Week 5: Coordination Chemistry 1 Overview Getting Started Techniques Procedure FAQ Full Lab Manual Introduction & Goals Chemistry & Background Key Questions Prelab Problems Safety Procedure In Your Write-up Experiments Index ChemLab Home		Chemistry & Background In addition to naturally occurring insoluble minerals, the transition metals form water-soluble salts. These salts give rise to an aqueous chemistry of transition metal cations which has tremendous scope. When crystallized from water, the cations of these salts characteristically have a number of water molecules closely associated with them. Some examples of salts of known molecular structures and the hydrated cations which they contain are listed below: Salts Hydrated form of cation Co(NO3)2•6H2O [Co(H2O)6]2+ Fe(NH4)2 (SO4)2•6H2O [Fe(H2O)6]2+ Fe(NO3)3•9H2O [Fe(H2O)6]3+ Cr(NO3)3•9H2O [Cr(H2O)6]3+ In these particular cases, six water molecules are closely bound to the metal even though additional water molecules may be present in the salt. The spatial arrangement of the six water molecules around the metal is presented below: The drawing emphasizes the coplanar arrangement of four oxygens around the metal. The dashed lines indicate bonds behind the plane of the paper, while the bold lines denote bonds coming out of the plane of the paper. Two more oxygens are above and below the plane defined by the 4 coplanar oxygen atoms. All of the metal-oxygen bonds are equivalent, however, and any pair of opposite water molecules could be chosen as the two to draw above and below a plane which would include the other four oxygens and the metal. All metal-oxygen bonds in a particular complex are the same length (~2 Å or 2 x 10-8 cm) and the oxygen-metal-oxygen bond angles are 90° for pairs of adjacent oxygens and 180° for opposite pairs. The geometry about the metal is described as octahedral since the oxygens can be regarded as positioned at the six equivalent vertices of a regular octahedron which is centered about the metal. The water molecules present in these cations are examples of ligands. "Ligand" is a general term for a neutral molecule or anion which is bonded to a metal. Metal-ligand bonding is often referred to as coordination of the metal by the ligand. The array of ligands is said to constitute the metal's coordination sphere. Finally, the entire ensemble of metal and ligands is called a transition metal complex or coordination complex. Although the above examples of cationic aqua (water) complexes are drawn from solid salts, there is overwhelming evidence that transition metal cations have similar structures in solution. Therefore, the common practice of writing the formulas of ions such as Cr3+ and Fe2+ provides an inadequate description. These two ions, for instance, exist in aqueous solution as the octahedral complexes [Cr(H2O)6]3+ and [Fe(H2O)6]3+. The number of water molecules coordinated to a particular cation in solution is not always known, but it is certain that all cations of charges +2 or higher are present as aqua complexes. Octahedral coordination by six water molecules appears the most common structure for +2 and +3 cations. Note that a positively charged species like [Cr(H2O)6]2+ requires negatively charged counter ions to balance its charge. Some typical anions include nitrate (NO3-), chloride (Cl-), and sulfate (SO42-) and a typical complex is written [Cr(H2O)6][NO3]2. A wide variety of neutral molecules other than water, as well as a large number of anions, can function as ligands. Common examples of ligands and complexes which they form are given below. The list includes both octahedral cases and also complexes with coordination numbers lower than six. Ligands Complexes NH3 [Cr(NH3)6]3+ [Cu(NH3)4]2+ Cl- [PtCl6]2- [CoCl4]2- CN- [Fe(CN)6]3-, [Fe(CN)6]4 - [Ag(CN)2]- CO Cr(CO)6       (These two are Ni(CO)4       neutral molecules!) The chemistry of transition metal complexes, covered in detail in Chem 64, is of enormous importance in biological and geological systems as well as in industrial applications. An example of singular importance is hemoglobin, an iron complex which is the prime biological oxygen carrier. Although some details are still in question, the gross features of the immediate coordination sphere of iron are known. Five coordination positions are occupied by nitrogens which are a part of the hemoglobin protein structure. The sixth position is occupied by oxygen in oxyhemoglobin and is thought to be vacant in deoxyhemoglobin. The extreme toxicity of carbon monoxide (CO) is associated with its ability to coordinate to iron more strongly than does oxygen, thereby inhibiting hemoglobin's oxygen carrying function. The Cobalt Ammines The most extensively studied class of octahedral transition metal compounds are cobalt(III) complexes in which ammonia (or other neutral molecules, closely related to ammonia, called amines) occupy some or all of the six coordination positions. The (III) in the name is a way of indicating the +3 oxidation state of the Co3+ ion. These complexes played a decisive role in early formulations of the structure of transition metal compounds and they continue to be important model systems for contemporary research into the properties of complex ions. The first and simplest cobalt ammine complex ion, [Co(NH3)6]3+, was prepared in 1798. Alfred Werner, a German chemist, studied the cobalt ammines extensively in the late 19th and early 20th centuries. He correctly interpreted his observations as requiring an octahedral geometry of the ligands about the metal. Modern transition metal chemistry has evolved from his work for which he was awarded the Nobel Prize in chemistry in 1913. The intensity of ongoing research interest in cobalt ammine complexes is measured by the fact that a recent Chemical Abstracts cumulative index to the chemical literature has about 5000 entries referring to articles on the subject over a five year period. This laboratory experiment involves the preparation of aquapentaammine-cobalt(III) as a nitrate salt, [Co(NH3)5(H2O)] [NO3]3: Note the spelling of the complex name. There are molecules which, as a class, are called amines, but the ammonia as a ligand is called ammine in the chemical's name. Water as a ligand is called aqua (formerly aquo) in the name. Once a successful synthesis has been carried out, a number of reactions of the complex will be explored that will establish the purity of the product and characterize some of its chemical behavior in weeks two, three, and four. Complexes of amines with cobalt(III) are nearly always prepared from a cobalt(II) salt, the amine, and a reagent which will convert cobalt(II) to cobalt(III). The procedure used here is typical, with hydrogen peroxide serving as the reagent (called an "oxidizing agent" for its ability to remove an electron) and ammonia as the amine. Here is the stoichiometric net reaction for this synthesis: 2 HNO3 + 2 [Co(H2O)6] [NO3]2(s) + H2O2 + 10 NH3 2 [Co(NH3)5(H2O)] [NO3]3(s) + 12 H2O The oxidation-reduction half-reactions consist of the oxidation of cobalt (II) to cobalt (III): 2 Co+2 2 Co+3 + 2 e- And the reduction of the hydrogen peroxide: 2 H+ + H2O2 + 2 e- 2 H2O The purpose of each reagent in the mixture is described below. About the Reagents Cobalt Nitrate ([Co(H2O)6] [NO3]2 or Co(NO3)2•6H2O) is a typical hydrated cobalt(II) salt which consists of octahedral Co(H2O)62+ cations and nitrate anions in the solid. It is extremely soluble in water and the solubility increases rapidly with increasing temperature. Ammonia is an aqueous solution of NH3 gas (sometimes laboratory bottles bear the old-fashioned label "ammonium hydroxide" or NH4OH). The solution is basic due to the following equilibrium: NH3(aq) + H2O NH4+ + OH- The 6 M solutions used here should be treated with respect. Aqueous ammonia is quite volatile, losing ammonia gas. In low concentrations, exposure to the gas causes irritation to the eyes and throat. At higher concentrations it is extremely toxic. Although the 6 M reagent does not present a serious hazard, ammonia solutions should be covered or kept in the hood throughout the procedure. Dilute ammonia solutions are familiar as household cleaning agents. They derive their effectiveness from the degradative action of basic solutions on natural fats. Ammonia provides the ammine ligands for your complex Ammonium nitrate (NH4NO3) is a common salt which requires no special precautions in standard laboratory applications. It is a major product of the chemical industry, largely because of its use as a source of nitrogen in fertilizers. When exposed to extreme heat and/or pressure it is explosive. Several major industrial disasters have occurred when fires have initiated the explosion of large stocks of ammonium nitrate. The truck bomb of the Oklahoma City bombing was filled ammonium nitrate. Ammonium nitrate is added to the reaction mixture to provide nitrate ions and to increase the concentration of NH3 (aq) in solution. According to LeChatelier's Principle, addition of NH4+ ions will push the equilibrium between NH3 and NH4+ to the reactant, or NH3, side of the above equation. Hydrogen peroxide (H2O2) is another ubiquitous chemical. Its commonplace applications include use as a disinfectant and as a bleaching agent. Industrially it is widely used for bleaching, particularly in the processing of paper pulp. Although concentrated solutions, such as 30% by weight, cause severe burns, the 3% solutions used here do not. The concentration of 3% H2O2 is 0.9 M. (% concentrations are almost always expressed as weight percents.) In your synthesis reaction, hydrogen peroxide is the oxidizing agent that oxidizes the cobalt (II) reactant to the cobalt (III) product. Concentrated nitric acid (HNO3, 16 M) is the most hazardous reagent used in this experiment. It is a typical strong acid in its capacity to burn skin and other tissue. It has the added peculiarity of turning exposed skin yellow. Gloves are recommended, although they provide only temporary protection from this strong acid--be careful! Although pure nitric acid solutions are colorless, the laboratory reagent is sometimes significantly discolored due to the presence of NO2 and N2O4, which are decomposition products of nitric acid when exposed to light. Nitric acid provides the H+ that appears in the reduction half-reaction. Ethyl alcohol (ethanol, C2H5OH) should require no introduction. A 95% ethanol-5% water solution is the standard laboratory form. It is that mixture of water and ethanol that has a boiling point, (78.15 °C), which is lower than that of pure ethanol (78.3 °C), pure water (100 °C), or any other mixture. The residual water, therefore, cannot be removed by simple distillation. The solutions are quite volatile and easily ignited. Exposure to open flame should be avoided. Alcohol is included in the reaction mixture to allow the product to precipitate. The product complex is soluble in water, but insoluble in ethanol.
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