The compound \(\mathrm{CH}_{3} \mathrm{CH}=(\mathrm{CHCHO}\) has a strong absorption in the ultraviolet at \(46950 \mathrm{~cm}^{-1}\) and a weak absorption at \(30000 \mathrm{~cm}^{-1}\) Justify these features in terms of the structure of the molecule.

When we talk about π→π* transitions, we are discussing a specific type of electronic transition that occurs in molecules containing pi (π) bonds — these are the bonds formed by the sideways overlap of p orbitals. During UV light absorption, electrons in the π bonding orbitals can be promoted (excited) to π* antibonding orbitals, which require a particular amount of energy.

In our molecule, CH₃CH=(CHCHO, the presence of the C=C double bond signifies a locus for such a transition. The strong absorption peak at 46950 cm⁻¹ in the UV spectrum is the fingerprint of a π→π* transition. This high-energy transition reflects the fact that π electrons are initially in a bonding state, and the jump to an antibonding (π*) state represents a significant alteration in electron density and energy within the molecule.

This type of transition is typically responsible for the vivid colors seen in organic compounds and is a crucial indicator while analyzing UV-Vis spectroscopy data for conjugated systems.

On the other side, we have the n→π* transitions. These occur when a nonbonding electron (n) absorbs energy and transitions to an antibonding pi (π*) orbital. Nonbonding electrons are found in atoms with lone pairs, such as the oxygen in the aldehyde group (CHO) of our example molecule.

The weaker absorption peak at 30000 cm⁻¹ corresponds to this lower energy transition. Why is it lower energy? Unlike π→π* transitions, where the electron departs from a bonding molecular orbital, n→π* transitions involve nonbonding electrons. These electrons are not involved in bonding and are already at a higher energy state compared to bonding electrons. Thus, the energy required to promote them to the antibonding π* state is less.

Understanding n→π* transitions is especially important for compounds that contain atoms with spare lone pairs subject to excitation. These absorbances are pivotal for elucidating the structure of organic compounds with oxygen, nitrogen, or sulfur atoms.

Diving deeper into the concept of UV light absorption, we discover that it's a fundamental process used in spectroscopy to study the electronic structure of molecules. When UV light is shone upon a sample, electrons within the molecule may absorb the energy and move to higher energy states. The specific wavelengths at which these absorptions occur can give us detailed insight concerning the molecule’s structure.

The energy of UV light absorbed is quantified by its wavelength or frequency. As seen in our problem, two distinct peaks at 46950 cm⁻¹ and 30000 cm⁻¹ are observed, each indicative of different transitions for electrons. The intensity of the peaks — strong for π→π* and weak for n→π* — also informs us about the probability and nature of these transitions.

Practically speaking, UV spectroscopy is an analytical method widely used for characterizing the presence of conjugated systems, functional groups, and overall molecular architecture, making it a cornerstone technique in both academic and industrial chemistry settings.