Problem 64 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following compounds is expected to be coloured? (a) \(\mathrm{Ag}_{2} \mathrm{SO}_{4}\) (b) \(\mathrm{CuF}_{2}\) (c) \(\mathrm{MgF}_{2}\) (d) \(\mathrm{CuCl}\)

Step-by-Step Solution

Answer

\( \mathrm{CuF}_{2} \) (copper(II) fluoride) is expected to be colored.

Step 1: Identify Transition Metal Ions

Look for compounds that contain transition metal ions, as they often have coloration due to their electronic structure. Here, compounds (b) \( \mathrm{CuF}_{2} \) and (d) \( \mathrm{CuCl} \) both contain copper, a transition metal.

Step 2: Understand the Role of Transition Metals

Transition metals are expected to be colored because they have partially filled d-orbitals, which can absorb certain frequencies of light, leading to visible coloration. Therefore, examine whether the copper in \( \mathrm{CuF}_{2} \) and \( \mathrm{CuCl} \) contributes to the compound's color.

Step 3: Analyze the Oxidation State of Copper

In \( \mathrm{CuF}_{2} \), copper is in the +2 oxidation state (\( ext{Cu}^{2+} \)), and this ion typically shows color due to the d-d transition. In \( \mathrm{CuCl} \), copper is in the +1 oxidation state (\( ext{Cu}^{+} \)), which generally does not show color as it has a filled d-shell.

Step 4: Conclude Based on Oxidation State

The compound \( \mathrm{CuF}_{2} \) is expected to be colored due to the presence of \( ext{Cu}^{2+} \), which has unpaired electrons leading to d-d transitions. \( \mathrm{CuCl} \) with \( ext{Cu}^{+} \) is less likely to be colored.

Key Concepts

Coloration in Compounds

In chemistry, particularly when exploring the beauty of transition metal compounds, coloration is a fascinating topic. Compounds are often colored due to the absorption of visible light. The frequencies of light absorbed depend on the electronic configuration of the metal ion. This usually happens in compounds containing transition metals, due to their unique electronic structures.

• Transition metal ions often have incompletely filled d-orbitals.
• This allows electrons to jump (transition) between different energy levels within the d-orbitals.
• During these electronic transitions, certain wavelengths of light are absorbed, and the remaining light gets transmitted or reflected, leading to the color we perceive.

For example, copper compounds can exhibit beautiful colors. In the exercise, compounds with copper like \( \mathrm{CuF}_{2} \) can be colored because of these transitions. Understanding how compounds get their colors helps in fields ranging from art restoration to developing new materials for electronic displays.

d-Orbitals

d-Orbitals are essential in understanding why transition metal compounds are often colored. These orbitals can hold a maximum of 10 electrons, and in transition metals, they are usually partially filled.
This partial filling results in:

• A wide array of possible electronic arrangements and distributions.
• Diverse ways in which electrons can be excited by light, leading to absorption of different wavelengths.

In transition metals, the d-orbitals are closely spaced in energy. This means only a small amount of energy, like that from visible light, is required to promote an electron from one d-orbital to another. These excitations and the subsequent relaxation cause the observed color in many compounds.
For instance, in \( \mathrm{CuF}_{2} \), the copper ion has a configuration that allows it to absorb certain light frequencies, resulting in its characteristic color.

Oxidation States

The oxidation state of an element in a compound can greatly influence its color. In transition metals, the oxidation state is crucial in determining whether a compound will exhibit coloration. Here's why:

• Different oxidation states alter the number of electrons present in the d-orbitals, thus changing how they interact with light.
• Higher oxidation states typically mean more d electrons are available for transition, whereas lower states might mean filled d-orbitals with little to no transition activity.

For example, in \( \mathrm{CuF}_{2} \), copper is in the +2 oxidation state, which means it has unpaired electrons that are free to participate in d-d transitions, often resulting in colored compounds. Conversely, in \( \mathrm{CuCl} \), the copper is in the +1 state. With fully filled d-orbitals, there's less tendency for transitions, making it much less likely to display coloration. Understanding oxidation states not only helps predict coloration but is also pivotal in understanding the chemical reactivity and stability of compounds.