Download 6. d and f-Block Elements and Coordination Chemistry

Chemistry 2810 Lecture Notes
6.
Dr. R. T. Boeré
Page95
d and f-Block Elements and Coordination Chemistry
Sc
Y
Ti V Cr Mn
Zr Nb Mo Tc
Hf Ta W Re
Rf Db Sg Bh
Fe Co Ni Cu Zn
Ru Rh Pd Ag Cd
Os Ir Pt Au Hg
Hs Mt
La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The transition elements comprise those metals in Groups 3 to 12, i.e. the d-block elements, as well as the so-called “inner
transition elements” which are the lanthanides and actinides, also know as the f-block elements. These are all metallic
elements, so they are often called the transition metals, which we take to be synonymous with the transition elements. The
transition metals include many of the most important structural metals for an industrial society, including iron (Fe),
chromium (Cr), nickel (Ni) and copper (Cu). The text-book describes how some of these metals are won from their ores. Our
focus in lecture is going to be on coordination complexes of these elements.
We could study the properties of the transition elements in crystalline solids. However, there are complexities. First of
all, many additional crystal structure types would have to be studied. Many of these structures involve a large amount of
covalent bonding, so that our simple ionic model is not completely adequate. But in reality, we could do so, and the decision
to move to soluble coordination compounds is a deliberate one. The study of the transition metals in solids have been
dominated by solid-state physicists, and is closely tied to important technical advances, such as the development of the first
Ruby lasers. Chemists have been more involved in studying soluble metal complexes, and today these have become extremely
important compounds for industrial chemical processes (e.g. in the petrochemical and pharmaceutical industries), as well as
in medicine an as model systems for many plant and animal proteins and enzymes.
There is a strong conceptual relationship between the metals in a solid and in solution. In a coordination compound, we
suround the metal ion by the same kind of donor atoms as are in the anions found in their crystalline compounds. Thus, to
mimic an oxide environment, we can simply use water, or another oxygen-containing species. Of course, the metal ions will
act as Lewis acids (both hard and soft acids are found among the transition elements, as well as everything in between), and
the donor atoms act as Lewis bases. Thus coordination compounds are nothing more or less than Lewis acid-base complexes.
You should review your tables of hard and soft acids and bases at this time.
If we want to mimic a crystalline metal halide (from the hard fluoride to the soft iodide), we can employ the halogen
anions as the donor atoms. Alternatively, we can employ donor atoms such as nitrogen or carbon that are rare in solids,
although nitrides and carbides of these metals do exist. In summary, a coordination compound is a Lewis acid/base complex
that solubilizes the metal ion while preserving the local environment found in crystalline solids of these metals. The
coordination number of the central ion, the geometry and the nature of the donor atom found in solids can often be mimicked
quite accurately in soluble complexes.
6.1
the d-block metals have multiple oxidation states
The transition elements are those which
have open d-subshells (BACK POINTER: atom
electron configurations).
There are several
important consequences which are due to these
d-orbitals. First of all, they are the cause of the
large number of common oxidation states for
transition metals, as seen in the diagram at the
left. The most common oxidation states are
given by the square entries. Notice the pattern
that the ends of the series, there are fewer
common oxidation states, whereas in the middle
(i.e. exactly where the d-orbitals become halffilled) there is the greatest variety. Removing d-electrons is relatively easy compared to s,p electrons, and this is the reason
that high-oxidation state forms of these elements can be obtained.
A second consequence of d-electrons is that coordination compounds are especially stable among the transition elements.
It is important to stress that all metallic and metalloid elements form coordination compounds, but they are definitely most
well established and most stable among the transition elements.
Chemistry 2810 Lecture Notes
6.2
Dr. R. T. Boeré
Page96
Introduction to coordination chemistry
A metal cation acts as a Lewis acid. A Lewis acid is a substance that can accept a pair of electrons from another atom to
form a new bond.
The electron pairs which are received come from surrounding groups which are called ligands. Ligands are typically
either neutral or anionic atoms or molecules. Ligands act as Lewis bases. A Lewis base is a substance that can donate a pair
of electrons to another atom to form a new bond. (BACK POINTER: review the Lewis theory which we have already covered.)
The combination of a metal cation and all its ligands is called a coordination complex. These difficult-sounding
definitions are much clearer from a few examples.
First, we construct the Lewis diagrams of the ligands H2O, Cl-, and CO
-
¨:
:Cl
:O:
:C O:
H
H
¨
is dissolved in water, the resulting Fe2+(aq) ion is in reality a coordination complex between
When a metal such as Fe2+
iron(II) and water:
Fe2+ + 6 H2O →
[Fe(OH2)6]2+ (here water is written OH2 to emphasize that the electron pair on oxygen of
water is acting as the Lewis base donor site to the metal; the charge on the resulting complex ion is simply the sum of the
charges of the metal ion and those on the ligands; square brackets are used to isolated the complex ion, and we refer to the six
water molecules as comprising the primary coordination sphere of the Fe(II) ion.
Copper(II) ions in the presence of a high concentration of chloride ion forms a chloro complex:
Cu2+ + 4 Cl- →
[CuCl4]2- (here the overall charge is -2 since that is the sum of the ligands and the metal
ion)
Nickel when finely divided reacts readily with an atmosphere of gaseous carbon monoxide to form the liquid nickel
tetracarbonyl:
Ni + 4 CO → [Ni(CO)4] (here the overall charge is 0, since the metal and the CO are all uncharged)
In summary, coordination compounds can be overall neutral, cationic or anionic, and the charge is the sum of all the
constituents. Charged complex compounds are called complex ions, and these can form salts with appropriate counter ions.
For example, the complexes mentioned above might be isolated as the salts [Fe(OH2)6]Br2 and K2[CuCl4]
You have already met several transition metal coordination complexes back during the Chemistry 2000 laboratory. For
example, in the first lab you prepared, and then analyzed, the complex salt K3[Fe(C2O4)3].3H2O. Later on, you determine the
relative number of thiocyanate and water ligands in complex ions of the type [Fe(SCN)n(H2O)6 - n](3-n)-. Several of the
compounds you may prepare in lab 6, the Thirteen Mystery Test Tubes, are in fact coordination compounds, while you may
also have met copper complexes [Cu(OH2)4]2+ and [Cu(NH3)4]2+. In the 2810 lab you will study the electronic spectra of a
number of simple aqua, chloro and ammine complexes of the d and f elements.
6.2.1 Common Ligands
There are literally thousands of ligands which have at times been used towards metals to form coordination complexes.
Some are just common small molecules or ions, such as water, ammonia and the halogen and pseudo-halogen ions. Such
species typically have a single donor atom available, and even if this donor has more than one electron lone pair, it is rare for
one atom to donate more than once to a given metal. The most common exception to this rule is the oxo ligand, O2–, for
which M=O double bonds are extremely common. In all such cases where a single donor atom is linked to the metal, the
ligand is said to be monodentate (literally, "single-toothed").
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page97
However, using simple
organic chemistry, it is fairly
easy to link together in one
molecule two or more donor
atoms linked by an inert
"backbone".
These more
complicated ligands are called
chelating ligands, from the
Greek word chele, meaning
“claw”. Chelating ligands can
have two, or more donor sites.
With two donors, they are
called bidentate ("doublytoothed"),
with
three,
The complex ion [Co(en)3]3+
The neutral complex [CoEDTA]
tridentate, etc.
Whereas
iron(III) has room in the primary coordination sphere for six water molecules in the complex ion [Fe(OH2)6]3+, it only has
room for three chelating bidentate oxalate ions in the complex ion [Fe(C2O4)3]3-. Can you see why? The pictures show some
computerr-drawn "ball and stick" diagrams of coordination complexes with chelating ligands. The structure of [Co(en)3]3+ is
that of a complex cation, which can be isolated from solution as a salt with any suitably charged anions. On the other hand,
[CoEDTA] where the cobalt is in the +3 oxidation state can be a neutral molecule (an hence is not particularly soluble in
water.) Alternatively, you should get used to interpreting simple line diagrams of typical coordination complexes, such as the
following examples:
3-
O
OH2
2+
NH 3
H2O
OH 2
OH2
H 3N
NH3
NH3
C
3+
OH2
O
OH2
C
O
C
O
O
Fe
Cu
Cu
H2O
2+
O
Fe
H2O
OH 2
O
OH2
C
O
O
O
C
C
O
O
Perspective line drawings of some typical coordination complexes
We identify the total number of Lewis base attachments to a given metal center as the coordination number. Thus in the
example discussed above, both [Fe(OH2)6]3+ and [Fe(C2O4)3]3- have a coordination number of six. In fact the iron(III) ion
does not particularly care whether the six oxygen donor atoms come from six monodentate water molecules or three
bidentate oxalate ions. One of the most amazing ligands is EDTA, which supplies up to six donor atoms to a single metal
ion, the ligand wrapping around the metal and encasing it. EDTA4- is a hexadentate ligand. In general, chelating ligands
have additional stability over non-chelating or monodentate ligands. One of the reasons that the oxalate ion displaces H2O
during the preparation of [Fe(C2O4)3]3- in the Chem2000 lab experiment is that oxalate is a chelating ligand.
Formula
Name
Abbreviation
H2 O
aqua
use formula
F–, Cl–, Br–, I–
fluoro, chloro,
bromo, iodo
use formula
ammine
use formula
Structure
:O:
H3 N
-
¨
Chg
monodentate
0
monodentate
–1
monodentate
0
monodentate
(through C)
monodentate
(through C)
monodentate
(often M=O)
0
H
H
-
: F̈ :
Donor number
:C̈l :
, ¨
¨
N
, etc.
H
H
H
CO
carbonyl
use formula
CN–
cyano
use formula
O2–
oxo
use formula
:C
O:
- :C
N:
-2
: Ö :
¨
–1
–2
Chemistry 2810 Lecture Notes
CO32–
carbonato
Dr. R. T. Boeré
Page98
O
use formula
C
C 2 O4
2–
oxalato
-O
O¨
ox
H2NCH2CH2NH2
ethyelenediamine
C12H8N2
phenanthroline
C10H12O8
ethylenediaminetetraacetato
-O
¨
en
¨O
O
H2 H2
C C
¨H
C
2
N
¨
-O C
2
bidentate
0
bidentate
0
hexadentate
4–
NH2
N
-O C
–2
-
¨
¨
¨
EDTA
bidentate
C
H2N
phen
–2
-
O
C
bidentate
N
2
H2
C
¨
N
-
¨
CO2
CO2
The table above lists a group of common ligands that you need to learn to be able to answer the assignment and test
questions for this course. You need to remember the name and abbreviation of each complex, the number of donor atoms, as
well as the overall charge on the ligand (typically either zero or some negative number; there do exist a very small number of
positively charged ligands, but these are extremely rare.) All questions will be restricted to complexes using only this small
set of ligands.
6.2.2 Naming Coordination Compounds
Just as there are rules for naming simple inorganic and organic compounds, coordination compounds are named
according to an established system. For example, the following compounds are named according to the rules outlined below.
Compound
[Ni(H2O)6]SO4
[Cr(en)2(CN)2]Cl
K[Pt(NH3)Cl3]
Systematic Name
Hexaaquanickel(II) sulfate
Dicyanobis(ethylenediamine)chromium(III) chloride
Potassium amminetrichloroplatinate(II)
As you read through the rules, notice how they apply to the examples above:
1. In naming a coordination compound that is a salt, name the cation first and then the anion. (This is how all salts are
commonly named).
2. When giving the name of the complex ion or molecule, name the ligands first, in alphabetical order, followed by the
name of the metal.
a. If a ligand is an anion whose name ends in -ite or -ate, the final e is changed to o (as in sulfate → sulfato or nitrite →
nitrito).
b. If the ligand is an anion whose name ends in -ide, the ending is changed to o (as in chloride → chloro or cyanide →
cyano).
c. If the ligand is a neutral molecule, its common name is usually used. The important exceptions to this rule are water,
which is called aqua, ammonia, which is called ammine, and CO, called carbonyl.
d. When there is more than one of a particular monodentate ligand with a simple name, the number of ligands is
designated by the appropriate pre-fix: di, tri, tetra, penta, or hexa. If the ligand name is complicated (whether
monodentate or bidentate), the prefix changes to bis, tris, tetrakis, pentakis, or hexakis, followed by the ligand name
in parentheses.
3. If the complex ion is an anion, the suffix -ate is added to the metal name.
4. Following the name of the metal, the oxidation number of the metal is given in Roman numerals. Complexes can be
considerably more complicated than those described in this chapter; then, even more rules of nomenclature must be
applied. The brief rules just outlined, however, are sufficient for the vast majority of complexes.
Here follow some examples of how the rules can be applied:
1. [Cu(NH3)4]SO4
The sulfate ion has a 2– charge, so the complex ion has a 2+ charge (that is, [Cu(NH3)4]2+ ).
Because NH3 is a neutral molecule, the copper ion is Cu 2+ . The compound’s name is therefore tetraamminecopper(II)
sulfate.
2. K2[CoCl4]
Two K+ ions occur in this compound, so the complex ion has a 2– charge ([CoCl4]2– ). Because four
Cl– ions occur in the complex ion, the cobalt center is Co2+ . Thus, the name of the compound is potassium
tetrachlorocobaltate(II).
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page99
Co(phen)2Cl2
This is a neutral compound. Because two Cl– ions and two neutral phen (phenanthroline) ligands
are bonded to a cobalt ion, the metal ion must be Co2+ . This means the compound name is dichlorobis(phenanthroline)
cobalt(II).
[Co(en)2(H2O)Cl]Cl2 Here the complex ion has a 2+ charge because it is associated with two uncoordinated Cl– ions. The
cobalt ion must be Co3+ because it is bonded to two neutral en (ethylenediamine) ligands, one neutral water, and one Cl– .
The name is aquachlorobis(ethylenediamine)cobalt(III) chloride.
K3[Fe(C2O4)3].3H2O First of all, the .3H2O are waters of crystallization. These are additional water molecules which
are incorporated into empty spaces in the crystals of the product, and are called hydrates. Thus the salt is a trihydrate.
(NB: many so-called hydrates of simple salts of metal ions are in fact metal aqua complexes. Thus CuSO4.5H2O is in fact
actually the complex salt [Cu(OH2)4]SO4.H2O. Just how many water molecules in a given hydrated salt are coordinated to
a metal ion must be established by experiment. It is not something you can know automatically. However, you should
from now on be aware that hydrated salts may in fact contain coordinated water molecules!). Complex salts obey the
same convention as simple salts in that the name of the cation is always given first, followed by the name of the anion.
Thus the first part of the name of the salt is potassium. Since K always gives a 1+ ion, its oxidation state is not specified.
The complex anion is named by giving the name of the ligand first, followed by that of the metal and the oxidation state
of the metal in Roman numerals. If the ligand is an anion whose name ends in ite or ate, the ending is changed to o.
Thus oxalate becomes oxalato. When there is more than one ligand of the same type (the normal situation) the number
is given by di, tri, tetra, penta, hexa, etc. Thus our example is trioxalato. If the complex ion is an anion, the suffix ate is
added to the metal name. Iron is a little bit odd in that we revert to the Latin name for anionic forms. Thus ironate is
never used; instead we call it ferrate. The oxidation state of the iron is +3, so this is given by (III). We are now ready to
name the complete salt which you will prepare next week in the laboratory: potassium trioxalatoferrate(III) trihydrate.
That’s quite a mouthful! You can see why on the whole chemists prefer a picture of a molecule to its name.
3.
4.
5.
6.2.3 Structure and Isomerism
The structure of coordination compounds is a combination of the coordination number (CN) and the possible geometries
of the attached ligands.
If CN = 2, the structure is always linear, for example in [AgCl2]- and [Au(CN)2]-.
Cl Cl
Ag
If CN = 4, the geometry may be one of two structures: square planar or tetrahedral. For example, the complexes
[Cu(OH2)4]2+ and [Cu(NH3)4]2+ are both tetrahedral, as indicated in the picture above. On the other hand, the complex
[Ni(CN)4]2- has the square-planar shape. Note that these shapes in general are not predictable by VSEPR theory. They must
be established by experiment. Isomerism is not possible in tetrahedral complexes unless all four attached groups are different.
This is an extremely rare situation for coordination compounds (but happens fairly frequently for sp3-hybridized carbon
atoms).
One of the consequences of square planar geometry is that cis and trans isomers are possible. Such geometrical isomers
only show up if there are two different types of ligands present on the molecules. A very famous example of this kind of
isomerism occurs for the formula [PtCl2(NH3)2]. The cis form of this neutral coordination complex is a potent anticancer
drug know in medicine as cisplatin. The trans form has no antitumor activity at all, and is useless as a drug.
If CN = 6, the structure is almost always octahedral. This is also the most common geometry for transition metal
coordination complexes. Octahedral complexes have a wide range of possible isomers, depending on the type of ligands
attached. If the general formula is MX4Y2, cis and trans geometrical isomers are possible. On the other hand, if the general
formula is MX3Y3, mer and fac geometrical isomers are possible. Multidentate ligands can change the general picture. For
example M(X∩ X)3 complexes, in which X∩ X represents a chelating bidentate ligand, have only one feasible geometry;
however such complexes do form optical isomers. This means that they are chemically indistinguishable, but rotate plane
polarized light in opposite directions. You will learn more about optical isomers in introductory organic chemistry classes.
You should be aware, however, that optical isomerism is rooted purely in the geometry of the molecule, not in the types of
elements involved. Chelating complexes of the general formula M(X∩ X)2Y2 can exist as cis and trans geometrical isomers.
The cis form of M(X∩ X)2Y2 exists as two optical isomers as well.
NH3
Cl
Pt
NH3
Cl
NH3
trans-[PtCl2(NH3)2]
Cl
Pt
NH3
Cl
H 3N
NH3
Pt
H 3N
NH3
Cl
2+
NH3
2+
Cl
H3 N
NH3
Pt
H3 N
Cl
Cl
+
Cl
H 3N
NH3
Pt
H 3N
cis-[PtCl2(NH3)4]2+
NH3
Pt
Cl
Cl
Cl
Cl
NH3
mer-[PtCl3(NH3)3]+
fac-[PtCl3(NH3)3]+
cis-[PtCl2(NH3)2]
trans-[PtCl2(NH3)4]2+
+
Cl
H 3N
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page100
All of these isomers are due to only to geometry.
You are encouraged to make some models and
prove to yourself that these differences are
possible, and that they are the only possible
variations. In addition, you should obtain the
point groups for each structure. The computergenerated ball-and-stick diagrams at right
illustrate cis and trans isomerism for sixcoordinate octahedral shapes. Note that the cis
isomer in this case involves a chelating ligand, so
it will exist in two optical isomers, i.e. it is an
cis isomer of [Co(en)2Cl2]+
trans isomer of [Co(en)2Cl2]+
example of M(X∩ X)2Y2. Remember from our
section on point group symmetry that optical isomers exist for point groups which do not have any form of Sn symmetry axis,
including a centre of symmetry or a mirror plane. Thus cis-[Co(en)2Cl2]+ belongs to the point group C2, which has E, C2
symmetry elements only.
6.2.4 Electronic Structure and Colour of Transition Metal Coordination Compounds
One thing distinguishes transition metal coordination compounds from all other coordination compounds, and that is
colour. Transition metal complexes are very often coloured, whereas the metals and metalloids of the s and p blocks form
colourless complexes. Consider the nitrate salts of Fe3+, Co2+, Ni2+, Cu2+ and Zn2+. All of these are metal aqua complexes of
the type [M(OH2)n]m+. Why are four of them coloured, while the last is colourless?
a)
Crystal field theory for octahedral and tetrahedral complexes
The origin of these phenomena can be understood by considering what happens to the d orbitals when a coordination
compound forms. Remember what the shapes of the five d orbitals are. Important here is also the orientation of these
orbitals:
The effect on the d orbitals depends on geometry. If the complex is octahedral,
with CN = 6, the most common situation, the ligands are found in the sites
indicated by the dark circles in the figure at right. A simple electrostatic model,
called the crystal field theory, assumes that there will be a significant electronelectron repulsion between the electron pair the ligands donate and any
electrons already in the metal d orbitals. The greatest repulsion is felt by those d
orbitals directed most closely at the ligands.
This leads to an alteration of the energies of the formerly degenerate d
orbitals, and in fact causes dx2-y2 and dz2 to become slightly higher in energy than
the set dxy, dxz, and dyz. The resulting pattern is shown below. We call the
separation between the two sets the crystal field splitting, ∆o, where the “o”
stands for octahedral.
In an octahedral field, electrons in the shaded d
orbitals experience greater repulsion than in the
unshaded d orbitals because of greater geometric
proximity to the ligand lone pairs.
Chemistry 2810 Lecture Notes
Dr. R. T. Boeré
Page101
The effect of an octahedral crystal field on the five metal d orbitals
For tetrahedral complexes, exactly the inverse situation is obtained.
Here the CN=4, and the ligands are located where the dark circles are in the
geometrical diagram at right.
The consequence is that the dxy, dxz, and dyz become higher in energy
than dx2-y2 and dz2. Here the crystal field splitting is called ∆T, standing for
the tetrahedral case. It is easy to show by geometry that the magnitude of
∆T < ∆o, so long as ligands of these same type are used in both cases.
In a tetrahedral field, electrons in the unshaded d orbitals
experience greater repulsion than in the shaded d
orbitals because of greater geometric proximity to the
ligand lone pairs.
Finally, for square planar complexes, a third possibility occurs. The pattern here is more complicated, with dxz, and dyz
lowest, followed by dz2, dxy, and then dx2-y2. The spacing between the last two is called ∆SP, the square-planar crystal field
splitting energy. We will not consider the square planar case, except to say that typically ∆SP is very large.
b)
The spectrochemical series
Changing the type of atoms in the ligand primarily affects the size of the crystal field splitting energies. This ligand
effect is called the spectrochemical series. A simplified form of this series is:
{small orbital splitting, ∆} Halides < C2O42- < H2O < NH3 = en < phen < CN- {large orbital splitting, ∆}
Ligands at the left of the series are often called weak field ligands, while those at the right are strong field ligands.
c)
Possible electron configurations for d-element complexes
The real importance of all this becomes clear when we start putting the electrons in. We need the electron configurations
of the cations, e.g. Fe2+ and Fe3+. These are [Ar]3d6 and [Ar]3d5, respectively. Consider the complexes [Fe(OH2)6]2+ and
[Fe(C2O4)3]4-, iron(II), and [Fe(OH2)6]3+ and [Fe(C2O4)3]3-, iron(III), all of which can be prepared under suitable conditions.
For these cases, and in fact for any configuration from d4 to d7, there are two possible ways of filling the metal orbitals. Let us
see how this works out by putting the electrons into these diagrams for octahedral geometry:
Chemistry 2810 Lecture Notes
d1
d4
Dr. R. T. Boeré
d2
Page102
d3
d5
d6
"high spin" forms
d8
d7
d4
d9
d10
d5
d6
"low spin" forms
d7
The high spin forms predominate for small orbital splitting, and with both water and oxalate ligands, most complexes will be
high spin. So the iron(II) and iron(III) complexes will have 5 and 4 unpaired electrons. On the other hand, if cyanide ligands
were present there would be large orbital splitting and the low spin forms would predominate, so that [Fe(CN)6]4- and
[Fe(CN)6]3- will have 1 and 0 unpaired electrons. Notice also that this last example is diamagnetic (for a definition, see Kotz
& Treichel, p. 358), while a free Fe3+ ion is definitely paramagnetic.
Exercise - Work out the possibilities that exist for tetrahedral crystal fields as a function of the d counts. For which can high
and low spin conditions exist (i.e. those cases where there is a net difference between the two distributions)? In fact, we need
only consider the high spin case, since tetrahedral crystal fields are too small to cause spin pairing.
high spin situation
d1
d4
d8
d2
low spin situation
d3
d5
d6
d9
d10
d1
d7
d4
d8
d2
d3
d5
d6
d9
d10
d7