Atomic Structure and Electron Configuration

Ground State Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. The Aufbau Principle states that electrons fill orbitals starting from the lowest energy level to higher ones.

• Aufbau Principle: Electrons occupy the lowest energy orbitals first.

• Condensed Electron Configuration: Uses the previous noble gas to abbreviate the configuration.

Example: Phosphorus (Z = 15): Ground state: Condensed:

Electronegativity

Electronegativity (EN) measures an atom's ability to attract electrons in a chemical bond. It increases across a period (left to right) and decreases down a group.

• Periodic Trend: Highest EN is at the top right of the periodic table (excluding noble gases).

• Example: Fluorine (F) is the most electronegative element in Group 7A.

Chemical Bonding and Lewis Structures

Octet Rule

Atoms tend to achieve eight electrons in their valence shell (octet) through chemical bonding.

• Valence Electrons: Electrons in the outermost shell involved in bonding.

• Shared Electrons: Electrons shared between atoms in a bond.

Example: In ethanol (CH3CH2OH), oxygen has 8 electrons (6 valence + 2 shared).

Formal Charge

Formal charge helps determine the most stable Lewis structure.

• Formula:

• Only -1, 0, or +1 are typical for atoms in stable molecules.

Example: In NH3, nitrogen has a formal charge of 0.

Drawing Lewis Dot Structures

Lewis structures represent the arrangement of electrons in molecules.

1. Count total valence electrons.

2. Place the least electronegative atom in the center.

3. Connect atoms with single bonds.

4. Complete octets for outer atoms, then central atom.

5. Assign formal charges as needed.

Example: COCl2 Lewis structure follows these steps.

Resonance Structures

Some molecules have multiple valid Lewis structures, called resonance structures. Electrons move between atoms via double-headed arrows.

• Resonance hybrid: The true structure is a weighted average of all resonance forms.

• Major contributors: Structures with full octets and minimal formal charges.

Example: NO2- has two resonance structures.

Hybridization and Molecular Geometry

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

Example: HCN: Carbon is sp hybridized.

Molecular Polarity

Molecular polarity depends on the shape and distribution of charge.

• Nonpolar Molecule: Symmetrical shape, no net dipole moment.

• Polar Molecule: Asymmetrical shape or lone pairs, net dipole moment.

Example: NF3 is polar due to lone pairs on nitrogen.

Functional Groups in Organic Chemistry

Functional Groups

Functional groups are specific groups of atoms within molecules responsible for characteristic chemical reactions.

• Examples: Alcohols (-OH), Aldehydes (-CHO), Ketones (C=O), Carboxylic acids (-COOH), Amines (-NH2), Ethers (R-O-R).

Atomic Structure and Quantum Mechanics

Atomic Structure

The atom consists of a nucleus (protons and neutrons) surrounded by electrons.

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

• Isotopes: Atoms with same Z but different A.

Example: Hydrogen isotopes: , (deuterium), (tritium).

Quantum Mechanics and Wave Function

Electrons behave as both particles and waves. The Heisenberg Uncertainty Principle states that position and momentum cannot both be known exactly.

• Orbitals: Regions of space where electrons are likely to be found (s, p, d, f).

• Wave Function (): Describes the probability of finding an electron.

Molecular Orbital Theory

Atomic orbitals combine to form molecular orbitals, which can be bonding or antibonding.

• Bonding Orbitals: Constructive interference, lower energy.

• Antibonding Orbitals: Destructive interference, higher energy.

Sigma and Pi Bonds

Sigma () and Pi () Bonds

Example: Ethylene (C2H4) has a double bond: 1 , 1 .

Octet Rule and Bonding Preferences

Octet Rule

Atoms are most stable when they have eight electrons in their valence shell, similar to noble gases.

• Atoms may form bonds or ions to achieve an octet.

• Exceptions: Hydrogen (2 electrons), Boron (6 electrons), expanded octets for elements in period 3 or higher.

Bonding Preferences

Structural Formulas and Isomerism

Condensed Structural Formula

Condensed formulas represent molecules using only text, showing connectivity but not all bonds explicitly.

Skeletal Structure

Skeletal structures simplify organic molecules by omitting carbon and hydrogen atoms explicitly.

• Each vertex represents a carbon atom.

• Hydrogens attached to carbon are implied.

Constitutional Isomers

Constitutional isomers have the same molecular formula but different connectivity.

• Identify by comparing atom connectivity and arrangement.

Resonance Structures and Hybrids

Resonance Structures

Resonance structures show different possible arrangements of electrons in a molecule.

• Curved arrows indicate electron movement.

• Major contributors have full octets and minimal formal charges.

Resonance Hybrid

The resonance hybrid is the true structure, representing a weighted average of all resonance forms.

Hybridization and Molecular Geometry

Hybridization Summary

Molecular Geometry (VSEPR Theory)

VSEPR theory predicts molecular shapes based on electron pair repulsion.

• Linear: 180° bond angle

• Trigonal planar: 120° bond angle

• Tetrahedral: 109.5° bond angle

Electronegativity and Dipoles

Electronegativity and Bond Polarity

Bond polarity arises from differences in electronegativity between atoms.

• Dipole Moment (): (charge × distance)

• Polar bonds have unequal sharing of electrons.

Molecular Dipoles

Net dipole moments depend on both bond polarity and molecular geometry.

• Solvents can be polar (e.g., water) or nonpolar (e.g., hexane).

Degree of Unsaturation

HD (Hydrogen Deficiency Index)

HD indicates the number of rings and/or multiple bonds in a molecule.

• Formula:

• C = number of carbons, N = number of nitrogens, H = number of hydrogens, X = number of halogens

Summary Table: Bond Types

Additional info:

• These notes cover foundational topics in General Chemistry, including atomic structure, electron configuration, chemical bonding, molecular geometry, resonance, and basic organic functional groups.

• Practice problems and examples are provided throughout to reinforce key concepts.