Atomic Structure and Important Elements

Overview of Atomic Structure

Understanding atomic structure is foundational for organic chemistry, as it determines how atoms bond and interact in molecules.

• Atom: Composed of protons (positive charge), neutrons (neutral), and electrons (negative charge).

• Protons: Define the element and its atomic number.

• Neutrons: Contribute to atomic mass but do not affect chemical behavior.

• Electrons: Occupy orbitals around the nucleus and participate in chemical bonding.

• Valence electrons: Electrons in the outermost shell, crucial for bonding.

Important Elements in Organic Chemistry

Organic chemistry primarily involves a select group of elements.

• "Big 4": Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N).

• Halogens: Group 7A: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I).

• Common heteroatoms: Phosphorus (P), Sulfur (S).

Electronic Structure

Orbitals and Electron Configuration

Electrons are arranged in orbitals, which are grouped into shells at increasing distances from the nucleus.

• Orbitals: Regions of space where electrons are likely to be found.

• Shells: Defined by principal quantum number "n" (n = 1, 2, 3...).

• Order of filling: Electrons fill the lowest energy orbitals first (Aufbau Principle).

• Hund's Rule: Orbitals of equal energy are singly occupied before pairing.

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

Electron Configuration Example

For carbon (C, atomic number 6):

Valence electrons correspond to the group number in the periodic table.

Bond Formation and the Octet Rule

Octet Rule

Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, mimicking the electron configuration of noble gases.

• Transfer or sharing: Atoms will transfer or share electrons to satisfy the octet rule.

• Exceptions: Hydrogen (H) wants 2 electrons; elements beyond the second row (e.g., Al) can expand their octet.

Lewis Structures

Lewis structures visually represent the bonding in covalent molecules.

• Bonds: Represented by lines (pairs of electrons shared between atoms).

• Lone pairs: Nonbonding electrons shown as dots.

• Radicals: Molecules with unpaired electrons.

Example: Water (H2O)

• Oxygen has 6 valence electrons; forms two bonds with hydrogen and has two lone pairs.

Counting Valence Electrons and Charges

• For each negative charge, add one electron; for each positive charge, subtract one electron.

• Charge calculation: Charge = valence electrons - (bonds + nonbonding electrons)

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is an atom's ability to attract electrons in a chemical bond.

• Pauling scale: Fluorine (F) is the most electronegative element.

• Electronegativity increases across a period (left to right) and decreases down a group.

Electronegativity Values (Pauling Scale)

Bond Polarity

Bond polarity arises when atoms with different electronegativities share electrons unequally, resulting in dipole moments.

• Nonpolar covalent bond: Electrons shared equally (e.g., H2).

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

• Ionic bond: Electrons transferred (large electronegativity difference).

Formal Charges

Calculating Formal Charge

Formal charge helps keep track of electron distribution in molecules.

• Formula:

• Used to identify the most stable Lewis structure.

Summary Table: Valence, Bonds, and Lone Pairs for Common Elements

Key Concepts and Rules

• Octet Rule: Atoms strive for 8 electrons in their valence shell.

• Exceptions: Hydrogen (2 electrons), elements with expanded octets.

• Electronegativity: Determines bond polarity and molecular properties.

• Formal Charge: Used to assess stability and reactivity of molecules.

Additional info:

• Some context and examples were inferred to clarify fragmented notes and ensure completeness.

• Tables were reconstructed to summarize key properties and electronegativity values.