital Picture, Resonance in Benzene, Aromatic Characters, Huckel’S Rule B

    Structure of Benzene, Orbital Picture, Resonance in Benzene, Aromatic Characters, Huckel’S Rule B

    Dr.Mrityunjay Banerjee, IPT, Salipur B PHARM 3rd SEMESTER BP301T PHARMACEUTICAL ORGANIC CHEMISTRY –II UNIT I UNIT I Benzene and its derivatives 10 Hours A. Analytical, synthetic and other evidences in the derivation of structure of benzene, Orbital picture, resonance in benzene, aromatic characters, Huckel’s rule B. Reactions of benzene - nitration, sulphonation, halogenationreactivity, Friedelcrafts alkylation- reactivity, limitations, Friedelcrafts acylation. C. Substituents, effect of substituents on reactivity and orientation of mono substituted benzene compounds towards electrophilic substitution reaction D. Structure and uses of DDT, Saccharin, BHC and Chloramine Benzene and its Derivatives Chemists have found it useful to divide all organic compounds into two broad classes: aliphatic compounds and aromatic compounds. The original meanings of the words "aliphatic" (fatty) and "aromatic" (fragrant/ pleasant smell). Aromatic compounds are benzene and compounds that resemble benzene in chemical behavior. Aromatic properties are those properties of benzene that distinguish it from aliphatic hydrocarbons. Benzene: A liquid that smells like gasoline Boils at 80°C & Freezes at 5.5°C It was formerly used to decaffeinate coffee and component of many consumer products, such as paint strippers, rubber cements, and home dry-cleaning spot removers. A precursor in the production of plastics (such as Styrofoam and nylon), drugs, detergents, synthetic rubber, pesticides, and dyes. It is used as a solvent in cleaning and maintaining printing equipment and for adhesives such as those used to attach soles to shoes. Benzene is a natural constituent of petroleum products, but because it is a known carcinogen, its use as an additive in gasoline is now limited. In 1970s it was associated with leukemia deaths.
  • Reactions of Aromatic Compounds Just Like an Alkene, Benzene Has Clouds of  Electrons Above and Below Its Sigma Bond Framework

    Reactions of Aromatic Compounds Just Like an Alkene, Benzene Has Clouds of  Electrons Above and Below Its Sigma Bond Framework

    Reactions of Aromatic Compounds Just like an alkene, benzene has clouds of electrons above and below its sigma bond framework. Although the electrons are in a stable aromatic system, they are still available for reaction with strong electrophiles. This generates a carbocation which is resonance stabilized (but not aromatic). This cation is called a sigma complex because the electrophile is joined to the benzene ring through a new sigma bond. The sigma complex (also called an arenium ion) is not aromatic since it contains an sp3 carbon (which disrupts the required loop of p orbitals). Ch17 Reactions of Aromatic Compounds (landscape).docx Page1 The loss of aromaticity required to form the sigma complex explains the highly endothermic nature of the first step. (That is why we require strong electrophiles for reaction). The sigma complex wishes to regain its aromaticity, and it may do so by either a reversal of the first step (i.e. regenerate the starting material) or by loss of the proton on the sp3 carbon (leading to a substitution product). When a reaction proceeds this way, it is electrophilic aromatic substitution. There are a wide variety of electrophiles that can be introduced into a benzene ring in this way, and so electrophilic aromatic substitution is a very important method for the synthesis of substituted aromatic compounds. Ch17 Reactions of Aromatic Compounds (landscape).docx Page2 Bromination of Benzene Bromination follows the same general mechanism for the electrophilic aromatic substitution (EAS). Bromine itself is not electrophilic enough to react with benzene. But the addition of a strong Lewis acid (electron pair acceptor), such as FeBr3, catalyses the reaction, and leads to the substitution product.
  • 3-MO Theory(U).Pptx

    3-MO Theory(U).Pptx

    Molecular Orbital Theory Valence Bond Theory: Electrons are located in discrete pairs between specific atoms Molecular Orbital Theory: Electrons are located in the molecule, not held in discrete regions between two bonded atoms Thus the main difference between these theories is where the electrons are located, in valence bond theory we predict the electrons are always held between two bonded atoms and in molecular orbital theory the electrons are merely held “somewhere” in molecule Mathematically can represent molecule by a linear combination of atomic orbitals (LCAO) ΨMOL = c1 φ1 + c2 φ2 + c3 φ3 + cn φn Where Ψ2 = spatial distribution of electrons If the ΨMOL can be determined, then where the electrons are located can also be determined 66 Building Molecular Orbitals from Atomic Orbitals Similar to a wave function that can describe the regions of space where electrons reside on time average for an atom, when two (or more) atoms react to form new bonds, the region where the electrons reside in the new molecule are described by a new wave function This new wave function describes molecular orbitals instead of atomic orbitals Mathematically, these new molecular orbitals are simply a combination of the atomic wave functions (e.g LCAO) Hydrogen 1s H-H bonding atomic orbital molecular orbital 67 Building Molecular Orbitals from Atomic Orbitals An important consideration, however, is that the number of wave functions (molecular orbitals) resulting from the mixing process must equal the number of wave functions (atomic orbitals) used in the mixing