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 with a net electric charge due to the loss or gain of electrons.

  • 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 several 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:

Condensed Electron Configuration: Uses the previous noble gas as a starting point. For 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: Among Group 7A elements, Cl is more electronegative than Br or I.

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 Calculation: Octet electrons = valence electrons + shared electrons

Example: In H3COH, oxygen has 6 valence electrons and 2 shared electrons, totaling 8 octet electrons.

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: 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 and F).

3. Connect atoms with single bonds.

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

5. If needed, form double or triple bonds to satisfy the octet rule.

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 in molecules with conjugated pi bonds.

• Double-Headed Arrows: Indicate resonance between structures.

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

Example: Nitrate ion (NO3-) has three resonance structures, each with a different oxygen atom double-bonded to nitrogen.

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 no net dipole moment.

• Polar Molecule: Has an asymmetrical shape or contains polar bonds that do not cancel out.

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

Functional Groups and Organic Molecules

Functional Groups

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

• Hydrocarbons: Alkanes, alkenes, alkynes, aromatic compounds (benzene).

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

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

Example: Alcohols contain an -OH group; carboxylic acids contain a -COOH group.

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: 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.