Atomic Structure and Electron Configuration

Atomic Structure

The atom is the fundamental unit of matter, composed of protons, neutrons, and electrons. Understanding atomic structure is essential for predicting chemical behavior.

• Atomic Number (Z): The number of protons in the nucleus of an atom. Determines the element's identity.

• Mass Number (A): The total number of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge.

• Cations: Positively charged ions (fewer electrons than protons).

• Anions: Negatively charged ions (more electrons than protons).

Example: Hydrogen Isotopes

• Protium (¹H): 1 proton, 0 neutrons

• Deuterium (²H): 1 proton, 1 neutron

• Tritium (³H): 1 proton, 2 neutrons

Example: Hydrogen Ions

• Proton (H⁺): 1 proton, 0 electrons

• Hydride (H⁻): 1 proton, 2 electrons

Electron Configuration

Electron configuration describes the distribution of electrons in atomic orbitals. It is governed by three main principles:

• Aufbau Principle: Electrons fill the lowest energy orbitals 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.

Ground State Electron Configuration: The arrangement of electrons in the lowest possible energy state.

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

Example: Phosphorus (Z = 15)

• Ground State: 1s2 2s2 2p6 3s2 3p3

• Condensed: [Ne] 3s2 3p3

Periodic Trends

Electronegativity

Electronegativity (EN): 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.

• Most Electronegative Element: Fluorine (F)

Example: The most electronegative Group 7A element is Cl (chlorine), but among all elements, F (fluorine) is the highest.

Chemical Bonding and the Octet Rule

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, similar to noble gases.

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

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

• Octet: 8 electrons in the valence shell (for most main group elements).

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

Formal Charge

Formal charge helps determine the most likely Lewis structure for a molecule.

• Formula:

• Sum of formal charges in a molecule equals the overall charge.

• Only allowable formal charges: -1, 0, +1.

Example: For the thiocyanate ion (NCS-), calculate formal charges for each atom using the formula above.

Lewis Dot 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

Some molecules have more than one valid Lewis structure, called resonance structures. The actual molecule is a resonance hybrid of these forms.

• Resonance involves the movement of electrons (not atoms).

• Double-headed arrows (↔) indicate resonance between structures.

• The resonance hybrid is a weighted average of all resonance forms.

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

Hybridization and Molecular Geometry

Hybridization

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: HCN has 2 electron groups around the central carbon, so it is sp hybridized.

Molecular Polarity

Molecular Polarity

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

• Nonpolar Molecule: Has a symmetrical (perfect) shape and even charge distribution.

• Polar Molecule: Has an asymmetrical shape or uneven charge distribution.

Example: Nitrogen trifluoride (NF3) is polar due to its trigonal pyramidal shape and lone pair on nitrogen.

Organic Chemistry: Functional Groups and Molecules

Functional Groups

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

Example: Identify functional groups in given molecules and classify them as hydrocarbons or containing heteroatoms.

Organic Molecules

Organic chemistry studies molecules containing carbon and hydrogen, often with other elements such as oxygen, nitrogen, sulfur, and halogens.

• Hydrocarbons: Molecules containing only carbon and hydrogen (alkanes, alkenes, alkynes, aromatics).

• Organic Molecule: Any molecule containing both carbon and hydrogen.

Example: Determine which molecules are organic and which are hydrocarbons from a given set.

Summary Table: Principles of Electron Configuration

Additional info: These notes provide a concise yet comprehensive overview of foundational topics in general chemistry, including atomic structure, periodic trends, bonding, molecular geometry, and basic organic chemistry. They are suitable for exam preparation and as a reference for problem-solving in introductory college chemistry courses.