Atomic Structure and Electron Configuration

Atomic Structure

The atom is the fundamental unit of matter, composed of protons, neutrons, and electrons. Understanding atomic structure is essential for predicting chemical behavior.

• 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).

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

Example: Hydrogen Isotopes

• Protium (¹H): 1 proton, 0 neutrons

• Deuterium (²H): 1 proton, 1 neutron

• Tritium (³H): 1 proton, 2 neutrons

Example: Hydrogen Ions

• Proton (H⁺): Hydrogen atom missing its electron

• Hydride (H⁻): Hydrogen atom with an extra electron

Electron Configuration

Electron configuration describes the distribution of electrons in atomic orbitals. It is governed by 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 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.

General Electron Configuration Notation:

• For example, the ground state electron configuration of Phosphorus (Z = 15):

• Condensed Electron Configuration: Uses the previous noble gas as a starting point. For Phosphorus:

Periodic Table Blocks: The periodic table is divided into s, p, d, and f blocks, which correspond to the type of atomic orbital being filled.

Periodic Trends

Electronegativity

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.

Chemical Bonding and the Octet Rule

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full set of eight valence electrons, resembling the electron configuration of noble gases.

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

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

• Octet: 8 electrons in the valence shell (except for hydrogen, which seeks 2).

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

Formal Charge

Formal charge is used to determine the most likely Lewis structure for a molecule or ion.

• Formula:

• Only allowable formal charges for an element are -1, 0, or +1.

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

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

Lewis Dot Structures

Lewis dot structures represent the arrangement of valence electrons among atoms in a molecule.

1. Count total valence electrons.

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

3. Connect atoms with single bonds.

4. Distribute remaining electrons to complete octets (or duets for hydrogen).

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

6. Check formal charges to ensure the most stable structure.

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.

• Resonance: Delocalization of electrons across multiple atoms or bonds.

• Double-Headed Arrows: Used to indicate resonance between structures.

• Resonance Hybrid: The actual structure is a hybrid of all resonance forms.

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

Hybridization and Molecular Geometry

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 2 electron groups around the central carbon, so it is sp hybridized.

Molecular Polarity

Molecular polarity arises from the distribution of electron density in a molecule.

• 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 the presence of a lone pair on nitrogen.

Organic Chemistry Basics

Functional Groups

Functional groups are specific groups of atoms within molecules that are responsible for the characteristic chemical reactions of those molecules.

• Hydrocarbons: Compounds containing only carbon and hydrogen (alkanes, alkenes, alkynes, aromatic compounds).

• Other Functional Groups: Alcohols, ethers, amines, aldehydes, ketones, carboxylic acids, esters, amides, acid chlorides, thiols.

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

Organic Molecules

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

Summary Table: Key Concepts

Additional info: Some explanations and examples have been expanded for clarity and completeness, following standard General Chemistry curriculum.