An Examination Of The Geometries Of Coordination Compounds

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A Coordination compound is 'a compound formed from a Lewis acid and a Brønsted base', a Lewis acid being an electron pair acceptor and a Brønsted base a proton acceptor. Thus the interaction of the Lewis acid metal centre in Ni(ClO4)2 with bronsted base NH3 to form a complex which is known as a coordination complex.

Ni(CLO4)2 + NH3  [Ni(NH3)6](CLO4)2..................................................(1)

Equation (1) describes the formation of a coordination complex.

A coordination complex is structure consisting of a central metal ion or atom and

Ligands surrounding central metal ion.

A coordination compound (or coordination complex) consists of a metal cation or neutral atom to which neutral or negatively charged ligands have bonded. The number of ligand atoms to which the metal center is directly bonded is the metal cation's coordination number (c.n.), and this number is always greater than the regular valence or oxidation number (o.n.) of the metal. The coordination complex can be negatively charged, for example, [AuCl 4 ] − , [PtCl 6 ] 2− , [Co(NO 2 ) 6 ] 3− , and [Fe(CN) 6 ] 3− ; neutral, for example, [Fe(CO) 5 ], [Ni(PF 3 ) 4 ], and [Rh(NH 3 ) 3 Cl 3 ]; or positively charged, for example, [Cu(NH 3 ) 4 ] 2+ , [Mn(H 2 O) 6 ] 2+ , and [Pt(NH 3 ) 5 Cl] 3+ . TiCl 4 and UF 6 are neutral molecules (in which o.n. = c.n.), they are not coordination compounds; whereas [AlCl 4 ] − and [FeF 6 ] 3− are coordination complexes in which the coordination number exceeds the oxidation number. The chromate anion, CrO 4 = , is not a coordination complex; the o.n. of the Cr atom is 6, but only four O atoms are bonded to it.

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TERMINOLOGY RELATED TO COORDINATION COMPOUNDS

1. Ligand (complexing agents)-ligand is an atom, ion or a molecule that binds to central metal ion to form complex. The bonding between metal and ligand generally involves formal donation of one or more of the ligand's electron pairs, which is called coordinate covalent bond.

Ligands are of following types:

1. Unidentate:-which contains only one lone pair of electrons,for ex-NH3

2. Bidentate:-which contains two lone pairs of electrons,for ex- ethlene diamine

3. Polydentate:-which contains more than two lone pair of electrons,

For ex-[EDTA]-4

Certain polydentate ligands are particularly good at linking together several metal ions and are refered to as polynucleating ligands

2. Coordination entity

A coordination entity constitutes a central metal atom or ion bonded to a fixed number of ions or molecules. For example, [CoCl3(NH3)3] is a coordination entity in which the cobalt ion is surrounded by three ammonia molecules and three chloride ions. Other examples are [Ni(CO)4], [PtCl2(NH3)2], [Fe(CN)6]4-, [Co(NH3)6]3+.

3. Central atom/ion

In a coordination entity, the atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement around it, is called the central atom or ion. For example, the central atom/ion in the coordination entities: [NiCl2(H2O)4],[CoCl(NH3)5]2+ and [Fe(CN)6]3-are Ni2+, Co3+ and Fe3+, respectively. These central atoms/ions are also referred to as Lewis acids.

4. Coordination polyhedron

The spatial arrangement of the ligand atoms which are directly attached to the central atom/ion defines a coordination polyhedron about the central atom. The most common coordination polyhedra are octahedral, square planar and tetrahedral. For example, [Co(NH3)6]3+ is octahedral, [Ni(CO)4] is tetrahedral and [PtCl4]2- is square planar. Fig. 9.1 shows the shapes of different coordination polyhedra.

5. Lone pairs: A lone pair is a valance electron pair without bonding or sharing with other atoms. They are found in the outermost electron shell of an atom, so lone pairs are a subset of a molecule's valence electrons. Thus, the number of lone electrons plus the number of bonding electrons equal the total number of valence electrons from a compound.

6. Coordination sphere:- The coordination sphere refers to a central atom or ion and an array of molecules or anions, the ligands, around. Molecules that are attached non covalently to the ligands are called the second coordination sphere.

7. Coordination number:- The coordination number (CN) of a metal ion in a complex can be defined as the number of ligand donor atoms to which the metal is directly bonded. In the case of [Co (NH3)5Cl]2+ this will be 6, the sum of one chloride and five ammonia ligands each donating an electron pair. An alternative definition of CN would be the number of electron pairs arising from the ligand donor atoms to which the metal is directly bonded.

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8. Chelates:- When a bi- or polydentate ligand uses two or more donor atoms to bind to a single metal ion, it is said to form a chelate complex . Such complexes tend to be more stable than similar complexes containing unidentate ligands due to ring structure which is difficult to break.

9. Oxidation number of central atom

The oxidation number of the central atom in a complex is defined as the charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom. The oxidation number is represented by a Roman numeral in parenthesis following the name of the coordination entity. For example, oxidation number of copper in [Cu(CN)4]3- is +1 and it is written as Cu(I).

10. Homoleptic and heteroleptic complexes

Complexes in which a metal is bound to only one kind of donor groups, e.g., [Co(NH3)6]3+, are known as homoleptic. Complexes in which a metal is bound to more than one kind of donor groups,e.g., [Co(NH3)4Cl2]+, are known as heteroleptic.

11.Dipolar bond:- A dipolar bond, also known as coordinate link, coordinate covalent bond[, dative bond, or semi polar bond, is a description of covalent bonding between two atoms in which both electrons shared in the bond come from the same atom.

Dipolar bond b/w NH3(ligand) and Cu(central metal ion).

12. Primary valency (Pv): This is non- directional and ionizable. In fact it is the positive charge on the metal ion.

13. Secondary valency (Sv) : This is directional and non- ionizable. It is equal to the number of ligand atoms co-ordinated to the metal (co-ordination number)

HISTORY

Alfred Werner (1866-1919), a Swiss chemist was the first to formulate his ideas about the structures of coordination compounds. He prepared and characterized a large number of coordination compounds and studied their physical and chemical behavior by simple experiment techniques. Werner proposed the concept of a primary valence and secondary valence for a metal ion. Binary compounds such as CrCl3.CoCl2 or PdCl2 have primary valence of 3, 2 and 2 respectively. In a series of compounds of cobalt (III) chloride with ammonia, it was found that some of the chloride ions could be precipitated as AgCl on adding excess silver nitrate solution in cold but some remained in solution.

Werner proposed the term secondary valence for the number of groups bound directly to the metal ion; in each of these examples the secondary valences are six.

The last two compounds in Table have identical empirical formula, CoCl3.4NH3, but distinct properties. Such compounds are termed as isomers.

The main postulates are:

1. In coordination compounds metals show two types of linkages (valences)-primary and secondary.

2. The primary valences are normally ionisable and are satisfied by negative ions.

3. The secondary valences are non ionisable. These are satisfied by neutral molecules or negative ions. The secondary valence is equal to the coordination number and is fixed for a metal.

4. The ions/groups bound by the secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordination numbers

GEOMETRIES OF COORDINATION COMPOUNDS

Molecular geometry or molecular structure is the three-dimensional arrangement of the atoms that constitute a molecule. It determines several properties of a substance including its reactivity, polarity, phase of matter, color, magnetism, and biological activity.

Werner was the first to describe the bonding features in coordination compounds. But his theory could not answer basic questions like:

(i)Why only certain elements possess the remarkable property of forming coordination compounds?

(ii) Why the bonds in coordination compounds have directional properties?

(iii) Why coordination compounds have characteristic magnetic and optical properties.

To rationalize how the shapes of atomic orbitals are transformed into the orbitals occupied in covalently bonded species, we need the help of two bonding theories:

Valence Bond (VB) Theory, the theory we will explore, describes the placement of electrons into bonding orbitals located around the individual atoms from which they originated

Molecular Orbital (MO) Theory places all electrons from atoms involved into molecular orbitals spread out over the entire species. This theory works well for excited species, and molecules like O2.

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VALANCE BOND THEORY

Valence bond theory describes a chemical bond as the overlap of atomic orbitals. In the case of the hydrogen molecule, the 1s orbital of one hydrogen atom overlaps with the 1s orbital of the second hydrogen atom to form a molecular orbital called a sigma bond. Attraction increases as the distance between the atoms gets closer but nuclear-nuclear repulsion becomes important if the atoms approach too close.

In a HCl molecule, the overlaps occur between the 1s-orbital of hydrogen and the singly occupied 3p-orbital of the chlorine atom. In H2O molecule, each O-H bond is formed as a result of an orbital overlap between the 1s-orbital of hydrogen and the singly occupied 2p-orbital of oxygen atom. The three singly occupied 2p-orbitals in nitrogen are involved in covalent bond formation with other atoms such as with hydrogen atoms in NH3 molecule.

In general, covalent bonds are formed as a result of an orbital between two singly occupied valence shell orbitals of two atoms. The two bonding electrons (with opposite spins) then occupy the overlapped-region of the orbitals.

There are 3 methods of showing the formulas of molecules. Molecular formulas show only the types and numbers of atoms in the molecule. Structural formulas show the atoms in their correct placement in the molecule and allow for distinguishing isomers. Electron-dot formulas or Lewis dot structures are similar to structural formulas but also include all of the non-bonding outer electrons.

Lewis dot structures of some compounds:

Valence Bond Theory predicts metal complex bonding arises from overlap of filled ligand orbital and vacant metal orbitals. Resulting bond is a coordinate covalent bond. A covalent bond is formed by the overlap of two orbitals, one from each bonding atom. Each orbital originally holds one electron, and after the orbitals overlap, a bond is formed that holds two electrons. In the formation of a coordinate covalent bond in a complex a ligand orbital containing two electrons overlaps an on occupied orbital on the metal atom.

Main Postulates

1) The metal ligand bond arises by donation of pair of electrons by ligands to the central metal atom.

2) To accommodate these electrons the metal ion must possess requisite number of vacant orbitals of comparable energy. These orbitals of the metal atom undergo hybridization to give hybrid orbitals. The basic premise of hybridization is that appropriate linear combinations of non-equivalent orbitals of an atom give sets of hybrid orbitals that are equivalent and have specific spatial orientations. For coordination compounds, the hybridizations involving s, p and d orbitals are important. Different linear combinations of s, p and d orbitals, like dsp2, dsp3 and d2sp3 yield various types of coordination compounds.

3) Sometimes the unpaired (n-1)d orbitals pair up before bond formation making     (n-1)d orbitals vacant. The central metal atom makes available number of d-orbitals equal to its co-ordination number.

4) The metal ligand bonds are thus formed by donation of electron pairs by the ligands to the empty hybridized orbitals. These bonds are equal in strength and directional in nature.

5) Octahedral, trigonal-bipyramidal or square pyramidal, square planar and tetrahedral complexes are formed as a result of d2sp3 (or sp3d2), sp3d, dsp2 and sp3 hybridization respectively.

On the basis of the valence bond theory it is usually possible to predict the geometry of coordination complexes.

Complex geometry can be linked to five main orbital hybridization processes.

The most common geometries are explained below:

Tetrahedral Geometry: Gives [CoCl4]2-three unpaired electrons, which makes it paramagnetic

and attracted by magnets.

Square Planar Geometry:

Gives [Ni (CN)4]2- all paired electrons, which makes it diamagnetic and

weakly repelled by magnets.

Octahedral sp3d2 Geometry:

Gives [CoF6]3- four unpaired electrons, which makes it paramagnetic and is called a high-spin complex.

Octahedral d2sp3

Geometry: Gives [Co(CN)6]3-paired electrons, which makes it diamagnetic and is called a low-spin complex.

The difference between sp3d2 and d2sp3 hybrids lies in the principal quantum number of the d

orbital.

• In sp3d2 hybrids, the s, p, and d orbitals have the same principal quantum number-High Spin.

• In d2sp3 hybrids, the principal quantum number of the d orbitals is one less than s and p orbitals-Low Spin .A complex's magnetic properties determine which hybrid is being used.

When inner d orbital is used in hybridisation inner orbital or low spin complex is formed

When outer d orbital is used in hybridisation outer orbital or high spin complex is formed.

CONCLUSION

If we can draw a Lewis structure for a species, and count electronic regions around central atom,

We can immediately determine:

the shape of the species about the central atom

the hybridization of the species based on the central atom

LIMITATIONS OF VALANCE BOND THEORY

While the VB theory, to a larger extent, explains the formation, structures

and magnetic behavior of coordination compounds, it suffers from

the following shortcomings:

(i)It involves a number of assumptions.

(ii)It does not give quantitative interpretation of magnetic data.

(iii)It does not explain the color exhibited by coordination compounds

(iv) It does not give a quantitative interpretation of the thermodynamic

Or kinetic stabilities of coordination compounds.

(v)It does not make exact predictions regarding the tetrahedral and Square planar structures of 4-coordinate complexes.

(vi) It does not distinguish between weak and strong ligands.

The properties of weak field ligands and string field ligands are explained by crystal field theory (CFT).

MAGNETIC PROPERTIES OF COORDINATION COMPOUNDS

The magnetic properties of coordination compounds can also be explained on the basis of electrons' present in d orbital. The magnetic moment of coordination compounds can be measured by the magnetic susceptibility experiments. The results can be used to obtain information about the structures adopted by metal complexes .A critical study of the magnetic data of coordination compounds of metals of the first transition series reveals some complications. For metal ions with up to three electrons in the d orbitals, like Ti3+ (d1); V3+(d2); Cr3+ (d3); two vacant d orbitals are available for octahedral hybridisation with 4s and 4p orbitals. The magnetic behavior of these free ions and their coordination entities is similar. When more than three 3d electrons are present, the required pair of 3d orbitals for octahedral hybridisation is not directly available (as a consequence of Hund's rule). Thus, for d4 (Cr2+, Mn3+), d5 (Mn2+, Fe3+), d6 (Fe2+, Co3+)cases, a vacant pair of d orbitals results only by pairing of 3d electrons which leaves two, one and zero unpaired electrons, respectively .The magnetic data agree with maximum spin pairing in many cases, especially with coordination compounds containing d6 ions. However, with species containing d4 and d5ions there are complications. [Mn (CN)6]3-has magnetic moment of two unpaired electrons while [MnCl6]3-has a paramagnetic moment of four unpaired electrons. [Fe(CN)6]3- has magnetic moment of a single unpaired electron while [FeF6]3- has a paramagnetic moment of five unpaired electrons. [CoF6]3- is paramagnetic with four unpaired electrons while [Co(C2O4)3]3- is diamagnetic. This apparent anomaly is explained by valence bond theory in terms of formation of inner orbital and outer orbital coordination entities. [Mn(CN)6]3-, [Fe(CN)6]3-and [Co(C2O4)3]3- are inner orbital complexes involving d2sp3 hybridisation,the former two complexes are paramagnetic and the latter diamagnetic. On the other hand, [MnCl6]3-, [FeF6]3- and [CoF6]3- are outer orbital complexes involving sp3d2 hybridisation and are paramagnetic corresponding to four, five and four unpaired electrons.