5.3 Molecular Spectroscopy: IR, UV-Vis & NMR

Molecular Spectroscopy: IR, UV-Vis & NMR

When electromagnetic radiation of an appropriate energy strikes a molecule, the molecule can absorb a photon and jump to a higher-energy state — but only if the photon's energy exactly matches the gap between two allowed molecular energy levels. Different regions of the EM spectrum probe different kinds of molecular motion and electronic structure, and each is a distinct diagnostic tool on the MCAT.

Infrared (IR) Spectroscopy

Infrared photons carry relatively low energy (roughly 0.5 eV or less) — too little to promote an electron to a higher orbital, but exactly the right energy to excite intramolecular vibrations and rotations: bond stretching, bending, and twisting. Each type of chemical bond has its own characteristic vibrational frequency, determined mainly by the bond's stiffness (bond order) and the masses of the atoms involved, similar in spirit to how a stiffer spring or lighter mass vibrates faster. IR spectra are conventionally plotted against wavenumber (cm⁻¹, proportional to frequency) rather than wavelength.

Some high-yield characteristic absorptions:

A useful pattern: triple bonds and O–H/N–H stretches occur at the highest wavenumbers, followed by the carbonyl region, then single-bond stretches and bending modes at lower wavenumbers.

Below roughly 1,500 cm⁻¹ lies the fingerprint region — a dense, complex forest of overlapping bending and skeletal vibration modes that is extremely difficult to assign to individual bonds, but is essentially unique to each compound (much like a human fingerprint). Two different molecules essentially never produce identical fingerprint regions, so this range is used to confirm a compound's identity by comparison to a reference spectrum, even when the peaks it contains can't be individually interpreted.

Visible Region: Absorption and Complementary Color

A compound appears colored when it absorbs certain wavelengths of visible light (400–700 nm) and transmits or reflects the rest; the color you perceive is the complementary color to the light absorbed. For example, a compound that absorbs primarily blue light (~450–495 nm) appears orange to the eye — orange and blue sit opposite each other on the color wheel. A compound that absorbs no visible light at all appears colorless/white, and one that absorbs essentially all visible wavelengths appears black.

Structural changes that alter a molecule's electronic environment shift which wavelength it absorbs, and therefore shift its observed color. This is exactly how acid-base indicators work: protonation or deprotonation of an indicator molecule changes its conjugation or electron distribution enough to shift its absorption maximum from one part of the visible spectrum to another (or into the UV, making the molecule appear colorless). Phenolphthalein, for instance, is colorless in acidic/neutral solution but turns pink in base once deprotonation extends conjugation across the molecule and shifts absorption into the visible range.

Beer–Lambert Law (Quantifying Absorption)

How strongly a solution absorbs light of a given wavelength is quantified by the Beer–Lambert law:

A = εlc

where A is absorbance (dimensionless; often measured as −log₁₀(I/I0) of transmitted vs. incident intensity), ε is the molar absorptivity (extinction coefficient, a property of the molecule at that wavelength), l is the path length of the cuvette (usually 1 cm), and c is the concentration of the absorbing species. Absorbance is linearly proportional to concentration when the law holds — the basis of spectrophotometric assays that convert an optical reading into a concentration (enzyme assays, protein concentration by A280, clinical chemistries). On the MCAT, expect passages that ask you to scale concentration from absorbance ratios, or to recognize that doubling path length or concentration doubles A (all else equal).

Ultraviolet (UV) Region

Ultraviolet photons carry more energy than visible photons (shorter wavelength, higher frequency) — enough to promote valence electrons between molecular orbitals. Two transitions dominate organic UV absorption:

• π → π* transitions: an electron in a bonding π orbital is promoted to the corresponding antibonding π* orbital. These occur in any molecule with π bonds, but are strongest and shift to longer wavelength in conjugated systems (alternating double and single bonds), because conjugation narrows the energy gap between the highest-occupied and lowest-unoccupied π orbitals.
• n → π* transitions: a non-bonding electron pair (n), such as a lone pair on the oxygen of a carbonyl group, is promoted into a π* antibonding orbital. These transitions are generally weaker (lower probability) than π → π* transitions but still diagnostic, especially for carbonyl-containing compounds.

The more extensive a molecule's conjugation, the smaller the HOMO–LUMO energy gap, and the longer the wavelength of light it absorbs — a bathochromic (red) shift. Push this far enough (many alternating double bonds) and the absorption maximum moves out of the UV range entirely and into the visible range, which is why highly conjugated biomolecules are colored even without any metal center involved: β-carotene (11 conjugated double bonds, appears orange) and the conjugated polyene chain of retinal, the light-sensing chromophore in rhodopsin that initiates vision, are classic MCAT examples. Metal-containing chromophores like the heme group in hemoglobin add additional visible absorption from transitions involving the metal's d-electrons, but the basic conjugation-lengthens-wavelength principle applies broadly across biological pigments.

¹H Nuclear Magnetic Resonance (NMR) Spectroscopy

Nuclear magnetic resonance (NMR) spectroscopy exploits the fact that certain atomic nuclei — including the proton, ¹H — behave like tiny magnets because they possess intrinsic nuclear spin. When placed in a strong external magnetic field (B0), these nuclear spins can align either with the field (lower energy) or against it (higher energy). Irradiating the sample with radiofrequency (RF) energy exactly matching this energy gap flips spins between the two states — the nucleus "resonates," absorbing energy at a specific frequency.

Crucially, the exact resonance frequency of a given proton is not fixed — it depends on the local electron density surrounding that proton, which partially shields it from the full strength of B0. Protons in electron-rich environments are more shielded and resonate at a slightly different frequency than protons in electron-poor environments (e.g., near an electronegative atom or a carbonyl group). This difference is reported as chemical shift (δ), measured in parts per million (ppm) relative to a reference compound (tetramethylsilane, TMS, defined as δ = 0). Because it's a ratio (ppm), chemical shift is independent of the specific magnetic field strength of the instrument used.

Chemically equivalent protons — related by molecular symmetry or free rotation, and experiencing the same electronic environment — resonate at the same chemical shift and produce a single combined signal. The area under each signal (integration) is directly proportional to the number of protons producing it, so integration ratios reveal how many of each type of proton a molecule contains, even without knowing the exact structure in advance.

Spin-Spin Splitting

Protons don't just respond to the local electron density — they also "feel" the small magnetic fields generated by nearby non-equivalent protons on adjacent carbons (typically up to three bonds away). This is spin-spin (J) coupling, and it splits a single expected peak into a multiplet of several closely spaced peaks.

The governing rule is the n + 1 rule: a proton (or equivalent set of protons) with n non-equivalent neighboring protons is split into n + 1 peaks.

• 0 neighboring protons → singlet (1 peak)
• 1 neighboring proton → doublet (2 peaks)
• 2 neighboring protons → triplet (3 peaks)
• 3 neighboring protons → quartet (4 peaks)

The relative heights of the peaks within a multiplet follow the ratios of Pascal's triangle (a doublet is 1:1, a triplet is 1:2:1, a quartet is 1:3:3:1). Splitting is caused only by protons on different, non-equivalent carbons — protons on the same carbon, or protons that are chemically equivalent to the one being observed, do not split its signal.

Worked Example: Applying the n+1 Rule

Predict the ¹H NMR splitting pattern of the two proton environments in bromoethane, CH₃–CH₂–Br.

Step 1 — Identify proton environments. There are two distinct sets of equivalent protons: the three methyl (CH₃) protons and the two methylene (CH₂) protons, each set equivalent within itself by free rotation and molecular symmetry, but different from each other because of their different distance from the electronegative bromine.

Step 2 — Count neighboring protons for each set.

• The CH₃ protons are adjacent to the CH₂ group, which has n = 2 protons → split into n + 1 = 3 peaks (a triplet).
• The CH₂ protons are adjacent to the CH₃ group, which has n = 3 protons → split into n + 1 = 4 peaks (a quartet).

Step 3 — Predict chemical shift ordering. The CH₂ protons sit closer to the electronegative bromine, which withdraws electron density and deshields them, so they resonate farther downfield (higher δ, typically ~3.4 ppm) than the CH₃ protons (typically ~1.7 ppm), which are shielded by comparison.

This triplet-quartet pattern (with the quartet downfield of the triplet, and integration ratio 2:3 for CH₂:CH₃) is one of the most recognizable signatures in ¹H NMR and appears whenever an ethyl group (–CH₂CH₃) is attached to an electronegative group — worth memorizing as a template you can recognize instantly on a passage.

Common MCAT Traps

• Assuming the fingerprint region can be assigned bond-by-bond. Only the higher-wavenumber characteristic-absorption region (above ~1,500 cm⁻¹) is reliably interpreted for specific functional groups; the fingerprint region is used for whole-spectrum comparison, not individual peak assignment.
• Confusing which color is absorbed vs. observed. The observed color is the complement of the absorbed color, not the same color — a solution that absorbs blue light appears orange, not blue.
• Reversing the conjugation-wavelength relationship. More conjugation narrows the HOMO-LUMO gap and shifts absorption to longer wavelength (toward visible/red), not shorter wavelength.
• Forgetting that equivalent protons don't split each other. Only non-equivalent neighboring protons contribute to the n in the n+1 rule; a proton never splits itself or its chemically equivalent partners.
• Assuming absorbance is independent of path length. Beer–Lambert (A = εlc) makes A proportional to both concentration and cuvette path length — doubling either doubles A when the law holds.