Atomic Structure and Isotopes

Atoms and Atomic Number

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.

• 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

• Cation: Positively charged atom (fewer electrons than protons).

• Anion: Negatively charged atom (more electrons than protons).

Example: The hydride ion (H-) has one more electron than the proton (H+).

Electron Configuration

Aufbau Principle, Pauli Exclusion, and Hund's Rule

Electron configuration describes the arrangement of electrons in an atom's orbitals.

• 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 occupy degenerate orbitals singly before pairing up.

Example: The ground state electron configuration of phosphorus (Z = 15) is:

Condensed Electron Configuration: Use the previous noble gas as a starting point. For phosphorus:

Periodic Table Blocks

• s-block: Groups 1A and 2A

• p-block: Groups 3A to 8A

• d-block: Transition metals (Groups 3B to 2B)

• f-block: Lanthanides and actinides

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 achieve eight valence electrons through chemical bonding.

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

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

Example: In H3COH, oxygen has 6 valence electrons and 2 shared electrons, completing its octet.

Formal Charge

Definition and Calculation

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

• Formula:

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

• Allowed formal charges are typically -1, 0, or +1 for main group elements.

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).

3. Connect atoms with single bonds.

4. Complete octets for surrounding atoms (except hydrogen, which only needs 2 electrons).

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 occurs when there is at least one pi bond and delocalized electrons.

• Double-headed arrows () indicate resonance between structures.

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

Example: The nitrate ion (NO3-) has three resonance structures, each with a different oxygen atom double-bonded to nitrogen.

Hybridization

Electron Groups and Hybrid Orbitals

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding.

• Electron Groups: Number of atoms bonded to the central atom plus lone pairs.

Example: In HCN, the central carbon 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 Classification

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

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

• 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.

• An organic molecule contains both carbon and hydrogen.

• Hydrocarbons are organic molecules containing only carbon and hydrogen.

Example: Ethanol (C2H5OH) is an organic molecule; benzene (C6H6) is a hydrocarbon.

Summary Table: Principles of Electron Configuration

Additional info: Some content was inferred and expanded for clarity and completeness, including the full electron configuration example, the summary tables, and the explanation of functional groups and organic chemistry context.