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Chapter 23 – Transition Metals and Coordination Chemistry

Many compounds of transition metals are colored (used in paints and to stain glass; produce color in

gemstones).

23.1 The Transition Metals

Most metals, including transition metals, are found in solid inorganic compounds known as minerals.

Minerals are named by common, not chemical, names.

The properties of the transition metals are similar to each other (and very different to the properties of the main

group metals).

The similarities in properties come from similarities in valence electron configuration.

Electron Configuration

For 1st & 2nd transition series = ns2(n−1)dx

1st = [Ar]4s23dx; 2nd = [Kr]5s24dx

For 3rd & 4th transition series = ns2(n−2)f14(n−1)dx

Some individuals deviate from the general pattern by “promoting” one or more s electrons into the underlying d

to complete the subshell

– Group 1B

Form ions by losing the ns electrons first, then the (n – 1)d

Irregular Electron Configurations

Due to sublevel splitting, the 4s sublevel is lower in energy than the 3d; and therefore the 4s fills before the 3d.

But the difference in energy is not large.

Some of the transition metals have irregular electron configurations in which the ns only partially fills before

the (n−1)d or doesn’t fill at all; therefore, their electron configuration must be found experimentally.

Expected Found Experimentally

Cr = [Ar]4s23d4 Cr = [Ar]4s13d5

Cu = [Ar]4s23d9 Cu = [Ar]4s13d10

Mo = [Kr]5s24d4 Mo = [Kr]5s14d5

Ru = [Kr]5s24d6 Ru = [Kr]5s14d7

Pd = [Kr]5s24d8 Pd = [Kr]5s04d10

This leads to some exceptions on the overall trend.

Half-filled and filled d subshells are specially stable.

Transition metal ions lose e-‘s from s orbital first:

Fe: [Ar]4s23d6 Fe2+: [Ar]3d6 Fe3+: [Ar]3d5

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Atomic Radii

Why this trend?

1. Increasing nuclear charge tends to make atoms smaller.

2. However, the strongest (and, therefore, shortest) metallic bonds are found in the center of the transition

metals.

First these two factors cooperate, but then they counteract each other.

Traits of Transition Metals

Because most transition metals have only partially occupied d subshells, the metals and/or their compounds

often:

exhibit multiple oxidation states (vary by 1)

Highest oxidation state is the same as the group number for Groups 3B to 7B.

Because most transition metals have only partially occupied d subshells, the metals and/or their compounds

often:

exhibit multiple oxidation states (vary by 1).

Are pigmented.

Because most transition metals have only partially occupied d subshells, the metals and/or their compounds

often:

exhibit multiple oxidation states (vary by 1).

Are pigmented.

Have magnetic properties.

Magnetism

Diamagnetic substance – all 𝑒−′𝑠 are paired (weakly repelled from a magnetic field).

Paramagnetic – at least one 𝑒− is unpaired in the electronic structure (attracted to a magnetic field).

23.2 Transition-Metal Complexes

When a monatomic cation combines with multiple monatomic anions or neutral molecules it makes a complex

ion.

The attached anions or neutral molecules are called ligands.

The charge on the complex ion can then be positive or negative, depending on the numbers and types of ligands

attached.

When a complex ion combines with counter-ions to make a neutral compound it is called a coordination

compound.

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The primary valence is the oxidation number of the metal.

The secondary valence is the number of ligands bonded to the metal:

– coordination number

Coordination numbers range from 2 to 12, with the most common being 6 and 4:

𝐶𝑜𝐶𝑙3 ∙ 6𝐻2𝑂 = [𝐶𝑜(𝐻2𝑂)6]𝐶𝑙3

Swiss chemist Alfred Werner (1893):

Some molecules/ions are bonded directly to the metal and are part of the “coordination sphere”.

The rest are just counterions (present to balance the charge).

Assume a coordination number of 6 for 𝐶𝑜.

Table 23.3

The first compound splits into 4 ions: 3 free 𝐶𝑙− (not bonded to the metal).

The second compound: 2 free 𝐶𝑙−.

- Test conductivity (# ions it splits into).

- Add a ppt reagent to test for # chloride ions (like 𝐴𝑔𝑁𝑂3).

Compounds 3 and 4 have 1 free 𝐶𝑙−: different arrangements.

Q. Pd tends to form complexes with coordination number 4. A compound has the composition 𝑃𝑑𝐶𝑙2 ∙ 3𝑁𝐻3.

a) What is the correct formula?

b) if excess 𝐴𝑔𝑁𝑂3 (𝑎𝑞) is added, how many moles of 𝐴𝑔𝐶𝑙(𝑠) will ppt per formula unit?

The Metal–Ligand Bond

Reaction between a Lewis acid (the metal) and Lewis base (the ligand).

The new complex has distinct physical and chemical properties.

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Metal ions in water are attached to several water ligands:

𝑁𝑖(𝑎𝑞)2+ ⇒ 𝑁𝑖(𝐻2𝑂)6 (𝑎𝑞)

2+

The coordination number depends on the metal and the ligands.

Q. What is the oxidation number of the metal in:

[𝑅ℎ(𝑁𝐻3)5𝐶𝑙](𝑁𝑂3)2

Q. A complex contains: 𝐶𝑟3+, 4𝐻2𝑂, 2𝐶𝑙−. Formula?

Common Geometries

Coordination Number

Shape Model Example

2

Linear

[𝐴𝑔(𝑁𝐻3)2]

+

4

Square Planar

[𝑃𝑑𝐶𝑙4]

2−

4

Tetrahedral

[𝑍𝑛(𝑁𝐻3)4]

2+

6

Octahedral

[𝐹𝑒(𝐻2𝑂)6]

3+

23.3 Common Ligands in Coordination Chemistry

Some ligands are polydentate (i.e. take up 2 or more coordination sites on the metal). They are also known as

chelating agents. (Table 23.4)

Complexes containing polydentate ligands are especially stable. Their formation is entropically favorable:

𝐶𝑜(𝐻2𝑂)6 (𝑎𝑞)2+ + 𝐸𝐷𝑇𝐴4− ⇌ 𝐶𝑜(𝐸𝐷𝑇𝐴)2− + 6𝐻2𝑂

2 particles 7 particles

more disorder

𝐸𝐷𝑇𝐴4− can be used to “tie up” free metal ions in solution. (Added to some canned foods salad dressings, etc.

Treatment for lead poisoning.)

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Many transition metals (V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Mo) are components of “metalloproteins” in the body.

In the metalloprotein, the metal forms a complex with groups in the protein acting as ligands.

Chelates in Biological Systems

Common ligand: porphine.

Complex of porphine and metal is known as porphyrin.

Variations possess diff. metals, diff. groups attached to porphine.

This type of complex is a component of myoglobin (stores oxygen),

hemoglobin (transports oxygen in blood) and

chlorophyl (needed for photosynthesis in plants).

The iron in hemoglobin carries O2 and CO2 through the blood.

Carbon monoxide and cyanide are poisonous because they will bind more tightly to the iron than will oxygen.

Read “The Battle for Iron in Living Systems” (p. 987)

23.4 Nomenclature and Isomerism in Coordination Chemistry

1. Name the cation and then the anion.

2. The name of the complex is one single word.

Ligands first (in alphabetical order), and the metal is last.

Prefixes are not considered for the alphabetical order.

3. Anionic ligands end in the letter o, but electrically neutral ligands ordinarily bear the name of the molecules;

exceptions are: (Table 23.5)

For example:

[𝐹𝑒(𝐶𝑁)2(𝑁𝐻3)2(𝐻2𝑂)2]+

diamminediaquadicyanoiron(III) ion

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4. Greek prefixes (di-, tri-, tetra-, etc.) are used to indicate the number of each kind of ligand when more than

one is present.

If the ligand contains a Greek prefix or is polydentate, the prefixes bis-, tris-, tetrakis-, etc. are used and the

ligand name is placed in parentheses.

[𝐶𝑜(𝑒𝑛)3]𝐵𝑟3 tris(ethylenediamine)–cobalt(III) bromide

5. If the complex is an anion, its name ends in -ate.

Some metals use their Latin names:

Fe ferrate Sn stannate Ag argenatate

Cu cuprate Pb plumbate Au aurate

[𝐶𝑜𝐶𝑙4]2− 𝐾4[𝐹𝑒(𝐶𝑁)6]

tetrachlorocobalate(II) Potassium hexacyanoferrate(II)

6. The oxidation number of the metal is given in parentheses in Roman numerals following the name of the

metal.

[𝑁𝑖(𝑁𝐻3)6]𝐵𝑟2 Hexaamminenickel(II) bromide

[𝐶𝑜(𝑒𝑛)2(𝐻2𝑂)(𝐶𝑁)]𝐶𝑙2 Aquacyanobis(ethylendiamine)cobalt(III) chloride

𝑁𝑎2[𝑀𝑜𝑂𝐶𝑙4] Sodium tetrachlorooxomolybdate(IV)

More Examples

[𝑁𝑖(𝐻2𝑂)6]𝑆𝑂4 Hexaaquanickel(II) sulfate

[𝐶𝑟(𝑒𝑛)2(𝐶𝑁)2]𝐶𝑙 Dicyanobis(ethylenediamine)cromium(III) chloride

𝐾[𝑃𝑡(𝑁𝐻3)𝐶𝑙3] Potassium amminetrichloroplatinate(II)

𝑁𝑎2[𝐹𝑒𝐸𝐷𝑇𝐴] Sodium ethylendiaminetetraacetatoferrate(III)

𝐾[𝐴𝑔(𝐶𝑁)2] Potassium dicyanoargentate(I)

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Q. Name or give the formula of:

1. [𝐶𝑢(𝑁𝐻3)4]𝐶𝑙2

2. 𝐾2[𝐶𝑜𝐶𝑙4]

3. [𝐶𝑜(𝑒𝑛)2(𝐻2𝑂)𝐶𝑙]𝐵𝑟2

4. Lithium tetrahydroxostannate (II)

5. Sodium diaquedioxalatoaurate (III)

Isomerism

Isomers: Compounds with the same molecular formula, but different structures.

1. Structural isomers: differ in connectivity of atoms. 𝐻 − 𝑁 = 𝐶 = 𝑂 𝑣𝑠 𝑁 ≡ 𝐶 − 𝑂 − 𝐻

for complexes, 2 subcategories:

A. Linkage isomers – the ligand is bound to the metal by a different atom.

B. Coordiantion-sphere isomers differ in what ligands are bound to the metal and which fall outside the

coordination sphere.

𝐸𝑥𝑎𝑚𝑝𝑙𝑒: 𝐶𝑟𝐶𝑙3(𝐻2𝑂)6

[𝐶𝑟(𝐻2𝑂)6]𝐶𝑙3

Coordination-Sphere Isomers [𝐶𝑟(𝐻2𝑂)5𝐶𝑙]𝐶𝑙2 ∙ 𝐻2𝑂

[𝐶𝑟(𝐻2𝑂)4𝐶𝑙2]𝐶𝑙 ∙ 2𝐻2𝑂

2. Stereoisomers – same order of connections, but different arrangement in space.

A. Geometric isomers - the ligands have a different spatial relationship.

More polar, more soluble No anticancer activity.

in 𝐻2𝑂. Anticancer drug.

No cis-trans isomers for tetrahedral complexes.

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Facial – three identical ligands Meridional – three identical ligands

facing each other. form an arc around the center.

B. Optical isomers:

Enantiomers – nonsuperimposable mirror images (like our hands).

A molecule that is nonsuperimposable on its mirror image is chiral (a chiral molecule would be

superimposable).

Common type of chiral molecule:

Tetrahedral geometry with 4 different substituents.

Square planar compounds have no optical isomerism (since they are flat).

Optical isomers rotate polarized light in opposite directions (dextro or levorotatory).

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Q. What type of isomerism is possible? Draw the isomers.

a. [𝐹𝑒(𝐻2𝑂)3𝑂3] b. [𝑅𝑢(𝐶2𝑂4)3]

4−

c. [𝑃𝑡(𝐶𝑁)2𝐶𝑙2]2− (square planar)

23.5 Color and Magnetism in Coordination Chemistry

Transition metal ions show many intense colors in host crystals or solution.

The color of light absorbed by the complexed ion is related to electronic energy changes in the structure of the

complex:

the electron “jumping” from a lower energy state to a higher energy state.

E = ℎ𝜈 =ℎ𝑐

𝜆

The observed color is the complementary color of the one that is absorbed.

If something absorbs all 𝜆’s of visible light, it looks black.

If something absorbs no 𝜆’s of visible light, it looks colorless or white.

Q. The complex [Ti(H2O)6]3+ has an absorbance maximum at 500 nm, which corresponds to green light. What

color is this complex?

Aside: Review of orbital shapes and orientations

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23.6 Crystal-Field Theory

Ligands are assumed to behave like ⊝ point charges.

The ligands exert an electric field, and the energies of the d orbitals on the metal are altered (due to 𝑒− − 𝑒−

repulsions).

In an octahedral field : 6 ligands on 𝑥, 𝑦, 𝑧 axes.

d orbitals are degenerate (same E) on the free metal ion.

No longer degenerate on the octahedral field.

Why does the particular split?

Because of repulsions between ⊝ charges, d electrons on metal are more stable in orbitals pointing between

ligands (𝑑𝑥𝑦, 𝑑𝑦𝑧 , 𝑑𝑥𝑧).

Less stable in orbitals pointing directly toward ligands (𝑑𝑥2𝑦2 , 𝑑𝑧2).

The size of the crystal field splitting energy (∆𝐸 between d levels = ∆), depends on the kinds of ligands and

their relative positions on the complex ion, as well as the kind of metal ion and its oxidation state:

Strong ligand ⇒ large splitting (big ∆𝐸)

include CN─ > NO2─ > en > NH3

Weak ligand ⇒ small splitting (small ∆𝐸)

include H2O > OH─ > F─ > Cl─ > Br─ > I─

The colors of complex ions are due to electronic transitions between the split d sublevel orbitals.

The wavelength of maximum absorbance can be used to determine the size of the energy gap between the split

d sublevel orbitals:

𝐸𝑝ℎ𝑜𝑡𝑜𝑛 = ℎ𝜈 =ℎ𝑐

𝜆= ∆

The spectrochemical series ranks ligands in order of their ability to increase the energy gap between d orbitals.

weak Increasing ∆ strong

I─< Br─ < Cl─ < F─ < OH─ < H2O < NH3 < en < NO2─ < CN─ < CO

CO is a very strong ligand. Poisonous! It binds to 𝐹𝑒2+ in hemoglobin much more tightly than 𝑂2 does (blood

cannot get oxygen around).

The stronger the crystal-field strength of the ligand, the larger the energy gap between d orbitals, and the shorter

the wavelength of light absorbed by the complex.

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Electron Configurations in Octahedral Complexes

Recall that when transition metals form cations, s electrons are lost first. Thus, 𝑇𝑖3+ is a 𝑑1 ion, 𝑉3+ is a 𝑑2 ion,

and 𝐶𝑟3+ is a 𝑑3 Ion.

If one to three electrons add to the d orbitals in an octahedral complex ion, Hund’s rule applies. (The first three

electrons go into different d orbitals with their spins parallel.)

We have a choice for the placement of the fourth electron:

If it goes into a higher-energy orbital, then there is an energy cost (Δ).

If it goes into a lower-energy orbital, there is a different energy cost (called the spin-pairing energy “P” due to

pairing with the electron already present).

Weak-field ligands tend to favor adding electrons to the higher-energy orbitals (high-spin complexes) because

Δ is less than the spin-pairing energy. 𝑃 > ∆

Strong-field ligands tend to favor adding electrons to lower-energy orbitals (low-spin complexes) because Δ is

greater than the spin-pairing energy. 𝑃 < ∆

For 𝑑4 − 𝑑7 ions (octahedral field), both high and low spin arrangements are possible.

For 𝑑1 − 𝑑3 and 𝑑8 − 𝑑10, only one possible way of arranging 𝑒− in each.

Can determine # unpaired 𝑒− experimentally.

Paramagnetic – at least 1 unpaired 𝑒−.

Diamagnetic – all 𝑒−’s paired.

Q. Out of the following complex ions, one is high spin and the other is low spin:

𝐶𝑜(𝑁𝐻3)63+, 𝐶𝑜𝐹6

3−

Which is which? Draw the splitting diagram for each. How many unpaired 𝑒− in each?

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Another factor that influences crystal-field splitting is the size of d orbitals.

3𝑑 < 4𝑑 < 5𝑑

Larger orbitals lead to more effective overlap and stronger interactions so 4d and 5d complexes are always low–

spin.

Tetrahedral Complexes

By using the same arguments as for the octahedral case, we can derive the relative orbital energies for d orbitals

in a tetrahedral field.

Since the ligands interact more strongly with the planar orbitals in the tetrahedral geometry, their energies are

raised.

This reverses the order of energies compared to the octahedral geometry.

All tetrahedral complexes are high spin because of reduced metal orbital – ligand interaction.

Square Planar Geometry and Crystal Field Splitting

The most complex splitting pattern.

Square-planar complexes can be thought of as follows: Start with an octahedral complex and remove two

ligands along the z-axis.

As a consequence the four planar ligands are drawn in towards the metal.

Most d8 metal ions form square-planar complexes (𝑃𝑡2+, 𝑃𝑑2+, 𝐼𝑟+, 𝐴𝑢3+).

The majority of complexes are low-spin (i.e., diamagnetic).

Q. 𝑁𝑖(𝑁𝐻3)42+ is paramagnetic, while 𝑁𝑖(𝐶𝑁)4

2− is diamagnetic. Draw the splitting diagram for each complex.

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Q. 𝐹𝑒(𝐻2𝑂)6 (𝑎𝑞)2+ is light blue-green. Would you expect the 𝐹𝑒2+ ion in this complex to have a high-spin or

low-spin configuration?

Q. An octahedral complex of 𝐶𝑟3+ with unknown ligand A looks bluish-green. A complex of 𝐶𝑟3+ with

unknown ligand B looks purple. Which is the stronger ligand, A or B? Explain. Draw splitting diagrams.

Q. 𝑍𝑛2+ compounds are colorless. Explain why.