Atomic Structure and Electron Configuration

Atoms, Isotopes, and Ions

Atoms are the fundamental units of matter, consisting of protons, neutrons, and electrons. The atomic number (Z) is the number of protons in the nucleus, while the mass number (A) is the sum of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge.

• Cations: Positively charged ions (fewer electrons than protons).

• Anions: Negatively charged ions (more electrons than protons).

Example: Hydrogen has three isotopes: protium (1H), deuterium (2H), and tritium (3H).

Electron Configuration

Electron configuration describes the distribution of electrons in atomic orbitals. The Aufbau Principle states that electrons fill the lowest energy orbitals first. The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of quantum numbers. Hund's Rule states that electrons fill degenerate orbitals singly before pairing.

• Ground State Electron Configuration: The arrangement of electrons in the lowest possible energy state.

• Condensed Electron Configuration: Uses the previous noble gas to abbreviate the configuration.

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

Periodic Trends: Electronegativity

Electronegativity (EN): A measure of an atom's ability to attract electrons in a chemical bond. Electronegativity increases from left to right across a period and from bottom to top within a group.

• Most electronegative element: Fluorine (F)

• Periodic Trend: Increases across a period, decreases down a group.

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

Chemical Bonding and Molecular Structure

Octet Rule

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, oxygen has 6 valence electrons and forms 2 shared (bonding) pairs to complete its octet.

Formal Charge

Formal charge helps determine the most stable Lewis structure for a molecule. It is calculated as:

Formal Charge Formula:

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

• Only allowable formal charges are -1, 0, or +1 for most main group elements.

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

Lewis Dot Structures

Lewis structures represent the arrangement of valence electrons among atoms in a molecule. Steps to draw 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 outer atoms, then central atom.

5. Use double or triple bonds if needed to satisfy octets.

6. Check formal charges for correctness.

Example: Draw the Lewis structure for COCl2.

Resonance Structures

Some molecules have more than one valid Lewis structure, called resonance structures. These structures differ only in the placement of electrons, not atoms.

• Double-sided arrows (↔) indicate resonance.

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

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

Example: Nitrate ion (NO3-) has three resonance structures, each with a different N=O double bond.

Hybridization

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: HCN has a linear geometry and sp hybridization.

Molecular Polarity

Molecular polarity arises from the distribution of electron density in a molecule. A molecule is polar if it has a net dipole moment.

• Nonpolar Molecule: Symmetrical shape, no net dipole (e.g., CO2).

• Polar Molecule: Asymmetrical shape or presence of lone pairs on the central atom (e.g., H2O).

Example: Nitrogen trifluoride (NF3) is polar due to the lone pair on nitrogen.

Organic Chemistry Basics

Functional Groups

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: Aldehydes, ketones, carboxylic acids, esters, amides, acid chlorides.

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

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

Organic Molecules and Hydrocarbons

Organic molecules contain both carbon and hydrogen. Hydrocarbons are organic molecules consisting solely of carbon and hydrogen.

• Alkanes: Single bonds only (saturated hydrocarbons).

• Alkenes: At least one double bond.

• Alkynes: At least one triple bond.

• Aromatic: Contain benzene rings.

Example: Methane (CH4) is the simplest alkane.

Summary Table: Principles of Electron Configuration

Additional info: These notes provide a concise overview of foundational concepts in general chemistry, including atomic structure, periodic trends, bonding, molecular geometry, and basic organic chemistry. For further study, refer to your textbook for detailed examples and practice problems.