Principles of Chemistry: Atomic Orbitals and Periodicity

Electron Density

Probability distribution of finding an electron.

Radial Node

Region in an orbital with zero electron density.

Angular Node

Region in an orbital where probability density is zero.

Hydrogen-like Species

Atoms with one electron, similar to hydrogen.

Energy Level Diagram

Graphical representation of orbital energies.

Avogadro's Constant (NA)

6.022×10²³ mol⁻¹; number of particles per mole.

Bohr Radius (a0)

53 pm; distance for maximum electron probability in H.

Quantum Mechanics

Theory describing behavior of particles at atomic scale.

Wavefunction (ψ)

Mathematical function describing quantum state of a system.

Electron Contour Diagrams

Visual representations of electron density in orbitals.

Energy of 1s Orbital (E1s)

−2.179×10⁻¹⁸ J per atom for hydrogen.

Nuclear Charge (Z)

Total positive charge from protons in nucleus.

Shielding Effect

Core electrons reduce effective nuclear charge on outer electrons.

Ionic Structures

Arrangements of ions in solid crystalline forms.

Quantum Theory Beginnings

Foundation of modern physics explaining atomic behavior.

Periodic Table

Arrangement of elements based on atomic structure.

Chemical Reactions

Processes where substances transform into new products.

e-density plot

Graph showing electron density distribution in orbitals.

2s orbital

Penetrates 1s core, higher effective nuclear charge.

Zeff

Effective nuclear charge experienced by electrons.

Orbital stability

Energy level of an atomic orbital; lower is more stable.

Penetration

Ability of an orbital to approach the nucleus.

Orbital energy sequence

ns < np < nd < nf based on penetration.

Aufbau principle

Electrons fill lowest energy orbitals first.

Pauli exclusion principle

No two electrons can have identical quantum numbers.

Hund's rule

Maximize number of unpaired spins in degenerate orbitals.

Degenerate orbitals

Orbitals with the same energy level.

Electron configuration

Distribution of electrons among atomic orbitals.

Valence electrons

Electrons in outermost orbitals, involved in bonding.

Core electrons

Electrons in filled inner orbitals, not involved in bonding.

Ionization energy (Ei1)

Energy required to remove an outer electron.

Atomic radius

Size of the outermost occupied atomic orbital.

Spectral lines

Visible lines resulting from electron transitions.

Laporte selection rule

Rule governing allowed transitions in spectroscopy.

Bohr model limitations

Fails to explain spectra of non-hydrogen atoms.

Group

Vertical column in the periodic table; similar properties.

Period

Horizontal row in the periodic table; same principal energy level.

Transition metals

Elements with partially filled d orbitals.

Electron shielding

Reduction of effective nuclear charge by inner electrons.

Exchange energy

Energy associated with electron spin arrangements.

Core shielding

Core electrons reduce the nuclear charge felt by valence electrons.

n quantum number

Principal quantum number indicating energy level.

l quantum number

Azimuthal quantum number indicating orbital shape.

Atomic properties periodicity

Trends in atomic properties across the periodic table.

Electron pairing

Two electrons occupying the same orbital with opposite spins.