In aqueous solution, water competes effectively with bromide ions for coordination to \(\mathrm{Cu}^{2+}\) ions. The hexaaquacopper(II) ion is the predominant species in solution. However, in the presence of a large concentration of bromide ions, the solution becomes deep violet. This violet color is due to the presence of the tetrabromocuprate(II) ions, which are tetrahedral. This process is reversible, and so the solution becomes light blue again on dilution with water. (a) Write the formulas of the two complex ions of copper(II) that form. (b) Is the change in color from violet to blue on dilution expected? Explain your reasoning.

The hexaaquacopper(II) ion is a vibrant example of the beauty and complexity of coordination compounds. It consists of a central copper(II) ion surrounded by six water molecules, each acting as a ligand. These water molecules are bonded to the copper ion through coordinate covalent bonds, where the oxygen atom of each water molecule donates a pair of electrons to the copper ion. The chemical formula for this stunning blue complex is \(\mathrm{[Cu(H_2O)_6]^{2+}}\), reflecting the +2 oxidation state of copper and the neutral charge of the six coordinated water molecules. The hexaaquacopper(II) ion is predominantly found in aqueous solutions of copper(II) salts and exemplifies the central concepts of coordination chemistry, such as ligand exchange and coordination numbers.

In practice, this ion imparts the characteristic light blue color to solutions of copper(II) salts. When isolated, this ion forms blue crystals that can be studied and analyzed for their structural properties. Moreover, the ability of the hexaaquacopper(II) ion to exchange its water ligands for other potential ligands, like bromide ions, makes it an excellent subject to explore reaction mechanisms and dynamic equilibria in chemistry.

The tetrabromocuprate(II) ion, with its deep violet color, tempts the curious minds of chemists seeking to understand the interplay of metal ions and ligands. This ion features a central copper ion, as the name suggests, with a +2 charge, but instead of water molecules, it is surrounded by four bromide ions. The bromide ions, each carrying a -1 charge, act as ligands and coordinate with the copper ion to form a tetrahedral complex. The resulting formula is denoted as \(\mathrm{[CuBr_4]^{2-}}\), which communicates the overall charge of -2 due to the four bromide ions and the charge compensating copper ion at the center.

Interestingly, the formation of tetrabromocuprate(II) ions demonstrates the versatility of copper(II) ions in forming various complexes with different ligands and geometries. It also plays a pivotal role in signaling the presence of bromide ions in a solution. The color transition from light blue to violet upon adding bromide ions serves as a qualitative test, allowing chemists to infer the composition of the solution.

Coordination chemistry is an enchanting field that delves into the compositions, structures, and properties of complex ions or coordination compounds. These compounds consist of a central metal ion, such as copper(II) in our previous examples, surrounded by a cohort of molecules or ions called ligands. The ligands are not merely passive onlookers; they actively donate electron pairs to the metal ion to form coordinate covalent bonds, which define the structure and properties of the complex ion. This fascinating interplay gives rise to various coordination numbers (the total number of bonds formed between the metal ion and its ligands), geometries (shapes like octahedral or tetrahedral), and types of ligands ({

• neutral molecules like water
• charged ions like bromide

}), each bringing a unique twist to the properties of the complex.

Coordination chemistry is not just academically stimulating but also industrially significant, with applications in medicinal chemistry, materials science, catalysis, and environmental chemistry. Whether it is the formulation of cancer-treating drugs or the synthesis of high-performance materials, coordination compounds play a critical role in scientific advancements.

Le Chatelier's principle is like a dance of equilibria, beautifully choreographed to maintain balance in chemical reactions. This principle is grounded in the concept that if a dynamic equilibrium system is subjected to a change in concentration, temperature, or pressure, it will adapt to minimize that change and restore equilibrium. When we dilute a solution containing complex ions, for instance, we are effectively reducing the concentration of certain ligands. According to Le Chatelier's principle, the system will respond by shifting the equilibrium to favor the formation of products that consume the ligands in lower concentrations.

In the case of our color-changing copper(II) solution, adding water reduces the concentration of bromide ions. Consequently, the equilibrium shifts away from the formation of the violet tetrabromocuprate(II) ions towards the production of more hexaaquacopper(II) ions, restoring the light blue color. Exploring this principle allows students not only to predict the behavior of chemical systems but also to understand and control chemical processes used in various industrial operations.