Electron configurations reveal the arrangement of electrons in atoms, shaping the periodic table's structure. This fundamental concept explains element properties, bonding behavior, and periodic trends, connecting atomic structure to chemical reactivity and spectroscopic observations.

Understanding electron configurations unlocks the periodic table's secrets, from ionization energies to magnetic properties. It's the key to predicting chemical behavior, making sense of spectral lines, and grasping why elements react the way they do.

Electron configurations and the periodic table

Principles of electron configurations

• Electron configurations represent electron arrangements in atomic orbitals following Aufbau principle, Hund's rule, and Pauli exclusion principle
• Aufbau principle dictates electrons fill lowest energy orbitals first (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p)
• Hund's rule states electrons occupy orbitals of equal energy individually before pairing
• Pauli exclusion principle allows maximum of two electrons with opposite spins per orbital
• Shorthand notation uses noble gas cores to simplify writing for elements with many electrons (e.g. [Ne] for sodium's core electrons)
• Exceptions occur due to stability of half-filled and fully-filled subshells (chromium: [Ar]4s13d5 instead of [Ar]4s23d4)

Periodic table organization and electron configurations

• Table structure reflects electron configurations with s, p, d, f blocks corresponding to subshell being filled
• Group numbers indicate number of valence electrons determining chemical behavior (Group 1: 1 valence electron, Group 18: 8 valence electrons)
• Periods represent highest principal quantum number (n) in element's electron configuration
• Transition metals (d-block) and inner transition metals (f-block) involve filling d and f subshells
• Table organization allows quick determination of electron configurations without memorization
• Atomic radius trends explained by number of energy levels and effective nuclear charge (increases across period, decreases down group)

Electron configurations for ions

• Cations form by removing electrons from highest energy level (Na+: 1s22s22p6)
• Anions form by adding electrons to lowest available energy level (Cl-: 1s22s22p63s23p6)
• Transition metal ions often lose electrons from s orbital before d orbital (Fe2+: [Ar]3d6)
• Main group ions typically achieve noble gas configurations (Al3+: 1s22s22p6)

Electron configurations and element properties

Ionization energy and electron affinity

• Ionization energy increases across periods and decreases down groups
  • Explained by distance of valence electrons from nucleus and electron shielding effects
  • First ionization energy of lithium lower than beryllium due to stable half-filled 2s orbital
• Electron affinity influenced by stability of noble gas configurations or half-filled/fully-filled subshells
  • Halogens have highest electron affinities due to achieving noble gas configurations
  • Noble gases have low electron affinities due to stable full outer shells

Metallic and nonmetallic properties

• Metallic character determined by ease of losing electrons (increases down group, decreases across period)
• Nonmetallic character determined by ease of gaining electrons (opposite trend of metallic character)
• Transition metals exhibit variable oxidation states due to availability of both ns and (n-1)d electrons
• Metalloids (boron, silicon, germanium) have intermediate properties between metals and nonmetals

Magnetic and spectroscopic properties

• Magnetic properties predicted by presence of unpaired electrons in electron configurations
  • Paramagnetism occurs in atoms with unpaired electrons (oxygen: two unpaired electrons in 2p orbitals)
  • Diamagnetism occurs in atoms with all paired electrons (neon: fully filled 2p orbitals)
• Spectroscopic properties and color often explained by electron transitions within specific configurations
  • Transition metal complexes exhibit colors due to d-d transitions (copper sulfate solution: blue color)
  • Flame tests produce characteristic colors based on valence electron transitions (sodium: yellow flame)

Valence electrons and chemical bonding

Valence electron concepts

• Valence electrons occupy outermost shell of atom and participate in chemical bonding and reactions
• Number of valence electrons quickly determined from group number for main group elements
• Transition metals include both ns and (n-1)d electrons as valence electrons
• Octet rule guides many bonding interactions (atoms tend to achieve 8 valence electrons)
• Exceptions to octet rule include expanded octets (sulfur in SF6) and incomplete octets (boron in BF3)

Valence electrons and bonding theories

• Valence bond theory explains covalent bonding through orbital overlap
  • Sigma bonds form by head-on overlap (H-H bond in hydrogen molecule)
  • Pi bonds form by side-by-side overlap (C=C double bond in ethylene)
• Molecular orbital theory provides more advanced explanation of bonding
  • Bonding and antibonding molecular orbitals form from atomic orbital combinations
  • Bond order calculated from number of electrons in bonding and antibonding orbitals

Electronegativity and bond types

• Electronegativity related to valence electron behavior predicts bond types
• Ionic bonds form between elements with large electronegativity difference (sodium chloride)
• Covalent bonds form between elements with similar electronegativities (hydrogen molecule)
• Polar covalent bonds form between elements with intermediate electronegativity difference (water)
• Metallic bonds involve delocalized valence electrons in a "sea of electrons" (copper metal)