Atomic Structure

Atoms and Isotopes

Atoms are the fundamental units of matter, consisting of protons, neutrons, and electrons. The atomic number (Z) is the number of protons in the nucleus, while the mass number (A) is the sum of protons and neutrons.

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

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

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

Example: Hydrogen has three isotopes: protium (1H), deuterium (2H), and tritium (3H).

• Ions: Atoms that have gained or lost electrons. Cations are positively charged (lost electrons), anions are negatively charged (gained electrons).

Electron Configuration Principles

Electrons occupy orbitals according to three main principles:

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

• Pauli Exclusion Principle: No two electrons in the same 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 fill degenerate (equal energy) orbitals singly before pairing up.

Electron Configuration

Ground State Electron Configuration

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

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

• Condensed Electron Configuration: Uses the previous noble gas as a starting point, then adds the remaining electrons.

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

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

Octet Rule and Valence Electrons

Octet Rule

The octet rule states that main group elements tend to gain, lose, or share electrons to achieve eight valence electrons, resembling the electron configuration of a noble gas.

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

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

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

Formal Charge

Calculating Formal Charge

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

• Formula:

Formal Charge = Valence Electrons − (Nonbonding Electrons + ½ Bonding Electrons)

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

• Acceptable formal charges are typically −1, 0, or +1.

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

Lewis Dot Structures

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 for the 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 (phosgene).

Resonance Structures

Resonance in Polyatomic Ions

Some molecules or ions have more than one valid Lewis structure, called resonance structures. These differ only in the placement of electrons, not atoms.

• Resonance structures are connected by double-headed arrows.

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

• Resonance occurs when there are delocalized electrons (usually in π bonds or lone pairs).

Example: The nitrate ion (NO3−) has three resonance structures, each with a different N–O double bond.

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 (hydrogen cyanide) has a linear geometry and sp hybridization at the central carbon.

Molecular Polarity

Polarity of Molecules

Molecular polarity arises from the distribution of electron density and the shape of the 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.

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

Functional Groups in Organic Chemistry

Definition and Types

Functional Group: A specific group of atoms within a molecule responsible for characteristic chemical reactions.

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

• 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

Organic Molecules and Hydrocarbons

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.

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

Example: Methane (CH4) is a hydrocarbon; ethanol (C2H5OH) is an organic molecule but not a hydrocarbon.

Summary Table: Principles of Electron Configuration

Additional info: Some content was inferred and expanded for clarity and completeness, including full definitions, examples, and tables for hybridization and molecular polarity.