THE ELECTRON CONFIGURATION

Ground State Electron Configuration

Electron configuration describes the distribution of electrons among the atomic orbitals (1s, 2s, 2p, etc.) of an atom, following specific rules.

• Aufbau Principle: Electrons fill the lowest energy orbitals first before occupying higher energy orbitals.

• Condensed Electron Configuration: Uses the previous noble gas to simplify the notation for electron configuration.

Example: Write the ground state and condensed electron configurations for phosphorus (Z = 15).

Additional info: The electron configuration for phosphorus is ; condensed: .

ELECTRONEGATIVITY

Definition and Periodic Trend

Electronegativity (EN) is a measure of an atom's ability to attract electrons in a chemical bond.

• Periodic Trend: Electronegativity increases from left to right across a period and increases going up a group in the periodic table.

Example: The most electronegative Group 7A element is chlorine (Cl).

Additional info: Fluorine (F) is the most electronegative element overall.

OCTET RULE

Valence Electrons and Shared Electrons

The octet rule states that atoms tend to form compounds in ways that give them eight valence electrons, achieving a noble gas configuration.

• Valence Electrons: Electrons in the outermost shell of an atom, involved in bonding.

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

Example: In H3COH (methanol), oxygen has 6 valence electrons and forms 2 shared (bonding) pairs, completing its octet.

FORMAL CHARGE

Definition and Calculation

Formal charge is used to determine the most likely Lewis structure for a molecule by assigning charges to atoms based on electron distribution.

• Formula:

• Only allowable formal charges are -1, 0, or +1 for most main group elements.

• The sum of all formal charges in a molecule equals the overall charge.

Example: Calculate the formal charges for each atom in the thiocyanate ion (NCS-).

DRAWING LEWIS DOT STRUCTURES

Steps for Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons among atoms in a molecule.

1. Count total valence electrons.

2. Place the least electronegative atom in the center (except hydrogen).

3. Connect atoms with single bonds.

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

5. If octets are incomplete, form double or triple bonds as needed.

6. Check formal charges to ensure the best structure.

Example: Draw the Lewis structure for COCl2.

RESONANCE STRUCTURES

Definition and Representation

Resonance structures are two or more valid Lewis structures for a molecule that differ only in the placement of electrons.

• Resonance involves the movement of pi electrons or lone pairs.

• Double-sided arrows () indicate resonance between structures.

• The actual structure is a resonance hybrid, a composite of all resonance forms.

Example: Draw all resonance structures for the nitrate ion, NO3-.

HYBRIDIZATION

Electron Groups and Hybrid Orbitals

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

• Electron Groups: Number of atoms bonded to the central atom plus lone pairs.

Example: Determine the hybridization and number of unhybridized orbitals in HCN.

MOLECULAR POLARITY

Polar and Nonpolar Molecules

Molecular polarity arises from the distribution of electron density in a molecule.

• Nonpolar Molecule: Has a symmetrical shape and equal sharing of electrons.

• Polar Molecule: Has an asymmetrical shape or unequal sharing of electrons, resulting in a dipole moment.

Example: Is nitrogen trifluoride (NF3) polar or nonpolar? (Answer: Polar)

FUNCTIONAL GROUPS

Definition and Types

A functional group is a specific group of atoms within a molecule responsible for its characteristic reactions.

• Hydrocarbons: Alkanes, alkenes, alkynes, aromatic compounds (benzene).

• With Carbonyls: Aldehyde, ketone, acid chloride, amide, carboxylic acid, ester.

• Without Carbonyls: Alkyl halide, amine, alcohol, ether, thiol.

Additional info: Recognizing functional groups is essential for predicting chemical reactivity.

ATOMIC STRUCTURE AND ISOTOPES

Atomic Number, Mass Number, and Isotopes

The atom is the basic unit of matter, consisting of protons, neutrons, and electrons.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons.

• Isotopes: Atoms with the same number of protons but different numbers of neutrons.

Example: Hydrogen isotopes: Protium (1H), Deuterium (2H), Tritium (3H).

IONS

Cations and Anions

Ions are atoms or molecules with a net electric charge due to the loss or gain of electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Example: Proton (H+), Hydride (H-).

PRINCIPLES OF ELECTRON CONFIGURATION

Aufbau Principle, Pauli Exclusion Principle, Hund's Rule

• Aufbau Principle: Electrons fill orbitals starting with the lowest energy first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; each orbital holds a maximum of two electrons with opposite spins.

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

Additional info: Violations of these principles can be identified in electron configuration diagrams.