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Transcript
VSEPR Theory Valence Shell Electron Pair Repulsion Chapter 10 ◆ Molecular Geometry and Chemical Bonding Theory ◆ Different geometrical arrangements of regions of electron density around a molecule’s central atom would result; dependent on the number of electron pairs. VSEPR Theory and Molecular Geometries ◆ Imagine 2, 3, 4, 5, or 6 balloons tied together: If we could see electron pairs, what would this look like? What 3–D shape would result? VSEPR Theory and Molecular Geometries: The Balloon Model ◆ In a polyatomic covalent species, electron pairs arrange themselves in such a way as to minimize electron – electron repulsion. Now consider that each balloon represents the space occupied by an electron pair (or region of electron density; electron domain). What is the geometrical arrangement of the electron domains around a central atom? ◆ ◆ What shapes result? What angles are created between the balloons by their natural tendency to stay as far away from each other as possible? Electron Domain Geometry Trigonal Bipyramidal Geometry 2 types of positions around the central atom: A – axial positions E – equitorial positions note the angles: ∠ A to E: 90° ∠ E to E: 120° Electron Domain Geometry Determine the of the Number of Electron Domains: # e– domains electron domain around central atom geometry angles between electron domains ◆ draw the Lewis structure ◆ count the e– domains around the central atom 2 linear 180° lone pair of electrons = 1 electron domain 3 trigonal planar 120° 4 tetrahedral 109.5° 5 trigonal bipyramidal 90° & 120° 6 octahedral 90° single bond = 1 electron domain multiple bond = 1 electron domain CO2 SF4 central atom: C 2 double bonds ∴ 2 e– domains e– domain geometry: linear central atom: S 4 single bonds & 1 lone pair ∴ 5 e– domains e– domain geometry: trigonal bipyramidal BF3 XeF4 central atom: B 3 single bonds ∴ 3 e– domains e– domain geometry: trigonal planar central atom: Xe 4 single bonds & 2 lone pair ∴ 6 e– domains e– domain geometry: octahedral Molecular Geometry Molecular Geometry ◆ ◆ What shape is created by the atoms that make up a molecule? How, geometrically, are outer atoms arranged around the central atom? OR ◆ How, geometrically, are the bonding e– domains arranged around the central atom? ◆ When the e– domains around the central atom are all bonding domains, the molecular geometry is the same as the electron domain geometry. Correlation Table for Molecular Geometries linear bent tetrahedral Correlation Table for Molecular Geometries ◆ Where do you “put” the lone pairs of electrons in trigonal bipyramidal geometry? recall: the goal is to minimize e– pair – e– pair repulsion ◆ a lone (nonbonding) pair of e–’s is bigger than a shared (bonding) pair of e–’s ∴ lone pairs of e–’s occupy positions that are farthest away from other electron domains ◆ in trigonal bipyramidal geometry, equitorial positions are favored for lone pairs of electrons Correlation Table for Molecular Geometries see-saw t-shaped ◆ Where do you “put” the lone pairs of electrons in octahedral geometry? linear square pyramidal central atom: I 5 single bonds & 1 lone pair ∴ 6 e– domains e– domain geometry: octahedral molecular geometry: square pyramidal F–I–F bond angles: 90° central atom: Xe 2 single bonds & 3 lone pairs ∴ 5 e– domains e– domain geometry: trigonal bipyramidal molecular geometry: linear F–Xe–F bond angle: 180° square planar central atom: Cl 3 single bonds & 2 lone pairs ∴ 5 e– domains e– domain geometry: trigonal bipyramidal molecular geometry: t-shaped F–Cl–F bond angles: 90° ◆ ◆ Why are observed bond angles not always the same as the predicted bond angles? Lone pairs of electrons are bigger and occupy more space, therefore they have the effect of compressing bond angles slightly. Bond Polarity vs. Molecular Polarity ◆ ◆ Covalent bond polarity is determined by the difference in electronegativity between 2 atoms. Molecular polarity is determined by the arrangement of bond dipoles around the central atom. ◆ ◆ Why are observed bond angles not always the same as the predicted bond angles? Multiple bonds occupy more space than single bonds, therefore they compress bond angles slightly. Molecular Polarity: Behavior of Polar Molecules in the Presence of an Electric Field ◆ Bond Polarity vs. Molecular Polarity bond dipoles are vector quantities: have a magnitude and a direction Molecular Polarity: Considering Bonds Dipoles and a Molecule’s Electron Density Map symmetrically arranged, equal bond dipoles can cancel each other out CO2: 2 polar C–O bonds arranged 180° apart ∴ CO2 is nonpolar H2O: the arrangement of the 2 polar H–O bonds, and the 2 lone pairs of e–’s on O give H2O a net dipole moment, μ > 0 ∴ H2O is polar Geometric Isomers and Polarity * if dipole moment equals zero (i.e. a zero dipole moment, the molecule is nonpolar if dipole moment is not equal to zero (i.e. a nonzero dipole moment), the molecule is polar Valence Bond Theory (VBT) Recall Covalent Bond Formation Between 2 H Atoms Lewis theory (Ch. 9): covalent bond forms when an electron pair is shared between 2 atoms. Valence Bond Theory: concentration of electron density between atoms occurs when a valence atomic orbital on one atom overlaps with the valence atomic orbital on the other atom. ◆ ◆ overlap of orbitals allows 2 e–’s to be shared in the common space between nuclei the greater the orbital overlap, the stronger the bond H2: Consider BeF2 Lewis structure predicts (correctly) that BeF2 is linear with 2 identical Be–F bonds. ◆ Consider H–Cl ◆ ◆ How do we describe the bonding in terms of valence bond theory? Which orbitals on Be and F overlap? Why does this overlap result in linear geometry and identical bonds? sp Hybrid Orbitals: BeF2 Hybrid Atomic Orbitals Atomic orbitals can mix together in a process called hybridization: ◆ ◆ hybridization results in formation of new hybrid atomic orbitals: new shape new spatial orientations new, degenerate energy 2p E e– promotion hybridization ↑↓ 2p ↑ 2s the total number of orbitals remains constant – orbital conservation # hybrid orbitals formed = # atomic orbitals mixed sp2 Hybrid Orbitals: BF3 ↑ sp hybrid orbitals shape of hybrid orbitals sp3 Hybrid Orbitals: CH4 Valence Bond Theory Description of H2O Hybridization Involving d orbitals: sp3d and sp3d2 Hybrid Orbitals XeF2 and XeF4 Prepare and Orbital Diagram for XeF2: Orbital Diagram for XeF4: Geometry Created by Large Lobes of Hybrid Orbitals Sigma (σ) vs. Pi (π) Bonds: Explaining Single, Double, and Triple Bonds There is a specific correlation between: ◆ the number of electron domains around an atom ◆ the electron domain geometry around that atom ◆ the hybridization of the atom: In a σ bond, e– density is concentrated symmetrically around the internuclear axis. ◆ results from the head-to-head overlap of orbitals atomic orbitals (ex. s–p or p–p overlap) OR hybrid orbitals (ex. sp3–sp3 or sp3d2–sp3 overlap) Sigma (σ) vs. Pi (π) Bonds: Explaining Single, Double, and Triple Bonds In a π bond, e– density is concentrated in 2 parallel regions above and below the internuclear axis. ◆ ◆ ◆ Sigma (σ) vs. Pi (π) Bonds: Explaining Single, Double, and Triple Bonds ◆ a single bond is one sigma bond (2 e–’s) results from the parallel overlap (or side-to-side overlap) of unhybridized p orbitals 2 regions of electron density, but still 1 bond – still just 1 pair of e–’s for atoms to participate in π bonding, they must have at least one unhybridized p orbital π bonds require atoms with sp2 or sp hybridization bond multiplicity a double bond is one sigma bond + one pi bond (4 e–’s) a triple bond is one sigma bond + two pi bonds (6 e–’s) ◆ ◆ a σ bond must form before a π bond forms can have a σ bond without a π bond cannot have a π bond without a σ bond the amount of overlap in a π bond is typically less than the amount of overlap in a σ bond ∴ a π bond is weaker than a σ bond Bonding Description for Ethene, C2H4 ◆ ◆ ◆ ◆ ◆ σ Bonding Description for Ethene, C2H4 each C is sp2 hybridized each C has 1 unhybridized p orbital that can participate in π bonding each C–H σ bond is the result of sp2–1s overlap the C–C σ bond is the result of sp2–sp2 overlap the C–C π bond is the result of 2p–2p overlap π Bonding Description for Ethene, C2H4 Bonding Description for Ethene, C2H4 : Experimental Evidence in Support of π Bonding Bonding Description for Acetylene, C2H2 ◆ ◆ ◆ ◆ ◆ σ Bonding Description for Acetylene, C2H2 each C is sp hybridized each C has 2 unhybridized p orbitals that can participate in the formation of 2 π bonds each C–H σ bond is the result of sp–1s overlap the C–C σ bond is the result of sp–sp overlap each C–C π bond is the result of 2p–2p overlap π Bonding Description for Acetylene, C2H2 σ and π Bonding Description for Acetylene, C2H2 Using Valence Bond Theory to Interpret Resonance Structures and Delocalized Bonding Parallel, unhybridized p orbitals on each O are aligned for continuous overlap and ∴ delocalization of π electron density evenly over all 3 atoms. HCN CO32— PF4— Orbital Overlap Bonding Description Hybridization of Central Atom Polar or Nonpolar? Bond Angles Molecular Geometry # Nonbonding Domains # Bonding Domain Template for Complete Structure, Geometry, Polarity, Hybridization and Bonding Description Chapters 9 & 10; Lewis Structures, Resonance, VSEPR Theory, and Valence Bond Theory Electron Domain Geometry Using Valence Bond Theory to Interpret Resonance Structures and Delocalized Bonding: Benzene, C6H6 Complete Molecular Structure-Bonding Template: Total # of Electron Domains at Central Atom Chem 1711: Lewis Structure (including resonance if appropriate) Each O is sp2 hybridized; the σ bonding picture is: Total # Valence Electrons ◆ Recall ozone, O3 and its Lewis structure: Chemical Formula ◆ Using Valence Bond Theory to Interpret Resonance Structures and Delocalized Bonding 89! ! )(+5&,K65-$! +6'0)/5-!Q'4R651'K5-$! XD$!=9! %&02!'1!2&,!)*+,(-+0.!3,*),2.4!0.*-/G!(0.Q*/!02*)!= D9! ! -'/,56$! @9! ◆ What is the hybridization B9! +,+65&,K65-$! of each numbered C? C9! +6'0)/5-!4-5/56$! Molecular Orbital (MO) Theory Extension of quantum mechanics: 8F:G! 8! >H:G! I! LC;G! IM$A! consider the wave functions of electrons and N=OG! 844-R!Y5-,/(,!Q)/K!+&,)6R!S+.)^!+&6,,^!)6!*)26!,-,(+6)/!45'639$!SCP514-,!IM$<9! a resulting orbitals extending over a whole molecule ◆ How many σ and π HN>G! Q)/K'/0!T!Q)/K'/0!+&,)6',3 ! bonds are in this molecule? instead of just a single atom ! ! :;! 2,2.0&,G.0+! #M$! D).!15/R!3'015!5/K!4'!Q)/K3!56,!'/!+&,!1)-,(2-,!4'(+26,K!Q,-).7! molecular orbitals: <;! 2.'3*/0+!O+0/0.! a molecular orbital (MO) is to an molecule as an =;! +'/,0.! b ◆ How many sp3 hybridized ?;! 2.'3*/0+!O4.0)'G0+! atomic orbital (AO) is to an atom atoms are there in this molecule? ! ! ◆ What is the approximate BJ;G! ,53R! H–C–C angle (a)? ! @;! aQ,/2! ! ◆ How many sp2 hybridized :>5A! <! BC5A! D! atoms are ?EFA! there in)*G,.02,! this molecule?H@FA! DI$D! 6<JA! B.,G'(2!2&,!)*+,(-+0.!3,*),2.4!'/!0!)*+,(-+,$! ! ◆ What are the approximate C6BA! Q*/G'/3!R!)*+,(-+0.!3,*),2.4 bond angles (a and b)? ! ! ! each MO holds 2 spin-paired e–’s when full each MO has a specific, characteristic E each MO has a specific shape and spatial orientation ! Molecular Orbitals for H2 !"#$%&'( !"#$%&'( Construct a MO Diagram for H2 $, !"#"$%&'(")*+,$-)*+ "-.$./' !"#"$%&'(")*+,$-)*+$,"-.$./' *0 ** MO Diagrams for 2nd Row Elements MO Diagrams for 2nd Row Elements MO Diagrams for 2nd Row Elements: Molecular Orbitals Formed From p Orbitals Interaction of s and p atomic orbitals causes the energy levels of some MO’s to shift: MO Diagrams for 2nd Row Elements: N2 (after consideration of 2s–2p interaction) Experimental Evidence That O2 Is Paramagnetic: MO Diagrams and Interpretation of Magnetic Properties MO Diagrams for 2nd Row Elements σ2p* σ2p* π2p π2p* σ2p π2p * π2p σ2p σ2s* σ2s* σ2s σ2s An Example of a MO Diagram for a Transition Metal Complex: