Lecture on Color and Paramagnetism of Complex Ions and Coordination Compounds

Introduction

• Presenter: Chad from Chad's Prep
• Focus on science prep for high school, college, MCAT, DAT, and OAT exams.
• This lesson is part of the general chemistry playlist.
• Discussion on color and magnetism in complex ions and coordination compounds.

Color in Complex Ions and Coordination Compounds

• Nature of Color
  • Transition metals often associated with bright colors due to electronic transitions.
  • Color arises from the absorption of light that promotes electrons between orbitals.
• Orbital Energy Differences
  • Low energy orbital with an electron can transition to a higher energy orbital.
  • The energy gap corresponds to visible light, unlike the large gap in oxygen which corresponds to ultraviolet light.
• Color Wheel & Complementary Colors
  • Absorbed light isn't the color seen; complementary color appears.
  • For example, if red light is absorbed, the compound appears green.
• Factors Influencing Color
  • D-orbital splitting: Smaller gaps tend to absorb red light; larger gaps absorb violet/blue light.
  • Complex structure (octahedral, tetrahedral, square planar) affects electron transitions.

Conditions for Color

• Requirements
  • Must have electrons in d-orbitals and available higher energy orbitals.
  • Square planar complexes always colored due to electron configuration.
• Exceptions
  • D0 (no electrons) or D10 (completely filled) tend to be colorless.
• Influence of Spin
  • Low spin and high spin states affect the energy gap and thus color.

Magnetism: Paramagnetism and Diamagnetism

• Definitions
  • Paramagnetic: Unpaired electrons, leads to attraction in a magnetic field.
  • Diamagnetic: All electrons are paired, leads to slight repulsion.
• Electron Spin and Magnetism
  • Electrons have spin (up or down); paired spins cancel out.
  • Unpaired electrons result in a net spin and paramagnetism.

Evaluating Complexes

• Sc3+
  • D0, likely colorless, diamagnetic (all paired electrons).
• Cu+
  • D10, likely colorless, diamagnetic (full orbitals).
• Fe2+ (low spin octahedral)
  • D6, colored, diamagnetic (paired electrons).
• Fe2+ (high spin octahedral)
  • D6, colored, paramagnetic (unpaired electrons).
• Cr3+
  • D3, likely colored, paramagnetic (unpaired electrons).

Practical Examples and Spectrochemical Series

• Crystal Field Theory
  • Split of d-orbitals influenced by surrounding ligands.
  • Weak field ligands lead to high spin; strong field leads to low spin.
• Spectrochemical Series
  • Ligands ranked by their field strength, affecting the d-orbital splitting.

Conclusion

• Color and magnetism of complexes influenced by electron configurations and ligand fields.
• Importance of understanding the spectrochemical series for predicting properties of complexes.
• Encouragement to explore further resources and courses offered by Chad's Prep.