Atomic Structure and Electron Configuration

Atomic Structure

The atom is the fundamental unit of matter, composed of protons, neutrons, and electrons. Understanding atomic structure is essential for predicting chemical behavior.

• Atomic Number (Z): The number of protons in the nucleus of an atom. Determines the element's identity.

• Mass Number (A): The sum of protons and neutrons in the nucleus.

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A).

• Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge. Cations are positively charged, anions are negatively charged.

Example: Hydrogen has three isotopes: Protium (1 proton), Deuterium (1 proton, 1 neutron), and Tritium (1 proton, 2 neutrons).

Electron Configuration

Electron configuration describes the distribution of electrons in atomic orbitals. It is governed by three main principles:

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; each orbital holds a maximum of two electrons with opposite spins.

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

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): A measure of an atom's ability to attract electrons in a chemical bond.

• Periodic Trend: Electronegativity increases from left to right across a period and increases going up a group.

• Most Electronegative Element: Fluorine (F) 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 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.

Example: In H3COH (methanol), 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 or ion.

• Formula:

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

• Acceptable formal charges are typically -1, 0, or +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.

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 complete octets (or duets for hydrogen).

5. Form double or triple bonds if necessary to satisfy the octet rule.

6. Check formal charges to ensure the most stable 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.

• Resonance structures are connected by double-headed arrows ().

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

• Resonance occurs when there are delocalized electrons (e.g., in NO3-).

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: The number of atoms bonded to a central atom plus the number of lone pairs.

• Types of Hybridization:

  • 2 electron groups: sp (linear geometry)

  • 3 electron groups: sp2 (trigonal planar geometry)

  • 4 electron groups: sp3 (tetrahedral geometry)

Example: In HCN, the central carbon is sp hybridized.

Molecular Polarity

Molecular polarity arises from the distribution of electron density in a molecule.

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

• Polar Molecule: Has an asymmetrical shape or contains polar bonds that do not cancel out.

• Central atom with lone pairs often leads to polarity.

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

Organic Chemistry: Functional Groups and Molecules

Functional Groups

A functional group is a specific group of atoms within a molecule responsible for its characteristic chemical reactions.

• Hydrocarbons: Alkanes, alkenes, alkynes, and aromatic compounds (benzene).

• Other Functional Groups: Alcohols, ethers, amines, aldehydes, ketones, carboxylic acids, esters, amides, acid chlorides, and thiols.

Example: Alcohols contain an -OH group; carboxylic acids contain a -COOH group.

Organic Molecules

Organic chemistry studies molecules containing carbon and hydrogen, often with other elements such as oxygen, nitrogen, sulfur, and halogens.

• Organic Molecule: Contains both carbon and hydrogen.

• Hydrocarbon: Contains only carbon and hydrogen.

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

Summary Table: Principles of Electron Configuration

• Aufbau Principle: Fill lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in the same atom can have identical quantum numbers.

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

Example: Electron configuration diagrams can be used to identify violations of these principles.

Additional info: The study notes above expand on the provided slides and images, filling in missing context and definitions for clarity and completeness. All key concepts are explained with examples and academic context suitable for General Chemistry students.