Atomic Structure

Atoms and Isotopes

Atoms are the basic units of matter, consisting of a nucleus (protons and neutrons) surrounded by electrons. The atomic number is equal to the number of protons, while 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: Hydrogen (1 proton), Deuterium (1 proton, 1 neutron), and Tritium (1 proton, 2 neutrons).

• Ion: An atom with a net charge due to loss or gain of electrons.

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

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

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

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.

Example: Electron configuration diagrams showing correct and incorrect applications of these principles.

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:

Periodic Table Blocks

The periodic table is divided into blocks based on the type of orbital being filled:

• s-block: Groups 1A and 2A

• p-block: Groups 3A to 8A

• d-block: Transition metals

• f-block: Lanthanides and actinides

Electronegativity

Definition and Trends

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 and 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.

Formal Charge

Calculating Formal Charge

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

• Formula:

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

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

Example: Calculate formal charges for N, C, and S in the thiocyanate ion (NCS-).

Lewis Dot Structures

Drawing Lewis Structures

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

1. Determine total number of valence electrons.

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

3. Add electrons to complete octets (except hydrogen, which only needs 2 electrons).

4. Place remaining electrons on the central atom.

5. If atoms lack octets, form double or triple bonds.

6. Check formal charges to confirm the best structure.

Example: Draw the Lewis Dot 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.

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

• Double-sided arrows indicate resonance between structures.

• The true 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 bonds and lone pairs around the central atom.

Example: HCN has 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: Molecule with a perfect shape and no lone pairs on the central atom.

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

Example: Nitrogen trifluoride (NF3) is polar due to lone pairs on the central atom.

Functional Groups

Definition and Classification

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

• Hydrocarbons: Alkanes, alkenes, alkynes, aromatic compounds

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

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

Example: Identify organic molecules and hydrocarbons from given structures.

Organic Chemistry Introduction

Organic Molecules and Hydrocarbons

Organic chemistry studies molecules containing carbon and hydrogen, often found in biological systems. Hydrocarbons are organic molecules composed solely of carbon and hydrogen.

• Organic Molecule: Contains both carbon and hydrogen.

• Hydrocarbon: Organic molecule with only carbon and hydrogen.

Example: Identify organic and hydrocarbon molecules from a list.

Summary Table: Key Concepts

Additional info: Some content was expanded for clarity and completeness, including definitions, examples, and tables for functional groups and hybridization.