Atomic Structure

Atoms and Isotopes

The atom is the fundamental unit of matter, composed of protons, neutrons, and electrons. Each element is defined by its atomic number, which equals the number of protons in its nucleus. 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: Protium (1 proton), Deuterium (1 proton, 1 neutron), and Tritium (1 proton, 2 neutrons).

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

• Positively charged ions are called cations; negatively charged ions are anions.

Example: Hydrogen ions: 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 fill degenerate orbitals singly before pairing.

Electron Configuration

Ground State Electron Configuration

The ground state electron configuration describes the distribution of electrons among the orbitals (1s, 2s, 2p, etc.) using the 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: [Ne]

Periodic Table Blocks

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.

• Fluorine is the most electronegative element.

Example: Among Group 7A elements, Cl is more electronegative than Br or I.

Octet Rule

Valence 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, totaling 8 octet electrons.

Formal Charge

Calculation and Application

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

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

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

Formula: Formal Charge = Valence Electrons – (Nonbonding Electrons + ½ Bonding Electrons)

Example: For the thiocyanate ion (NCS-), calculate the formal charge for each atom using the formula above.

Lewis Dot Structures

Drawing Rules

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

1. Count total valence electrons.

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

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

4. Place remaining electrons on the central atom.

5. If octets are incomplete, 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 or ion that differ only in the placement of electrons.

• Resonance involves the movement of pi electrons or lone pairs.

• Double-sided arrows indicate resonance between structures.

• The true structure is a resonance hybrid of all possible 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 a central atom.

Example: HCN has a linear geometry and sp hybridization.

Molecular Polarity

Polar and Nonpolar Molecules

Molecular Polarity refers to the distribution of electrical charge over the atoms in a molecule.

• Nonpolar Molecule: Molecule with a symmetrical shape and no net dipole moment.

• Polar Molecule: Molecule with an asymmetrical shape or differing electronegativities, resulting in a net dipole moment.

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

Functional Groups

Definition and Classification

Functional Group: A specific group of atoms within a molecule responsible for characteristic chemical reactions.

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

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

• Without Carbonyls: Alkyl halides, amines, alcohols, ethers, thiols.

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

Organic Chemistry Overview

Organic Molecules and Hydrocarbons

Organic chemistry studies molecules containing carbon and hydrogen, often with other elements. Organic molecules are defined by the presence of both carbon and hydrogen. Hydrocarbons contain only carbon and hydrogen.

• Example: Benzene (C6H6) is a hydrocarbon; ethanol (C2H5OH) is organic but not a hydrocarbon.

Additional info: Organic chemistry is foundational for understanding biological molecules and industrial compounds.

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