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.

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

Example: Hydrogen Isotopes

• Protium (¹H): 1 proton, 0 neutrons

• Deuterium (²H): 1 proton, 1 neutron

• Tritium (³H): 1 proton, 2 neutrons

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.

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.

• Periodic Trend: Electronegativity increases from left to right across a period and increases going up a group.

• Most Electronegative Element: Fluorine (F)

Example: The most electronegative Group 7A element is Cl (Chlorine), but among all elements, F (Fluorine) is the highest.

Bonding and the Octet Rule

Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, 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 (bonding) pairs, achieving an octet.

Formal Charge

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

• Formula:

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

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

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

Lewis Dot Structures

Lewis 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 H, which is never central).

3. Connect atoms with single bonds.

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

5. Form double/triple bonds if needed to satisfy octet rule.

6. Check formal charges to ensure the best structure.

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

Resonance Structures

Some molecules have more than one valid Lewis structure, called resonance structures. The actual molecule is a resonance hybrid of these forms.

• Resonance: Delocalization of electrons across multiple atoms.

• Double-Headed Arrows: Indicate resonance between structures.

• Resonance Hybrid: The true structure, a blend of all resonance forms.

Example: 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 and lone pairs around a central atom.

Example: HCN (hydrogen cyanide) has a linear geometry and sp hybridization.

Molecular Polarity

Polarity

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

• Nonpolar Molecule: Symmetrical shape, even charge distribution (e.g., CO2).

• Polar Molecule: Asymmetrical shape or uneven charge distribution (e.g., H2O).

• Perfect Shape: Central atom has no lone pairs and all surrounding atoms are identical.

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

Functional Groups and Organic Molecules

Functional Groups

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

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

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

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

Example: Identify functional groups in a given molecule (e.g., alcohol in ethanol, carboxylic acid in acetic acid).

Organic Molecules

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: Determine which molecules are organic and which are hydrocarbons from a given set.

Additional info:

• Some context and definitions were expanded for clarity and completeness.

• Tables were recreated in HTML for hybridization and molecular polarity.

• Examples were inferred or expanded based on standard general chemistry curriculum.