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 arrangement of electrons in an atom's orbitals. This determines chemical properties and reactivity.

• 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.

Notation: Electron configurations are written as a sequence of subshells (e.g., 1s2 2s2 2p6).

Condensed Electron Configuration: Uses the previous noble gas in brackets to simplify notation (e.g., [Ne] 3s2 3p3 for phosphorus).

Example: Phosphorus (Z = 15)

• Ground State: 1s2 2s2 2p6 3s2 3p3

• Condensed: [Ne] 3s2 3p3

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: 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 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.

Example: In H3COH, oxygen has 6 valence electrons and forms 2 shared pairs (bonds), achieving an octet.

Formal Charge

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

• Formula:

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

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

Example: Calculate formal charges for each atom in the thiocyanate ion (NCS-).

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 always terminal).

3. Connect atoms with single bonds.

4. Distribute remaining electrons to complete octets (except for H, which only needs 2 electrons).

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. The actual molecule is a resonance hybrid of these forms.

• Resonance: Delocalization of electrons across multiple atoms.

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

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

Example: Draw all resonance structures for the nitrate ion, NO3-.

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, so the central atom is sp hybridized.

Molecular Polarity

Molecular polarity depends on the shape of the molecule and the distribution of electronegative atoms.

• Nonpolar Molecule: Has a symmetrical (perfect) shape and even charge distribution.

• Polar Molecule: Has an asymmetrical shape or uneven charge distribution.

Example: Nitrogen trifluoride (NF3) is polar due to its trigonal pyramidal shape and lone pair on nitrogen.

Organic Chemistry Fundamentals

Functional Groups

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

• Hydrocarbons: Compounds containing only carbon and hydrogen.

• Alkane: Single bonds only (C–C).

• Alkene: Contains at least one C=C double bond.

• Alkyne: Contains at least one C≡C triple bond.

• Aromatic: Contains a benzene ring.

• Alcohol: Contains an –OH group.

• Ether: Contains a C–O–C linkage.

• Amine: Contains a nitrogen atom bonded to carbon(s).

• Carbonyl Compounds: Contain a C=O group (includes aldehydes, ketones, carboxylic acids, esters, amides, acid chlorides).

Organic Molecules and Hydrocarbons

Organic molecules are defined as compounds containing both carbon and hydrogen. Hydrocarbons are a subset containing only carbon and hydrogen.

• Example: Identify which molecules are organic and which are hydrocarbons based on their structure.

Summary Table: Key Concepts

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