Atomic Structure and Electron Configuration

Atoms and Isotopes

Atoms are the fundamental units of matter, consisting of protons, neutrons, and electrons. The atomic number is the number of protons in an atom, while the mass number is the sum of protons and neutrons. Isotopes are atoms of the same element with different numbers of neutrons.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Number of protons + neutrons.

• Isotopes: Atoms with the same atomic number but different mass numbers.

• Ions: Atoms that have gained or lost electrons. Cations are positively charged, anions are negatively charged.

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

Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's 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.

• Aufbau Principle: Electrons occupy the lowest energy orbitals available.

• Pauli Exclusion Principle: Each orbital holds a maximum of two electrons with opposite spins.

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing.

Condensed Electron Configuration: Uses the previous noble gas to simplify notation.

Example: Phosphorus (Z = 15): Ground state: 1s2 2s2 2p6 3s2 3p3 Condensed: [Ne] 3s2 3p3

Periodic Trends

Electronegativity

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

• Periodic Trend: Increases across a period (left to right) and 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.

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

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. It is calculated as:

Formal Charge Formula:

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

• Allowed 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 dot structures represent the arrangement of valence electrons in molecules. There are systematic steps to draw the best structure:

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/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.

• Resonance: Delocalization of electrons across multiple structures.

• Double-sided arrows (↔) indicate resonance between structures.

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

Example: Nitrate ion (NO3-) has three resonance structures.

Hybridization and Molecular Geometry

Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding. The number of electron groups (bonds and lone pairs) determines the hybridization:

Example: HCN has 2 electron groups, so the central atom is sp hybridized.

Molecular Polarity

Molecular polarity arises from differences in electronegativity and molecular geometry. A molecule is polar if it has a net dipole moment.

• Nonpolar molecule: Symmetrical shape, even charge distribution.

• Polar molecule: Asymmetrical shape or uneven charge distribution.

Example: Nitrogen trifluoride (NF3) is polar due to its lone pair and asymmetrical shape.

Organic Chemistry Basics

Functional Groups

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

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 made solely of carbon and hydrogen.

• Hydrocarbons: Alkanes, alkenes, alkynes, and aromatics.

• Organic molecule: Any molecule with both carbon and hydrogen atoms.

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

Summary Table: Key Concepts