Valence Bond Theory
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Transcription 12.02.15
Lecture 12A • 02/15/12 We’re going to continue our discussion of conjugation. If we go back and talk about electrons and how they act like waves, [there’s] something known as the particle in a box. You’ve got some kind of box where imagine that the walls are infinitely tall. It’s a one-dimensional system, where the electron, all it can do, is go left and right between the two walls of the box. The electron’s a wave; it can’t exist outside the box, so whatever function, whatever wave we use to describe the electron, has to have a value of zero at one end of the box and zero at the other end of the box. Graphically, what do the solutions look like? We have something like this, where we’re talking about energy potential; it’s infinite at either end of the box, and zero in between. We have an electron that’s bouncing around inside of it. An electron’s a wave, and there’s function that describes that wave. The only way it can fit in this box and physically make sense is if the wave starts and stops and the ends of the box. It turns out that – long, long, long story short – one of the solutions for the Schrödinger equation is just a sine function – actually, it’s part of an exponential version of a sine function. It’s something, at least, we can draw a pretty picture of. That’s why you’ll start with this problem in a discussion of quantum mechanics, because it is solvable, it is only on in dimension, and it’s more humanly possible to discuss. -
Aromaticity Sem- Ii
AROMATICITY SEM- II In 1931, German chemist and physicist Sir Erich Hückel proposed a theory to help determine if a planar ring molecule would have aromatic properties .This is a very popular and useful rule to identify aromaticity in monocyclic conjugated compound. According to which a planar monocyclic conjugated system having ( 4n +2) delocalised (where, n = 0, 1, 2, .....) electrons are known as aromatic compound . For example: Benzene, Naphthalene, Furan, Pyrrole etc. Criteria for Aromaticity 1) The molecule is cyclic (a ring of atoms) 2) The molecule is planar (all atoms in the molecule lie in the same plane) 3) The molecule is fully conjugated (p orbitals at every atom in the ring) 4) The molecule has 4n+2 π electrons (n=0 or any positive integer Why 4n+2π Electrons? According to Hückel's Molecular Orbital Theory, a compound is particularly stable if all of its bonding molecular orbitals are filled with paired electrons. - This is true of aromatic compounds, meaning they are quite stable. - With aromatic compounds, 2 electrons fill the lowest energy molecular orbital, and 4 electrons fill each subsequent energy level (the number of subsequent energy levels is denoted by n), leaving all bonding orbitals filled and no anti-bonding orbitals occupied. This gives a total of 4n+2π electrons. - As for example: Benzene has 6π electrons. Its first 2π electrons fill the lowest energy orbital, and it has 4π electrons remaining. These 4 fill in the orbitals of the succeeding energy level. The criteria for Antiaromaticity are as follows: 1) The molecule must be cyclic and completely conjugated 2) The molecule must be planar. -
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). -
Alkenes and Alkynes
02/21/2019 CHAPTER FOUR Alkenes and Alkynes H N O I Cl C O C O Cl F3C C Cl C Cl Efavirenz Haloprogin (antiviral, AIDS therapeutic) (antifungal, antiseptic) Chapter 4 Table of Content * Unsaturated Hydrocarbons * Introduction and hybridization * Alkenes and Alkynes * Benzene and Phenyl groups * Structure of Alkenes, cis‐trans Isomerism * Nomenclature of Alkenes and Alkynes * Configuration cis/trans, and cis/trans Isomerism * Configuration E/Z * Physical Properties of Hydrocarbons * Acid‐Base Reactions of Hydrocarbons * pka and Hybridizations 1 02/21/2019 Unsaturated Hydrocarbons • Unsaturated Hydrocarbon: A hydrocarbon that contains one or more carbon‐carbon double or triple bonds or benzene‐like rings. – Alkene: contains a carbon‐carbon double bond and has the general formula CnH2n. – Alkyne: contains a carbon‐carbon triple bond and has the general formula CnH2n‐2. Introduction Alkenes ● Hydrocarbons containing C=C ● Old name: olefins • Steroids • Hormones • Biochemical regulators 2 02/21/2019 • Alkynes – Hydrocarbons containing C≡C – Common name: acetylenes Unsaturated Hydrocarbons • Arene: benzene and its derivatives (Ch 9) 3 02/21/2019 Benzene and Phenyl Groups • We do not study benzene and its derivatives until Chapter 9. – However, we show structural formulas of compounds containing a phenyl group before that time. – The phenyl group is not reactive under any of the conditions we describe in chapters 5‐8. Structure of Alkenes • The two carbon atoms of a double bond and the four atoms bonded to them lie in a plane, with bond angles of approximately 120°. 4 02/21/2019 Structure of Alkenes • Figure 4.1 According to the orbital overlap model, a double bond consists of one bond formed by overlap of sp2 hybrid orbitals and one bond formed by overlap of parallel 2p orbitals. -
II. Nomenclature Rules for Alkenes 1. the Parent Name Will Be the Longest
1 Lecture 9 II. Nomenclature Rules For Alkenes 1. The parent name will be the longest carbon chain that contains both carbons of the double bond. Drop the -ane suffix of the alkane name and add the –ene suffix. Never name the double bond as a prefix. If a double bond is present, you have an alkene, not an alkane. alkane + -ene = alkene 2. Begin numbering the chain at the end nearest the double bond. Always number through the double bond and identify its position in the longest chain with the lower number. In the older IUPAC rules the number for the double bond was placed in front of the stem name with a hyphen. Under the newer rules, the number for the double bond is placed right in front of “ene”, with hyphens. We will use the newer rules for specifying the location of pi bonds. 1 2 3456 H3CCHCH CH 2 CH2 CH3 hex-2-ene (newer rules) 2-hexene (older rules) 3. Indicate the position of any substituent group by the number of the carbon atom in the parent (longest) chain to which it is attached. CH 1 2 345 3 H3CCHCHCHCH2 CH CH3 6 7 CH3 Numbering is determined by the double bond, not the branches, because the double bond has 5,6-dimethylhept-3-ene (newer rules) higher priority than any alkyl branch. 5,6-dimethyl-3-heptene (older rules) 4. Number cycloalkenes so that the double bond is 1,2 (number through the double bond). Number in the direction about the ring so that the lowest number is used at the first point of difference. -
Bsc Chemistry
____________________________________________________________________________________________________ Subject Chemistry Paper No and Title 2 and Physical Chemistry-I Module No and Title 27 and Valence Bond Theory I Module Tag CHE_P2_M27 SUBJECT PAPER : 2, Physical Chemistry I MODULE : 27, Valence Bond Theory I ____________________________________________________________________________________________________ TABLE OF CONTENTS 1. Learning Outcomes 2. Introduction 3. Valence Bond Theory (VBT) 3.1 Postulates of VBT 3.2 VBT of Hydrogen molecule 4. Summary 1. SUBJECT PAPER : 2, Physical Chemistry I MODULE : 27, Valence Bond Theory I ____________________________________________________________________________________________________ 1. Learning Outcomes After studying this module, you shall be able to: Learn about electronic structure of molecules using VBT Understand the postulates of VBT The VBT of Hydrogen molecule 2. Introduction Now, at this point we know that the Schrodinger equation cannot be solved exactly for multi-electron system even if the inter-nuclear distances are held constant. This is due to the presence of electron-electron repulsion terms in the Hamiltonian operator. Consequently, various approaches have been developed for the approximate solution of Schrodinger equation. Of these, Molecular Orbital Theory (MOT) and Valence Bond Theory (VBT) have been widely used. These two approaches differ in the choice of trial/approximate wave-function which is optimized via variation method for the system under consideration. Of these, we have already discussed MOT in earlier modules. In this module, we will take up the formalism of VBT in detail. 3. Valence Bond Theory Valence Bond Theory was the first quantum mechanical treatment to account for chemical bonding. This theory was first introduced by Heitler and London in 1927 and subsequently by Slater and Pauling in 1930s. -
I 1.9 Aromotic Compounds
5t8 CHAPTERI I CarbonChains and Rings PRACTICE EXERCISEII.S \ Draw structural formulas for the following alkenes.If a compound has geometric isomers, draw both the cls and the transforms. (a) l-pentene (b) 2-hexene (c) 2-methyl-2-hexene (d) 2,3-dimethyl-2-butene Alkynes Organic compounds containing carbon-carbon triple bonds are called alkynes. Like alkenes,alkynes are unsaturated compounds. This is the car- bon-carbon triple bond found in alkynes: -C:C- Allqmes are quite rare in nature. The simplest alkyne is the gasethlme, C2H2. The common name for ethyne is acetylene-the fuel burned in oxyacety- lene torches used in welding. The single bonds that extend from the car- bons involved in the carbon-carbon Figurel l .l0 triple bond of ethpre are separatedby (Fig. Ball-and-stickmodel of ethyne the maximum bond angle, IB0 degrees f f .10). Ethyne is a linear (commonname acetylene). molecule. The development of an orbital picture for ethyne requires us to go through the same steps that we used for methane and ethene: hybridiza- tion of carbon atomic orbitals and orbital overlap. As with methane and ethene,remember that each orbital representsone electron.The descrip- tion of ethyne that best flts the experimental data is obtained if the 2s atomic orbital of carbon is mixed with only one of the three available 2p atomic orbitals (Fig. f 1.11).The result of this mixing is two sp atomic orbitals; two 2p orbitals remain unused. Overlap of one of the sp hybrid orbitals of one carbon with that of another carbon producesa carbon-car- bon sigma bond. -
Valence Bond Theory
Valence Bond Theory • A bond is a result of overlapping atomic orbitals from two atoms. The overlap holds a pair of electrons. • Normally each atomic orbital is bringing one electron to this bond. But in a “coordinate covalent bond”, both electrons come from the same atom • In this model, we are not creating a new orbital by the overlap. We are simply referring to the overlap between atomic orbitals (which may or may not be hybrid) from two atoms as a “bond”. Valence Bond Theory Sigma (σ) Bond • Skeletal bonds are called “sigma” bonds. • Sigma bonds are formed by orbitals approaching and overlapping each other head-on . Two hybrid orbitals, or a hybrid orbital and an s-orbital, or two s- orbitals • The resulting bond is like an elongated egg, and has cylindrical symmetry. Acts like an axle • That means the bond shows no resistance to rotation around a line that lies along its length. Pi (π) Bond Valence Bond Theory • The “leftover” p-orbitals that are not used in forming hybrid orbitals are used in making the “extra” bonds we saw in Lewis structures. The 2 nd bond in a double bond nd rd The 2 and 3 bonds in a triple bond /~harding/IGOC/P/pi_bond.html • Those extra bonds form only after the atoms are brought together by the formation of the skeletal bonds made by www.chem.ucla.edu hybrid orbitals. • The “extra” π bonds are always associated with a skeletal bond around which they form. • They don’t form without a skeletal bond to bring the p- orbitals together and “support” them. -
Visualizing Molecules - VSEPR, Valence Bond Theory, and Molecular Orbital Theory
Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Visualizing Molecules - VSEPR, Valence Bond Theory, and Molecular Orbital Theory Nestor S. Valera 1 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Visualizing Molecules 1. Lewis Structures and Valence-Shell Electron-Pair Repulsion Theory 2. Valence Bond Theory and Hybridization 3. Molecular Orbital Theory 2 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. The Shapes of Molecules 10.1 Depicting Molecules and Ions with Lewis Structures 10.2 Valence-Shell Electron-Pair Repulsion (VSEPR) Theory and Molecular Shape 10.3 Molecular Shape and Molecular Polarity 3 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 10.1 The steps in converting a molecular formula into a Lewis structure. Place atom Molecular Step 1 with lowest formula EN in center Atom Add A-group Step 2 placement numbers Sum of Draw single bonds. Step 3 valence e- Subtract 2e- for each bond. Give each Remaining Step 4 atom 8e- valence e- (2e- for H) Lewis structure 4 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Molecular formula For NF3 Atom placement : : N 5e- : F: : F: : Sum of N F 7e- X 3 = 21e- valence e- - : F: Total 26e Remaining : valence e- Lewis structure 5 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. SAMPLE PROBLEM 10.1 Writing Lewis Structures for Molecules with One Central Atom PROBLEM: Write a Lewis structure for CCl2F2, one of the compounds responsible for the depletion of stratospheric ozone. -
Valence Bond, Molecular Orbital Theory, Imf's
Unit 3 Review 2 VALENCE BOND, MOLECULAR ORBITAL THEORY, IMF’S Valence Bond vs. Molecular Orbital Theory • Valence Bond Theory • Molecular Orbital Theory • Hybrid Orbitals: sp, sp2, sp3, sp3d, sp3d2 • Diagrams • Sigma and pi bonds • Answers three fundamental quesHons regarding • Focuses on the electron “clouds” your molecule: bond order, magneHsm, HOMO- LUMO gap • Depends on the number of electron-dense regions • Sigma and pi bonds and anH-bonds • Focuses on the electrons as “waves” • Depends on the number of electrons Valence Bond Theory • Valence Bond Theory explains the behavior of the valence orbitals of bonding atoms by suggesHng that pure atomic orbitals can form energeHcally favorable hybrid orbitals • The conclusions of Valence Bond Theory are: 1. The hybridizaon of a bonding atom depends on the number of electron regions (bonding regions plus lone-pair regions) 2. Bonds can be made between pure and hybrid orbitals 3. Bonds can be classified as sigma or pi depending on the orientaon of the orbitals’ overlap (sigma – end-on-end; pi– off axis / side-by-side) Valence Bond Theory Pure Atomic Orbitals C 2px 2py 2pz 2s Valence Bond Hybrid Orbitals H 25% s ; 75% p H C sp H H + 4H 1s What you form: 4 End-on-End overlaps (sigma bonds) between the pure 1s of H and the hybrid sp³ of C σ1s-sp³ Your hybrid orbital shape (sp3,sp2,sp,etc.) absolutely depends on the molecule / surrounding atoms H σ1s-sp³ H C H H Pure Atomic Orbitals C 2px 2py 2pz 2s Valence Bond Orbitals O 33% s ; 67% p 2p C sp H H + 2H + O 2p 1s sp What you form: End-on-End overlap -
Chemical Bonds Bonding in Hydrogen
Chemical Bonds Up until this time, we have used a very simple model for bonding where lines represent a pair of shared electrons in a chemical bond and dots represent unshared electrons. You've seen how two orbitals from different atoms that point directly at each other can overlap in in a sigma bond and orbitals that are parallel can overlap to make a pi bond. In this unit we will extend ideas about bonding in diatomic molecules. Outline • First Row: Hydrogen • Second Row: Nitrogen • Electronegativity Effects • Homework Bonding in Hydrogen Linear Combination of Atomic Orbitals (LCAO) The orbitals from atoms in a molecule combine to make molecular orbitals. There are several rules for the construction of molecular orbitals from atomic orbitals. 1. Only orbitals that have the same symmetry can combine. 2. The number of orbitals can't change, so the number of atomic orbitals must equal the number of molecular orbitals. 3. The average energy of the orbitals can't change, so the average energy of molecular orbitals must equal the average energy of the atomic orbitals that made them up. Hydrogen Molecular Orbital Diagram Atomic hydrogen has 1 electron in a 1s orbital. Of course, there are 2s, 2p, 3s, 3p, etc. empty orbitals at higher energy. Let's just consider the 1s orbitals. Remember that an orbital is a mathematical function that describes the probability of finding an electron in space. Like all mathematical functions, it can have positive values and negative values. The figure on the left represents the function corresponding to +1s on the x,y,z coordinates while the figure on the right represents -1s function. -
Physical Organic 2015-2016 Coloured Lecture 2 Aromaticity
DAMIETTA UNIVERSITY CHEM-204: PHYSICAL ORGANIC CHEMISTRY LECTURE 2 Dr Ali El-Agamey Aromaticity—Hückel’s Rule HydrogenationUnusual of just oneStability double bond in benzene is endothermic! => Chapter 163 Resonance Energy • This difference between the predicted and the observed value is called the resonance energy. Chapter 164 Reactivity of naphthalene as compared to benzene Reactivity Stability of the As R.E. increases decreases compound increases 5 The Criteria for Aromaticity—Hückel’s Rule Four structural criteria must be satisfied for a compound to be aromatic. [1] A molecule must be cyclic. To be aromatic, each p orbital must overlap with p orbitals on adjacent atoms. [2] A molecule must be planar. All adjacent p orbitals must be aligned so that the π electron density can be delocalized. Since cyclooctatetraene is non-planar, it is not aromatic, and it undergoes addition reactions just like those of other alkenes. [3] A molecule must be completely conjugated. Aromatic compounds must have a p orbital on every atom. [4] A molecule must satisfy Hückel’s rule, and contain a particular number of π electrons. Hückel's rule: Benzene is aromatic and especially stable because it contains 6 π electrons. Cyclobutadiene is antiaromatic and especially unstable because it contains 4 π electrons. Considering aromaticity, a compound can be classified in one of three ways: 1. Aromatic—A cyclic, planar, completely conjugated compound with 4n + 2 π electrons. 2. Antiaromatic—A cyclic, planar, completely conjugated compound with 4n π electrons. 3. Not aromatic (nonaromatic)—A compound that lacks one (or more) of the following requirements for aromaticity: being cyclic, planar, and completely conjugated.