Structure Determines Properties

Introduction

The physical and chemical properties of organic molecules are fundamentally determined by their atomic structure and the nature of chemical bonds. Understanding atomic structure, electron configuration, and bonding principles is essential for predicting molecular behavior in organic chemistry.

Atoms, Electrons, and Orbitals

Subatomic Particles and Atomic Symbols

• Proton: Mass = kg; Charge = C

• Neutron: Mass = kg; Charge = C

• Electron: Mass = kg; Charge = C

The atomic number (Z) is the number of protons in the nucleus. The mass number (M) is the total number of protons and neutrons.

Electrons as Waves

• Electrons in atoms and molecules exhibit wave-like behavior rather than behaving solely as particles.

• The behavior of electrons is described by the wavefunction (), which is a solution to the Schrödinger equation:

• The orbital is a region in space where the probability of finding an electron is high.

• The shape of an orbital reflects the probability density at a point , which is related to .

• The energy of an orbital reflects the stability of an electron within it.

Quantum Numbers

Atomic orbitals are defined by four quantum numbers:

• Principal quantum number (n): Determines the energy level of the orbital.

• Orbital quantum number (l): Determines the shape of the orbital (e.g., s, p, d, f).

• Magnetic quantum number (ml): Determines the orientation of the orbital in space.

• Spin quantum number (s): Describes the spin of the electron ( or ).

Pauli Exclusion Principle: No two electrons in the same atom can have the same set of quantum numbers.

s and p Atomic Orbitals

• s orbitals: Spherically shaped, begin at , one per shell, lower in energy than other orbitals in the same shell.

• p orbitals: Dumbbell-shaped, begin at , three per shell (2px, 2py, 2pz), all have the same energy and contain a node at the nucleus.

Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

Ionic Bonds and Electron Transfer

Coulomb’s Law and Ionic Bonding

• Ions are atoms or molecules with electric charge.

• Oppositely charged ions attract each other according to Coulomb’s Law:

• Ionic bonds result from the electrostatic attraction between cations and anions.

• Ionic compounds form tightly packed lattices, are high-melting solids, and are soluble in polar solvents.

Ionic Bonding in Organic Chemistry

• Carbon atoms rarely form ions; thus, ionic bonds are uncommon in organic molecules.

• Ionic bonds do appear in salts of C, N, O, and H, especially when these atoms are nucleophilic.

Covalent Bonds, Lewis Structures, and the Octet Rule

The Lewis Model of Covalent Bonding

• Covalent bonds involve the sharing of electrons between two atoms.

• Atoms share electrons to achieve a more stable electron configuration, often resembling the nearest noble gas.

• Electrons may not be shared equally, resulting in polar covalent bonds.

Lewis Symbols and Structures

• Electrons are represented as dots (for lone pairs) or lines (for shared pairs/bonds).

• Not all electrons are involved in bonding; nonbonding electrons are shown as dots.

• Lewis structures help visualize electron arrangement and bonding in molecules.

Covalent Bonding Examples

• Hydrogen (H2): Each hydrogen shares one electron, achieving a full shell analogous to helium.

• Fluorine (F2): Each fluorine shares one electron, achieving a full shell analogous to neon.

The Octet Rule

• Second-row elements achieve stability by having eight electrons in their valence shell (1s, 2s, 2p).

• Octet Rule: Atoms share electrons until they achieve an octet.

• Shared electrons are counted for both atoms involved in the bond.

For example, oxygen in H2O has eight electrons in its valence shell.

Multiple Bonds

• More than two electrons can be shared between atoms, forming multiple bonds:

• Double bond: Four electrons shared (e.g., ethylene, C2H4).

• Triple bond: Six electrons shared (e.g., ethyne, C2H2).

Polar Covalent Bonds, Electronegativity, and Bond Dipoles

Electronegativity

• Electronegativity is the ability of an atom to attract electrons in a chemical bond.

• Electronegative atoms attract electrons strongly and tend to gain electrons.

• Electropositive atoms attract electrons weakly and may lose electrons.

Periodic Trends in Electronegativity

• Electronegativity increases from left to right across a period and from bottom to top within a group.

Polar Covalent Bonds

• When atoms of different electronegativities share electrons, the bond is polarized.

• The greater the difference in electronegativity, the more polarized the bond.

• Polar bonds have a dipole moment pointing from the positive to the negative end.

Electrostatic Potential Maps

• Electrostatic potential (ESP) maps show the distribution of charge over a molecule.

• Nonpolar molecules have symmetric charge distributions; polar molecules have regions of partial positive and negative charge.

Dipole Moments Table

The following table summarizes dipole moments for selected bonds:

Formal Charge

Definition and Calculation

• Formal charge is the difference between the number of valence electrons in a neutral atom and the number assigned to it in a Lewis structure.

• Formula:

• To calculate FEC: count all unshared electrons and half of all bonding electrons.

• Refer to the periodic table for the group number (valence electrons) of the neutral atom.

Examples of Formal Charge

• Oxygen in H2O: Oxygen has 6 valence electrons; in H2O, it is assigned 6 electrons (4 from lone pairs, 2 from bonds), so FC = 6 - 6 = 0.

• Nitrogen in NH4+: Nitrogen has 5 valence electrons; in NH4+, it is assigned 4 electrons, so FC = 5 - 4 = +1.

Recognizing formal charge patterns is important for understanding reactivity and stability in organic molecules.

Additional info:

• Understanding atomic structure and bonding principles is foundational for all subsequent topics in organic chemistry, including reaction mechanisms and molecular properties.