In-Depth Notes on Atomic Structure and the Periodic Table

Proposed by Niels Bohr, linking the atom's electron to photon emission.
Electrons circle the nucleus in allowed paths (orbits).
Energy levels increase with distance from the nucleus.
Emission occurs when an electron falls to a lower energy level, releasing a photon.
Absorption occurs when energy is added to an atom, moving an electron to a higher energy level.

The ground state is the lowest energy state of an atom.
An excited state has higher potential energy than the ground state.
Passing electric current through hydrogen gas results in emission of a characteristic pinkish glow.
A narrow beam of emitted light separated into four specific colors through a prism, known as hydrogen’s line-emission spectrum.

Louis de Broglie described electrons as waves confined to the nucleus's space.
Electron waves exist only at specific frequencies, correlating with quantized energies of Bohr’s orbits (E = hν).
Electrons can diffract (bend) like light waves when passing edges or openings.
Interference occurs when waves overlap.

Proposed by Werner Heisenberg; the attempt to locate an electron with a photon disrupts its course.
Impossible to simultaneously determine both the position and velocity of an electron.

Developed by Erwin Schrödinger in 1926, treating electrons as waves.
Along with Heisenberg's principle, it forms the foundation of modern quantum theory.
Electrons do not travel in neat orbits but exist in regions called orbitals (3D space indicating probable location).

Quantum numbers specify orbital properties and electron characteristics:
- Principal quantum number (n): main energy level occupied by the electron.
- Angular momentum quantum number (l): shape of the orbital.
- Magnetic quantum number (m): orientation of an orbital around the nucleus.
- Spin quantum number: (+1/2 or -1/2), indicating electron spin states.

Aufbau principle: electrons occupy the lowest-energy orbital first.
Pauli exclusion principle: no two electrons can have the same set of four quantum numbers.
Hund's rule: orbitals of equal energy are singly occupied before pairing occurs; same spin for singly occupied orbitals.
Ground-state electron configuration: lowest-energy arrangement of electrons for elements (e.g., Potassium: K = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹).
Irregularities in Chromium (Cr) and Copper (Cu): stable configurations favor full or half-full d-orbitals.

Orbital notation displays unoccupied and occupied orbitals.
Noble-gas configuration uses electron arrangement of noble gases to represent elements more succinctly.
- Example: Rubidium (Rb) = [Kr]5s¹.

Dmitri Mendeleev noticed patterns in elements arranged by increasing atomic mass; similar properties reappeared periodically.
He left gaps in the table, predicting undiscovered elements like Scandium, Germanium, and Gallium.

Henry Moseley discovered arranging elements by atomic number fits patterns better than by atomic mass.
The Periodic Law states that the properties of elements are periodic functions of their atomic numbers.

Arranged by atomic number; similar properties in columns (groups).
Organized in horizontal rows (periods); length determined by the number of electrons filling sublevels.

s block: Groups 1 (alkali metals) and 2 (alkaline-earth metals).
- Group 1: ns¹; properties include softness and high reactivity.
- Group 2: ns²; properties include higher density and strength than Group 1.
d block: Groups 3–12 (Transition metals), typical metallic properties.
p block: Groups 13–18; varying properties within metals, metalloids, and nonmetals.
f block: Includes lanthanides (reactive) and actinides (radioactive).

Defined as the distance from nucleus center to the outer electron cloud; measured in picometers (pm).
Smaller across a period due to increased nuclear charge pulling electrons closer; larger down a group due to increasing electron cloud size.

Energy required to remove an electron from a neutral atom (first ionization energy, IE₁).
General trend: increases across periods due to increasing nuclear charge; decreases down groups as electrons are further from the nucleus.

Energy change when an electron is added to a neutral atom.
Trends: Generally increases across periods; decreases down groups due to increased distance from nucleus.

Cations (positive ions) decrease atomic radius (e.g., Na⁺ < Na); anions (negative ions) increase atomic radius (e.g., Cl⁻ > Cl).
Ionic radii decrease across a period and gradually increase down a group.

A measure of an atom's ability to attract electrons in a compound.
Generally increases across periods and decreases down groups.
The difference in electronegativity helps determine bond type and compound polarity.