• Result of absorption of certain wavelengths of visible light
  - see mix of colors that are not absorbed

  -white light - no visible wavelengths absorbed - see clear or white crystal

• Chromophore: color-causing element
  common chromophores are elements with d-orbitals = transition metals (e.g., Fe²⁺, Fe³⁺Ti⁴⁺, Cr³⁺), and lanthanides
  For transition & lanthanide elements, energy levels of d-orbital electron shell transitions are within visible light wavelengths, which allows these wavelengths to be absorbed.

Color-Causing Mechanisms

1. "Crystal Field Transitions" or "Crystal Field Splitting"
   • most abundant type of color causing mechanism

   • occurs in elements with d-electrons (transition metals)
     • in unbonded or "free" ion all 5 d-orbitals have the same energy level (not within visible color wavelengths)

     • but in bonded ion - energy levels "split" into separate levels with higher and lower energies
       • electron orbitals whose axes are aligned with anion positions are shifted to higher energy level (repulsion causes excited electron state)

       • electron orbitals whose axes are not aligned with anion positions remain at lower energy level

       • difference in energy between split levels is now within visible color wavelengths - color wavelengths can now be absorbed (lower level electron absorbs photon and jumps to the higher level)

   
   • effect is most pronounced for octahedrally coordinated cations - maximizes the alignment between electron shell lobes and anion positions

   • if coordination is a distorted octahedron - further splitting occurs resulting in more energy level changes

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   Effect of Valence on Color
   Energy needed for an electron to make the transition between d orbitals is less when the number of paired electrons within the 5 orbitals is the same before and after the transition. This is called a "Spin Allowed" Transition.
   e.g.,

   For Fe²⁺ there are 6 d electrons spread among 5 d orbitals
   

   => 4 of the orbitals have unpaired electrons and the fifth has a pair of electrons.

   Absorption of a photon would simply cause the second electron in the pair to pair up with one of the unpaired electrons. The result would still be 4 unpaired and 1 paired.

   i.e., before transition: 4 UNPAIRED, 1 PAIRED

   after transition: 4 UNPAIRED, 1 PAIRED
   • This would result in the absorption of lower energy wavelengths e.g., yellows and reds => minerals would show blue - green colors

   Energy needed for an electron to make the transition between d orbitals is more when the number of paired electrons within the 5 orbitals is different after the transition. This is called a "Spin Forbidden" Transition.

   e.g.,

   For Fe³⁺ there are 5 d electrons spread among 5 d orbitals
   

   => initially all 5 of the orbitals have UNPAIRED electrons, and there would be NO PAIRS

   Absorption of a photon would cause one of these unpaired electrons to jump up and pair with one of the other unpaired electrons. The result would be 3 UNPAIRED and 1 PAIRED.

   i.e., before transition: 5 UNPAIRED, 0 PAIRED

   after transition: 3 UNPAIRED, 1 PAIRED
   • This would result in the absorption of higher energy wavelengths e.g., greens and blues => minerals would show yellow - red colors

   Effect of Coordination on Color
   Higher coordination numbers (more surrounding ions) results in longer distances between the central ion (e.g., Fe²⁺) and the coordinating ion (e.g., oxygen). This greater distance results in lower energy levels for the cation's electrons => absorb lower energy wavelengths compared to the same cation with fewer (closer) coordinating ions.
   e.g., Fe²⁺ in 6CN absorbs higher energy wavelengths than Fe²⁺ in 8CN

   Effect of Bond Strength on Color

   Strong bonds result in higher energy electron states => absorb higher energy photons
   e.g., Ionic bonds in corundum result in absorption of blues and greens in ruby (gem form of corundum)

   Weaker bonds result in lower energy electron states => absorb lower energy photons
   e.g., Covalent bonds in beryl result in absorption of reds and yellow in emerald (gem form of beryl)

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2. Molecular Orbital Transitions
   • are the transfer of electrons from one cation to another

     Same rules as above for Spin Allowed and Spin Forbidden Transitions (see "Effect of Valence on Color") except that electrons are transferring between cations not just changing orbitals within one cation.

     For example, Ti⁴⁺ gives an electron to Fe
     i.e., Ti⁴⁺ & Fe²⁺ ---> Ti³⁺ & Fe³⁺

     Cation	# Unpaired	# Paired
     BEFORE TRANSITION
     Ti4+ there are 0 d electrons spread among 5 d orbitals	0	0
     Fe2+ there are 6 d electrons spread among 5 d orbitals	5	1
     AFTER TRANSITION
     Ti3+ there is 1 d electron spread among 5 d orbitals	1	0
     Fe3+ there are 5 d electrons spread among 5 d orbitals	5	0

     i.e., Before transition: 1 PAIRED
     After transition: 0 PAIRED

     Therefore this is a Spin-Forbidden Transition => higher energy wavelengths are absorbed

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3. Color Centers or Farbe Centers
   are lattice defects which result in a "hole" surrounded by (+) charges into which electrons enter to satisfy the missing (-) charge
   

   In the instance of a mineral like halite, with strong ionic bonds, the energy gaps (difference in energy between electron sites) is too great for photons in the visible range to be absorbed.

   Farbe Centers provide new lower energy "sublevels" into which electrons can drop. The energy gaps between these sublevels and the main levels are within the visible light range.