Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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First Thoughts

Coordination Number: To address this question, we must ﬁrst identify the ligands

and the central metal to which they are attached. In part c, the formula is written

as a neutral compound, not only as the coordination complex. The species in

brackets represents the coordination complex. It is only the coordination complex

that we should consider when answering a question about the coordination

number. Remember that oxalate (ox) is a bidentate ligand.

Geometry: Once we know the coordination number of the complex, we can

determine the geometry of the complex. For coordination number 2 or 6, the

geometry is linear or octahedral, respectively. For four-coordinate complexes,

either tetrahedral or square planar geometry is expected. Remember that metals

with eight d orbital electrons typically form square planar complexes.

Solution

Coordination Number a. 6 b. 2 c. 4 d. 6

Geometry a. Octahedral b. Linear c. Square planar d. Octahedral

Further Insights

Remember that the “coordination” of a ligand to a metal is a covalent bond, and the

same concepts that you learned earlier for covalent bonds still apply. In short, the

metal and ligand mutually attract an electron pair.

PRACTICE 20.2

Determine the coordination number and geometry for each of the following

compounds.

a. [Cr(CO)

] b. [Fe(H

]

c. K

[TiCl

]d.[Cu(NH

)

]

See Problems 17, 18, and 23–26.

20.5 Isomers

Organic molecules (Chapter 12) are not unique in their ability to form isomeric

structures. Coordination compounds can also form isomers. These types of iso-

mers are divided into two main categories on the basis of whether there is a

change in the structure or in the geometry of the complex. Structural differences

are observed in the linkage isomers, ionization isomers, and coordination sphere

isomers. Differences in geometric isomers arise from the nonspeciﬁc nature of

the formation of a coordinate covalent bond

between a ligand and a metal center.

Some ligands coordinate to a metal

through one of several donor atoms. For exam-

ple, the thiocyanate ion (SCN

−

) can coordinate

to a chromium(III) ion either through the

nonbonded pair of electrons on the nitrogen

or through the nonbonded pair of the elec-

trons on the sulfur atom:

878 Chapter 20 Coordination Complexes

Geometric

isomers

Coordination-

sphere isomers

Ionization

isomers

Linkage

isomers

Isomers

Stereoisomers

Structural

isomers

Isomers in coordination compounds.

Cr(H

NCS

Cr(H

SCN

Video Lesson: Structures of

Coordination Compounds and

Isomers

These are called linkage isomers. They differ only in the atom that participates in

the coordinate covalent bond with the metal.

Another isomer that is common among the coordination compounds is the

ionization isomer. These are species in which a ligand and an ion (a counter ion)

exchange roles, as with the cobalt complexes shown in Figure 20.15. Note that the

counter ion and a ligand have changed positions in the two compounds. It is that

change that makes these two compounds ionization isomers.

Coordination sphere isomers contain different ligands in the coordination

spheres of cations and anions. In the coordination sphere isomers, both the cation

and anion are coordination complexes. Examples are shown in Figure 20.16. Note

that each of these complexes includes a cation that is a coordination complex and

20.5 Isomers 879

[Co(NH

)

]Cl [Co(NH

)

Cl]SO

FIGURE 20.15

Ionization isomers.

[Pt(NH

)

][CuCl

]

[Co(NH

)

][Cr(NH

)(NO

)

] [Cr(NH

)

(NO

)][Co(NH

)(NO

)

]

[Cu(NH

)

][PtCl

]

FIGURE 20.16

Examples of coordination sphere isomers.

an anion that is a coordination complex. The metal ions in each case have

switched location from the cation to the anion.

Geometric isomers, which we ﬁrst discussed in Section 12.3, are substances in

which all of the atoms are attached with the same connectivity, or bonds, but

the geometric orientation differs. Square planar four-coordinate complexes of

the general composition MA

and octahedral complexes of the composition

, where M is the metal and A and B are different ligands, can exist as geo-

metric isomers labeled either cis or trans. As shown in Figure 20.17 for the plat-

inum complex [Pt(NH

)

], the cis isomer has two identical ligands next to

each other. The

trans isomer has two identical ligands on opposite sides of the

metal. Octahedral complexes can exist as cis or trans isomers in complexes such

as [Co(NH

)

]

(Figure 20.18). In fact, in the days when Alfred Werner

(1866–1919) ﬁrst deﬁned the nature of coordination complexes (for which he

won the 1913 Nobel Prize in chemistry), he noted the two isomers of

[Co(NH

)

]Cl. He called one the Praseo complex and the other the Violeo

complex. These two isomers were identiﬁed easily, and named to reﬂect their

identiﬁcation, because the difference in their color was quite evident. To be sure,

the color of a coordination complex is markedly inﬂuenced by the geometric

arrangement of the ligands around the metal center. Table 20.2 gives a summary

of the various isomers common among coordination compounds.

880 Chapter 20 Coordination Complexes

Isomers of Coordination Complexes

Isomer Example Explanation

Linkage Ligand binds through different donor atoms

Ionization Anion and ligand interchanged

Coordination sphere Distribution of coordinating ligands differs

Geometric Orientation of

ligands around the

metal center differs

TABLE 20.2

Co(NH

)

Co(NH

)

ONO

Pt(NH

)

Cl Br Pt(NH

)

Br Cl

Co(en)

Cr(ox)

Cr(en)

Co(ox)

cis

trans

FIGURE 20.17

Cis and trans isomers are geometric isomers. In the cis isomer shown here, the two chloro ligands

are on the same side of the complex. The chloro ligands are on opposite sides in the trans complex.

HERE’S WHAT WE KNOW SO FAR

■

Coordination complexes are present in simple metal ions in solution and as

the reaction center in many biological molecules.

■

A Lewis base that donates a lone pair of electrons to a metal to form a coordi-

nate covalent bond forms a coordination complex.

■

The coordination numbers of various metal centers are commonly observed

to be 2, 4, or 6.

■

Common coordination geometries are linear, tetrahedral, square planar, and

octahedral.

■

A variety of isomer classes are observed for metal complexes.

20.6 Formulas and Names

Coordination compounds are quite varied in their structure, bonding, ability

to form isomers, and other features. For this reason, assigning speciﬁc names to

these compounds must be done with care. Just as in naming binary compounds

(Chapter 2) and organic compounds (Chapter 12), IUPAC rules will guide us in

constructing the proper name.

Formulas

Coordination complexes (the metal center and the ligands) can be neutral, an-

ionic, or cationic in nature. To ensure continuity from one structure to another,

we often write the formula of a coordination compound in accordance with a set

of rules. Those rules are outlined in Table 20.3.

Nomenclature

Systematic names for coordination compounds are constructed in accordance

with a set of rules. Those rules are outlined in Table 20.4. Let’s use these rules to

name a few examples. Consider that we are interested in naming the neutral com-

plex [PtCl

(NH

)

]. Step 1 from Table 20.4 indicates that we should name the

compound using the rules in step 3. The ligands are named alphabetically before

20.6 Formulas and Names 881

cis-Diamminedichloroplatinum(II)

ClCl

trans-Diamminedichloroplatinum(II)

NCo

cis



FIGURE 20.18

Cis and trans isomers are also

possible in octahedral complexes.

NCo

trans



Rules for Writing

Formulas for Coordi-

nation Compounds

1. The metal is written ﬁrst,

followed by the ligands.

2. Anionic ligands are written

before the neutral ligands,

each in alphabetical order.

3. The complex is enclosed in

brackets.

4. Polyatomic ligands, such as

or NO

−

, are enclosed

in parentheses.

5. Ionic compounds containing

coordination complexes are

written in the traditional way:

cation on the left, anion on

the right.

TABLE 20.3

we name the metal. Thus the ammine ligand is named before the chloro ligand.

Moreover, there are two of each of these ligands, so we should write down

diamminedichloro

without any spaces. Then, in step 3d, we write the name of the metal and its

charge in parentheses immediately afterward. Because the charge on the entire

complex is zero, and the complex is neutral, we do not add a sufﬁx to the name:

diamminedichloroplatinum(II)

The name we have created for the formula indicates the number and type of each

ligand, and the metal and its oxidation state. If we knew the three-dimensional

arrangement of the atoms in this complex, we could also include that in-

formation in the name. For example, if we knew that the structure indi-

cated a cis arrangement of the atoms, we could designate that by writing

cis-diamminedichloroplatinum(II). Without that knowledge, though, we can

provide only the name of the formula.

If we wish to name [Co(NH

)

]Cl

, we can do so by following the rules. In

step 1, we note that the compound is made up of a cationic complex and some

anionic counter ions. Step 2 in Table 20.4 directs us to write the name of the com-

plex (the cation). The ligands are named ﬁrst, like this:

hexaammine

Then in step 3d we name the metal and its charge. Again, we do not add a sufﬁx,

because the complex is a cation:

hexaamminecobalt(III)

Finally, steps 6 and 7 tell us to add the counter ions by naming them as we did in

Chapter 3. Note that we don’t add a preﬁx to their name.

hexaamminecobalt(III) chloride

An anionic complex is not much different. Suppose we wished to name

[NiCl

]. In step 1, we note that the anion is a complex and the cation is a

counter ion. Accordingly, we name the cation ﬁrst.

potassium

882 Chapter 20 Coordination Complexes

Rules for Naming Coordination Compounds

1. Name the cation,and then the anion. If the complex is neutral, name it using step 3.

2. Is the cation a complex ion?

3. Ye s – name it using these rules, and then skip to step 5. No – skip to step 4.

a. Name the ligands ﬁrst using Table 20.1.

b. Name the ligands in alphabetical order. Anionic ligands often end in “o,” as in bromo,

hydroxo, and sulfato.

c. If more than one ligand of the same type is present, a preﬁx indicates the number of

units. For simple ligands, the preﬁxes di-, tri-, tetra-, penta-, and hexa- are used. For

complex ligands, the preﬁxes bis-, tris-, tetrakis-, pentakis-, and hexakis- are used.

d. Name the metal last, with its oxidation number in parentheses in Roman numerals.

(1) If the complex is negatively charged, the name of the metal ends in -ate.

(2) If the complex is positively charged or neutral, no sufﬁx is added to the metal name.

4. Name the cation using the conventions described in Chapter 3.

5. Is the anion a complex ion?

6. YES – Name it using the rules in step 3, and then stop. NO – Go to step 7.

7. Name the anion using the conventions described in Chapter 3, and then stop.

TABLE 20.4

Video Lesson: Naming

Coordination Compounds

20.7 Color and Coordination Compounds 883

The complex anion is named using step 3. The ligands go ﬁrst (note that there are

four of them):

potassium tetrachloro

And the metal is part of an anion, so it gets the sufﬁx -ate:

potassium tetrachloronickelate(II)

NiCl

Co(NH

)

Co(NO

)

(NH

)

Examples of coordination compounds. See if you can name them all.

EXERCISE 20.3 Naming Compounds

Give the name for each of these compounds.

a. [Cr(NH

)

(en)

]SO

b. (NH

)

[Fe(CN)

]

First Thoughts

Naming coordination compounds requires us to identify the oxidation states of the

metal center in addition to understanding and knowing the charges on the individ-

ual ligands and ions. To answer this question accurately, separate the compounds

into their two halves (cation and anion). For the half containing the metal ion,

determine its oxidation state, and then use the rules in Table 20.4 to name it.

Solution

a. diamminebis(ethylenediamine)chromium(II) sulfate

b. ammonium hexacyanoferrate(III)

Further Insight

The names for these compounds have a different sound to them, compared to the

simpler compounds we worked with in Chapter 3. Just remember to follow the rules

in Table 20.4, and be sure that you know the names of the ligands from Table 20.1.

PRACTICE 20.3

Name each of these compounds.

a. [CrCl

(NH

)

]

b. Na

[Ni(CN)

]

Draw the formula for each of the following coordination compounds.

c. calcium hexaﬂuoroferrate(II)

d. tetraamminedicarbonylmanganese(II) sulfate

See Problems 37–40.

20.7 Color and Coordination Compounds

Many painters create images that both catch our eye and induce a feeling in our

mind. The goal of the painter is often to portray more than just a staged scene.

Van Gogh’s painting Starry Night (see Figure 20.19) evokes a certain sense of

wonder. The effect is generated by the brush strokes and the colors applied to the

canvas. Are chemical compounds responsible for the color of paint? The answer

is a resounding yes. Colors used in paintings are often constructed from ancient

Video Lesson: Color and

Transition Metals

formulations that include compounds of transition metals. One of the general

characteristics of transition metals is the color of their compounds. Compare

table salt (NaCl), baking soda (NaHCO

), chalk (CaCO

), and Epsom salts

(MgSO

) to ruby (with Cr

), emerald (with Cr

), and rust (Fe

· nH

O, or

hematite). Chromium was so named because of the variety and brilliance of the

colors of its compounds.

What causes compounds that contain transition metals to

have such striking colors?

Transition Metals and Color

Because color is a common feature of transition metal compounds, it seems rea-

sonable that there must be some common characteristics that give rise to their

colors. The presence of color in most transition metal compounds can be attrib-

uted to the presence of partially ﬁlled d orbitals in the compounds and to the in-

ﬂuence of the coordination environment on the energies of the d orbitals. Let’s

see how the color of transition metal compounds is related to fundamental

atomic structure principles.

We perceive color when our eye detects light rays that differ from the ordinary

distribution of those present in white light. For example, a sweater appears blue

if white light strikes it, colors complementary to blue are absorbed, and the re-

maining light is reﬂected. A glass of fruit punch may appear red if white light

strikes it and the colors complementary to red are absorbed. The remaining light,

which appears red to our eye, is transmitted through the liquid and gives the liq-

uid its red color. Color, then, is the array of light rays that our eyes observe being

reﬂected from or transmitted through an object.

There are many objects that can absorb or transmit certain wavelengths of

light but not be “colored.” This occurs when the light being absorbed or trans-

mitted is outside of the visible region (400–700 nm) of the electromagnetic spec-

trum. Recall from Chapter 6 that ultraviolet (UV) light has shorter wavelengths,

typically deﬁned as light in the wavelength range of 200–400 nm. A photon of UV

light carries a lot of energy and may be damaging to substances it strikes. Infrared

light has longer wavelengths than visible light.

Given our understanding of light (Chapter 6), we can begin to ask questions

about how a compound can appear to be colored.

For example, what characteris-

tic change occurs in a substance when light is absorbed?

Answering this question

will enable us to understand why lime (CaO) is colorless but rust (Fe

) is col-

ored. Several fundamental principles are involved.

884 Chapter 20 Coordination Complexes

FIGURE 20.19

Van Gogh’s Starry Night uses the colors

of transition metals to evoke the feeling

of magic in the night sky.

Application

HEMICAL

ENCOUNTERS:

Transition Metals

and Color

750 nm

650 nm

580 nm

430 nm

490 nm

400 nm

560 nm

The color wheel indicates the range of

wavelengths that corresponds to each

color.

Absorption of a photon of visible light by a compound results in the

excitation of an electron from a low energy orbital to a higher energy or-

bital within the substance. In order for the photon to be absorbed, the

difference in orbital energy levels must match the energy of the photon

absorbed. Finally, the excited electron must be able to be excited to the

higher energy orbital—that is, the higher energy orbital must not be full. For ex-

ample, when visible light strikes rust, an electron in a lower energy orbital ab-

sorbs a photon of blue light. The electron is excited to a higher energy orbital.

(Usually the high-energy electron dissipates its energy as heat and relaxes back to

its original energy level, ready to absorb again!) Orange light is reﬂected, and that

is what our eyes detect. Rust is orange-colored.

In many transition metal compounds, of which rust is one, the same process

occurs. What electronic transitions take place to allow the absorption of visible

light? The d orbitals of the transition metal are often arranged such that they

have a slight difference in energy. One example is shown in Figure 20.20. This

difference in energy is roughly that of a photon in the visible region of the elec-

tromagnetic spectrum. If the d orbitals are partially ﬁlled, as they are in the tran-

sition metal ions, a photon can cause an electron in the lower energy orbitals to

jump to an empty (or partially ﬁlled) higher energy d orbital.

Let’s use this information to compare the colors of rust (Fe

) and chalk

(CaCO

). The iron(III) ion in rust has the electron conﬁguration [Ar]3d

. It con-

tains partially ﬁlled d orbitals as its valence orbitals. These d orbitals on the metal,

surrounded by a ﬁeld of oxide ions, are split into different energy levels in ways

that we will discuss in the next section. When visible light strikes rust, some wave-

lengths of light are absorbed, and an electron is excited to a higher empty energy

level. The calcium ion in lime is isoelectronic with argon. It does not have par-

tially ﬁlled d orbitals, cannot absorb visible light energy, and reﬂects all of the vis-

ible light to our eyes. It appears to be white, as shown in Figure 20.21.

20.7 Color and Coordination Compounds 885

400450500550600650750

IR UV

700

Wavelength (nm)

RYGBVIO

The visible spectrum.

Light

absorbed

FIGURE 20.20

A slight difference in the energy of the

ﬁve d orbitals on a transition metal in a

complex allows visible light to be ab-

sorbed. This causes an electron to be ex-

cited to the higher-energy d orbitals.

Lime Rust

FIGURE 20.21

The color of lime and that of rust are a result of

the absorption of speciﬁc wavelengths of light

in the visible region of the electromagnetic

spectrum. Calcium ions in lime do not possess

electrons in d orbitals and cannot absorb light

in the visible region. Iron ions in rust have d or-

bital electrons and do absorb visible light.

xy,

–y

Crystal Field Theory

We can understand the nature of the d orbital splitting from careful analysis of

geometry and d orbital shape. In a free, gaseous metal atom or ion, there is no dif-

ference in energy among the orbitals of the same sublevel. We say that these

orbitals, such as the 3d orbitals on gaseous iron, are degenerate. But something

happens when we bring other atoms (such as the negatively charged anions) up

close to the metal. They cause distortions in the orbitals of the metal. In a six-

coordinate octahedral complex, for example, ions approach the central metal

from each end of the three coordinate axes. Their orientation causes the orbitals

on the metal to lose their degenerate nature.

Let’s look just at the xy plane in an octahedral complex. The octahedral lig-

ands are oriented along the x axis and the y axis. The d orbitals on a metal at

the center of our plane are oriented in speciﬁc ways. Which electron would

have lower energy, an electron in a d

−y

orbital pointed directly at a negatively

charged ligand or an electron in a d

orbital pointed between the negatively

charged ligands? Using Figure 20.22 as a graphical guide, we can answer this

question. Because like charges repel, an electron in a d

−y

orbital would be

higher in energy. We would represent this on an energy-level diagram as follows:

886 Chapter 20 Coordination Complexes

–y

–

–––

–

–––

–

FIGURE 20.23

d orbitals in an octahedral crystal ﬁeld.

Ligands interacting

with a d

–y

2 orbital

Ligands interacting

with a d

orbital

FIGURE 20.22

As ligands approach the d orbitals, they

affect the relative energy levels of the

5d orbitals. Some, like the d

– y

orbital,

have an increased energy due to the

repulsion of similar charges.

Octahedral Crystal Field Splitting

For a metal center in an octahedral crystal ﬁeld, an electron in a d

−y

orbital

would be at higher energy than an electron in a d

orbital. We can extend this

analysis to three dimensions and include all d orbitals as shown in Figure 20.23.

Note that the d

−y

and d

orbitals are pointed right at the negatively charged lig-

ands, but the d

, d

, and d

orbitals are pointed between the axes. The d

and d

orbitals will be at lower energy than the d

−y

and d

orbitals. Based on

Video Lesson: Crystal Field

Theory

the symmetry of the system, the d

, d

, and d

orbitals are at the same energy

level. Although it is less obvious, the d

−y

and d

orbitals are at the same energy

level, which is higher than that of the d

, d

, and d

orbitals.

In an octahedral array of anions or ligands, the d orbital energies of the cen-

tral metal are split into a lower energy set (d

) that we label the t

set and

an upper energy set (d

−y

, d

) that we label the e

set, as shown in Figure 20.24.

The energy difference between the t

orbitals and the e

orbitals in the octahedral

ﬁeld, the

crystal ﬁeld splitting energy, is given the symbol ∆

, where the subscript

“o” indicates the splitting energy for an octahedral complex. The magnitude of

∆

, the difference in energy between the t

and e

orbitals, depends on the nature

of the central metal and the nature of the ligands.

Tetrahedral Crystal Field Splitting

A tetrahedral ligand ﬁeld can be viewed as one with the metal ion at the center of

a cube and ions at alternate corners of a cube, with the coordinate axes going

through the centers of the faces of the cube as shown in Figure 20.25. In this

situation, none of the d orbitals of a central atom are pointed at the ions. The d

, and d

lobes are closer to the ions than the d

−y

and d

lobes. Thus the d

, and d

orbitals are at a higher energy than the d

−y

and d

orbitals. As

shown in Figure 20.25, the tetrahedral crystal ﬁeld splitting energy, ∆

, where the

subscript “t”indicates a tetrahedral complex, is qualitatively the inverse of the oc-

tahedral ﬁeld splitting diagram. We label the lower orbital set “e” and the upper

orbital set “t”. As the comparison in Figure 20.26 shows, the magnitude of the

splitting, ∆

, is estimated to be

⁄

9 the size of ∆

, so the magnitude of tetrahedral

crystal ﬁeld splitting is smaller than that of octahedral ﬁeld splitting.

20.7 Color and Coordination Compounds 887

Tetrahedral crystal field splitting

∆

Octahedral crystal field splitting

∆

FIGURE 20.24

Crystal ﬁeld splitting of metal d orbitals for

octahedral and tetrahedral geometries.

Square Planar Crystal Field Splitting

We can deﬁne the square planar geometry as one with the

four ligands in the xy plane as in an octahedral complex,

but with the two opposing ligands on the z axis com-

pletely removed. The crystal ﬁeld effect in the xy plane

remains the same, but the effect in the z direction is re-

moved. This results in stabilization of the d orbitals with

z axis components. The resulting crystal ﬁeld splitting

diagram is shown in Figure 20.27. The difference in en-

ergy between d

and d

− y

is still ∆

. The other energy

differences are considerably smaller.

–y

–

FIGURE 20.25

Tetrahedral crystal ﬁeld splitting (arrows show the distance from orbital lobe to ions).

∆

–y

FIGURE 20.26

The value of ∆

⁄9 that of

∆

, the distance in energy

from the d

–

to the d

orbital.

FIGURE 20.27

Square planar coordination complex.

∆

–y +y

–x

Ligand arrangement of a

square planar complex

d Orbital energies in a

square planar comple

–y