Inorganic Pharmaceutical Chemistry
Lecture No. 5
Date : 22/11 /2012
Dr. Mohammed Hamed
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Coordination compounds
Coordination compounds have been recognized by chemists as a special type of
compound since late in the eighteenth century . Throughout the Nineteenth
century , a wide variety of "complex inorganic compounds" were prepared by
reacting metal salts with neutral organic or inorganic~oolecules, such as
ethylene or ammonia, or with halide Ions and other simple anions, such as
cyanide and nitrite . In 1892 Alfred Werner, proposed the Coordination Theory.
Werner's Coordination Theory is one of the great landmarks of chemistry and
has proved extremely fruitful in foster inga myriad of developments that have
taken coordination chemistry in to such diverse areas as the use of
organometallic compounds in synthesis and catalysis , new insights in to the role
of metal ions in biochemical reactions , and investigations in to metal-metal
bonding and the formation of clusters .where a number of metal centers come
together in one compound. Today, coordination compounds play key roles in
medicine, agriculture , materials science , electrochemistry , catalysis , and
industrial organic synthesis. Definitions and Terminology
Werner's Coordination Theory
1.
In co-ordination compounds, central metal atoms exhibit primary valence
and secondary valence.
primary valence is ionizable. Secondary valence is not ionizable. The primary
valence corresponds to oxidation state.The secondary valence corresponds to
coordination number. (the central metal ion and ligands are not ionizable)
2. Every metal atom has a fixed number of secondary valences (coordination
number(s)).
3. The metal atom tends to satisfy both its primary valence as well as its
secondary valence. Primary valence is satisfied by negative ions (metal ion has a
positive charge) whereas secondary valence (coordination number) is satisfied
either by negative ions or by neutral molecules. (In certain case a negative ion
may satisfy both types of valences).
4. The coordination number or secondary valencies are always directed towards
the fixed positions in space and this leads to definite geometry of the
coordination compound
A coordination compound is one in which a central atom or ion is joined to
one or more ligands through what is called a coordinate covalent bond in which
both of the bonding electrons are supplied by the ligand. In such a complex the
central atom acts as an electron-pair acceptor (Lewis acid — think of H+ which
has no electrons at all, but can accept a pair from something like Cl –) and the
ligand as an electron-pair donor (Lewis base ). The central atom and the ligands
coordinated to it constitute the coordination sphere. Thus the salt
[Co(NH3)5Cl]Cl2 is composed of the complex ion [Co(NH3)5Cl]2+ and two Cl–
ions; components within the square brackets are inside the coordination sphere,
whereas the two chloride ions are situated outside the coordination sphere.
These latter two ions could be replaced by other ions such as NO3– without
otherwise materially changing the nature of the salt.
The central atoms of coordination complexes are most often cations (positive
ions), but may in some cases be neutral atoms, as in nickel carbonyl Ni(CO)4.
Ligands composed of ions such as F– or small molecules such as H2O
or CN– possess more than one set of lone pair electrons, but only one of these
pairs can coordinate with a central ion. Such ligands are said to be monodentate
(―one tooth‖.) Larger ligands may contain more than one atom capable of
coordinating with a single central ion, and are described as polydentate. Thus
ethylenediamine (shown below) is a bidentate ligand. Polydentate ligands whose
geometery enables them to occupy more than one coordinating position of a
central ion act as chelating agents and tend to form extremely stable complexes
known as chelates.
The coordination sphere of a coordination compound or complex consists
of the central metal atom/ion plus its attached ligands. The coordination sphere
is usually enclosed in brackets when written in a formula.
A complex is a structure composed of a central metal atom or ion, generally
a cation, surrounded by a number of negatively charged ions or neutral
molecules possessing lone pairs. A complex may also be called a coordination
compound or metal complex.
The ions or molecules surrounding the metal are called ligands. A ligand that is
bound to a metal ion is said to be coordinated with the ion. The process of
binding to the metal ion with more than one coordination site per ligand is called
chelation. Compounds that bind avidly to form complexes are thus called
chelating agents (for example, EDTA).
Simple ligands like water or chlorine form only one link with the central atom
and are said to be monodentate. Some ligands are capable of forming multiple
links to the same metal atom, and are described as bidentate, tridentate etc.
EDTA is hexadentate, which accounts for the great stability of many of its
complexes.
Typically, the chemistry of complexes is dominated by interactions between s
and p orbitals of the ligands and the d (or f) orbitals of the metal ions. Because
of this, simple octet bonding theory fails in the case of complexes and to
understand the chemistry of these systems, a deeper understanding of chemical
bonding rules is necessary.
Naming Coordination Compounds
The coordination compounds are named in the following way. (At the end of
this tutorial we have some examples to show you how coordination compounds
are named.)
A. To name a coordination compound, no matter whether the complex ion is the
cation or the anion, always name the cation before the anion. (This is just like
naming an ionic compound.)
B. In naming the complex ion:
1. Name the ligands first, in alphabetical order, then the metal atom or ion.
Note: The metal atom or ion is written before the ligands in the chemical
formula.
2. The names of some common ligands are listed in Table 1.
For anionic ligands end in "-o"; for anions that end in "-ide"(e.g. chloride), "ate" (e.g. sulfate, nitrate), and "-ite" (e.g. nirite), change the endings as follows: ide
-o; -ate
-ato; -ite
-ito
For neutral ligands, the common name of the molecule is used e.g.
H2NCH2CH2NH2 (ethylenediamine). Important exceptions: water is called
‗aqua‘, ammonia is called ‗ammine‘, carbon monoxide is called ‗carbonyl‘, and
the N2 and O2 are called ‗dinitrogen‘ and ‗dioxygen‘.
Table 1. Names of Some Common Ligands
Anionic
Ligands
Names
Neutral Ligands Names
Br-
bromo
NH3
ammine
F-
fluoro
H2O
aqua
O2-
oxo
NO
Nitrosyl
OH-
Hydroxo
CO
Carbonyl
CN-
cyano
O2
dioxygen
C2O42-
oxalato
N2
dinitrogen
CO32-
carbonato
C5H5N
pyridine
CH3COO-
acetato
H2NCH2CH2NH2 ethylenediamine
2. Greek prefixes are used to designate the number of each type of ligand in
the complex ion, e.g. di-, tri- and tetra-. If the ligand already contains a
Greek prefix
one binding site) the prefixes bis-, tris-, tetrakis-, pentakis-, are used instead.
The numerical prefixes are listed in Table 2.
Table 2. Numerical Prefixes
Number Prefix
Number Prefix
Number Prefix
1
mono
5
penta
9
(pentakis)
nona
(ennea)
2
di (bis)
6
hexa
(hexakis)
10
deca
3
tri (tris)
7
hepta
11
undeca
4
tetra
(tetrakis)
8
octa
12
dodeca
3. After naming the ligands, name the central metal. If the complex ion is a
cation, the metal is named same as the element. For example, Co in a
complex cation is call cobalt and Pt is called platinum.. If the complex ion
is an anion, the name of the metal ends with the suffix –ate.. For example,
Co in a complex anion is called cobaltate and Pt is called platinate. For
some metals, the Latin names are used in the complex anions e.g. Fe is
called ferrate (not ironate).
Table 3: Name of Metals in Anionic Complexes
Name of Metal
Name in an Anionic Complex
Iron
Ferrate
Copper
Cuprate
Lead
Plumbate
Silver
Argenate
Gold
Aurate
Tin
Stannate
6. Following the name of the metal, the oxidation state of the metal in the
complex is given as a Roman numeral in parentheses.
C. To name a neutral complex molecule, follow the rules of naming a complex
cation. Remember: Name the (possibly complex) cation BEFORE the
(possibly complex) anion.
For historic reasons, some coordination compounds are called by their common
names. For example, Fe(CN)63and Fe(CN)64are named ferricyanide and
ferrocyanide respectively, and Fe(CO)5 is called iron carbonyl.
VALENCE BOND THEORY (VBT) & HYBRIDIZATION
The valence bond theory was proposed by Heitler and London to explain the
formation of covalent bond quantitatively using quantum mechanics. Later on,
Linus Pauling improved this theory by introducing the concept of
hybridization.
The main postulates of this theory are as follows:
* A covalent bond is formed by the overlapping of two half filled valence
atomic orbitals of two different atoms.
* The electrons in the overlapping orbitals get paired and confined between the
nuclei of two atoms.
* The electron density between two bonded atoms increases due to overlapping.
This confers stability to the molecule.
* Greater the extent of overlapping, stronger is the bond formed.
* The direction of the covalent bond is along the region of overlapping of the
atomic orbitals i.e., covalent bond is directional.
* There are two types of covalent bonds based on the pattern of overlapping as
follows:
(i) σ-bond: The covalent bond formed due to overlapping of atomic orbital
along the inter nucleus axis is called σ-bond. It is a stronger bond and
cylindrically symmetrical.
Depending on the types of orbitals overlapping, the σ-bond is divided into
following types:
σs-s bond:
σp-p bond:
σs-p bond:
(ii) π-bond: The covalent bond formed by sidewise overlapping of atomic
orbitals is called π- bond. In this bond, the electron density is present above and
below the inter nuclear axis. It is relatively a weaker bond since the electrons are
not strongly attracted by the nuclei of bonding atoms.
Note: The 's' orbitals can only form σ-bonds, whereas the p, d & f orbitals can
form both σ and π-bonds.
Examples
1) H2 molecule:
* The electronic configuration of hydrogen atom in the ground state is 1s 1.
* In the formation of hydrogen molecule, two half filled 1s orbitals of hydrogen
atoms overlap along the inter-nuclear axis and thus by forming a σs-s bond.
2) Cl2 molecule:
* The electronic configuration of Cl atom in the ground state is [Ne]3s2 3px2
3py2 3pz1.
* The two half filled 3pz atomic orbitals of two chlorine atoms overlap along the
inter-nuclear axis and thus by forming a σp-p bond.
3) HCl molecule:
* In the ground state, the electronic configuration of hydrogen atom is 1s1.
* And the ground state electronic configuration of Cl atom is [Ne]3s 2 3px2 3py2
3pz1.
* The half filled 1s orbital of hydrogen overlap with the half filled 3p z atomic
orbital of chlorine atom along the inter-nuclear axis to form a σs-p bond.
4) O2 molecule:
* The electronic configuration of O in the ground state is [He] 2s2 2px2 2py1 2pz1.
* The half filled 2py orbitals of two oxygen atoms overlap along the internuclear axis and form σp-p bond.
* The remaining half filled 2pz orbitals overlap laterally to form a πp-p bond.
* Thus a double bond (one σp-p and one πp-p) is formed between two oxygen
atoms.
5) N2 molecule:
* The ground state electronic configuration of N is [He] 2s2 2px1 2py1 2pz1.
* A σp-p bond is formed between two nitrogen atoms due to overlapping of half
filled 2px atomic orbitals along the inter-nuclear axis.
* The remaining half filled 2py and 2pz orbitals form two πp-p bonds due to
lateral overlapping. Thus a triple bond (one and two) is formed between two
nitrogen atoms.