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Causes of color
Contents
1 Basic
2 Advanced
2.1 The Crystal Field Theory
2.1.1 Allochromatic colors caused by transition metal impurities
2.1.2 Idiochromatic colors caused by transition metal constituents
2.1.3 Color centers
2.1.3.1 Electron color centers
2.1.3.2 Hole color centers
2.2 The Molecular Orbital Theory
2.2.1 Electronegativity
2.2.2 Charge transfer
2.3 The Band Theory
2.4 The Physical Properties Theory
3 Sources
4 External links
Basic
In general color is caused by the absorption of certain wavelengths of light by a substance (as a gemstone) while permitting other wavelengths to pass through the substance unaltered. The net result of the wavelengths that pass through the gem give the final color to the gemstone.
In order to see color we need at least 3 variables:
Light
A substance
Vision (the eye)
When any of them is absent, we can't see color. For instance, your red sweater will have no color in the dark as there is no light for the sweater to absorb.
Absorption of all wavelenghts execpt red
White light is a mixture of the 7 spectral colors (visible in a rainbow): red, orange, yellow, green, indigo, blue and violet. Each of these colors travel at a certain wavelength (about 700nm for red and 400nm for violet) and carry a specific amount of energy.
When a substance absorbs all the spectral colors (the colors of the rainbow) except red, the residual color is red. Therefore the net result will be red. If all colors except red and blue are absorbed, the residual color (net result) will be a purple gemstone.
In total there are about 16 million combinations that can produce color.
Absorption of light usually results in transformation of light energy to heat. That is why frigidaires are white; white substances don't absorb light so they are very effective at keeping heat away.
Coloring elements
Element Colors
Titanium Blue
Vanadium Green / Color change
Chromium Red - Green
Manganese Pink
Iron Red - Green - Yellow
Cobalt Blue
Nickel Green
Copper Green - Blue
Absorption of wavelengths (colors) happens because certain elements absorb the energy of those wavelengths. The main elements that absorb color are the "transition metal" elements. These transition metals have partially filled d-shells (opposed to fully filled shells) and some of the electrons in the d-shells are unpaired.
When a wavelength carries enough energy to raise an unpaired electron to a higher energy state, the energy of that wavelength is completey absorbed by the electron and the light energy is usually transformed to heat when the electron falls down to its original (ground) state. This means that that particular wavelength (or color) is removed from the spectrum, or better "absorbed".
The energy required to raise an unpaired electron to a higher energy state (and so cause absorption) is far less than the energy required to raise a paired electron. The energy in visible light is not enough to raise a paired electron, but it can raise the unpaired electrons that reside in the d-shells of the transition metal elements. That is why the transition elements are also named the "coloring elements" or "coloring agents". Besides the transition metal elements, some rare earth metals also act as coloring elements.
The coloring agents can either be part of the ideal chemical makeup or occur as impurities in the crystal.
We divide minerals into two groups, depending on this.
Idiochromatic - minerals that are colored by a coloring agent that is part of the chemical formula (for instance malachite)
Allochromatic - the coloring agents are not part of the ideal chemical composition (such as emerald and ruby)
Although the theory of transition metal elements serves as a good basic understanding for the causes of color, there may be other mechanisms at work. Some elements work together to form color (the molecular orbital theory) while in other cases the absence of an electron or an element at a particular place ("site") in the crystal creates a color center.
What all the theories have in common is that some energy from incoming light (wavelengths) is absorbed and the residual colors (wavelengths) determine the final color of the gemstone (or any other substance).
Advanced
The causes of color can be divided into 4 different theories:
The Crystal Field Theory
Transition metal compounds (malachite, almandine) - idiochromatic
Transition metal impurities (ruby, emerald, citrine, jade) - allochromatic
Color centers (amethyst, maxixe-beryl)
The Molecular Orbital Theory
Charge transfer (sapphire, iolite)
The Band Theory
Insulators (glass)
Conductors (metals)
Semiconductors (galena)
Doped semiconductors (diamond)
The Physical Properties Theory
Dispersion (Fire in diamond)
Scattering (moonstone, cat's eyes, stars)
Interference (iridescence, opal)
Diffraction (opal)
The Crystal Field Theory
The crystal field theory describes colors originating from the excitation of electrons in transition elements (either idiochromatic or allochromatic) and color centers.
When a transition metal ion has a partially filled d-shell, the electrons in the outer d-shell orbit the nucleus unpaired (or atleast some of the electrons do). The surrounding ions of the crystal lattice create a force (a "crystal field") around such a transition element and the strength of those fields determine which energy levels are available for the unpaired electrons. Such a system of energy levels depends on the strength and nature of the bonding in the lattice aswell as on the valence state of the transition element. These are different in every crystal.
As energy and energy levels are quantized, the electrons need a specific amount of energy to "jump" from its groundstate to a higher energy level. The complex calculations that determine which energy levels are available for the electron to excite also provide a few selection rules which exclude some levels for excitation.
Absorption in ruby
In ruby, Cr3+ substitutes some of the Al3+ ions in the Al2O3 lattice. As the chromium is not part of the ideal make-up, ruby is said to be allochromatic. The crystal field around the chromium impurity makes a few, quantized, energy levels available to the unpaired electrons. These are presented as levels B, C and D. However selection rules determine that level B is not available for excitation in this case.
Levels C and D correspond with energies of respectively 2.23 eV and 3 eV. The energy needed to jump to level C (2.23 eV) corresponds with yellow-green light and level D (3 eV) corresponds with violet light.
This means that when white light enters a ruby, yellow-green and violet light will be absorbed by the unpaired electrons and these electrons now have sufficient energy to be excited to levels C or D. The residual colors that are not absorbed determine the red color of the ruby.
The same selection rules also forbid the excited electron to fall back to their ground states (A), but must instead first fall back to B. When the electrons in level B deexcite to their ground state, emmision of red light occurs (fluorescence) which gives an extra glow to the already red color caused by the absorption of the yellow-green and violet portions of white light.
Schematic overview of energy levels in ruby and emerald
For emerald, which allochromatic color is also caused by a Cr3+ impurity, the mechanism is similar but the crystal field from the surrounding elements has less strength and causes a shift in absorption bands. The D level is lowered to 2.8 eV and the C level is lowered to 2.05 eV but the B level is almost the same (1.82 instead of 1.79 eV).
The result is that emerald absorbs most violet and red portions of visible light, leaving a dominant blue-green transmission with a red fluorescence.
Alexandrite, a variety of chrysoberyl, is also colored by a Cr3+ impurity. The absorption scheme of alexandrite is between that of ruby and emerald and the intensity of the incident light determines the color of the alexandrite. Natural daylight is richer in blue-green while incandescent light has more red in its spectrum. This causes the alexandrite to be blue-green (emerald like) in daylight and red-violet (amethyst like) in incandescent light.
Vanadium (V3+) causes the same "alexandrite" color change effect in natural and synthetic corundum.
Allochromatic colors caused by transition metal impurities
The mere occurence of a transition metal impurity does not neccessarily cause the color in a gemstone. Apart from the need of a specific valency the ion must have to be responsible for absorption, other mechanisms (such as color centers and charge-transfer) may be more dominant.
Like certain impurities may be responsible for different colors, as red and green for chromium, other impurities may also cause the same color (like vanadium in emerald).
Idiochromatic colors caused by transition metal constituents
The crystal field theory described above also applies to minerals that have a transition metal ion in its ideal chemical composition.
Color centers
Unpaired electrons on non-transition metal ions may also produce colors in certain circumstances. This can either happen due to misplacement of an ion (and an unpaired electron taking its vacancy), or due to the displacement of an electron from (natural or artificial) radiation.
In both cases a "color center" is created and the unpaired electron can be raised to higher energy levels through the absorption of incident light as with unpaired electrons of the transition metal ions. In the first case (an electron substituting a misplaced ion), the color center is an "electron hole center" and in the latter case it creates a "hole color center".
Electron color centers
Electron color center in fluorite
Fluorite is most often used to describe the mechanism of an electron color center. The purple color of fluorite is caused by the absense of a fluorine (F-) ion and an electron is trapped in the vacancy it leaves behind.
There are various reasons why the fluorine ion is missing from a particular site in the crystal lattice. Among those are an excess of calcium and radiation. either during or after growth of the crystal. This creates a so called "F-center" (or "Farbe center" - Farbe is the German word for color) and a free electron from the pool of unpaired electrons in the crystal (see Band Theory) is trapped in the vacancy. This unpaired electron can then be raised to ,the now available, higher energy levels by the absorption of energy in photons and similar crystal field rules of absorption and fluorescence, as described above, are in effect.
The fluorine ion is usually displaced and creates an interstitual in the lattice, a so called "Frenkel defect", meaning that there is an ion at a particular site in the lattice where it normally would not be. This displaced fluorine ion does not play a role in the development of color (only the vacancy it leaves behind does).
The term "electron color center" refers to the fact that there is a "free" electron where it normally wouldn't be.
Situation A in the image shows the ideal configuration of fluorite, while situation B shows the electron that is trapped in the vacancy left behind by the displaced fluorine ion.
Hole color centers
Hole color center in smokey quartz
Hole color centers are usually illustrated by smokey quartz as in the image on the right.
In quartz (SiO2) some silicon ions with a valence state of 4+ are substituted by aluminium ions with a valency of 3+. In order to keep electroneutrality, a hydrogen atom (or Na+) will be present nearby. This causes the forces on the electrons of the oxygen atoms to be weakened and radiation (X-ray, gamma rays etc.) can remove one of the weaker bonded electrons of the oxygen atoms. This leaves a hole (one electron is missing) and different energy levels become available to the now unpaired electron on the oxygen ion.
The crystal field theory now applies for the remaining, unpaired, oxygen electron and the resulting color is a smokey brown to which smokey quartz owes its name. The displaced electron will be trapped at other sites in the crystal lattice and it doesn't contribute to the color making scheme.
The substitutional aluminum ion acts as a "precursor" and is vital to the mechanism.
In amethyst the workings are similar, but the precursor is ferric iron (Fe3+) and produces the typical purple color.
The term "hole color center" refers to the missing electron, leaving behind a hole.
If the crystal is heated (around 400° C. for smokey quartz and ± 450° C. for amethyst), the displaced electron is freed from it's trap and returns to its original site, traveling as paired electrons again. The color of the crystal will then return to its original color (usually yellow ("citrine") or green ("prasiolite") for amethyst and colorless for smokey quartz). After re-irradiation, the hole color center can be reproduced and aslong as the crystal is not overheated, this process can be repeated infinitly.
The process of heating the crystal and destroying the hole color center is named "bleaching".
The Molecular Orbital Theory
The molecular orbital theory describes the paths (orbitals) electrons travel when multiple atoms (two or more) combine chemically. In order for atoms to combine to molecules, they must share or exchange electrons.
Most diagrams and textbooks, used to explain orbitals in fundamental chemistry (and gemology), show the orbitals or movements of electrons in 2-dimensional rings. Although that suffies for a basic understanding, the orbitals are 3-dimensional and electrons move in complex clouds.
An indepth explanation of those orbitals can be very useful in describing this theory but it is a very complex topic and goes beyond the realm of gemology.
Instead we will focus on the different types of bonds that can occur between atoms.
Ionic bonding
Covalent bonding
While both these bonds have different characteristics, mostly they both play a role in the chemical bonding of gemstones and are directly related to the electronegativity of the different elements that make up the chemical formula of a gemstone.
It are only the electrons in the outer shells of atoms that play a role in both ionic or covalent bonding.
Electronegativity
Electronegativity scale
Electronegativity refers to the force at which certain elements attract eachother. One of the ways the periodic table of elements is arranged is by their electronegativity. In general, the elements in the upper right corner have a stronger electronegativity than the ones in the lower left corner.
As different elements try to combine to form molecules, the elements that have the highest electronegativity are the most greedy and take presidence over the elements with a lower electronegativity.
When there is a large electronegativity difference between two elements, the element with the higher electronegativity will grab the electrons from the other element and this causes a very strong bond between those two elements.
The "greedy" element will have an extra negative charge (because it grabbed negative charged electrons from the other element) and thus polarization between the two elements occurs. This polarization (one element will have a positive charge and the other a negative charge) acts like a magnet and the two elements will be very close to eachother. In other words, they form a very strong bond and there will be an exchange (donation and receiving) of electrons. This is refered to as ionic bonding.
However if the electronegativity difference is very low or the electronegativity is the same, the bonds are not as strong and the electrons in the outer shells of the atoms (elements) are shared between the different elements and those electrons can travel freely (but according to certain rules) between the two elements. This type of loose bonding is named covalent bonding and the atoms that make up the molecule are spaced further apart than in ionic bondings.
Covalent bonding plays a large role in the formation of color as the energy required to exchange the electrons causes absorption. Charge transfer is a common term to describe the continues exchange of electrons in covalent bondings when energy is applied to the molecule.
The main difference between ionic and covalent bondings is that in ionic bondings there is a donation/receiving of electrons, while in covalent bondings there is a sharing of electrons.
Charge transfer
gemologyproject.com/wiki/images/6/61/Charge_transfer2.png
Charge transfer in Iolite
Different elements (especially the transition elements) can excist in different valency states. When those charged atoms (ions) form covalent bondings, some electrons in the outer shells can travel between those two ions. This results in a charge transfer between those two ions and can only happen through absorption of energy. When the energy required for that transition is equivalent to energy in the visible range (or near that) of light, it will result in color.
In the case of iolite it is believed that the color is caused by the charge transfer between Fe2+ and Fe3+ ions. The Fe2+ ion has one more electron in its outer shell than the Fe3+ ion. That electron is attracted to the Fe3+ ion and when the electron orbits the Fe3+ ion, that ion will become an Fe2+ ion. The original Fe2+ ion will then become an Fe3+ ion and the process repeats.
The usual notation for such a transfer is: <math>Fe^{2+} + Fe^{3+} \rightarrow Fe^{3+} + Fe^{2+}</math>
For sapphire the workings are similar, but the transition is between Fe2+ and Ti4+. Although there are many other elements that could play a role in the coloration of sapphires, the general notation for (blue) sapphire is: Fe2+ + Ti4+ → Fe3+ + Ti3+
The Band Theory
The band theory was originally developed to explain electrical conductivity in metals. Later this theory was further developed with quantum mechanics and with that it could also explain colors in materials (such as gemstones).
Electrical conductivity depends on the free travel of negative charged electrons through a material. In order for an the electron to travel through a solid (like a metal or a crystal), there should be room for the electron to travel.
Imagine you are in a room full of people and you would want to walk to the other end of the room. In a fully packed room you could not cross it as there are too many people in it. Now imagine that there is a balcony above the room where there are no people. When some of the people would have enough energy in there body to jump onto the balcony, those people now have freedom to move. While those people move to the balcony, they leave spaces behind in the room so the rest of us have more space to move around.
Of course for people to jump onto the balcony, they need the energy to do so. If they don't have that energy, they will just fall back. So there is a gap between the room and the balcony they need to bridge.
The same thing applies to electrons in a solid.
The band (gap) theory illustrated
In previous theories the focus was on single or two charged atoms (ions). Some electrons are in filled energy levels around the atoms and are very localized in the sense that they have a very strong binding force to the nucleus. Yet the electrons in the outer (unfilled) shells can play a role in binding with other atoms, either through ionic or covalent bondings. Those electrons are named valence electrons and are loosly bound to the nucleus. In reality there is a very large number of valence electrons in a crystal (a multitude of billions per cubic millimeter).
If we would throw all those looser bound valence electrons in a pool of electrons we would get an imaginary pool of valence electrons. That "pool" is named the valence band.
As the valence band if completely filled with electrons, no travel of electrons is possible. So there can be no electrical conductivity. If the electrons would have enough energy to break from this valence band and enter the conduction band (the balcony), then electrical conductivity will be possible. In various materials the band gap (the distance between the groundfloor and the balcony) differs.
Depending on the energy required for an electron to overcome this gap, we divide materials into three catergories:
Insulators (large gap with no electrical conductivity)
Conductors (small to no gap to overlapping energy levels with a good electrical conductivity)
Semiconductors (intermediate between insulators and conductors)
In insulators the gap is so large that under normal conditions (room temperature) there is not enough energy for an electron to bridge the large gap between the valence and the conduction band. As a consequence materials with a large gap will be electrical insulators aswell as colorless. The last part (colorless) is logical because the solid can not absorb energy from visible light. Glass is an example of a good insulator.
In conductors (such as metals) there is a very small band gap or no band gap atall (usually the band and conduction gaps overlap). Therefor electrons can easily be in the conduction band. As little energy is required for an electron to go from the valence band to the conduction band, these materials will be opaque because all of the energy in visible light will be absorbed. The ideal situation would result in an opaque black material, but other forces may turn it into a colored opaque material (such as yellow for gold).
Semiconductors are intermediate between the two states mentioned above and we divide semiconductors into two categories:
Intrinsic semiconductors (pure elements)
Extrinsic semiconductors (doped elements)
Doping means that there are impurities with a different valence state in the crystal lattice, like in allochromatic gemstones.
Lewis dot structure of diamond.
Intrinsic semiconductors with a large band gap usually behave as insulators (such as diamond). Every carbon atom in the diamond structure has 4 valence atoms and they try to combine with 4 other carbon atoms to form a tetrahedral molecular structure. This means that in a pure diamond structure the valance band is fully occupied and the conduction band is empty (so no electrical conduction is present). Yet if we could find a means to make room in the valence band and/or to occupy the conduction band, then the diamond could be an electrical conductor.
For this we would need an impurity to substitute for a carbon atom in the "ideal" diamond structure.
Two types of doped semiconductors
If we would replace a carbon (C) atom with a nitrogen (N) atom, which has 5 valency electrons instead of 4, then we would have one electron too many. That excess electron would create an extra energy level inside the band gap below the conduction band. Some electrons in this extra electron energy level can now be excited into the conduction band. This type of a doped semiconductor is named an N-type semiconductor (N after negative charge from the extra electrons). Yellow type IIa diamond is an example of this type.
The opposite works as well, if we introduce boron (B) in the lattice, which has 3 valence electrons, the diamond would have a positive charge and it would create a hole in the energy gap (another term for a band gap). An electron from the valence band is excited to fill the hole and that creates a hole in the valence band. A neighboring electron in the valence band will fill that hole and the hole seems to be moving through the valence band, which creates conductivity. This type is named a P-type semiconductor (P after the positive charge created by the hole). Blue type IIb diamond is an example of this. Irradiated blue diamond does not conduct, so a distinction can be made between them.
The Physical Properties Theory
Although the previous theories describe the causes of colors originating from electron interactions, there are a few other ways colors can be created. Mostly they show as patches of colors caused by dispersion, scattering, interference and/or diffraction.
All these phenomena are explained in other chapters.
Sources
Kurt Nassau (1978) - The origins of color in minerals
The Physics and Chemistry of Color, 2nd Edition,Kurt Nassau (2001)