How Do You Tell If a Molecule Is Paramagnetic Or Diamagnetic
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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. -
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. -
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. -
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. -
Chapter 5 the Covalent Bond 1
Chapter 5 The Covalent Bond 1. What two opposing forces dictate the bond length? (Why do bonds form and what keeps the bonds from getting any shorter?) The bond length is the distance at which the repulsion of the two nuclei equals the attraction of the valence electrons on one atom and the nucleus of the other. 3. List the following bonds in order of increasing bond length: H-Cl, H-Br, H-O. H-O < H-Cl < H-Br. The order of increasing atom size. 5. Use electronegativities to describe the nature (purely covalent, mostly covalent, polar covalent, or ionic) of the following bonds. a) P-Cl Δχ = 3.0 – 2.1 = 0.9 polar covalent b) K-O Δχ = 3.5 – 0.8 = 2.7 ionic c) N-H Δχ = 3.0 – 2.1 = 0.9 polar covalent d) Tl-Br Δχ = 2.8 – 1.7 = 1.1 polar covalent 7. Name the following compounds: a) S2Cl2 disulfur dichloride b) CCl4 carbon tetrachloride c) PCl5 phosphorus pentachloride d) HCl hydrogen chloride 9. Write formulas for each of the following compounds: a) dinitrogen tetroxide N2O4 b) nitrogen monoxide NO c) dinitrogen pentoxide N2O5 11. Consider the U-V, W-X, and Y-Z bonds. The valence orbital diagrams for the orbitals involved in the bonds are shown in the margin. Use an arrow to show the direction of the bond dipole in the polar bonds or indicate “not polar” for the nonpolar bond(s). Rank the bonds in order of increasing polarity. The bond dipole points toward the more electronegative atom in the bond, which is the atom with the lower energy orbital. -
Carbon Carbon Spsp Hybrids : Acetylene and the : Acetylene And
CarbonCarbon spsp hybridshybrids:: AcetyleneAcetylene andand thethe TripleTriple bondbond CC2HH2 isis HH--CCCC--HH FormForm spsp onon eacheach CC leavingleaving 2p2px,, 2p2py “unused”“unused” H - sp C sp + +-+ sp C sp H PredictsPredicts linearlinear structure.structure. AboveAbove areare allall σσ bondsbonds H-C-C-H Uses up 2 valence e- for each C in sp (and 2 from 1s H). Leaves 2 valence e - for each C unused. 2py 2py Put last valence 2px e- into ππ orbitals sp formed from H C C H 2px, 2py 2px - - 2 e in πx 2 e in πy Gives two π bonds connecting carbon atoms. π bonds are at right angles to each other. Total C-C bonds are now 3, one σσ, 2 π Short Comparison of Bond Order, Bond Length, Bond Energy C-C C-C C-C Molecule Bond Order Bond Length Bond E, kcal/mole σσ Ethane, C2H6 1 (1 ) 1.54Å 83 σσ Ethylene, C2H4 2 (1 , 1π) 1.35Å 125 σσ Acetylene, C2H2 3 (1 , 2π) 1.21Å 230 ConjugationConjugation andand DelocalizationDelocalization ofof ElectronsElectrons andand BondsBonds Although energy of π* in ethylene < σ*, conjugated polylenes have even lower energy π* levels. These absorb light at longer wavelength- sometimes even in visible (human eye’s light perception). Conjugated polyenes: C=C-C-C=C-C 2 essentially independent double bonds. C=C-C=C-C=C polyene DoubleDouble bondsbonds ALTERNATE!ALTERNATE! Could draw ππ bonds between any 2 C’s ππ* bonds C C C C C C drawn are not unique! ψψ This gives delocalized structure MO = (Const)[2py(1) + 2py(2) + 2py(3) + 2py(4) +2py(5) + 2py(6) +………]. -
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 -
Bond Orders of the Diatomic Molecules†
RSC Advances PAPER View Article Online View Journal | View Issue Bond orders of the diatomic molecules† Cite this: RSC Adv.,2019,9,17072 Taoyi Chen and Thomas A. Manz * Bond order quantifies the number of electrons dressed-exchanged between two atoms in a material and is important for understanding many chemical properties. Diatomic molecules are the smallest molecules possessing chemical bonds and play key roles in atmospheric chemistry, biochemistry, lab chemistry, and chemical manufacturing. Here we quantum-mechanically calculate bond orders for 288 diatomic molecules and ions. For homodiatomics, we show bond orders correlate to bond energies for elements within the same chemical group. We quantify and discuss how semicore electrons weaken bond orders for elements having diffuse semicore electrons. Lots of chemistry is effected by this. We introduce a first-principles method to represent orbital-independent bond order as a sum of orbital-dependent bond order components. This bond order component analysis (BOCA) applies to any spin-orbitals that are unitary transformations of the natural spin-orbitals, with or without periodic boundary conditions, and to non-magnetic and (collinear or non-collinear) magnetic materials. We use this BOCA to study all À Creative Commons Attribution-NonCommercial 3.0 Unported Licence. period 2 homodiatomics plus Mo2,Cr2, ClO, ClO , and Mo2(acetate)4. Using Manz's bond order equation with DDEC6 partitioning, the Mo–Mo bond order was 4.12 in Mo2 and 1.46 in Mo2(acetate)4 with a sum of bond orders for each Mo atom of 4. Our study informs both chemistry research and education. As Received 5th February 2019 a learning aid, we introduce an analogy between bond orders in materials and message transmission in Accepted 10th May 2019 computer networks.