Physical Chemistry II CHEM 3340 Molecular Spectroscopy: Electronic Transitions PDF

Physical Chemistry II CHEM 3340 Molecular Spectroscopy: Electronic Transitions Ultraviolet and visible spectra The energies needed to change the electron distribution is much larger than those of vibrations and rotations: order of 100 kJ/mol This corresponds to the visible (Blue, 470 nm responds wit...

Physical Chemistry II CHEM 3340 Molecular Spectroscopy: Electronic Transitions Ultraviolet and visible spectra The energies needed to change the electron distribution is much larger than those of vibrations and rotations: order of 100 kJ/mol This corresponds to the visible (Blue, 470 nm responds with vibration Frank-Condon principle: Because nuclei are so much more massive than electrons, and electronic transition takes place faster than the nuclei can respond Transition from ground state to turning point: vertical transition The molecule can be in other vibrational states: the absorption occurs at several different frequencies Electronic Spectrum of SO2(g) Absorption spectra of polyatomic molecules MO theory: there are several MOs: bonding: σ, π anti-bonding: σ*, π* non-bonding: n Several transitions are possible however, most common: π* ← π σ* ← n Strongest transition occurs from HOMO to LUMO Some groups have well-de ned absorptions: chromophores fi Chromophores Groups with characteristic optical absorption: chromophores The exact position of transition depends on the rest of the molecule Two important chromophores Carbon-carbon double bond π* ← π transitions In ethylene: unconjugated double bond: λ=180 nm, ultraviolet When double bond is part of conjugated chain: Particle in a box -> wavelengths are longer -> it can shift into visible Indicative of extent of conjugation Substances with Color colorless naphtalene Conjugations extending many bonds orange colorless antrhacene tetracene Carbonyl group In addition to π* ← π π* ← n transitions (290 nm) n orbital is practically a p orbital on O (lone pair) π*: antibonding CO orbital close 290 nm, not very strong Spectrum of complex compounds of transition metal elements Observation: transition metals in solution often do not have colors, however, vivid colors are obtained in certain coordination compounds Example: Cu(NH3)42+, (dark blue) Unexpected because the d orbitals are degenerate: their energies are the same Ligand Field Theory When a ligand approaches an ion, the energies of the d orbitals can be different due to the presence of the electrons on the ligand In octahedral complexes the energies of (dz2, dx2y2) orbitals increase, eg orbitals the energies of (dxy, dyz, dzx) decrease: t2g orbitals Energy difference: ligand eld splitting parameter, ∆0 fi Spectrum of 3+ Ti(H2O)6 500 nm Transition between t2g and eg can produce visible spectrum Different ligands have different ∆O values -> different colors Analytical importance of UV/Vis spectroscopy Extremely important quantitative method Use Lambert-Beer law to calculate concentration of molecules A=εcl Use wavelength where only the analyte has large absorption coef cient, ε fi The fates of electronically excited states What happens after a molecule absorbs a photon and gets in an excited state? ◊ Non-radiative decay (internal conversion): the molecule collides with other molecules; the energy transferred into the vibration/translation/ rotation of the surrounding molecules (thermal motion, heat) ◊ Radiative decay: the molecule discards its excitation energy as a photon ◊ The excited molecule can take part in chemical reaction: Photochemistry on exam Radiative Decay Two types of radiative decays: ◊ Fluorescence: spontaneous emission of radiation occurs within a few nanoseconds after the exciting radiation is extinguished -> fast conversion ◊ Phosphorescence: The spontaneous emission may persist for a long period of time (seconds, or even hours) -> similar to an energy reservoir that slowly leaks Fluorescence The molecule absorbs a photon and gets into an excited state (vertical transition) The excited molecule is in a vibrational state v>0; the molecule gives up energy to surrounding molecules as it steps down to v=0 The surrounding molecules might not be able to accept the large energy difference needed to return back to the ground electronic state The molecule undergoes spontaneous emission (vertical transition): radiation This radiation ◊ corresponds to lower energy (lower frequency, longer wavelength) than that of the incident radiation ◊ Has vibrational structure similar to the ground electronic state Fluorescence Spectrum Fluorescence spectrum: similar in structure to the absorption spectrum, but ◊ Larger wavelength ◊ Vibrational characteristics are different Fluorescence quenching: the intensity of uorescence depends on the capability of solvent molecules to accept the transition energy: solvents with many energy levels (water) can effectively ‘quench’ uorescence (decrease intensity) fl fl Fluorescent dyes: absorb UV light and emit visible light (vivid orange/green colors) Fluorescence spectroscopy: images of biological cells by selectively attaching uorescent molecules to proteins, nucleic acids, membranes -> measuring distribution of uorescence fl fl Fluorescence Microscopy Lung carcinoma cell Phosphorescence: Mechanism The electronic ground state and the excited state is singlet (↑↓) The excited electronic state can exhibit a triplet state as well (↑↑); this triplet state has similar geometry at a point where the singlet/triplet energies cross Intersystem crossing: non-radiative transition from a singlet to a triplet state The molecule is stuck in a triplet state after internal conversion because the triplet to singlet transition to the ground electronic state is ‘forbidden’: weak emission Jablonsky Diagram Simpli ed portray of the energy levels Naphtalene fi Phosphorescence Example: ZnS Usually occurs with heavy elements (S) Phosphorous does not exhibit phosphorescence