Inorganic Pharmaceutical Chemistry. Coordination compounds

Transcription

1 Inorganic Pharmaceutical Chemistry Lecture No. 5 Date : 22/11 /2012 Dr. Mohammed Hamed 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).

2 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(NH 3 ) 5 Cl]Cl 2 is composed of the complex ion [Co(NH 3 ) 5 Cl] 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 NO 3 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 H 2 O 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

3 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. H 2 NCH 2 CH 2 NH 2 (ethylenediamine). Important exceptions: water is called aqua, ammonia is called ammine, carbon monoxide is called carbonyl, and the N 2 and O 2 are called dinitrogen and dioxygen.

4 Table 1. Names of Some Common Ligands Anionic Ligands Names Neutral Ligands Names Br - bromo NH 3 ammine F - fluoro H 2 O aqua O 2- oxo NO Nitrosyl OH - Hydroxo CO Carbonyl CN - cyano O 2 dioxygen C 2 O 4 2- CO 3 2- oxalato N 2 dinitrogen carbonato C 5 H 5 N pyridine CH 3 COO - acetato H 2 NCH 2 CH 2 NH 2 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 (pentakis) 2 di (bis) 6 hexa (hexakis) 9 nona (ennea) 10 deca 3 tri (tris) 7 hepta 11 undeca 4 tetra (tetrakis) 8 octa 12 dodeca

5 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 Iron Copper Lead Silver Gold Tin Name in an Anionic Complex Ferrate Cuprate Plumbate Argenate Aurate 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) 6 3 and Fe(CN) 6 4 are named ferricyanide and ferrocyanide respectively, and Fe(CO) 5 is called iron carbonyl.

6 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: σ s-s bond:

7 σ p-p bond: σ s-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) H 2 molecule: * The electronic configuration of hydrogen atom in the ground state is 1s 1.

8 * 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 σ s-s bond. 2) Cl 2 molecule: * The electronic configuration of Cl atom in the ground state is [Ne]3s 2 3p x 2 3p y 2 3p z 1. * The two half filled 3p z atomic orbitals of two chlorine atoms overlap along the inter-nuclear axis and thus by forming a σ p-p bond. 3) HCl molecule: * In the ground state, the electronic configuration of hydrogen atom is 1s 1. * And the ground state electronic configuration of Cl atom is [Ne]3s 2 3p x 2 3p y 2 3p z 1. * 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 σ s-p bond.

9 4) O 2 molecule: * The electronic configuration of O in the ground state is [He] 2s 2 2p x 2 2p y 1 2p z 1. * The half filled 2p y orbitals of two oxygen atoms overlap along the internuclear axis and form σ p-p bond. * The remaining half filled 2p z orbitals overlap laterally to form a π p-p bond. * Thus a double bond (one σ p-p and one π p-p ) is formed between two oxygen atoms. 5) N 2 molecule: * The ground state electronic configuration of N is [He] 2s 2 2p x 1 2p y 1 2p z 1. * A σ p-p bond is formed between two nitrogen atoms due to overlapping of half filled 2p x atomic orbitals along the inter-nuclear axis. * The remaining half filled 2p y and 2p z orbitals form two π p-p bonds due to lateral overlapping. Thus a triple bond (one and two) is formed between two nitrogen atoms.

10