Atomic Structure and Isotopes

Atoms and Subatomic Particles

An atom is the basic unit of matter, composed of protons, neutrons, and electrons.

• 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).

Example: Hydrogen has three isotopes: Protium (1 proton), Deuterium (1 proton, 1 neutron), and Tritium (1 proton, 2 neutrons).

• Ions: Atoms that have gained or lost electrons.

• Cations: Positively charged ions (lost electrons).

• Anions: Negatively charged ions (gained electrons).

Example: The proton (H+) and hydride (H-).

Electron Configuration 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.

The Electron Configuration

Ground State Electron Configuration

The ground state electron configuration describes the distribution of electrons among the orbitals (1s, 2s, 2p, etc.) of an atom using the Aufbau Principle.

• Electrons fill from lower to higher energy orbitals: ...

• Condensed Electron Configuration: Uses the previous noble gas to simplify notation. For example, phosphorus (Z = 15):

Periodic Table Blocks: s-block (Groups 1-2), p-block (Groups 13-18), d-block (transition metals), f-block (lanthanides/actinides).

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.

• Fluorine is the most electronegative element.

Example: Among Group 7A elements, Cl is more electronegative than Br or I.

Octet Rule

Valence and Shared Electrons

The octet rule states that main group elements tend to form bonds until they are surrounded by eight valence electrons.

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

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

• Octet: 8 electrons (sum of valence and shared electrons).

Example: In H3COH, oxygen has 6 valence electrons and forms 2 bonds to complete its octet.

Formal Charge

Definition and Calculation

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

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

• Formula:

Formal Charge = Valence Electrons - (Nonbonding Electrons + 1/2 Bonding Electrons)

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

Lewis Dot Structures

Drawing Rules

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.

Exceptions: Hydrogen only needs 2 electrons; halogens usually form one bond.

Resonance Structures

Definition and Representation

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

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

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

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

• Resonance is shown by placing a dashed line where delocalized electrons are shared.

Example: Nitrate ion (NO3-) has three resonance structures.

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: HCN has a linear geometry and sp hybridization.

Molecular Polarity

Polar and Nonpolar Molecules

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

• Nonpolar Molecule: Has a symmetrical (perfect) shape and no net dipole moment.

• Polar Molecule: Has an asymmetrical shape or contains polar bonds that do not cancel out.

• Central atom with lone pairs often leads to polarity.

Example: Nitrogen trifluoride (NF3) is polar due to the lone pair on nitrogen.

Functional Groups

Definition and Types

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

• Hydrocarbons: Alkanes, alkenes, alkynes, aromatic rings.

• With Carbonyls: Aldehydes, ketones, carboxylic acids, esters, amides, acid chlorides.

• Without Carbonyls: Alcohols, ethers, amines, thiols, alkyl halides.

Example: Alcohols contain an -OH group; carboxylic acids contain a -COOH group.

Organic Chemistry Basics

Definition and Examples

Organic chemistry is the study of carbon-containing compounds, especially those found in living organisms.

• Organic molecules contain both carbon and hydrogen.

• Hydrocarbons: Compounds containing only carbon and hydrogen.

Example: Identifying organic molecules and hydrocarbons from a list of structures.

Summary Table: Principles of Electron Configuration

Additional info: These notes cover 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.