In this multi-week experiment, you will synthesize a compound and then analyze it to determine its empirical formula. The substance you will prepare is a vividly colored coordination compound of copper. Based on the reagents used in the synthesis procedure, you can assume that the final compound contains copper(II), ammonia, sulfate, and water. In your analysis, you goal is to figure out the mole ratio of each component in the final compound. In other words, you will solve for x, y, z and a in the formula Cux(NH3)y(SO4)z • aH2O. Based on the way the formula is written you can assume that ammonia and sulfate are acting as ligands and counter-ions, respectively. The water, in contrast, is water of hydration. That is, it is incorporated into the crystal lattice of the solid compound in a noncovalent manner, usually by hydrogen bonds, and with a specific stoichiometry.
To prepare your own sample of the coordination compound, you’ll start with solid copper(II) sulfate pentahydrate, CuSO4 • 5H2O. Once dissolved in water, the copper ions take on water molecules as neutral ligands: copper(II) ions exist as the hexaquacopper(II) complex ion in aqueous solution. After adding concentrated ammonia, NH3 ligands displace the water molecules covalently bound in the original copper complex, and a dramatic color change occurs. In your lab report, you will be able to explain this color change using crystal field theory and the relative energies of copper’s d orbitals in the two complexes.
The copper(II) ammonia complex is a water-soluble ion; in order to precipitate and isolate the final product, an ionic compound, you need to decrease the solubility of the compound. A convenient way of doing this is to add a large amount of ethanol, CH3CH2OH, to the aqueous solution. Ethanol is miscible with water but is much less polar, and, as the amount of ethanol in the mixture increases, the solubility of ionic compounds decreases. After the addition of ethanol, your coordination compound will appear as a crystalline solid, and the synthesis procedure ends with filtration, rinsing, and drying of the visually stunning product. The next several parts of the lab involve analysis of your copper(II) coordination compound to determine its empirical formula.
Gravimetric Analysis of Sulfate
The amount of sulfate in your product will be determined with a simple gravimetric analysis, a method that depends on the accurate weighing of a certain sample. You will selectively react the sulfate in the copper(II) compound using an aqueous solution of lead(II) acetate. The entire sulfate in the copper(II) complex sample should precipitate as PbSO4 during this metathesis reaction. After careful filtering and drying of the precipitate, you can use the mass of PbSO4 to calculate the amount of sulfate in the sample of the copper(II) compound.
Volumetric Analysis of Ammonia
You will measure the ammonia content of the product with a simple but colorful titration. A standard HCl solution will be provided to titrate all the ammonia contained in a sample of your copper(II) compound. Since this is a strong acid-weak base titration, the equivalence point occurs at pH < 7, and methyl orange is a good indicator to use. Methyl orange changes from orange-yellow to red at the endpoint of the titration; however, due to the brightly colored copper(II) compounds produced during the titration you should observe several interesting color changes. The solution starts out deep blue becomes a blue-green suspension, and then a pea-green suspension when 70% of the HCl has been added. By the time 85% of the acid has been added, the suspension is golden yellow. The endpoint occurs when the suspension becomes a pumpkin-orange solution. A red-orange solution is passed the endpoint. Spectrophotometric...