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Structure and Properties of Organic Molecules

This chapter explores the formation of covalent bonds, the involvement of orbitals, the shapes of organic molecules, and how bonding and shape affect their properties. It also focuses on hybridization, bond rotation, and the hybridization of heteroatoms.

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Structure and Properties of Organic Molecules

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  1. Chapter 2: Structure andProperties of Organic Molecules • Previously: • Atomic/electronic structure • Lewis structures • Bonding • Now: • How do atoms form covalent bonds? • Which orbitals are involved? • What are the shapes of organic molecules? • How do bonding and shape affect properties?

  2. Which electrons are involved in bonding? Valence electrons • Where are valence electrons? In atomic orbitals • Bonds are formed by the combination of atomic orbitals • Linear combination of atomic orbitals (LCAO)

  3. Linear Combination of Atomic Orbitals • Two theories • Molecular Orbital Theory • Atomic orbitals of two atoms interact • Bonding and antibonding MO’s formed • Skip this stuff • Valence Bond Theory (Hybridization) • Atomic orbitals of the same atom interact • Hybrid orbitals formed • Bonds formed between hybrid orbitals

  4. Let’s consider carbon… • How many valence electrons? 4 • In which orbitals? 2s22p2 • So, both the 2s and 2p orbitals are used to form bonds • How many bonds does carbon form? • All four C-H bonds are the same • i.e. there are not two types of bonds from the two different orbitals • How do we explain this? Hybridization

  5. Hybridization • The s and p orbitals of the C atom combine with each other to form hybrid orbitals before they combine with orbitals of another atom to form a covalent bond

  6. sp3 hybridization 4 atomic orbitals → 4 equivalent hybrid orbitals s + px + py + pz → 4 sppp = 4 sp3 • Orbitals have two lobes (unsymmetrical) • Orbitals arrange in space with larger lobes away from one another (tetrahedral shape) • Each hybrid orbital holds 2e-

  7. Formation of methane • The sp3 hybrid orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds • Each C–H bond has a strength of 438 kJ/mol and length of 110 pm • Bond angle:each H–C–H is 109.5°, the tetrahedral angle.

  8. Motivation for hybridization? • Better orbital overlap with larger lobe of sp3 hybrid orbital then with unhybridized p orbital • Stronger bond • Electron pairs farther apart in hybrid orbitals • Lower energy

  9. Another example: ethane • C atoms bond to each other by overlap of an sp3 orbital from each • Three sp3 orbitals on each C overlap with H 1s orbitals • Form six C–H bonds • All bond angles of ethane are tetrahedral

  10. Both methane and ethane have only single bonds • Sigma (s) bonds • Electron density centered between nuclei • Most common type of bond • Pi (p) bonds • Electron density above and below nuclei • Associated with multiple bonds • Overlap between two p orbitals • C atoms are sp2 or sp hybridized

  11. Bond rotation • Single (s) bonds freely rotate • Multiple (p) bonds are rigid

  12. sp2 hybridization 4 atomic orbitals → 3 equivalent hybrid orbitals + 1 unhybridized p orbital s + px + py + pz → 3 spp + 1 p = 3 sp2 + 1 p • Shape = trigonal planar (bond angle = 120º) • Remaining p orbital is perpendicular to the plane

  13. Formation of ethylene (C2H4) • Two sp2-hybridized orbitals overlap to form a s bond • Two sp2 orbitals on each C overlap with H 1s orbitals • Form four C–H bonds • p orbitals overlap side-to-side to form a  bond • sp2–sp2s bond and 2p–2p bond result in sharing four electrons and formation of C-C double bond • Shorter and stronger than single bond in ethane

  14. sp hybridization 4 atomic orbitals → 2 equivalent hybrid orbitals + 2 unhybridized p orbital s + px + py + pz → 2 sp + 2p • Shape = linear (bond angle = 180º) • Remaining p orbitals are perpendicular on y-axis and z-axis

  15. Formation of acetylene (C2H2) • Two sp-hybridized orbitals overlap to form a s bond • One sp orbital on each C overlap with H 1s orbitals • Form two C–H bonds • p orbitals overlap side-to-side to form two  bonds • sp–sps bond and two p–p bonds result in sharing six electrons and formation of C-C triple bond • Shorter and stronger than double bond in ethylene

  16. Summary of Hybridization

  17. Hybridization of Heteroatoms • Same theory • Look at number of e- groups to determine hybridization • Lone pairs will occupy hybrid orbital • Ammonia: • N’s orbitals (sppp) hybridize to form four sp3 orbitals • One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H • H–N–H bond angle is 107.3° • Water • The oxygen atom is sp3-hybridized • The H–O–H bond angle is 104.5°

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