Structure, Bonding, and Polarity in Organic Molecules

Chapter 1: Structure and Bonding

1.11 Describing Chemical Bonds: Molecular Orbital Theory

Molecular Orbital (MO) Theory provides a quantum mechanical explanation for covalent bond formation. It describes how atomic orbitals from different atoms combine to form molecular orbitals, which can be occupied by electrons to create bonds.

• Molecular Orbitals: When two hydrogen 1s atomic orbitals combine, they form two molecular orbitals: a lower-energy bonding MO and a higher-energy antibonding MO.

• Bonding MO: This orbital is lower in energy and is filled with electrons, stabilizing the molecule.

• Antibonding MO: This orbital is higher in energy and remains unfilled in the ground state of H2.

• Node: A region where the probability of finding an electron is zero, present in the antibonding MO.

Example: In H2, the two electrons occupy the bonding MO, resulting in a stable covalent bond.

1.12 Drawing Chemical Structures

Organic chemists use several conventions to represent molecules, ranging from detailed to highly simplified forms. Understanding these conventions is essential for interpreting and drawing organic structures.

• Condensed Structures: Carbon–hydrogen and carbon–carbon single bonds are not explicitly shown. For example, CH3 indicates a carbon with three hydrogens.

• Skeletal Structures: Even simpler than condensed structures, skeletal structures omit carbon and hydrogen atoms bonded to carbon. Each intersection or end of a line represents a carbon atom, and hydrogens are implied to complete carbon's valence.

• Rules for Skeletal Structures:

  1. Carbon atoms are not shown; they are implied at line ends and intersections.

  2. Hydrogen atoms bonded to carbon are not shown; their number is inferred by carbon's valence (4).

  3. Atoms other than carbon and hydrogen are explicitly shown.

Example: 2-Methylbutane can be represented as a full structure, a condensed structure (CH3CH2CH(CH3)2), or a skeletal structure.

Table 1.3: Line-bond and Skeletal Structures for Some Compounds

Compound	Line-bond structure	Skeletal structure
Isoprene, C5H8	All atoms and bonds shown	Lines only, vertices and ends represent carbons
Methylcyclohexane, C7H14	All atoms and bonds shown	Hexagon with a line for methyl group
Phenol, C6H6O	All atoms and bonds shown	Benzene ring with -OH group

Chapter 2: Polar Covalent Bonds; Acids and Bases

2.1 Polar Covalent Bonds and Electronegativity

Most chemical bonds are not purely ionic or covalent but exist on a continuum. Polar covalent bonds occur when electrons are shared unequally between atoms, resulting in partial charges.

• Bond Polarity: Arises from differences in electronegativity (EN), the ability of an atom to attract shared electrons.

• Partial Charges (δ+ and δ−): The more electronegative atom acquires a partial negative charge (δ−), while the less electronegative atom becomes partially positive (δ+).

• Electronegativity Trends: Fluorine is the most electronegative element (EN = 4.0), cesium the least (EN = 0.7). Metals have low EN, nonmetals (especially on the right of the periodic table) have high EN. Carbon's EN is intermediate (2.5).

Example: In methanol (CH3OH), the C–O bond is polar covalent due to the difference in EN between carbon and oxygen.

Inductive Effect

The inductive effect refers to the shifting of electrons in a σ bond in response to the electronegativity of nearby atoms. This effect can stabilize or destabilize molecules depending on the direction of electron shift.

• Electron Donating: Metals like lithium and magnesium donate electrons inductively.

• Electron Withdrawing: Nonmetals like oxygen and nitrogen withdraw electrons inductively.

2.2 Polar Covalent Bonds and Dipole Moments

The dipole moment (μ) quantifies the separation of charge in a molecule and is a measure of molecular polarity. It is calculated as the product of the magnitude of the charge (Q) and the distance (r) between the charges.

• Formula:

$\mu = Q \times r$

• Units: Dipole moments are measured in debyes (D), where 1 D = $3.336 \times 10^{-30}$ C·m.

• Example Calculation:

$\mu = (1.60 \times 10^{-19} \text{ C})(100 \times 10^{-12} \text{ m}) \left( \frac{1 \text{ D}}{3.336 \times 10^{-30} \text{ C} \cdot \text{m}} \right) = 4.80 \text{ D}$

Table 2.1: Dipole Moments of Some Compounds

Compound	Dipole moment (D)	Compound	Dipole moment (D)
NaCl	9.00	NH3	1.47
CH2O	2.33	CH3NH2	1.31
CH3Cl	1.87	CO2	0
H2O	1.85	CH3CH3	0
CH3OH	1.70	Benzene	0
CH3SH	1.52

Example: Water (H2O) has a dipole moment of 1.85 D, indicating significant polarity, while carbon dioxide (CO2) has a dipole moment of 0 D due to its linear, symmetrical structure.