Electron Configuration and Effective Nuclear Charge

Electron Configuration and Energy Levels

• Introduction to Electron Configuration

  • Discusses the concept of energy levels and their relation to electron configuration.

  • Emphasizes that for neutral atoms, the number of electrons equals the number of protons, which is the atomic number.

• Basic Principles of Electron Configuration

  • Hydrogen, having one electron, has the configuration of 1s1.

• Filling Orbitals

  • Fill orbitals based on energy levels:

  • Helium, as a two-electron system, has the configuration 1s2.

  • Lithium has one additional electron, resulting in 1s2 2s1.

  • Proceeding in order through the periodic table, configurations expand as follows:

    • Beryllium: 1s2 2s2

    • Boron: 1s2 2s2 2p1.

• Degenerate Orbitals

  • Discusses the concept of degenerate orbitals where orbitals have the same energy level.

  • Hund's Rule

  • States that electrons will occupy degenerate orbitals singly with parallel spins before pairing.

  • Example: Carbon has two p orbitals that must fill singly (1s2 2s2 2p2).

• Electron Spin and Magnetic Properties

  • Explains how electron spins relate to their magnetic properties.

  • Spin orientations are either up or down, and the Schrodinger equation describes these orientations in terms of spin states.

  • Magnetism

  • Paragmagnetic atoms (e.g., Nitrogen) have unpaired electrons and will exhibit interaction in magnetic fields.

  • Diamagnetic atoms (e.g., Beryllium) have all electrons paired and do not exhibit magnetic interactions.

Ground State Electron Configurations

• Determining Electron Configurations

  • For an atom with 17 electrons (like Chlorine), the configuration can be methodically counted:

  • 1s2 2s2 2p6 3s2 3p5.

• Periodic Trends Related to Electron Configurations

  • Discusses trends such as ionization energy and electron affinity as they relate to electron configurations.

  • Periodic Table Trends:

  • Atomic radius increases down a group and decreases across a period.

  • Ionization energy generally increases across a period from left to right and decreases down a group.

  • Effective nuclear charge (Z_eff) concept introduced, which defines the net positive charge experienced by electrons.

Effective Nuclear Charge and Shielding

• Understanding Effective Nuclear Charge

  • The effective nuclear charge is the net positive charge felt by an electron in an atom after subtracting the shielding effect of inner shell electrons.

  • This relationship affects the stability of valence electrons and influences electron configurations.

• Trends in Periodic Table

  • As one moves across a period from lithium to beryllium, boron, etc., core electrons remain constant, while valence electrons increase, altering Z_eff.

  • Comparison of Lithium, Beryllium, and Boron:

  • Focuses on core electrons (remains the same) vs. increasing valence electrons as one moves right across the periodic table.

Ionization Energy and Electron Affinity

• Differences in Ionization Energy and Electron Affinity

  • Ionization Energy results from removing an electron from an atom.

  • Electron Affinity is the energy change that occurs when an electron is added to a neutral atom.

• Periodic Trend Observations

  • Both ionization energy and electron affinity increase across a period and decrease down a group, influenced by electron configurations.

Group Trends**

• Effective Nuclear Charge Stability

  • Explains that effective nuclear charge remains relatively consistent down a group, while atomic radius and potential energy are factors that impact stability.

  • As radius increases, potential energy of valence electrons decreases, impacting stability significantly.

• Example with Copper

  • Unique electron configuration requires consideration of the d and s orbitals for total understanding of its electronic structure.

  • Potential questions on exams may involve properties of configurations alongside group trends in the periodic table.