Studies of the colors and magnetic properties of transition-metal complexes have played an important role in the development of modern models for metal–ligand bonding. We discussed the various types of magnetic behavior of the transition metals in Section 23.1, and we discussed the interaction of radiant energy with matter in Section 6.3. Let's briefly examine the significance of these two properties for transition-metal complexes before we develop a model for metal–ligand bonding.


In Figure 23.3 we saw the diverse range of colors seen in salts of transition-metal ions and their aqueous solutions. In general, the color of a complex depends on the identity of the metal ion, on its oxidation state, and on the ligands bound to it. FIGURE 23.24, for instance, shows how the pale blue color characteristic of [Cu(H2O)4]2+ changes to deep blue as NH3 ligands replace the H2O ligands to form [Cu(NH3)4]2+.


Is the equilibrium binding constant of ammonia for Cu(II) likely to be larger or smaller than that of water for Cu(II)?

FIGURE 23.24 The color of a coordination complex changes when the ligand changes.

FIGURE 23.25 Two ways of perceiving the color orange. An object appears orange either when it reflects orange light to the eye (left), or when it transmits to the eye all colors except blue, the complement of orange (middle). Complementary colors lie opposite to each other on an artist's color wheel (right).

For a substance to have color we can see, it must absorb some portion of the spectrum of visible light.  (Section 6.1) Absorption happens, however, only if the energy needed to move an electron in the substance from its ground state to an excited state corresponds to the energy of some portion of the visible spectrum.  (Section 6.3) Thus, the particular energies of radiation a substance absorbs dictate the color we see for the substance.

When an object absorbs some portion of the visible spectrum, the color we perceive is the sum of the unabsorbed portions, which are either reflected or transmitted by the object and strike our eyes. (Opaque objects reflect light, and transparent ones transmit it.) If an object absorbs all wavelengths of visible light, none reaches our eyes and the object appears black. If it absorbs no visible light, it is white if it is a solid or colorless if it is a liquid. If it absorbs all but orange light, the orange light is what reaches our eye and therefore the color we see.

An interesting phenomenon of vision is that we also perceive an orange color when an object absorbs only the blue portion of the visible spectrum and all the other colors strike our eyes. This is because orange and blue are complementary colors, which means that the removal of blue from white light makes the light look orange (and the removal of orange makes the light look blue).

Complementary colors can be determined with an artist's color wheel, which shows complementary colors on opposite sides (FIGURE 23.25).

SAMPLE EXERCISE 23.7 Relating Color Absorbed to Color Observed

The complex ion trans-[Co(NH3)4Cl2]+ absorbs light primarily in the red region of the visible spectrum (the most intense absorption is at 680 nm). What is the color of the complex?


Analyze We need to relate the color absorbed by a complex (red) to the color observed for the complex.

Plan For an object that absorbs only one color from the visible spectrum, the color we see is complementary to the color absorbed. We can use the color wheel of Figure 23.25 to determine the complementary color.

Solve From Figure 23.25, we see that green is complementary to red, so the complex appears green.

Comment As noted in Section 23.2, this green complex was one of those that helped Werner establish his theory of coordination (Table 23.3). The other geometric isomer of this complex, cis-[Co(NH3)4Cl2]+, absorbs yellow light and therefore appears violet.


A certain transition-metal complex ion absorbs at 630 nm. Which color is this ion most likely to be—blue, yellow, green, or orange?

Answer: blue

The amount of light absorbed by a sample as a function of wavelength is known as the sample's absorption spectrum. The visible absorption spectrum of a transparent sample can be determined using a spectrometer, as described in the “A Closer Look” box on page 564. The absorption spectrum of the ion [Ti(H2O)6]3+ is shown in FIGURE 23.26. The absorption maximum is at 500 nm, but the graph shows that much of the yellow, green, and blue light is also absorbed. Because the sample absorbs all of these colors, what we see is the unabsorbed red and violet light, which we perceive as red-violet.


How would this absorbance spectrum change if you decreased the concentration of the [Ti(H2O)6]3+ in solution?

FIGURE 23.26 The color of [Ti(H2O)6]3+. A solution containing the [Ti(H2O)6]3+ ion appears red-violet because, as its visible absorption spectrum shows, the solution does not absorb light from the violet and red ends of the spectrum. That unabsorbed light is what reaches our eyes.

Magnetism of Coordination Compounds

Many transition-metal complexes exhibit paramagnetism, as described in Sections 9.8 and 23.1. In such compounds the metal ions possess some number of unpaired electrons. It is possible to experimentally determine the number of unpaired electrons per metal ion from the measured degree of paramagnetism, and experiments reveal some interesting comparisons.

Compounds of the complex ion [Co(CN)6]3– have no unpaired electrons, for example, but compounds of the [CoF6]3– ion have four unpaired electrons per metal ion. Both complexes contain Co(III) with a 3d6 electron configuration.  (Section 7.4) Clearly, there is a major difference in the ways in which the electrons are arranged in these two cases. Any successful bonding theory must explain this difference, and we present such a theory in the next section.


What is the electron configuration for

a. the Co atom and

b. the Co3+ ion? How many unpaired electrons does each possess? (See Section 7.4 to review electron configurations of ions.)