# 0.9 Transition metals

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## Objectives

• To synthesize a transition metal complex of cobalt three, Co(III), and ethylenediamine.
• To characterize the resulting metal complex spectroscopically.
• To understand concept of limiting reactant.

Your will be determined according to the following:

• prelab (10%)
• lab report form (80%)
• TA points (10%)

## Introduction

The transition metals are the largest“group”(classification) of elements from the periodic table. These can be found in nature as ores or in its elemental form, such as gold. All transition metals have more than one oxidation state. Most transition metals (TMs) can complex with other species (called ligands in“TM Complex”jargon) by giving their electrons to them, forming a complex. These ligands, which are the nearest neighbor atoms to the metal center, constitute the inner (or first) coordination sphere. Complexes may be either neutral or charged and have distinctive properties that may be quite unlike those associated with their constituent molecules and ions, each of which is capable of independent existence. An example of a charged complex is ferricyanide, $\left[\text{Fe}\left(\text{CN}{\right)}_{6}{\right]}^{-3}$ . The ${\text{Fe}}^{+3}$ and ${\text{CN}}^{-}$ ions found in the ferricyanide complex ion exist as independent species and in other compounds. The transition metals are well known for forming a large number of complex ions. In this experiment we will synthesize a transition metal complex containing cobalt, Co(III), and ethylenediamine.

## Stereochemistry

The most common coordination numbers (the number of individual ligands bound) are two, four, and six, with geometries illustrated in Fig 1:

Fig 1. Common geometries for complex ions. (A) linear, (B) square planar, (C) tetrahedral, and (D) octahedral

Complexes of Cu(I), Ag(I), Au(I) and some of Hg(II) form linear structures (A) such as $\text{Cu}\left(\text{CN}{\right)}_{2}^{-}$ , $\text{Ag}\left({\text{NH}}_{3}{\right)}_{2}^{+}$ , etc. Four-fold coordination (C) is not too common with transition metals, and the square planar geometry (B) occurs in complexes of Pd(II), Pt(II), Ni(II), Cu(II), and Au(III). Six-fold coordination (D) is the most common and in fact the one we will study in this laboratory exercise.

A ligand that is capable of occupying only one position in the inner coordination sphere by forming only one bond to the central atom is called a monodentate (“one tooth”) ligand. Examples are ${F}^{-}$ , ${\text{Cl}}^{-}$ , ${\text{OH}}^{-}$ , ${H}_{2}O$ , ${\text{NH}}_{3}$ and ${\text{CN}}^{-}$ . If the ligand has two groups that are capable of bonding to the central atom, it is called a bidentate ("two teeth") ligand, and so forth. An example of a bidentate ligand is ethylenediamine $\left({\text{CH}}_{2}{\text{NH}}_{2}{\text{CH}}_{2}{\text{NH}}_{2}\right)$ , which is commonly abbreviated "en". Both nitrogen atoms in "en" can bond to the central atom in a complex at the same time.

Complex ion salts with the same chemical formulas often behave differently because the same number of atoms can be arranged into different forms called isomers. Hydrate isomerism is illustrated by the following example: There are three distinct compounds with the formula $\text{Cr}\left({H}_{2}O{\right)}_{6}{\text{Cl}}_{3}$ . One of these, violet in color, reacts immediately with ${\text{AgNO}}_{3}$ to precipitate all of the chlorines as AgCl. The second is light green but only⅔of the chlorine is precipitated as AgCl. The third compound is dark green and only⅓of the chlorine is precipitated as AgCl. The last compound has only one reactive Cl, so apparently two chlorines in this compound are bonded tightly to the Cr and are not available for reaction. We might thus write this compound as $\left[{\text{CrCl}}_{2}\left({H}_{2}O{\right)}_{4}\right]\cdot \left({H}_{2}O{\right)}_{2}$ , where the species within the brackets are regarded as ligands bonded fairly strongly to the central chromium, and this species would behave as a single ion in solution. i.e., in aqueous solution,

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