Chemistry 105
The Chemistry of Color

Course Lecture Topic Information

Chem 105 Web Site
Dr. Kimberly Lawler-Sagarin
Elmhurst College

Discussion 11:
Rubies are red and sapphires are blue

This discussion will take place May 2 - May 8

Assignments for the discussion board

  • Q1: Based on your student number, answer the following question (see web links to answer this question. Include information about bonding in these compounds, examples and physical properties):

    • students 1,2,3,4: What is a covalent network solid?
    • students 5,6,7,8: What is a metallic solid?
    • students 9,10,11,12: What is a molecular solid (can also be called a Van der Waals solid)?
    • students 13,14,15: What are ionic solids?
  • Q2: For this question, look up information about the causes of color in the following gemstones (see links at the end of the page to get started with some basic information):
    • Students 1,6,11: Emeralds
    • Students 2,7,15: Diamonds
    • Students 3,8,14: Choose from: Quartz, Amethyst or Citrine
    • Students 4,9,13: Choose from two beryl-based gems: Aquamarine or Heliodor (also called golden beryl)
    • Students 5,10,12: Choose from: Tourmaline or Garnet

  • Q3: Comment on the answer of another student, either by adding to the answer or asking a followup question. -OR- Answer another student's question.

Text Readings These are the same as last week. Most of the new material is contained in the virtual lecture notes, but I draw heavily on last weeks material.
  • Check out these two websites on the structure of solids:
    Topic Background

    This week, we continue to explore transition metals and d electrons in more detail.

    Electronic Structure of Transition Metals

    The "d" Orbitals

    Transition metals mark the region of the periodic table where the d subshell begins to be filled. Most transition metals have a partially filled d subshell. This partially filled shell is responsible for many of the physical and chemical properties of the transition metals.

    The d subshell consists of 5 orbitals, so can hold a maximum of 10 electrons. Here are images of the five d orbitals.

    image of dz2 orbital image of dxy orbital image of dxz orbital
    dz2 dxy dxz
    image of dyz orbital image of dx2-y2 orbital
    dyz dx2-y2

    Compounds formed by Transition Metal

    Transition metals are found in a wide variety of compounds. I'll mention two common types here that are relevant for our discussion of color.

    • Coordination compounds - a small molecule or ion consisting of a metal surrounded by a number of ligands. Ligands are simply chemical species that will form bonds to a metal center. A generic formula for such a compound would be MLn - a metal surrounded by "n" ligands (denoted "L"). Some common small ligands include H2O, NH3, Cl-, OH-, CO and CN-. However, there are also much larger ligands, such as cyclopentadiene (C5H5) and tri-t-butylsiloxide ( OSi(C(CH3)3)3 ).

    • Ionic solids - these are solids consisting of a repeating array of ions held together by ionic bonds. The positive and negative ions pack together in the solid by alternating with one another in a repeating pattern - rather like the tiling of a floor. (see web reading links for examples). In an ionic solid, a transition metal will be positively charged, and will generally be surrounded by a number of negative ions. For example, in TiO2 (consisting of Ti4+ and O2- ions), Ti4+ ions are surrounded by six O2- ions. (These O2- ions are shared by other Ti4+ ions, so that the 2:1 ratio of oxygen to titanium is preserved.)

    d Orbitals and Color

    Transition metal compounds are often very brightly colored. This is primarily a result of the energies of the d orbitals. In an isolated transition metal atom, the d orbitals all have the same energy. However, once other chemical species surround a transition metal atom, this is no longer true. The d orbitals in such compounds will have different energies. The five orbitals are said to be "split".

    The pattern of splitting and the energy differences observed depends on two things: (1) the type of species surrounding the metal atom and (2) the geometry or arrangement in space of these surrounding species.

    Electrons will generally fill the lowest energy orbitals first, thus there is often a small energy gap between the the filled and unfilled orbitals. It turns out that the energy differences between the various d orbitals is actually quite small. Transition metal complexes can absorb light in the visible region of the spectrum by exciting an electron from one of the lower energy d orbitals to a higher energy empty d orbital.

    Common Geometries for Transition Metal Complexes

    There are many possible geometries for transition metal complexes. Three common geometries are:

    • Octahedral - this is a structure adopted by many ML6 complexes (a metal with six ligands) and is also a transitional metal environment found widely in solids. The octahedral arrangement consists of six ligands placed evenly around a central atom or ion (the transition metal in our case). Neighboring ligands make L-M-L angles of 90 degrees from one another. If we consider a Cartesian coordinate system where the transition metal is in the center, the size ligands lie directly on the x, y and z axes.

    • Tetrahedral - this is the familiar tetrahedral structure we find for many main ground compounds. Many ML4 compounds have this structure. A perfect tetrahedron consists of a central atom or ion (the transition metal in this case) surrounded by four ligands at 109.5 degree angles from one another.

    • Square Planar - In this arrangement, four ligands radiate out from the central transition metal in one plane, forming a square. The angle between neighboring ligands is 90 degrees. This is essentially an octahedral arrangement with two ligands removed.

    Each of these geometries gives rise to a different d orbital splitting pattern.

    Crystal Field Theory

    Crystal field theory is a simple way to describe the d orbital splitting patterns found in transition metal complexes. To introduce this theory, we will consider the case of an octahedral transition metal complex, ML6

    . For simplicity, we will arrange our coordinate system so that the ligands lie on the x, y and z axes, as shown below.

    Now, imagine an isolated transition metal atom or ion. This atom or ion will have five d orbitals with equivalent energies:

    Now imagine six ligands approaching the transition metal. They will bond to the metal by sharing their electrons with the metal center, often by sharing their lone pair electrons.

    From the metal's point of view, negatively charged electrons are approaching it. As these ligands approach, the metal's d orbitals that have lobes pointing directly toward the ligands are raised in energy. Why? In a simplified view of crystal field theory, these orbitals are now less likely to contain the transition metal's electrons, as these electrons would be repelled by the negative charge of the electrons coming in from the ligands. This results effectively results in raising the energies of these orbitals.

    In an octahedral complex, the dx2-y2 and dz2 orbitals have large lobes pointing toward the ligands, thus these two orbitals will rise in energy. The other three will remain at their original location.

    This leads to a "three below two" pattern of d orbital energies. Electrons will generally occupy the bottom orbitals only.

    A similar analysis of a tetrahedral ML4 complex will give just the opposite pattern. Because of the different geometry, the dxy, dxz and dyz orbitals are the orbitals pointing more directly toward the ligands. Thus, these three orbitals are pushed up in energy forming a "two below three" pattern. Furthermore, because tetrahedral complexes only have four ligands compared to the size ligands found in octahedral complexes, the splitting in tetrahedral complexes is not as large as it is in octahedral complexes.


    Sapphires are precious gemstones made of the mineral corundum. Corundum is a very hard mineral - it rates a 9 on the hardness scale (diamond has a hardness of 10). Corundum is a form of aluminum oxide, an ionic solid with the formula Al2O3, formed from Al3+ ions and O2- ions. Pure corundum is colorless. You may wonder then what makes sapphires blue!

    It may surprise you to learn that sapphires get their blue color from impurities. Blue sapphires contain a small amount of Fe2+ (iron 2+) and Ti4+ (titanium 4+) ions in their crystal lattice. These ions replace a very small number of the Al3+ ions in corundum. It is the iron and titanium ions that give rise to the beautiful deep blue color of sapphires. Only a few hundredths of a percent of the titanium and iron impurities are needed to produce the bright color.

    The mechanism that produces the color is called charge transfer. When light of appropriate frequency is absorbed, an Fe2+ ion transfers an electron to a neighboring Ti4+ ion. In doing so, the iron ion is oxidized, whereas the titanium ion is reduced. The iron goes from Fe2+ to Fe3+, and titanium goes from Ti4+ to Ti3+:

    Fe2+ + Ti4+ + light --> Fe3+ + Ti3+

    The Fe3+ + Ti3+ combination lies higher in energy, however the difference in energy is small, falling into the visible range. A schematic absorption spectrum for sapphire is shown below. Note the absorption occurs primarily on the red end of the spectrum.

    Sapphires actually come in colors other than blue. Such sapphires are called "fancy sapphires" and are found in all different shades of pink, purple, yellow, orange and green. Different impurities are responsible for the different colors. Sapphires which are red fall into a different category - they are called rubies.


    Yes, rubies are actually red sapphires! Like sapphires, rubies are also made of the mineral corundum. However, rubies contain the impurity Cr3+. This chromium impurity is responsible for the bright red color of rubies. In order to achieve a bright red color, about 1% of the aluminum ions must be replaced by chromium ions.

    These chromium ions are found in an octahedral environment - they are surrounded by six oxide ions. (Recall, Al2O3 is formed from Al3+ ions and O2- ions.) This will lead to the "three below two" d orbital splitting pattern we saw earlier.

    A neutral chromium atom has 6 valence electrons and has the following electron configuration:


    The arrangement above is actually lower in energy than the [Ar]4s23d4 arrangement you might expect, because half filled subshells are especially stable. This is partially due to the fact that it costs energy to put two electrons into the same orbital - they are both negatively charges, so they repel one another.

    When chromium loses three electrons to form Cr3+, it will have a total of three valence electrons. These three electrons will go into the lower energy d orbitals.

    When exposed to light of an appropriate wavelength, electrons are promoted into the higher energy d orbitals. A schematic absorption spectrum for ruby is show below. Note that ruby absorbs in the violet-blue and yellow-green regions. Thus, there is strong red transmission, giving rise to the strong red color.

    Interestingly, rubies also undergo fluorescence. Once the Cr3+ ion is in an excited state caused by the absorption of violet light, it may lose some energy as heat, falling down to an intermediate energy level. It then returns to the ground state by fluorescence. The fluoresced light is also red, adding to the stone's brilliance.

    Gemstone Links

Chemistry 105 Home
Blackboard Home
Elmhurst Chemistry Home
Elmhurst College Home

 K. Lawler-Sagarin, 2004-2005