Question

Write equations to represent the following observations. (a) A quantity of \(\mathrm{CuCl}_{2}(\mathrm{s})\) is dissolved in concentrated HCl(aq) and produces a yellow solution. The solution is diluted to twice its volume with water and assumes a green color. On dilution to ten times its original volume, the solution becomes pale blue. (b) When chromium metal is dissolved in \(\mathrm{HCl}(\mathrm{aq}), \mathrm{a}\) blue solution is produced that quickly turns green. Later the green solution becomes blue-green and then violet.

Short answer

The chemical equations representing the reactions are: \n(a) \n1. \( \mathrm{CuCl}_{2}(\mathrm{s}) + 2\mathrm{HCl}(\mathrm{aq}) \rightleftharpoons [CuCl4]^{2-}(\mathrm{aq}) + 2H^+(\mathrm{aq}) \)\n2. \([CuCl4]^{2-}(aq) \rightleftharpoons Cu^{2+}(aq) + 4Cl^-(aq)\) and \(Cu^{2+}(aq) + 4H2O(l) \rightarrow [Cu(H2O)4]^{2+}(aq)\) \n(b) \n1. \(\mathrm{Cr(s)}+ 2\mathrm{HCl(aq)} \rightarrow \mathrm{CrCl2(aq)} + \mathrm{H2(g)}\)\n2. \(\mathrm{Cr^{2+}(s)}+ \mathrm{6H2O(l)} \rightarrow [\mathrm{Cr(H2O)6}]^{2+}(aq)\)\n3. \([\mathrm{Cr(H2O)6}]^{2+}(aq) + 2 \mathrm{H2O(l)} \leftrightarrow [\mathrm{Cr(H2O)4(OH)2}]^{+}(aq) + 4 \mathrm{H2O(l)}\) and \n4. \([\mathrm{Cr(H2O)4(OH)2}]^{+}(aq) \rightarrow [\mathrm{Cr(H2O)3(OH)3]\)

Step-by-step solution

Formation of CuCl2 Complex Ion

Copper (II) chloride reacts with hydrochloric acid to form a CuCl4 complex ion which gives the solution a yellow color. This equilibrium process can be represented by the equation: \n\( \mathrm{CuCl}_{2}(\mathrm{s}) + 2\mathrm{HCl}(\mathrm{aq}) \rightleftharpoons [CuCl4]^{2-}(\mathrm{aq}) + 2H^+(\mathrm{aq}) \)

Dilution and color change of CuCl2 Complex Ion

When the yellow solution of [CuCl4]2- complex is diluted, the equilibrium is disturbed and shifts back towards reactants according to Le Chatelier’s principle, producing Cu(H2O)4 complex ion. This change of complex ion through dilution leads to the change of the solution color, which can be represented by the following equation: \n\([CuCl4]^{2-}(aq) \rightleftharpoons Cu^{2+}(aq) + 4Cl^-(aq)\) followed by \(Cu^{2+}(aq) + 4H2O(l) \rightarrow [Cu(H2O)4]^{2+}(aq)\)

Reaction of Chromium in Hydrochloric Acid

When chromium metal reacts with hydrochloric acid, the chromium is oxidized to Cr2+ with the release of hydrogen gas. This can be represented by the equation: \n\(\mathrm{Cr(s)}+ 2\mathrm{HCl(aq)} \rightarrow \mathrm{CrCl2(aq)} + \mathrm{H2(g)}\)

Formation and change of Chromium Complex Ions

The Cr2+ ion in solution readily reacts with water to form a series of coordination complexes that are responsible for the color changes observed. The main reactions are: \n\(\mathrm{Cr^{2+}(s)}+ \mathrm{6H2O(l)} \rightarrow [\mathrm{Cr(H2O)6}]^{2+}(aq)\) (which is blue), and \n\([\mathrm{Cr(H2O)6}]^{2+}(aq) + 2 \mathrm{H2O(l)} \leftrightarrow [\mathrm{Cr(H2O)4(OH)2}]^{+}(aq) + 4 \mathrm{H2O(l)}\) (which is green) and \n\([\mathrm{Cr(H2O)4(OH)2}]^{+}(aq) \rightarrow [\mathrm{Cr(H2O)3(OH)3]\) (aqua color) which eventually becomes violet as [\mathrm{Cr(H2O)3(OH)3}] decomposes.

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle explains how a change in conditions affects chemical equilibrium. When a system at equilibrium is disturbed by changes in concentration, temperature, or pressure, it will shift to counteract the change and restore equilibrium. This principle helps in predicting the direction of a reaction shift when such external changes occur.

In the context of the copper chloride solution, the color changes from yellow to green to pale blue indicate shifts in equilibrium. Initially, the copper (II) chloride forms a yellow [CuCl4]^{2-} complex ion. When water is added to the solution, it gets diluted.
According to Le Chatelier's Principle, the equilibrium shifts to favor the dissociation of the yellow complex back into its components. This results in a shift towards the formation of a green color from a different complex formed, ultimately stabilizing as a pale blue when the concentration of complex ions is significantly reduced.

Complex Ions

Complex ions are formed when metal ions bind with surrounding molecules or ions, known as ligands, through coordinate covalent bonds. These complexes often exhibit unique colors depending on the ligands attached to the central metal ion.

For instance, copper (II) can form different colored complexes. In the initial step, copper chloride reacts with hydrochloric acid to form a [CuCl4]^{2-} complex, which is yellow. As this solution gets diluted, equilibrium shifts. Eventually, a different complex, [Cu(H2O)4]^{2+}, forms as water becomes the primary ligand.
This complex is largely responsible for the bluish tint seen in more diluted solutions.

• Formation: Metal ions and ligands create bonds, forming stable compounds with potential for different electronic structures.
• Color variation: Depends on metal and ligand identities, affecting how light absorption (and thus color) changes.

In the case of chromium, similar changes occur where the metal interacts initially as Cr^{2+} and progresses through various hydrated complexes, accordingly changing solution colors.

Color Change in Solutions

Color changes in solutions are interesting visual cues that indicate chemical changes or interactions. These changes are often due to the formation or change of complex ions within the solution, significantly altering how they absorb and reflect light.

Different complex ions will absorb distinct wavelengths of light, consequently impacting the perceived color of the solution:

• The initial yellow color in the diluted copper chloride solution signifies [CuCl4]^{2-} complex presence.
• A green hue appears as equilibrium shifts back, leading to the formation of different copper ions.
• Further dilution emphasizes the prominent blue color from the [Cu(H2O)4]^{2+} complex as chloride ligands are replaced by water.

In chromium solutions, varying complex formations lead to an exciting sequence of colors transforming from blue to green, and eventually shifting to violet.
The entire sequence underscores chemical equilibrium's sensitivity to changing solute concentration, providing a vivid platform for learning and observing chemistry in action.