Atomic Structure and Isotopes

Atoms and Subatomic Particles

The atom is the fundamental unit of matter, composed of protons, neutrons, and electrons. The arrangement and number of these subatomic particles determine the identity and properties of each element.

• Atomic Number (Z): The number of protons in the nucleus of an atom. It defines the element.

• Mass Number (A): The total number 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: protium (1H), deuterium (2H), and tritium (3H).

• Ions: Atoms that have gained or lost electrons. Cations are positively charged (loss of electrons), and anions are negatively charged (gain of electrons).

Example: D+ (deuterium ion) and 14C2− (carbon ion).

Electron Configuration

Principles of Electron Configuration

Electron configuration describes the distribution of electrons in atomic orbitals. Three main principles govern this arrangement:

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an 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 for phosphorus (Z = 15) is:

$$1s^2 2s^2 2p^6 3s^2 3p^3$$

The condensed electron configuration uses the previous noble gas:

$$[Ne] 3s^2 3p^3$$

Periodic Table Blocks

• s-block: Groups 1A and 2A (alkali and alkaline earth metals)

• p-block: Groups 3A to 8A (main group elements)

• 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) is the most electronegative element.

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 main group elements tend to form bonds until they are surrounded by eight valence electrons, achieving a noble gas configuration.

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

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

Example: In H3COH (methanol), oxygen has 6 valence electrons and forms 2 shared pairs (bonds), completing its octet.

Formal Charge

Calculating Formal Charge

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

• Formula:

$$\text{Formal Charge} = \text{Valence Electrons} - (\text{Nonbonding Electrons} + \text{Bonding Electrons}/2)$$

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

• Acceptable formal charges are typically -1, 0, or +1.

Example: In 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 for all atoms.

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

3. Connect atoms with single bonds.

4. Complete octets for outer atoms, then for the central atom.

5. If octets are incomplete, form double or triple bonds as needed.

6. Check formal charges to ensure the best structure.

Example: Draw the Lewis structure for COCl2 (phosgene).

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 (not atoms).

• 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 weighted average of all resonance forms.

Example: The nitrate ion (NO3−) has three resonance structures, each with a different N=O double bond location.

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 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 depends on the shape of the molecule and the distribution of electronegative atoms.

• Nonpolar Molecule: Has a symmetrical (perfect) shape and even charge distribution.

• Polar Molecule: Has an asymmetrical shape or uneven charge distribution, resulting in a dipole moment.

Example: Nitrogen trifluoride (NF3) is polar due to its trigonal pyramidal shape and lone pair on nitrogen.

Functional Groups

Definition and Types

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

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

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

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

Example: Ethanol contains an alcohol functional group; acetone contains a ketone group.

Organic Chemistry Overview

Organic Molecules and Hydrocarbons

Organic chemistry studies molecules containing carbon and hydrogen, often with other elements such as oxygen, nitrogen, sulfur, and halogens.

• Organic Molecule: Contains both carbon and hydrogen.

• Hydrocarbon: Contains only carbon and hydrogen.

Example: Methane (CH4) is a hydrocarbon; ethanol (C2H5OH) is an organic molecule but not a hydrocarbon.

Summary Table: Key Concepts