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 the nucleus, while the mass number 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 (lost electrons), and anions are negatively charged (gained 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 simplify notation.

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. Electronegativity increases from left to right across a period and from bottom to top within a group.

• Periodic Trend: Highest in the top right of the periodic table (excluding noble gases).

Example: The most electronegative Group 7A element is chlorine (Cl).

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 2 shared electrons, totaling 8 electrons (octet satisfied).

Formal Charge

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

Formula:

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

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

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. 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. Distribute remaining electrons to satisfy the octet rule.

5. Form double or triple bonds if necessary to complete octets.

6. Check formal charges to ensure the best 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.

• Double-sided arrows indicate resonance between structures.

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

• Resonance occurs when there are delocalized electrons (pi bonds 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, so the central atom 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.

• Nonpolar Molecule: Has a symmetrical (perfect) shape and no net dipole.

• 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.

Organic Chemistry Basics

Functional Groups

A functional group is a specific group of atoms within a molecule responsible for its characteristic reactions. Recognizing functional groups is essential for understanding organic reactivity.

• Hydrocarbons: Alkanes, alkenes, alkynes, and aromatic compounds.

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

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

Example: Identify the functional groups in a given molecule, such as an alcohol (-OH) or a ketone (C=O).

Organic Molecules and Hydrocarbons

Organic molecules contain both carbon and hydrogen. Hydrocarbons are organic molecules consisting solely of carbon and hydrogen atoms.

• Alkanes: Single bonds only.

• Alkenes: At least one double bond.

• Alkynes: At least one triple bond.

• Aromatic: Contain benzene rings.

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

Summary Table: Principles of Electron Configuration

Additional info: These notes provide a concise overview of foundational concepts in general chemistry, including atomic structure, periodic trends, bonding, molecular geometry, and basic organic chemistry. Mastery of these topics is essential for success in college-level chemistry courses.