General Chemistry Study Notes: Atomic Structure, Bonding, and Molecular Properties

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Atomic Structure

Atoms and Isotopes

The atom is the fundamental 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 cations; negatively charged ions are anions.

Example: 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.

Example: Electron configuration diagrams show how electrons fill orbitals according to these rules.

Electron Configuration

Ground State Electron Configuration

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

Aufbau Principle: Start from 1s, fill lower energy orbitals before higher energy ones.
Condensed Electron Configuration: Use the previous noble gas to abbreviate the configuration.

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

Block	Elements
s-block	Groups 1A, 2A
p-block	Groups 3A-8A
d-block	Transition metals
f-block	Lanthanides & 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.

Example: The most electronegative Group 7A element is Cl (Chlorine).

Octet Rule

Valence Electrons 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

Calculating Formal Charge

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/ion.

Formula: Formal Charge = Valence Electrons – (Nonbonding Electrons + 1/2 Bonding Electrons)

Example: Calculate formal charges for N, C, and S in the thiocyanate ion (NCS-).

Lewis Dot Structures

Drawing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules.

Count total valence electrons.
Place the least electronegative atom in the center (except H and halogens).
Connect atoms with single bonds.
Complete octets for surrounding atoms.
Place remaining electrons on the central atom.
If octets are not complete, form double/triple bonds as needed.
Check formal charges for stability.

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, a composite 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 bonds and lone pairs around the central atom.

Electron Groups	Hybridization	Geometry
2	sp	Linear
3	sp2	Trigonal Planar
4	sp3	Tetrahedral

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: Hydrocarbons and molecules with perfect symmetry.
Polar Molecule: Molecules with an asymmetric distribution of electrons.

Electron Groups	Lone Pairs	Polarity
2	0	Nonpolar
3	1	Polar
4	2	Polar

Example: Nitrogen trifluoride (NF3) is polar due to lone pairs and asymmetry.

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.

Functional Group	Structure
Alcohol	R–OH
Aldehyde	R–CHO
Ketone	R–CO–R'
Carboxylic Acid	R–COOH
Amine	R–NH2
Ether	R–O–R'

Organic Chemistry Overview

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 molecules and hydrocarbons from given structures.

Additional info:

Some context and examples were inferred to provide complete academic explanations.
Tables were recreated to summarize block classification, hybridization, molecular polarity, and functional groups.