Remembering General Chemistry: Electronic Structure and Bonding

What Makes Carbon So Special?

Carbon's unique ability to form four covalent bonds by sharing electrons makes it central to organic chemistry. Atoms to the left of carbon in the periodic table tend to lose electrons, while those to the right tend to gain electrons. Carbon, however, shares electrons, allowing for a vast diversity of stable compounds.

The Structure of an Atom

Atomic Structure and Isotopes

An atom consists of a nucleus (containing protons and neutrons) surrounded by an electron cloud. The atomic number equals the number of protons, which defines the element. Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

• Protons: Positively charged particles in the nucleus.

• Neutrons: Neutral particles in the nucleus.

• Electrons: Negatively charged particles in orbitals around the nucleus.

• Isotopes: Same atomic number, different mass numbers (e.g., 12C, 13C, 14C).

The Distribution of Electrons in an Atom

Electronic Configurations and the Periodic Table

Electrons occupy atomic orbitals in shells around the nucleus. The arrangement follows the Aufbau principle (lowest energy orbitals filled first), Pauli exclusion principle (maximum two electrons per orbital), and Hund’s rule (electrons fill degenerate orbitals singly first).

• Aufbau principle: Electrons fill orbitals of lowest energy first.

• Pauli exclusion principle: No more than two electrons per orbital.

• Hund’s rule: Electrons occupy empty degenerate orbitals before pairing.

Stability of Electron Configurations

Atoms are most stable when their outer shell is filled or contains eight electrons (octet rule). Elements in the first column (e.g., Li, Na) lose electrons to achieve a filled shell, while elements on the right (e.g., F, Cl) gain electrons.

• Example: Lithium loses one electron to form Li+; fluorine gains one electron to form F-.

Hydrogen: A Special Case

Hydrogen can either lose its single electron to form a proton (H+) or gain an electron to form a hydride ion (H-).

Bonding and the Octet Rule

Covalent Bond Formation

Atoms can achieve filled outer shells by sharing electrons, forming covalent bonds. The number of bonds an atom forms is determined by the number of electrons needed to complete its octet (or duet for hydrogen).

• Example: Two hydrogen atoms share electrons to form H2; two fluorine atoms share electrons to form F2.

Bonding Patterns of Common Elements

The typical number of bonds formed by common elements in organic chemistry:

• Carbon: 4 bonds

• Nitrogen: 3 bonds, 1 lone pair

• Oxygen: 2 bonds, 2 lone pairs

• Halogens (F, Cl, Br, I): 1 bond, 3 lone pairs

• Hydrogen: 1 bond

Expanded Octets

Elements like phosphorus and sulfur can have more than eight electrons in their valence shell due to available d orbitals.

Bond Polarity and Electronegativity

Nonpolar and Polar Covalent Bonds

Bonds between atoms with similar electronegativities are nonpolar covalent, while those with different electronegativities are polar covalent. The greater the difference, the more polar the bond.

• Nonpolar covalent: Electrons shared equally (e.g., C–C, H–H).

• Polar covalent: Electrons shared unequally (e.g., H–Cl, H–O).

• Ionic: Electrons not shared; atoms held by electrostatic attraction (e.g., Na+Cl-).

Dipole Moments and Electrostatic Potential

The dipole moment of a bond depends on the difference in electronegativity and the distance between charges. Electrostatic potential maps visually represent electron density and polarity in molecules.

Lewis Structures and Formal Charges

Drawing Lewis Structures

Lewis structures show how atoms are bonded and the location of lone pairs and formal charges. The sum of bonds and lone pairs for C, N, O, and halogens typically equals four.

• Formal charge: Calculated as (valence electrons) – (nonbonding electrons + 1/2 bonding electrons).

• Example: Methane (CH4) has no formal charges; methyl cation (CH3+) has a positive charge on carbon.

Molecular Representations

Kekulé, Condensed, and Skeletal Structures

Organic molecules can be represented in several ways:

• Kekulé structures: Show all atoms and bonds, but omit lone pairs.

• Condensed structures: Omit some or all bonds, grouping atoms together.

• Skeletal structures: Show only the carbon skeleton and heteroatoms; hydrogens on carbons are implied.

Atomic and Molecular Orbitals

s and p Atomic Orbitals

Atomic orbitals are regions of space where electrons are likely to be found. The s orbital is spherical, while p orbitals are dumbbell-shaped and have two lobes with opposite phases.

Bond Formation: Sigma (σ) and Pi (π) Bonds

Sigma bonds are formed by end-to-end overlap of orbitals, while pi bonds are formed by side-to-side overlap of p orbitals. Molecular orbitals result from the combination of atomic orbitals.

Hybridization and Molecular Geometry

sp3, sp2, and sp Hybridization

Hybridization explains the observed shapes of molecules:

• sp3 hybridization: Four equivalent orbitals, tetrahedral geometry (e.g., methane, CH4).

• sp2 hybridization: Three equivalent orbitals, trigonal planar geometry (e.g., ethene, C2H4).

• sp hybridization: Two equivalent orbitals, linear geometry (e.g., ethyne, C2H2).

Bond Strength, Bond Length, and s Character

The more s character in a hybrid orbital, the shorter and stronger the bond. For example, sp bonds (50% s character) are shorter and stronger than sp2 (33%) or sp3 (25%).

Summary Table: Typical Bonding Patterns

Learning Objectives

• Write ground-state electronic configurations for elements H through Ca.

• Describe bond polarity and determine dipole directions.

• Represent organic compounds using Lewis, Kekulé, condensed, and skeletal structures.

• Assign lone pairs and calculate formal charges.

• Determine hybridization of C, N, or O from molecular formula.

• Describe how hybridization determines molecular geometry.

• Explain how hybridization affects bond strength and length.

• Describe how bond order affects bond length and strength.