Atomic Structure

Atoms and Isotopes

The atom is the fundamental unit of matter, composed of protons, neutrons, and electrons. Each element is defined by its atomic number, which equals the number of protons in its nucleus. The mass number is the sum of protons and neutrons.

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

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

• Isotopes: Atoms of the same element with different numbers of neutrons.

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

• Ion: An atom or molecule 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: D+ (Deuterium ion): 1 proton, 1 neutron, 0 electrons. 14C2- (Carbon ion): 6 protons, 8 neutrons, 8 electrons.

Electron Configuration Principles

Electrons occupy regions of space called orbitals. The arrangement of electrons in an atom follows 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 quantum numbers; each orbital holds a maximum of two electrons with opposite spins.

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

Example: Electron configuration diagrams can be analyzed to determine which principle is being violated.

Electron Configuration

Ground State Electron Configuration

The ground state electron configuration describes the distribution of electrons among the orbitals of an atom using the Aufbau Principle.

• Aufbau Principle: Electrons fill from lowest to highest energy orbitals.

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

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

• Most Electronegative Element: Fluorine (F).

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

Octet Rule

Valence Electrons and Shared Electrons

The Octet Rule states that most main group elements tend to achieve eight electrons in their valence shell through chemical bonding.

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

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

Example: In H3COH, oxygen has 6 valence electrons and 2 shared electrons, achieving an octet.

Formal Charge

Calculating Formal Charge

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

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

• The sum of formal charges equals the overall charge of the molecule or ion.

Formula:

Lewis Dot Structures

Drawing Lewis Structures

Lewis Dot Structures represent the arrangement of valence electrons in molecules.

1. Determine total valence electrons.

2. Place the least electronegative atom in the center (except H and F).

3. Add electrons to surrounding atoms to complete octets.

4. Place remaining electrons on the central atom.

5. If octets are not complete, form double or triple bonds.

6. Check formal charges for stability.

Example: COCl2 Lewis structure.

Resonance Structures

Definition and Representation

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

• Movement of electrons occurs in pi bonds or lone pairs.

• Double-sided arrows indicate resonance.

• The actual structure is a resonance hybrid.

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 bonds and lone pairs around the central atom.

Example: HCN has a linear geometry and sp hybridization.

Molecular Polarity

Polar and Nonpolar Molecules

Molecular polarity arises from differences in electronegativity and molecular geometry.

• Nonpolar Molecule: Hydrocarbons or molecules with perfect symmetry and no lone pairs on the central atom.

• Polar Molecule: Molecules with an asymmetric shape or lone pairs on the central atom.

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

Functional Groups

Definition and Classification

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

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

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

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

Example: Organic molecules contain both carbon and hydrogen. Hydrocarbons contain only carbon and hydrogen.

Additional info:

• Some content is inferred for completeness, such as specific examples and expanded definitions.

• Tables are reconstructed to summarize block classification, formal charge calculation, hybridization, molecular polarity, and functional groups.