Chapter_3_ELECTRON_CONFIGURATIONS

Modern Atomic Theory

• The modern atomic theory offers the most accurate depiction of the atom based on various atomic models.

Nucleus

• The nucleus is at the center of an atom, containing:

  • Protons: Positively charged particles.

  • Neutrons: Neutral particles (no charge).

• Electrons orbit around the nucleus but do not do so in defined paths as initially thought.

Subatomic Particles and Energy Levels

• Electrons exist outside the nucleus in regions defined as energy levels with quantized energy values.

  • These energy levels influence the physical and chemical properties of elements.

  • Electrons occupy specific energy regions and can be excited to higher energy states by absorbing energy (thermal, electrical, etc.).

Bohr Model of the Atom

• Proposed by Niels Bohr, likened electrons to planets orbiting the sun:

  • Electrons inhabit specific energy levels but this model is primarily applicable to the hydrogen atom.

  • Important concepts:

    • Ground State: Lowest energy state of an electron.

    • Excited State: Higher energy states when electrons absorb energy.

    • When excited electrons return to ground state, they emit energy as photons.

Quantum Mechanical Model

• Developed by Erwin Schrödinger; key to modern quantum mechanics.

• Electrons exist in regions of probability called orbitals:

  • An orbital is defined as a space where there is a high probability of finding an electron.

  • Principal energy levels are denoted by the variable n (e.g., n=1 for the first level).

Sublevels and Orbitals

• Each principal energy level contains sublevels (s, p, d, f):

  • Energy Level 1: 1s

  • Energy Level 2: 2s, 2p

  • Energy Level 3: 3s, 3p, 3d

  • Energy Level 4: 4s, 4p, 4d, 4f

• Orbitals can hold a maximum of 2 electrons with opposite spins (spin-up and spin-down).

Electron Configuration

• The arrangement of electrons in an atom is known as electron configuration and greatly influences an element's properties.

• Electrons fill orbitals from the lowest energy level to the highest:

  • The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s...

• Each orbital must be filled with one electron before pairing up (Hund's Rule).

Example Configurations

• Hydrogen: 1s¹

• Helium: 1s²

• Lithium: 1s² 2s¹

• Carbon: 1s² 2s² 2p²

• Use the periodic table for configurations for elements up to atomic number 18 (argon).

Abbreviated Notation

• Use the nearest noble gas for shorthand notation:

  • Potassium: [Ar] 4s¹

  • Bromine: [Ar] 4s² 3d¹⁰ 4p⁵

Trends in Periodic Properties

• Atomic Size: Increases down a group and decreases across a period (left to right).

• Ionization Energy: Energy required to remove an electron; increases across a period and decreases down a group.

• Metallic Character: Trends regarding the reactivity of metals decrease across a period and increase down a group.

Summary of Periodic Table

• Horizontal rows are periods related to energy levels.

• Vertical columns are known as groups/families correlating to outermost (valence) electrons.

• Valence electrons determine bonding and reactions of atoms.

Conclusion

• Key takeaways include understanding the modern interpretation of atomic structure through quantum mechanics, filling order of electron configurations, and the periodic trends affecting the physical and chemical properties of elements.

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