Atomic Structure and Electron Configuration

Atoms, Isotopes, and Ions

Atoms are the fundamental units of matter, composed 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 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 (lost electrons), 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.

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

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

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

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: 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.

Example: In H3COH, oxygen has 6 valence electrons and forms 2 shared (bonding) pairs to complete its octet.

Formal Charge

Formal charge is used to 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.

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

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 in a molecule. The best structure minimizes formal charges and satisfies the octet rule.

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. Complete octets for outer atoms, then central atom.

5. Use double/triple bonds if needed to satisfy octets.

6. Check formal charges.

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 structures are connected by double-headed arrows.

• The actual structure is a resonance hybrid of all possible structures.

• Resonance involves the movement of pi 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.

• Electron Groups: Number of atoms bonded to the central atom plus lone pairs.

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

Molecular Polarity

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: Symmetrical shape, even distribution of charge.

• Polar Molecule: Asymmetrical shape, uneven distribution of charge.

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

Organic Chemistry Fundamentals

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.

• 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 and Hydrocarbons

Organic molecules are compounds containing both carbon and hydrogen. Hydrocarbons are organic molecules consisting solely of carbon and hydrogen.

• Alkanes: Single bonds only.

• Alkenes: At least one double bond.

• Alkynes: At least one triple bond.

• Aromatic: Contain benzene rings.

Example: Methane (CH4) is an alkane; ethene (C2H4) is an alkene.

Summary Table: Key Concepts