Atomic Structure

Atoms and Isotopes

The atom is the basic unit of matter, consisting of a nucleus (protons and neutrons) surrounded by electrons.

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

• Mass Number (A): The sum 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: Hydrogen (1 proton), Deuterium (1 proton, 1 neutron), Tritium (1 proton, 2 neutrons).

• Ions: Atoms that have gained or lost electrons. Positively charged ions are called cations; negatively charged ions are called anions.

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

Electron Configuration

Ground State Electron Configuration

Describes the distribution of electrons among atomic orbitals using the Aufbau Principle, Pauli Exclusion Principle, and Hund's Rule.

• 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 fill degenerate orbitals singly before pairing.

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

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

Periodic Table Blocks

• 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): 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 and Shared Electrons

The tendency of main group elements 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 H3C–O–H, oxygen has 6 valence electrons and 2 shared electrons, totaling 8 octet electrons.

Formal Charge

Calculation and Application

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: For the thiocyanate ion (NCS-): Calculate formal charges for N, C, and S using the formula above.

Lewis Dot Structures

Drawing Rules

Lewis structures represent the arrangement of valence electrons in molecules.

1. Count total valence electrons.

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

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

4. Place remaining electrons on the central atom.

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

6. Check formal charges to confirm the best structure.

Example: Draw the Lewis structure for COCl2.

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 between structures.

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

• Dashed lines are used to show delocalized electrons in the resonance hybrid.

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 2 electron groups around C, so its hybridization is sp.

Molecular Polarity

Polar and Nonpolar Molecules

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

• Nonpolar Molecule: Has a perfect shape and no lone pairs on the central atom (e.g., CO2).

• Polar Molecule: Has an asymmetrical shape or lone pairs on the central atom (e.g., H2O).

Example: 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

Organic Chemistry Basics

Organic Molecules and Hydrocarbons

Organic chemistry studies molecules containing carbon and hydrogen, often with other elements.

• Organic Molecule: Contains both carbon and hydrogen.

• Hydrocarbon: Contains only carbon and hydrogen.

Example: Identify organic and hydrocarbon molecules from a given set.

Additional info:

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

• Tables reconstructed for clarity and academic context.