inorganic pharmaceutical · pdf filevalency (v), and a valency that ... atoms or radicals;...

11/26/2016 Dr. Mohammed H. Said 1

In linear molecules, like CO2, the central atom has only two equivalent bonding

orbitals.

Draw the energy levels and name the orbitals formed in this hybridization.

5lecture Attention is now directed toward double and triple covalent bonds.


Similarly, in hydrogen cyanide, HCN, we assume that the carbon is sp-hybridized, since it is joined to only two other atoms,


and is hence in a divalent state. One of the sp-hybrid orbitals overlaps with the hydrogen 1s orbital, while the other overlaps end-to-end with one of the three unhybridized p orbitals of the nitrogen atom. This leaves us with two nitrogen p-orbitals which form two mutually perpendicular π bonds to the two atomic p orbitals on the carbon. Hydrogen cyanide thus contains one single and one triple bond. The latter consists of a σ bond from the overlap of a carbon sp hybrid orbital with a nitrogen p orbital, plus two mutually perpendicular π bonds deriving from parallel atomic p orbitals on the carbon and nitrogen atoms.


(3) Dipole Moment. The extent of polarization or polarity in a bond or molecule is frequently expressed as the dipole moment. In polar molecules , one portion is relatively positive and the other relatively negative due to the displacement of electrons in relation to the atomic centers. In diatomic molecules , this situation occurs because of the difference in electronegativities of the two atoms . Such a molecule has a positive pole and a negative pole separated by some distance (d) and is said to have a dipole moment. The magnitude of the moment is determined by the product of the magnitude of the charge at one end , and the distance (d) between it and the center of charge at the other end.

If the charge is unity ,equal to 4.80x10-10 e.s.u., and the distance is 1 A0 The dipole moment in Deby units, D., is

(4.80 x10-10) (1)/1 x 10-10 =4.80D

(2) Polarizing Power :-The influence that an atom ,ion, or molecule has to cause polarization in another species is determined by its polarizing power. This is generally viewed as the strength of the electrostatic filed for instance, a small cation can produce extreme polarization in an anion with

a large ionic radius or in an easily polarizable molecule .


Unsaturated compounds , due to the greater mobility of the electrons yield more possibilities for canonical forms .Using carbon dioxide as an example,

Coordination Compounds and Complexation

Basic theory Werner's Coordination Theory When aqueous ammonia is added to a solution of cobalt dichloride, CoCl2, a blue precipitate forms of the corresponding hydroxide, Co(OH)2, which dissolves on the addition of an excess of ammonia to give a solution that immediately begins to absorb oxygen and turn brown. From the oxidized solution, the following compounds can be isolated:



Number of ions present in solution per CoCl3

Moles of AgCl

precipitated per mole

Complexes No.

4 4 3 2 2

3 3 2 1 1

CoCl3·6NH3 CoCl3·5NH3·H2O

CoCl3·5NH3

CoCl3·4NH3 CoCl3·4NH3

(I)

(II)

(III) (IV) (V)

These properties, and those of many other compounds of a

similar kind, were brilliantly rationalized by Alfred Werner in

1893. In this year, at the age of only 26, he proposed what is

now referred to as his “coordination theory”,

Its principal postulates are as follows:

(1) An atom exhibits two types of valency, its ordinary valency (V), and a valency that determines the number of neighbouring atoms to which it is bound (the “coordination number”).


(2) An atom’s ordinary valency is satisfied by other

atoms or radicals; its coordination number is satisfied by

atoms, radicals, or molecules.

(3) Bonds to neighbouring atoms are directed towards

fixed positions in space.

On the basis of postulates (1) and (2), Werner formulated

compounds (I) (V) as shown diagrammatically below.

Ordinary valency bonds are designated by black lines

( ) and bonds between neighbouring atoms by green

lines ( ). The cobalt atoms have their ordinary valency

of three (as in CoF3) and are given a coordination

number of six. The other atoms are given coordination

numbers to match


Now if in these formulations the black lines are taken to be ionic bonds, and the green and double lines are taken to be bonds of a non-ionic character, the properties of compounds (I) - (V) given above are fully explained. Thus (I) and (II) would be expected to give three Cl– ions in solution, (III) only two, and (IV) and (V) only one. This leads to the customary formulations:

(I) [Co(NH3)6]Cl3 containing the [Co(NH3)6]3+ ion; (II) [Co(NH3)5H2O]Cl3 containing the [Co(NH3)5H2O]3+ ion;

(III) [CoCl(NH3)5]Cl2 containing the [CoCl(NH3)5]2+ ion; (IV), (V) [CoCl2(NH3)4]Cl containing the [CoCl2(NH3)4]+ ion.

The existence of two isomers of [CoCl2(NH3)4]Cl was explained by Werner on the basis of postulate (3). Indeed, the fact that only two isomers of this formula were known led him to propose that the arrangement of atoms around the cobalt atom is an octahedral one, since other arrangements would lead to more than two isomers.


The absence of a third isomer of [CoCl2(NH3)4]Cl does not of

course prove that the octahedral structure is correct. The third

isomer may simply be much less stable or more difficult to

isolate. Werner was able to prove, however, that the octahedral

structure is the correct one by preparing a compound that would

be expected to exist in optically active forms if the structure were

octahedral, but not if it was planar or a trigonal prism, and

showing that the compound can indeed be resolved into optically

active isomers. This conclusion has been verified by X-ray

crystallography.

Most metals can form some coordination compounds, but few

form as many as trivalent cobalt. Others that do include trivalent

chromium, and bi- and quadri-valent platinum. Trivalent

chromium and quadrivalent platinum have the same coordination

number and geometry as trivalent cobalt; bivalent platinum has a

coordination number of four with the groups around it arranged

in a square.

The complex part of a coordination compound may be a

cation, an anion, or an uncharged molecule. For example: complex cation [Co(NH3)6]3+

complex anion [CoF6]3-, [Co(CN)6]3–, [Co(NO2)6] 3–

neutral complex [CoCl3(NH3)3]0, [Co(NO2)3(NH3)3]0

Neutral complexes are frequently more soluble in nonpolar organic solvents than they are in water. When they do dissolve in water, they behave as non-electrolytes.


Valence bond theory In 1927, valence bond theory was formulated and it argues that a chemical bond forms when two valence electrons, in their respective atomic orbitals, work or function to hold two nuclei together, by virtue of effects of lowering system energies. Building on this theory, the chemist Linus Pauling published in 1931 what some consider one of the most important papers in the history of chemistry: "On the Nature of the Chemical Bond". In this paper, elaborating on the works of Lewis, and the valence bond theory (VB) of Heitler and London, and his own earlier works, Pauling presented six rules for the shared electron bond, the first three of which were already generally known:



1. The electron-pair bond forms through the interaction of an

unpaired electron on each of two atoms.

2. The spins of the electrons have to be opposed.

3. Once paired, the two electrons cannot take part in additional

bonds.

His last three rules were new:

4. The electron-exchange terms for the bond involves only one

wave function from each atom.

5. The available electrons in the lowest energy level form the

strongest bonds.

6. Of two orbitals in an atom, the one that can overlap the most

with an orbital from another atom will form the strongest bond,

and this bond will tend to lie in the direction of the concentrated

orbital.

THE SHAPES OF COMPLEX METAL IONS

6-coordinated complex ions

These are complex ions in which the central metal ion is

forming six bonds. In the simple cases we are talking

about, that means that it will be attached to six ligands.

These ions have an octahedral shape. Four of the

ligands are in one plane, with the fifth one above the

plane, and the sixth one below the plane.



4-coordinated complex ions These are far less common, and they can take up one of two different shapes. Tetrahedral ions These are the ones you are most likely to need for A' level purposes in the UK. There are two very similar ions which crop up commonly at this level: [CuCl4]2- and [CoCl4]2-. The copper(II) and cobalt(II) ions have four chloride ions bonded to them rather than six, because the chloride ions are too big to fit any more around the central metal ion A square planar complex Occasionally a 4-co-ordinated complex turns out to be square planar. There's no easy way of predicting that this is going to happen. The only one you might possibly come across at this level is cisplatin which is used as an anti-cancer drug. Cisplatin is a neutral complex, Pt(NH3)2Cl2. It is neutral because the 2+ charge of the original platinum(II) ion is exactly cancelled by the two negative charges supplied by the chloride ions

RULES FOR NAMING COORDINATION COMPLEXES

The name of the positive ion is written before the name of the negative ion.

The name of the ligand is written before the name of the metal to which it is coordinated.

Ligands are listed in the following order: negative ions, neutral molecules, and positive ions. Ligands with the same charge are listed in alphabetical order.

The Greek prefixes mono-, di-, tri-, tetra-, penta-, hexa-, and so on are used to indicate the number of ligands when these ligands are relatively simple. The Greek prefixes bis-, tris-, andtetrakis- are used with more complicated ligands.

The names of negative ligands always end in o, as in fluoro (F-), chloro (Cl-), bromo (Br-), iodo (I-), oxo (O2-), hydroxo (OH-), and cyano (CN-).

A handful of neutral ligands are given common names, such as aquo (H2O), ammine (NH3), andcarbonyl (CO).

The oxidation number of the metal atom is indicated by a Roman numeral in parentheses after the name of the metal atom.

The names of complexes with a net negative charge end in -ate. Co(SCN)42-, for example, is the tetrathiocyanatocobaltate(II) ion. When the symbol for the metal is derived from its Latin name,-ate is added to the Latin name of the metal. Thus, negatively charged iron complexes are ferrates and negatively charged copper complexes are cuprates.


List of common ion names

Monatomic anions: Cl− chloride

S2− sulfide

P3− phosphide

Polyatomic ions:

NH4+ ammonium

H3O+ hydronium

NO3− nitrate

NO2− nitrite

ClO− hypochlorite

ClO2− chlorite

ClO3− chlorate

AsO43− arsenate

C2O42− oxalate

CN− cyanide

ClO4− perchlorate SO32− sulfite SO42− sulfate HSO3− hydrogen sulfite (or bisulfite) HCO3− hydrogen carbonate (or bicarbonate) CO32− carbonate PO43− phosphate HPO42− hydrogen phosphate H2PO4− dihydrogen phosphate CrO42− chromate Cr2O72− dichromate BO33− borate SCN− thiocyanate MnO4− permanganate


Coding for the ligand

(old name) coded by ligand

aquo aqua H2O

ammino ammine NH3

hydroxy hydroxo OH-

chloro Cl-

fluoro F-

cyano CN-


Coding for the number of ligands

The normal prefixes apply if there is more than one ligand. no of ligands coded by

2 di

3 tri

4 tetra

5 penta

6 hexa


For negatively charged complex ions A negatively charged complex ion is called an anionic

complex. An anion is a negatively charged ion.

In this case the name of the metal is modified to show

that it has ended up in a negative ion. This is shown by

the ending -ate.

With many metals, the basic name of the metal is

changed as well - sometimes drastically!


Common examples include: metal changed to

cobalt cobaltate

aluminium aluminate

chromium chromate

vanadium vanadate

copper cuprate

iron ferrate


inorganic pharmaceutical · pdf filevalency (v), and a valency that ... atoms or radicals;...

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