Quantum Mechanics Overview

Quantum Behavior Overview

• Quantum mechanics describes the bizarre and often counterintuitive behavior of matter and energy at the atomic and subatomic levels.
• Quantum phenomena are governed by probabilistic rules, not deterministic ones.
• Quantization means that properties like energy and angular momentum can only take on specific discrete values, not continuous ranges.
• Superposition is a key concept; a quantum system can exist in multiple states simultaneously until measured.
• Entanglement is a phenomenon where two or more particles become linked, influencing each other's states regardless of distance.

Electron Orbitals and Quantum Numbers

• Electrons occupy orbitals, regions of space around the nucleus where there's a high probability of finding an electron.
• Four quantum numbers describe the properties of each electron:
  • Principal quantum number (n): Determines the energy level and overall size of the orbital. Higher n values correspond to higher energy levels.
  • Azimuthal quantum number (l): Determines the shape of the orbital. Values range from 0 to n-1. l=0 is s, l=1 is p, l=2 is d, and l=3 is f.
  • Magnetic quantum number (ml): Determines the orientation of the orbital in space. Values range from -l to +l.
  • Spin quantum number (ms): Determines the intrinsic angular momentum or spin of the electron. Can be +1/2 or -1/2.

Orbital Shapes

• s orbitals are spherical.
• p orbitals are dumbbell-shaped.
• d orbitals have more complex shapes, including cloverleaf and double-dumbbell shapes.
• f orbitals are even more complex in shape.

Electron Configurations

• Electron configurations describe the arrangement of electrons in atomic orbitals.

• Aufbau principle: Electrons fill lower energy levels first.

• Hund's rule: Within a subshell, electrons occupy each orbital singly before pairing up.

• Pauli exclusion principle: No two electrons in an atom can have the same four quantum numbers.

• Abbreviated electron configurations use the noble gas that precedes the element to represent the filled inner electron shells. Example: Magnesium [Ne] 3s².

Orbital Diagrams

• Orbital diagrams represent electron configurations using boxes to show orbitals and arrows to indicate the electrons' spin.
• Use arrows pointing up and down to represent the two possible spins ( +1/2 and –1/2) .

Magnetic Properties

• Magnetic properties depend on the electron configuration.
• Elements with unpaired electrons are paramagnetic, meaning they are attracted to magnetic fields.
• Elements with all electrons paired are diamagnetic, meaning they are repelled by magnetic fields.

Valence and Core Electrons

• Valence electrons are the electrons in the outermost shell of an atom; they are involved in chemical bonding.
• Core electrons are the inner-shell electrons; they are not typically involved in chemical reactions.

Electron Configurations of Ions

• Cations (positive ions) lose electrons. Their electron configuration reflects the loss of the outer electrons.
• Anions (negative ions) gain electrons. Their electron configuration reflects the addition of electrons to the outer shell.

Predicting Ion Charges

• Many elements tend to gain or lose electrons to achieve a noble gas electron configuration (8 valence electrons for most), thus achieving stability.
• Group trends on the periodic table can be used to predict ion charges.