Lecture Notes: Quantum Numbers and Electron Configuration

Importance of Electrons in Chemistry

• Electrons are crucial as they participate in chemical reactions.
• Understanding electron positions helps determine atom properties.
• Electrons don't have a fixed location; probability determines their placement.

Schrodinger's Contributions

• Developed equations to understand electron positions, known as Schrodinger equations.
• Solutions provide electron orbitals, representing probability regions for electrons.
• From Schrodinger's equations, three quantum numbers are derived:
  • n (principal quantum number): Indicates energy level/shell.
  • l (angular quantum number): Indicates orbital shape.
  • ml (magnetic quantum number): Indicates orientation of an orbital.
• ms (spin quantum number): Added later, indicates electron spin.

Quantum Numbers and Their Significance

Principal Quantum Number (n)

• Determines size and energy of an orbital.
• Corresponds to periodic table periods.
• Higher n means larger size and higher energy.

Angular Quantum Number (l)

• Determines shape of orbitals.
• Types of orbitals (subshells): s (sphere), p (dumbbell), d, and f.
• Based on l values: 0 (s), 1 (p), 2 (d), 3 (f).

Magnetic Quantum Number (ml)

• Specifies orientation in space, possible values range from -l to +l.
• Defines the number of orbitals per subshell.

Spin Quantum Number (ms)

• Electrons can have spins of +1/2 or -1/2.
• Maximum 2 electrons per orbital with opposite spins.

Concepts

Degenerate Orbitals

• Orbitals with the same energy level.
• Examples: p orbitals have three degenerate orbitals.

Pauli Exclusion Principle

• No two electrons in an atom can have the same set of four quantum numbers.

Hund's Rule

• Electrons occupy degenerate orbitals singly first, then pair up.

Aufbau Principle

• Electrons fill orbitals starting from the lowest available energy level.

Electron Configuration

Writing Electron Configurations

• Expresses electron arrangements in atoms.
• Notation: 1s² - 1 = principal quantum number, s = orbital type, ² = number of electrons.
• Abbreviated configurations use noble gases as shorthand.

Example Configurations

• Sodium: [Ne] 3s¹
• Selenium: [Ar] 4s² 3d¹⁰ 4p⁴

Transition Metals and Exceptions

• Transition metals have unique configurations due to d orbitals.
• Notable exceptions:
  • Chromium: [Ar] 4s¹ 3d⁵
  • Copper: [Ar] 4s¹ 3d¹⁰

Orbitals in Excited States

Ground vs. Excited State

• Ground State: Electrons in the lowest energy configuration.
• Excited State: Electrons absorb energy and jump to higher levels temporarily.

Paramagnetism and Diamagnetism

• Paramagnetic: Unpaired electrons are present, attracted to magnets.
• Diamagnetic: All electrons are paired, slightly repelled by magnets.

Shielding Effect

• Inner electrons reduce effective nuclear charge experienced by outer electrons.
• Shielding increases down a group but remains constant across a period.

Summary

• Understanding electron configurations and quantum numbers is crucial for predicting chemical properties and behaviors.
• This knowledge bridges chemistry and physics, explaining atomic and molecular structures related to the periodic table.