Organic Chemistry I - Structure of Organic Molecules - Quick Guides | Chemistry

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Study GuideOrganic Chemistry I–Structure of Organic Molecules1.Molecular OrbitalsWhen atoms come together to form a molecule, their electrons don’t just stay in the same atomicorbitals. Instead, the atomic orbitals combine to formmolecular orbitals. These new orbitals belongto the entire molecule, not to any single atom.Let’s understand how this works using the hydrogen molecule as an example.How Molecular Orbitals FormWhen two hydrogen atoms approach each other, their1s atomic orbitals overlap. From this overlap,two molecular orbitalsare created:1.A bonding molecular orbital2.An antibonding molecular orbitalThese two orbitals come from the same atomic orbitals but differ in how the wave functions interact.1.1Bonding Molecular Orbital (σBond)In one combination, the wave functions of the two atomic orbitalsadd together (in phase).•This causesmore electron density between the two nuclei.•More electron density between the nuclei means stronger attraction between the negativelycharged electrons and the positively charged nuclei.•This attraction pulls the atoms together andholds the molecule together.This low-energy, stable orbital is called abonding molecular orbital.Because it forms from end-to-end overlap of atomic orbitals, it is called aσ (sigma) bondingmolecular orbital.Key idea:High electron density between nuclei = strong bond.

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Study Guide1.2Antibonding Molecular Orbital (σ*Orbital)In the other combination, the wave functionssubtract from each other (out of phase).•This creates a regionwith no electron density between the nuclei, called anode.•Without electrons between them, the positively charged nucleirepel each other strongly.•This makes the molecule unstable.This higher-energy orbital is called anantibonding molecular orbital, written asσ*.Important note:The plus (+) and minus (–) signs shown in diagrams representwave phase, notelectrical charge.Figure 11.3Electron Spin and StabilityJust like in atomic orbitals,electrons in molecular orbitals must have opposite spinswhen theypair up.In a hydrogen molecule:•Both electrons occupy thelower-energyσbonding orbital.•Theσ*antibonding orbital remains empty.

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Study GuideAs long as a molecule hasmore electrons in bonding orbitals than in antibonding orbitals, themolecule isstable.1.4Why Helium Molecules Do Not ExistHelium atoms havetwo electrons each. If two helium atoms tried to bond:•Electrons would fill both thebondingandantibondingorbitals.•The stabilizing effect of bonding electrons would becancelled outby antibonding electrons.Because there isno overall energy advantage, helium molecules donotform.Bottom line:Equal bonding and antibonding electrons = no bond.Figure 2

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Study GuideWhat Is aσ (Sigma) Bond?Aσbondhas:•High electron densityalong an imaginary line connecting the nuclei•Electron density that comes fromend-to-end overlapof orbitalsσbonds can form from:•Two atomic orbitals•An atomic orbital and a hybrid orbital•Two hybrid orbitalsσbonds are thestrongest type of covalent bondbecause the overlap is very effective.Figure 3What Is aπ (Pi) Bond?Aπbondforms when orbitals overlapside-by-side, instead of head-on.•Electron density is foundabove and below the line connecting the nuclei•The molecular orbital formed is called aπmolecular orbitalBecause this overlap is less direct:•πbonds areweaker thanσbonds•Repulsion between nuclei reduces the quality of overlapπbonds usually formin addition to aσbond, such as in double and triple bonds.

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Study GuideFigure 4Key Takeaways•Atomic orbitals combine to formbonding and antibonding molecular orbitals•Bonding orbitalsincrease electron density between nuclei and stabilize molecules•Antibonding orbitalscreate nodes and destabilize molecules•σbonds form byend-to-end overlap•πbonds form byside-to-side overlap•Molecules are stable only whenbonding electrons outnumber antibonding electrons2.Hybridization of Atomic OrbitalsTo understand hybridization,very simple and familiar molecule:methane (CH₄).Careful experiments on methane show some important facts:•Allcarbon–hydrogen bond lengths are equal•AllH–C–H bond angles are equal•Each bond angle is about109.5° (roughly 110°)•All the bonds arecovalentThese observations tell us that all four bonds in methane areidentical. But when we look at carbon’selectron configuration, this creates an interesting problem.

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Study Guide2.1The Ground State of CarbonCarbon has an atomic number of 6. In itsground state(the lowest-energy, unexcited state), itselectrons are arranged like this:•1s²•2s²•2p²This means carbon hasonly two half-filled orbitals, which would allow it to form justtwo covalentbonds.But methane clearly hasfour bonds.So how does carbon do this?2.2The Excited State of CarbonWhen energy is added to a carbon atom,one electron from the 2s orbital is promoted to a 2porbital.Now carbon has:•Four half-filled orbitals•Each orbital can form a covalent bondThis excited state explains how carbon can form four bonds. However, there’s still a problem:•The2s orbital is smaller•The2p orbitals are larger

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Study GuideIf carbon used these orbitals directly, the bonds would havedifferent lengths, which contradictsexperimental evidence.Why Hybridization Is NeededTo makeall four bonds equal, carbon needsfour identical orbitals.Nature solves this problem through a process calledhybridization.What Is Hybridization?Hybridizationis the mixing (linear combination) of atomic orbitals on the same atom to formnew,equivalent hybrid orbitals.•The wave functions ofs and p orbitals combine•New orbitals form withdifferent shapes and energies•These hybrid orbitals explainequal bond lengths and anglessp³ Hybridization (Methane)In methane:•One s orbitalmixes withthree p orbitals•This producesfour identical sp³ hybrid orbitalsEach sp³ orbital has:•1 part s character•3 parts p characterThese four orbitals point toward the corners of atetrahedron, giving bond angles of109.5°, exactlywhat experiments show.This explains whyall C–H bonds in methane are equal.2.3Bond Angles and VSEPR TheoryBond angles are explained usingValence-Shell Electron-Pair Repulsion (VSEPR) theory.

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Study GuideAccording to VSEPR:•Electron pairs repel each other•They arrange themselves as far apart as possibleIn methane:•Four bonding pairs surround carbon•Maximum separation gives atetrahedral geometry•Bond angle =109°28′ (about 110°)sp² HybridizationCarbon can also formsp² hybrid orbitals.In this case:•One s orbitalcombines withtwo p orbitals•Three equivalent sp² hybrid orbitalsare formed•One p orbital remains unhybridizedKey features:•The three sp² orbitals lie inone plane•They are separated by120°•The unhybridized p orbital isperpendicularto the planeThis arrangement allows for the formation ofdouble bonds.sp HybridizationCarbon can also undergosp hybridization.Here:•One s orbitalmixes withone p orbital•Two sp hybrid orbitalsare formed•Two p orbitals remain unhybridized

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