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 Isotopes: Protium (1 proton), Deuterium (1 proton, 1 neutron), Tritium (1 proton, 2 neutrons).

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 degenerate orbitals singly before pairing up.

Condensed Electron Configuration: Uses the previous noble gas to abbreviate the configuration.

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: Increases across a period (left to right) and up a group.

• Most electronegative element: Fluorine (F).

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

Chemical 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: 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 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 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 octet rule.

6. Check formal charges for correctness.

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. The real structure is a resonance hybrid, a composite of all resonance forms.

• Resonance involves the movement of electrons, not atoms.

• Double-sided arrows () indicate resonance.

• Resonance hybrid is drawn with dashed lines to show delocalized electrons.

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 shape and equal sharing of electrons.

• Polar Molecule: Has an asymmetrical shape or unequal sharing of electrons.

• Perfect Shape: Central atom has no lone pairs and all surrounding atoms are identical.

Example: Nitrogen trifluoride (NF3) is polar due to 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, arenes (aromatic compounds).

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

• Without Carbonyls: Alkyl halides, amines, alcohols, ethers, thiols.

Example: Identify functional groups in a given molecule, such as alcohol (-OH), amine (-NH2), or carboxylic acid (-COOH).

Organic Molecules

Organic chemistry is the study of carbon-containing compounds, especially those found in living organisms. Organic molecules must contain both carbon and hydrogen. Hydrocarbons are organic molecules containing only carbon and hydrogen.

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

Summary Table: Principles of Electron Configuration

Additional info: Some content was inferred and expanded for clarity and completeness, such as the explicit electron configuration for phosphorus and the summary tables for hybridization and electron configuration principles.