CHEM 101 12/16

Electron Configuration Basics

• Electron Configuration: A way to represent the distribution of electrons in an atom's orbitals.

Chlorine Atom

• Atomic Number: 17 electrons.

• Electron Configuration:

  • 1s² 2s² 2p⁶ 3s² 3p⁵

Noble Gases and Reactivity

• Noble Gases: Nonreactive elements due to complete valence shells, e.g., Helium, Neon, Argon, Krypton.

• Helium: 1s², completely filled shell with 2 electrons.

Orbital Diagrams

• Visual representation of electrons in orbitals; fill orbitals based on energy.

• Energy Order:

  • 1s < 2s < 2p < 3s < 3p

• Filling Order:

  • First electron in an orbital has spin up or down (50/50).

  • Follow Hund's Rule: Fill degenerate orbitals singly before pairing.

Quantum Numbers

• Four Quantum Numbers: Describe electron's location and energy level.

  • Principal Quantum Number (n): Indicates energy level (e.g., n=3 for 3p).

  • Azimuthal Quantum Number (l): Indicates subshell type (s, p, d, f)

  • Magnetic Quantum Number (m_l): Orientation of orbital.

  • Spin Quantum Number (m_s): Direction of electron's spin (+1/2 or -1/2).

Chlorine's Last Electron

• Quantum Numbers for last electron in 3p orbital: (n=3, l=1, m_l=1, m_s=-1/2).

Transition Metals and Electron Configurations

Chromium (24 electrons)

• Electron Configuration:

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴ (uses noble gas notation: [Ar] 4s² 3d⁴).

• Observed Configuration:

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵ (4s electron promotes to 3d for stability).

Copper (29 electrons)

• Predicted Configuration:

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹.

• Observed Configuration:

  • 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰.

Periodic Table Characteristics

• S Block: Groups 1 & 2 (alkali and alkaline earth metals).

• P Block: Groups 13-18 (main group elements).

• Metal Character:

  • More metallic behavior decreases left to right and increases down a group.

• Metalloids: Found along the stair-step line separating metals and nonmetals.

Atomic Size and Trends

• Atomic Radius: Distance from nucleus to outermost electron shell.

  • Trend: Increases down a group, decreases across a period from left to right.

• Effective Nuclear Charge (Z_eff): The net positive charge experienced by valence electrons.

  • Typically less than actual nuclear charge due to shielding effects.

Ionization Energy

• Ionization Energy (IE): Energy required to remove an electron from an atom.

  • First IE: Removal of the first electron; increases across periods and decreases down groups.

  • Second IE: Removal of a second electron; generally requires more energy than the first due to increased positive charge on the remaining ions.

Electron Affinity

• Electron Affinity: Energy released when an atom gains an electron.

• Nonmetals (like halogens) tend to have high electron affinities, while metals typically do not.

Summary of Key Concepts

• Understand the significance of noble gases and their role in chemical stability due to full valence shells.

• Practice drawing orbital diagrams and writing electron configurations.

• Note the trends in the periodic table affecting metallic character, atomic radius, ionization energy, and electron affinity.