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) \(\mathrm{CuCl}\)

Short answer

Compounds (b) \( \mathrm{CuF}_2\\) and (d) \( \mathrm{CuCl}\\) are expected to be colored.

Step-by-step solution

Understand Color and Electron Configuration

The color of a compound is often due to the presence of partially filled d-orbitals, which allow for electronic transitions that can absorb visible light. These transitions are typically found in transition metal compounds with partially filled d-orbitals.

Examine Electrons in Silver Compounds

Silver in \(\mathrm{Ag}_2\mathrm{SO}_4\) has a completely filled d-orbital (d\^10 configuration), meaning there are no possible d-d transitions that can absorb visible light. Therefore, this compound is expected to be colorless.

Analyze the Copper Fluoride Compound

Copper in \(\mathrm{CuF}_2\) is in the +2 oxidation state, which corresponds to a d\^9 configuration, allowing for possible d-d transitions. These transitions can absorb visible light, imparting color to the compound.

Evaluate the Magnesium Compound

Magnesium in \(\mathrm{MgF}_2\) has no d-electrons (it’s not a transition metal), thus there are no d-d transitions to color the compound. Therefore, it is expected to be colorless.

Look at the Copper Chloride Compound

Similar to \(\mathrm{CuF}_2\), copper in \(\mathrm{CuCl}\) can also have d-d transitions due to its d\^9 configuration. These possible electronic transitions result in the compound being colored.

Conclusion: Identify the Coloured Compound

Based on our analysis, both \(\mathrm{CuF}_2\) and \(\mathrm{CuCl}\) can exhibit color due to d-d transitions from the copper ion's d\^9 configuration.

Key concepts

Transition Metals

Transition metals are unique elements found in the center of the periodic table. These metals, encompassing elements like copper, iron, and nickel, are known for their ability to form colorful compounds. What makes them so special is their d-orbitals. Transition metals often have partially filled d orbitals. As a result, they can engage in d-d transitions, which are electronic transitions within these d-orbitals.
This phenomenon is crucial for the absorption of visible light, leading to the beautiful colors often associated with transition metal compounds. Not all metals can display such properties, which is why non-transition metals, such as magnesium, do not exhibit colorful compounds under the same chemical context.

Electronic Configuration

The electronic configuration of an element describes how its electrons are arranged in atomic orbitals. It is essential for understanding how an element engages with light and forms bonds. Transition metals have in particular d-orbitals that are filled with electrons.
For example:

• In copper(II) fluoride (\( \mathrm{CuF}_2 \) ), copper is in a +2 oxidation state, which gives it a d\(^9\) configuration.
• This partially filled d-orbital allows the electrons to move between different energy levels within these d-orbitals.

This movement is called a "transition" and causes absorption of specific wavelengths of light. It allows \( \mathrm{CuF}_2 \) to exhibit color.

Oxidation States

Oxidation states refer to the charge of an atom within a compound due to electron loss or gain. This concept is vital for understanding electron distribution in compounds.

• For instance, copper, when forming copper(II) fluoride (\( \mathrm{CuF}_2 \) ), exhibits a +2 oxidation state. \( \mathrm{Cu}^{2+} \) suggests that copper has lost two electrons.
• This affects its electronic configuration, leaving copper with 9 electrons in the d-orbital or a d\(^9\) state.

Determining the oxidation state of metals in transition metal compounds is crucial. It helps predict possible electronic configurations and potential for d-d transitions, influencing if the compound will be colored.

Visible Light Absorption

The absorption of visible light is what lends color to compounds. When light strikes a compound, the energy of the light can cause electrons to jump from a lower energy level to a higher one. Specifically in transition metals, d-d transitions allow electrons to move between d-orbitals.

• This transition absorbs certain wavelengths of visible light, while other wavelengths are reflected or transmitted. This is perceived by our eyes as color.
• Copper(II) fluoride and copper chloride are good examples, where copper ions with a d\(^9\) configuration absorb certain wavelengths that generate their colored appearance.

Notably, compounds without d-d transitions, like \( \mathrm{Ag}_2\mathrm{SO}_4 \), do not absorb light in the visible spectrum, making them appear colorless.