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Electron Configurations and Periodicity in Chemistry

Learn about electron spin, configurations, and subshell occupancies in this chapter. Understand the Aufbau principle and how to write electron configurations using the periodic table. Discover how electron addition to subshells follows Hund's rule.

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Electron Configurations and Periodicity in Chemistry

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  1. Chapter 8: Electron configurations and periodicity Chemistry 1061: Principles of Chemistry I Andy Aspaas, Instructor

  2. Electron spin and configurations • Electrons may have only one of two possible spins • Quantum number ms = +1/2 or -1/2 • Similar to poles of a magnet • Electron configuration: distribution of electrons among available subshells • Notation lists one subshell after another, with number of electrons in each given in superscript • Ex. 1s22s1 means 2 electrons in 1s and 1 electron in 2s • Orbital diagram: drawing of orbitals and their occupations • Electrons indicated by arrows, either up or down indicates spin

  3. Orbital occupancy • Pauli exclusion principle: no 2 electrons in the same atom may have the same 4 quantum numbers • Each orbital may contain two electrons, but they must be opposite spin - same-spin pair in an orbital is forbidden • Each subshell can hold twice as many electrons as orbitals in the subshell

  4. Subshell occupancies

  5. Filling orbitals • Every atom has an infinite number of electron configrations, but the ground state (stablest) electron configuration can be predicted • Aufbau principle: electron configurations can be predicted by successively filling subshells with electrons in a specific order • # electrons in a neutral atom = atomic number

  6. Using the periodic table to fill orbitals • Use the row number to determine n for s and p sublevels • Section of the periodic table corresponds with the subshell that’s added • First two groups: s • Last 6 groups: p • Transition metals: d • Inner transition metals: f

  7. Available subshells • 1s • 2s, 2p • 3s, 3p, 3d • 4s, 4p, 4d, 4f • 5s, 5p, 5d, 5f • 6s, 6p, 6d, 6f • 7s, 7p, 7d, 7f • These subshells are allowed based on allowed quantum numbers • But, they are filled based on the order they appear in the periodic table • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 etc

  8. Writing electron configurations • Just follow the periodic table from the start, adding electrons in the appropriate sublevel • Electron configurations can be abbreviated using a noble-gas core (in brackets) • Ex. Ca = 1s22s22p63s23p64s2 or just [Ar]4s2

  9. Electron addition to subshells • Hund’s rule: lowest energy arrangement of electrons in a subshell is obtained by first putting in as many same-spin electrons as possible before pairs are made • Oxygen: 1s22s22p4 therefore has 2 unpaired electrons in the 2p subshell • Any atom with unpaired electrons is paramagnetic (attracted to a magnetic field due to unpaired electrons)

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