14 - Quantum Numbers & Electron Configuration

Modern View of the Atom

  • Transition from Bohr's model to Schrodinger's model.

  • Schrodinger introduced wave functions to describe electron orbitals.

  • Emphasizes the uncertainty principle: cannot know exact momentum and location of electrons simultaneously.

Uncertainty Principle

  • Impossibility to measure both momentum and location with precision.

  • Analogy: Moving car example illustrates difficulty in predicting exact future position.

  • Key takeaway: observing a particle can change its behavior.

Double Slit Experiment

  • Light behaves as both a wave and a particle (photons).

  • Observer's effect alters the outcome of experiments.

  • Importance of understanding the uncertainty principle and observer's effect.

Quantum Numbers

  • Quantum numbers (n, l, m_l, m_s) provide a way to locate electrons:

    • n: Principal quantum number (energy shell), integers 1 to ∞.

    • l: Azimuthal quantum number (subshell), ranges from 0 to n-1.

    • m_l: Magnetic quantum number (number of orbitals), ranges from -l to +l.

    • m_s: Spin quantum number, either +1/2 or -1/2.

Pauli Exclusion Principle

  • Each orbital can hold a maximum of 2 electrons with opposite spins.

Types of Orbitals

  • s Orbital: Spherical shape.

  • p Orbitals: Dumbbell shape, includes three orientations (px, py, pz).

Electron Configuration

  • Understanding how to write electron configurations using the periodic table.

  • Example for Oxygen: 1s² 2s² 2p⁴.

  • Noble gas notation simplifies writing of electron configurations.

Cations and Anions

  • Anion: Negative charge, gain of electrons.

  • Cation: Positive charge, loss of electrons. Determine the configuration by removing or adding electrons from the outermost shell.

Practical Tips

  • Regular practice with quantum numbers and electron configurations.

  • Use of periodic tables is crucial.

  • Be careful with details to avoid mistakes in calculations.