Unit 3 Review 2 VALENCE BOND, MOLECULAR ORBITAL THEORY, IMF S
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1 Unit 3 Review 2 VALENCE BOND, MOLECULAR ORBITAL TEORY, IMF S
2 Valence Bond vs. Molecular Orbital Theory Valence Bond Theory ybrid Orbitals: sp, sp 2, sp 3, sp 3 d, sp 3 d 2 Sigma and pi bonds Focuses on the electron clouds Depends on the number of electron- dense regions Molecular Orbital Theory Diagrams Answers three fundamental quesons regarding your molecule: bond order, magnesm, OMO- LUMO gap Sigma and pi bonds and an- bonds Focuses on the electrons as waves Depends on the number of electrons
3 Valence Bond Theory Valence Bond Theory explains the behavior of the valence orbitals of bonding atoms by suggesng that pure atomic orbitals can form energecally 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)
4 Valence Bond Theory Pure Atomic Orbitals C 2s 2px 2py 2pz Valence Bond ybrid Orbitals 25% s ; 75% p C sp s What you form: 4 End-on-End overlaps (sigma bonds) between the pure 1s of and the hybrid sp³ of C σ1s-sp³ Your hybrid orbital shape (sp3,sp2,sp,etc.) absolutely depends on the molecule / surrounding atoms σ1s-sp³ C
5 O O Pure Atomic Orbitals C 2s 2px 2py 2pz Valence Bond Orbitals 33% s ; 67% p C 2p sp O 1s What you form: End-on-End overlap (sigma bonds) between the pure 1s of and the hybrid sp² of C End-on-End overlap (sigma bond) between the hybrid sp² of O and the hybrid sp² of C Side-by-Side overlap (pi bond) between the pure 2p of O and the pure 2p of C sp2 2p σ1s-sp² σsp²-sp² π2p-2p σ1s-sp² σ1s-sp² C π2p-2p σsp²-sp²
6 Pure Atomic Orbitals C 2s 2px 2py 2pz Valence Bond ybrid Orbitals C C 2p sp C 1s What you form: 1 End-on-End overlap (sigma bond) between the pure 1s of and the hybrid sp of C 1 End-on-End overlap (sigma bond) between the hybrid sp of C and the hybrid sp of C 2 Side-on-Side overlaps (pi bonds) between the pure 2p of C and the pure 2p of C sp 50% s ; 50% p 2p σ1s-sp σsp-sp π2p-2p π2p-2p σ1s-sp C σsp-sp C σ1s-sp π2p-2p
7 Pure atomic Orbital P 3s 3px 3py 3pz 3d Valence Bond Orbitals What you form: + 5Cl sp3d sp3 5 End-on-End overlaps (sigma bonds) between the hybrid sp³d of P and the hybrid sp³ of Cl σsp³d-sp³ Pure atomic Orbital S 3s 3px 3py 3pz 3d Valence Bond Orbitals SCl₆ sp3d2 What you form: + 6Cl 6 End-on-End overlaps (sigma bonds) between the hybrid sp³d² of S and the hybrid sp³ of Cl σsp³d²-sp³ sp3
8 Valence Bond: sp 3 The sp 3 hybridizaon is observed when a central atom has 4 electron- dense regions opefully it is clear to see that this configuraon doesn t really help us make a bond This looks bezer for forming four sigma (end- on- end) bonds
9 Valence Bond: sp 3 Methane forms 4 sigma bonds to fill Carbon s octet
10 Valence Bond: sp 3 with 3 bonding regions Ammonia has sp 3 hybridizaon with only 3 bonding regions This is because ammonia sll has a tetrahedral electronic geometry, resulng in paired electrons in one of the sp 3 orbitals
11 Valence Bond: sp 2 The sp 2 hybridizaon is observed when a central atom has 3 electron- dense regions The carbon in this configuraon has both pure and hybrid orbitals. The hybrid orbitals will form sigma bonds, while the pure 2p orbital will form a pi bond with the pure 2p of oxygen Can you spot the lone pairs in oxygen in the valence bond orbitals?
12 Valence Bond: sp 2
13 Valence Bond: sp The sp hybridizaon is observed when a central atom has 2 electron- dense regions
14 Valence Bond: sp The sp hybridizaon is observed when a central atom has 2 electron- dense regions
15 Valence Bond: sp 3 d The sp 3 d hybridizaon is observed when a central atom has 5 electron- dense regions This can only occur with central atoms with 3p or greater electrons the same rules apply for expanded valence Lewis Structures
16 Valence Bond: sp 3 d 2 The sp 3 d 2 hybridizaon is observed when a central atom has 6 electron- dense regions This can only occur with central atoms with 3p or greater electrons the same rules apply for expanded valence Lewis Structures
17 Valence Bond Summary ybrids Types of Bonds # of Electron- Desnse Regions ybridiza8on 2 sp 3 sp 2 4 sp 3 5 sp 3 d 6 sp 3 d 2 Single Bonds: 1 sigma bond Double Bonds: 1 pi bond and 1 sigma bond Triple Bonds: 2 pi bonds and 1 sigma bond Note: is necessary to memorize this trend with bonds but not necessarily sufficient for the exam. Be able to determine the types of orbitals (pure, hybrid) that go into making the parcular bond.
18 QuesPon Which of the following shapes is possible with sp 3 d hybridizaon (choose all that apply)? a. square planar b. bent c. linear d. see- saw
19 Organic Molecules What is the hybridizaon of the circled central atoms? +
20 Organic Molecules What is the hybridizaon of the circled central atoms? Oxygen sp 3 Nitrogen sp 3 Carbon sp 2 + Phosphorus sp 3
21 Molecular Orbital Theory ATOMIC ORBITALS COMBINE TO FORM MOLECULAR ORBITALS
22 Molecular Orbital Theory For every bonding orbital, there is a higher energy an- bonding orbital Molecular Orbital Theory looks at individual electrons as waves. This theory looks at bonds as regions with a high probability of electron density Construc8ve interference results in a high electron density (bond) Low Energy Stable Destruc8ve interference results in repulsion (zero electron density; an- bond) igh Energy Unstable Note: our focus in this class will be on ydrogen, elium, row 2 diatomics, and their ions
23 Molecular Orbital Theory MO Diagram for ₂ σ1s* 1s 1s σ1s O₂ and F₂ B₂, N₂, and C₂ σ2p* σ2p* πx* πy* πx* πy* 2p 2p 2p 2p πx πy σ2p σ2p πx πy σ2s* σ2s* 2s 2s 2s 2s σ2s σ2s
24 O O₂ O LUMO: lowest unoccupied molecular orbital OMO: highest occupied molecular orbital σ2p* 2p πx* πx πy* πy 2p σ2p σ2s* 2s 2s 6 available valence electrons Conclusions from this diagram: Bond Order = ½ (Bonding electrons - Anti-bonding electrons) Bond Order = ½ (8-4) = 2 Oxygen is paramagnetic, meaning it contains unpaired electrons. Paramagnetic compounds are attracted to a magnetic field. σ2s
25 Molecular Orbital Theory
26 Molecular Orbital Theory: Conclusions Molecular Orbital Theory provides three conclusions that Valence Bond and VSEPR are unable to explain: 1. MO Theory can describe fraconal bond orders ascribed to charged molecules and resonance structures using an- bonding and bonding orbitals 2. MO Theory can idenfy the magnec properes of a given molecule (diamagnec, paramagnec) 3. MO Theory can predict the photon emission energy of excited electrons (OMO- LUMO)
27 Intermolecular Forces
28 Intermolecular Forces S t r e n g t h 3 main types of IMF s: ydrogen bonding (extra strong dipole- dipole) Determine relave strength based on number of hydrogen bonds or electronegave difference (example: O- vs. N- ) Dipole- dipole Determine relave strength based on electronegavity (Lewis Structures are important) Consider polarity when comparing similar molecules in the same period Dispersion forces (present in all molecules) Determine relave strength based on MW, polarizability, shape/ stackability Consider polarizability when comparing similar molecules with atoms in the same group (ex. alogen diatomics) The name of the game for IMF s: Predict the dominant IMF s a molecule and compare relave physical properes
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