Organic Chemistry Notes

Naming Organic Compounds

• Systematic nomenclature for naming organic compounds is introduced by IUPAC.
• Names are divided into three parts:
  • Describe the hydrocarbon framework.
  • Describe the functional groups.
  • Indicate where the functional groups are attached to the skeleton.
• The IUPAC naming system consists of a suffix, parent chain, and prefix.

Hydrocarbon Framework

• Number of carbon atoms in the chain and the name of the parent hydrocarbon:
  • 1: methane (CH_4)
  • 2: ethane (CH3-CH3)
  • 3: propane (CH3CH2CH_3)
  • 4: butane
  • 5: pentane
  • 6: hexane
  • 7: heptane
  • 8: octane
  • 9: nonane
  • 10: decane
• Cyclic derivatives:
  • cyclopropane
  • cyclohexane

Functional Groups

• The name of a functional group is combined with the relevant hydrocarbon framework.
  • Methane + alcohol → methanol (CH_3OH)
  • Ethane + aldehyde → ethanal
  • Cyclohexane + ketone → cyclohexanone
• Examples:
  • Butanoic acid
  • Heptanoyl chloride
  • Propene
  • Ethene

Locating Functional Groups

• Numbers are used to locate functional groups; carbon atoms are counted from one end of the framework, using the lowest possible number.
• Example: propan-1-ol
  • 3-carbon chain (propane)
  • Alcohol functionality at C1
• Example: 2-aminobutane
  • 4-carbon chain (butane)
  • Amine functionality at C2
• Application to branched hydrocarbon frameworks: treat as a derivative of the longest carbon chain.
  • 2-methylbutane
    • 4-carbon chain (butane)
    • Methyl functionality at C2
  • 2-ethylpropane (incorrect)
    • 3-carbon chain (propane)
    • Ethyl functionality at C2

Structure and Bonding

Fundamental Questions

• Why does carbon form 4 covalent bonds with H, N, O, S, Si (N forms 3, O forms 2, etc.)?
• Why do molecules adopt specific three-dimensional shapes?
• To answer these questions, models for structure and bonding in organic chemistry are needed.

Atomic Orbitals

• Electronic configuration of elements is based upon atomic orbitals.
• Quantum numbers determine the shape of the periodic table.
  • Principal Quantum number (shells)
    • 1: 1s (Each set of s holds 2 electrons)
    • 2: 2s, 3 x 2p (Each set of p holds 6 electrons)
    • 3: 3s, 3 x 3p, 5 x 3d (Each set of d holds 10 electrons)

Electronegativity

• Electronegativity is the measure of the relative ability of an atom to attract valence electrons.
  • A higher number indicates greater electronegativity.
• Trends in electronegativity:
  • Increases across a period (increased nuclear charge).
  • Decreases down a group.

Lewis Scheme

• The simplest model for bonding.
• "Bonding is the achievement of completely filled (or empty) valence shells by sharing (covalent bonds) or transfer of electrons (ionic bonds)."
  • Ionic Bonds: Formed between atoms at opposite ends of the periodic table with very different electronegativities; achieved by transfer of one or more electrons (e.g., LiF = Li^+F^-, Li = 1.0, F = 4.0).
  • Covalent Bonds: Common between atoms of similar electronegativity; results in a “sharing” of electrons (e.g., H-H).
• Bond Polarisation:
  • In H-F, the shared pair is not equally distributed but displaced toward F due to large difference in nuclear charge between H and F.
  • Results in bond polarisation (inductive, through-bond displacement) (H^{\delta+} - F^{\delta-}).

Carbon and Covalent Compounds

• Carbon is in the center of the periodic table and has no great differences in electronegativity with many common elements.
• It is difficult to form stable closed-shell compounds by ionization to C^{4+} or C^{4-}.
• Carbon therefore forms covalent compounds (e.g., methane CH4, ammonia NH3).
• Lewis Scheme:
  • A line = a shared pair = covalent bond.
  • An unshared or lone pair is represented by a pair of electrons.
• The Lewis Scheme does not tell us about the shape of molecules or how bonds are formed.

VSEPR and Molecular Orbitals

• VSEPR allows a simple method to predict structures based on the number of electron pairs in the outer valence shell and their repulsion.
• To understand why specific shapes are adopted, we need to consider the atomic orbitals that make up the molecule and how they combine to form molecular orbitals.

Linear Combinations of Atomic Orbitals

• Electrons in atomic orbitals are best described by wavefunctions.
• They can combine together either constructively (in-phase) or destructively (out-of-phase) like waves.

Atomic Orbital Overlap

• Atomic orbitals can combine in-phase or out-of-phase.
• s orbitals: spherically symmetric (1s, 2s, …).
• p orbitals: three p orbitals in each quantum shell, oriented along the x, y, and z axes, with two “lobes” of electron density separated by a nodal surface.

Representing Orbitals

• A simple way of representing these orbitals is to take a cross-section, giving a spherical s orbital and an “hour-glass” p orbital.
• At a nodal surface, there is a change in sign of the wavefunction.

Combining Atomic Orbitals - Homonuclear Diatomics

• Model H_2: Combine two 1s orbitals in-phase (add together) or out-of-phase (they cancel) in the space between the nuclei.
• In-phase combination: leads to increased electron density between nuclei and a bonding orbital.
• Out-of-phase combination: leads to decreased electron density between nuclei (anti-bonding).
• MO arising from in-phase combination is lower in energy than out-of-phase combinations.

Energy Level Diagram

• The two electrons from the constituent atoms are placed in the lowest energy MO, the bonding orbital, which results in net bonding between the atoms.
• Combining s orbitals gives rise to a cylindrically symmetric MO along the intermolecular axis, known as a sigma bond (\sigma).
• Anti-bonding orbitals are given a * (\sigma^*) to distinguish them.

Helium Dimer

• Consider He_2 - each He atom has two electrons (1s^2) - after constructing a MO diagram, both the bonding and anti-bonding orbitals are full, resulting in no net bonding.
• He2 does not exist, but what about He2^+?

Molecular Orbitals from Overlap of p Orbitals

• Electronic configuration of C = 1s^2 2s^2 2p^2
• On a carbon atom, there are three mutually perpendicular p orbitals with direction along the x, y, and z axes.
• End-on overlap results in a pair of MO (bonding / anti-bonding) that are \sigma symmetric.

Pi Orbitals

• Overlap of the two remaining p orbitals form MO that are not symmetric about the nuclear axis (they have a phase change) - this type of orbital is known as a \pi orbital.
• We can combine them either in-phase (bonding) or out-of-phase (anti-bonding).
• There are 2 mutually perpendicular pairs of p orbitals that can combine in this way, a pair of degenerate (equal energy) mutually perpendicular \pi bonding MOs and a pair of degenerate mutually perpendicular \pi^* MOs are formed.

Hybridisation

Hybridisation of Atomic Orbitals

• Allows the mixing of atomic orbitals on the same atom to create hybrid atomic orbitals that point directly to the atoms we wish to construct bonds with.
• For methane (CH_4), which adopts a tetrahedral shape, we need to use 4 valence AO on C (1 x 2s and 3 x 2p) to form 4 hybrid atomic orbitals that each interact with a H atom to give methane.

sp3 Hybrid Orbitals

• We combine the 2s and 3 x 2p orbitals in the valence shell of C to create 4 x sp^3 hybrid orbitals (these are an equal mix of 1 x s and 3 x p orbitals = 25% s; 75% p character).
• Carbon has 4 equivalent sp^3 hybrid orbitals that point to the corner of a tetrahedron.
• Each sp^3 hybrid orbital is considered to contain a single electron.

Formation of Methane

• 4 equivalent sp^3 hybrid orbitals point to the corner of a tetrahedron.
• They each interact with a 1s orbital of hydrogen to form an in-phase \sigma-bonding MO and an out-of-phase anti-bonding \sigma^*-MO.
• Each C-H \sigma-bond contains two electrons.
• The valence electrons fill the bonding MO to form tetrahedral methane.

Summary of Approach

• 1. Hybridise: take 4 AO on C and hybridise to create 4 sp^3 hybrid orbitals each containing one electron.
• 2. Combine with 4 x 1s orbitals to create 4 x C-H \sigma-bonds.

Extensions of Hybridization

Ethane (C2H6)

• Both C atoms are tetrahedral and sp^3 hybridised.
• Two sp^3 hybrid AO from the neighbouring C atoms overlap to form a C-C \sigma bond (and a vacant C-C \sigma^* bond).
• Each remaining sp^3 hybrid AO overlaps with 3 x H 1s atomic orbitals to form C-H \sigma bonds (and vacant C-H \sigma^* bonds).
• Rotation can occur about the C-C \sigma-bond.

Ethene (C2H4)

• Both C atoms are planar.
• Each C atom forms bonds to three other atoms.
• We combine the 2s and 2 x 2p orbitals in the valence shell of C to create 3 x sp^2 hybrid orbitals that are planar (these are an equal mix of 1 x s and 2 x p orbitals = 33% s; 67% p character); leaving an unchanged p orbital.