Atomic Structure and Isotopes

Atoms and Subatomic Particles

The atom is the fundamental 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 an atom's nucleus.

• 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: Positively charged ions (lost electrons).

• Anions: Negatively charged ions (gained electrons).

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

Electron Configuration 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.

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

The Electron Configuration

Ground State Electron Configuration

Describes the distribution of electrons among the orbitals of an atom in its lowest energy state.

• Aufbau Principle: Electrons fill from lower to higher 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): 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 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 (full s and p subshells).

Example: In H3COH, oxygen has 6 valence electrons and 2 shared electrons, totaling 8 octet electrons.

Formal Charge

Definition and Calculation

Formal charge helps determine the most likely 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 the formal charge for each atom using the formula above.

Lewis Dot Structures

Steps for Drawing Lewis Structures

1. Count total valence electrons.

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

3. Connect atoms with single bonds.

4. Add electrons to outer atoms to complete octets.

5. Place remaining electrons on the central atom.

6. If needed, form double or triple bonds to satisfy the octet rule.

7. Check formal charges to ensure 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.

• 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.

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 atoms and lone pairs attached to the central atom.

Example: HCN has 2 electron groups, so the central atom is sp hybridized.

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.

Example: Nitrogen trifluoride (NF3) is polar due to its lone pair and asymmetrical shape.

Functional Groups

Definition and Types

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

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

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

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

Example: Alcohols contain the –OH group; carboxylic acids contain the –COOH group.

Organic Chemistry Overview

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 are organic molecules containing only carbon and hydrogen.

Example: Methane (CH4) is a hydrocarbon; acetic acid (CH3COOH) 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 example electron configurations, explanations of resonance, and the summary tables.