3.8-3.9:

3.8: The periodic table and filling orbitals:

• Aufbau principal: the concept of building up ground-state atoms so that their electrons occupy the lowest-energy orbitals available

  • electrons placed in lowest-energy orbitals available

  • only other restriction is no more than two electrons from Pauli’s exclusion principle

• Electron configurations: the distribution of electrons among the orbitals of an atom or ion

• 1s^1

  • First 1 is principal quantum number (n), s indicates subshell (l), and superscripted 1 means that there is one electron in the 1s orbital

• Pauli exclusion states that if two electrons on in an orbital one has spin up and one has spin down

  • Two opposite electrons with opposite spin quantum numbers are spin-paired

• Core electrons: the electrons in the filled, inner shells in an atom or ion; they are not involved in chemical reactions

• Valence shell: the outermost occupied shell of an atom

  • Electrons in this shell are called valance electrons

• Effective nuclear charge (Zeff): the attraction toward the nucleus experienced by an electron in an atom; the positive charge on the nucleus is reduced by the extent to with other electrons in the atom shield the electron from the nucleus. 

• Degenerate orbitals: orbitals with the same energy

• Hund’s rule: the principle that the lowest-energy electron configuration of an atom has the maximum number of unpaired electrons, all of which have the same spin, in degenerate orbitals

• Orbital diagrams: a depiction of the arrangement of electrons in an atom or ion, using boxes to represent orbitals

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3.9: Electron configurations of ions:

• Stuff wants to become like noble gases

• Isoelectronic: describes atoms or ions that have identical electron configurations

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