Atomic Structure and Electron Configuration

Atoms and Isotopes

Atoms are the fundamental units of matter, consisting of protons, neutrons, and electrons. The atomic number (Z) is the number of protons in the nucleus, while the mass number (A) is the sum of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Ions: Atoms that have gained or lost electrons. Cations are positively charged (loss of electrons), and anions are negatively charged (gain of electrons).

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

Electron Configuration

Electron configuration describes the distribution of electrons in atomic 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.

• 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: 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: Highest in the top right (excluding noble gases); fluorine 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 main group elements tend to form bonds until they are surrounded by eight valence electrons, achieving a noble gas configuration.

• 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 (bonding) pairs to complete its octet.

Formal Charge

Formal charge helps determine the most stable Lewis structure for a molecule. It is calculated as:

Formal Charge Formula:

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

• Only allowable formal charges are -1, 0, or +1 for most main group elements.

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

Lewis Dot Structures

Lewis structures represent the arrangement of valence electrons among atoms in a molecule. Steps to draw:

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 stability.

Example: Draw the Lewis structure for COCl2.

Resonance Structures

Resonance structures are two or more valid Lewis structures for a molecule that differ only in the placement of electrons (not atoms). The real structure is a resonance hybrid, a composite of all resonance forms.

• Double-headed arrows (↔) indicate resonance.

• Resonance involves movement of π electrons or lone pairs.

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. The number of electron groups (bonding pairs + lone pairs) determines the hybridization:

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

Molecular Polarity

Polarity of Molecules

Molecular polarity arises from the distribution of electron density in a molecule. A molecule is polar if it has a net dipole moment due to differences in electronegativity and molecular geometry.

• Nonpolar Molecule: Symmetrical shape and/or identical surrounding atoms (e.g., CO2).

• Polar Molecule: Asymmetrical shape or different surrounding atoms (e.g., H2O).

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

Organic Chemistry: Functional Groups and Molecules

Functional Groups

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

Example: Identify which molecules are organic and which are hydrocarbons based on their functional groups.

Additional Info

• Atomic orbitals are regions of space where electrons are likely to be found.

• Electron configuration principles (Aufbau, Pauli, Hund) are essential for understanding chemical bonding and reactivity.