分子轨道理论,杂化

1Chemistry 5.12, Lecture #3, 2/10/03OutlineIII. Review of Molecular Orbital Theory (Read Chapter 2) A. Atomic Orbitals (2-1) B. Sigma-Bonding (2-2A) C. Pi-Bonding (2-2B) D. VSEPR Theory (2-4)IV. Hybridization/LCAO A. sp Hybridization (2-4A)

B. sp2 Hybridization (2-4B)

C. sp3 Hybridization (2-4C) 1. Rotation of Ethane versus Ethylene (2-3,7,8)

Problem Set #1 Due Today by 4pmProblem Set #2 Posted: Due Tuesday (2/18/03)

Suggested Problems: 2-28,29,31,33* Read 2-12,13: Survey of Organic Molecules

Electrons Are Waves!They Exist as 3-D Standing Waves (Orbitals).

atomic orbitals: unhybridized orbitals on an atom (s, p, d)Linear Combination of Atomic Orbitals (LCAO): Individual wave functions (orbitals) combine to form hybrid atomic orbitals (sp, sp2, sp3) and molecular orbitals (σ, σ*, π, π*)

Hybrid Atomic Orbital: Combination of atomic orbitals from the same atom

Molecular Orbital: Combination of atomic orbitals from different atomsConservation of Orbitals: When you add orbitals together, you always end up with the same number of orbitals that you started with.

III. Molecular Orbital Theory

󰻻㒆󰀃󰀃󰣪󱄺5.12,䆆󱑻3,2003󱑈2󱳜10󱮹󰟚󱄤䔼䘧⧚䆎󰻡д˄䯙䇏㄀Ѡゴ˅󰥳󱄤䔼䘧(2-1)³䬂(2-2A)±䬂(2-2B)Ӌ󱈖⬉󱄤󱇍Ѧ󱭹⧚䆎(VSEPR)(2-4)󱴖󰣪/󰥳󱄤䔼䘧㒓󱗻㒘󰧜（LCAO）󱴖󰣪(2-4A)󱴖󰣪(2-4B)󱴖󰣪(2-4C)Э⛋󰩠Э⛃󱮟䕀П󱇍↨(2-3,7,8)

Ҟ󰻽ϟ󰤜4⚍ϞѸд乬䲚󰀆1󱏗㕂д乬#2：ϟ󰨼ѠϞѸ(2003年2月18日)

󱥼㤤乬Ⳃ˖

䯙䇏󱳝󱴎󰟚󱄤㒉㾜

󰟚󱄤䔼䘧⧚䆎⬉󱄤󱰃ϔ⾡⊶ʽ󱅗Ӏҹϝ㓈偏⊶ⱘ󱔶󱓣󱄬󰳼󰀋䔼䘧󰀌

󰥳󱄤䔼䘧：原子中未杂化的轨道(s,p,d)󰥳󱄤䔼䘧㒓󱗻㒘󰧜˖⣀ゟⱘ䔼䘧⊶󰟑󱭄󰀋䔼䘧󰀌ⳌѦ㒘󰧜󱔶󱟤њ󰥳󱄤ⱘ󱴖󰣪䔼䘧(sp,sp2,sp3)ҹ󰦞󰟚󱄤䔼䘧(σ,σ∗,π,π∗)

󰟚󱄤䔼䘧˖ϡ󰧠󰥳󱄤ⱘ䔼䘧㒘󰧜䔼䘧󱅜󱘦˖䔼䘧󰦴󰡴ҹ󰧢󱔶󱟤󱮄䔼䘧ⱘ󱭄ⳂϢ󰦴󰡴󰠡ⱘ󱭄ⳂⳌ󰧠DŽ

󰥳󱄤󱴖󰣪䔼䘧˖Ⳍ󰧠󰥳󱄤ⱘ󰥳󱄤䔼䘧㒘󰧜2

A. Atomic Orbitals• The s- and p-orbitals are the most important in organic compounds.• Orbital pictures are actually electron probability clouds.

xyzxyzxyzxy

zspXpYpZ

s-orbitals: spherical, electrons held close to nucleus, one sign

p-orbitals: two lobes with opposite signs, electrons further from nucleus

• Remember, the sign of the orbital does not indicate charge. It represents the sign of the wavefunction and lets us think qualitatively about whether orbital interactions are constructive (bonding) or destructive (anti-bonding).

B. Sigma-Bonding (σ)• Sigma-bonding orbitals are cylindrically symmetrical molecular orbitals.• Electron density is centered along the axis of the bond.• Single bonds are sigma-bonds.

e.g. H2 is the simplest sigma-bond

H.H.H–Hσ (bonding)

σ* (anti-bonding)

nodebonding: (+/+ or –/–) electron

density centered between nuclei

anti-bonding: (+/–) generally has a node between nuclei

node: area of zero electron density

• In stable bonding situations, usually only the bonding orbitals (σ, π) are occupied.

󰥳󱄤䔼䘧s󰩠p䔼䘧󱰃󱳝󱴎󰣪󰧜⠽Ё󱳔䞡㽕ⱘ䔼䘧DŽ䔼䘧󰳒󱅲䰙󱰃⬉󱄤󰞴⥛ѥDŽ

s䔼䘧˖󰨜⧗󱔶󰟚󱏗ˈ⬉󱄤⾏Ḍ䖥ˈ󰤩ϔヺ󰧋p䔼䘧˖ϸ⪷󱓖󰧋ˈ⬉󱄤⾏Ḍ䕗䖰

󰟛䆄ˈ䔼䘧ⱘℷ䋳󰧋ϡҷ㸼⬉㥋DŽ󰀃󰀃󱅗󰦾ҷ㸼⊶󰟑󱭄ⱘヺ󰧋ˈ󱟥Ӏ󰧃ҹ󱅮󱗻󰴄⧚㾷Ў䔼䘧䯈ⱘⳌѦ԰⫼󱰃󰧺󱳝󰟽󰀋󱟤䬂Ϣ󰦡䬂󰀌DŽ

³䬂³󱟤䬂䔼䘧Ў󰳚󱷅󱔶󱇍⿄ⱘ󰟚󱄤䔼䘧DŽ⬉󱄤䲚Ё󰟚󱏗󰳼䬂䕈Ё󰼂DŽ󰤩䬂Ў³䬂DŽ

՟󰽖˖H2󱰃󱳔ㅔ󰤩ⱘ³䬂

󱟤䬂˖󰀋+/+󱟪󰀐󰀒󰀐)˖Ḍ䯈⬉󱄤ѥ󱆚󱑺󱳔󰻻DŽ󰦡䬂˖󰀋󰀎󰀒󰀐󰀌䗮󱐌󰳼Ḍ䯈󱄬󰳼ϔϾ㡖⚍DŽ㡖⚍˖⬉󱄤ѥ󱆚󱑺Ў䳊ⱘ󰤎󰶳DŽ

󰳼〇󱅮󱟤䬂ⱘ󱚙󰞉ϟˈ⬉󱄤ϔ㠀󰦾󰤴󱥂󱟤䬂䔼䘧˄³,±˅

㡖⚍˄󰦡䬂˅

˄󱟤䬂˅3

• Sigma bonds aren't necessarily between two s-orbitals.σnodeσ*

spX

s + ppXpXnodeσσ*p + p

• These are all examples of single bonds.C. Pi-Bonding (π)• Pi-bonding orbitals are not cylindrically symmetrical.• Electron density is located above and below the axis of the bond.• Double and triple bonds are pi-bonds.

node

ππ*

pYpY

• Double bond = σ + πe.g. Ethylene

CCHHHHCCHHHH

HHHH

σπCCHHHHCC

HH

HHor

to simplify drawing of orbitalsσ

π

πσ

π

π

³䬂ⱘ󱔶󱟤󱑊ϡ󱈔䰤ѢϸϾs䔼䘧䯈DŽҹϞ󰴛󱰃󰧃󱔶󱟤³䬂ⱘ՟󱄤DŽ㡖⚍㡖⚍

±䬂±䔼䘧ϡ󰨜󰳚󱷅󱔶󱇍⿄DŽ⬉󱄤䲚Ё󰟚󱏗󰳼䬂䕈ⱘϞ󱮍󰩠ϟ󱮍DŽ󰦠䬂󰩠ϝ䬂Ў±䬂DŽ

㡖⚍󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰦠䬂՟󰽖˖Э⛃

󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃󰀃䔼䘧ⱘㅔ󱯧⬏⊩