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What is Operando Spectroscopy? Spectroscopic characterization of catalysts under realistic reaction conditions with simultaneous real-time online analysis of reaction products SPECTROSCOPIC METHODS IN CATALYSIS COVERAGE Application of UV-Visible spectroscopy Application of Infra red spectroscopy Application of Resonance spectroscopy ELECTRONIC SPECTROSCOPY Using light absorption for changing the charge distribution about a molecule This is a lot of energy - often can break bonds Types of electronic transitions ORGANICS: Involving p, s, n electrons Saturated compounds s s (<150 nm), n s (<250 nm) Deep UV Double bonds/unsaturated systems p p , n p transitions : UV and visible (200-700 nm) Inorganics: Additionally, transitions between d orbitals split by presence of ligand field. Usually in visible d-d transition Charge transfer transition Electron moves between ligand and metal. One must act as donor and other as acceptor Light will be resonant with electronic energy gap at equilibrium nuclear geometry Electronic Spectra At equilibrium, molecule is in ground electronic state → lowest energy electronic state and typically in v=0. • Transitions to higher lying electronic states are accompanied by changes in v, J. • Excitation is accompanied by vibrational excitation, feels restoring force in excited state. Franck-Condon principle vertical transitions Electrons respond much faster than nuclear motion, therefore an excitation proceeds without a change to the nuclear geometry. Light will be resonant with electronic energy gap at equilibrium nuclear geometry. Table . The electronic spectral data of the complexes recorded in ethanol [frequency (cm-1)/ εmax (mol-1 dm3 cm-1)]. Complex ν(104 cm-1)/ ε μeff/μB π→π* π→π* π*→d d→d* d→d* [Mn(C5H7 O2)3] 4.90 40.8/175 36.4/24000 00 30.7/9500 24.8/950 17.5/100 [Mn(C5H7 O2)2L1] 4.76 41.6/310 30.2/12000 00 30.2/17500 24.3/100 17.6/70 0 [Mn(C5H7 O2)2L2] 4.81 42.0/225 35.9/14500 00 30.3/11000 24.8/150 17.5/170 0 Selection Rules for Electronic Spectra of Transition Metal Complexes. The Selection Rules governing transitions between electronic energy levels of transition metal complexes are: ΔS = 0 The Spin Rule Δl = +/- 1 The Orbital Rule (Laporte) The first rule says that allowed transitions must involve the promotion of electrons without a change in their spin. The second rule says that if the molecule has a centre of symmetry, transitions within a given set of p or d orbitals (i.e. those which only involve a redistribution of electrons within a given subshell) are forbidden. Relaxation of the Rules can occur through: a) Spin-Orbit coupling - this gives rise to weak spin forbidden bands b) Vibronic coupling - an octahedral complex may have allowed vibrations where the molecule is asymmetric. Absorption of light at that moment is then possible. c) π-acceptor and π-donor ligands can mix with the d-orbitals so transitions are no longer purely d-d. Types of transition Charge transfer, either ligand to metal or metal to ligand. These are often extremely intense and are generally found in the UV but they may have a tail into the visible. d-d, these can occur in both the UV and visible region but since they are forbidden transitions have small intensities. Expected Values The expected values should be compared to the following rough guide. For M2+ complexes, expect Δ = 7500 - 12500 cm-1 or λ = 800 - 1350 nm. For M3+ complexes, expect Δ= 14000 - 25000 cm-1 or λ = 400 - 720 nm. For a typical spin-allowed but Laporte (orbitally) forbidden transition in an octahedral complex, expect ε < 10 m2mol-1. Extinction coefficients for tetrahedral complexes are expected to be around 50-100 times larger than for octrahedral complexes. B for first-row transition metal free ions is around 1000 cm-1. Depending on the position of the ligand in the nephelauxetic series, this can be reduced to as low as 60% in the complex. Expected intensities of electronic transitions Expected intensities of electronic transitions Transition type Example Typical value of ε m2 mol-1 Spin forbidden, Laporte forbidden [Mn(H2O)6]2+ 0.1 Spin allowed (octahedral complex), Laporte forbidden [Ti(H2O)6]3+ 1 Spin allowed (tetrahedral complex), Laporte partially allowed by d-p mixing [CoCl4]2- 50 Spin allowed, Laporte allowed e.g. charge transfer bands [TiCl6]2- or MnO4- 1000 Basics of Light, EM Spectrum, and X-rays • Light can take on many forms. Radio waves, microwaves, infrared, visible, ultraviolet, X-ray and gamma radiation are all different forms of light. • The energy of the photon tells what kind of light it is. Radio waves are composed of low energy photons. Optical photons--the only photons perceived by the human eye--are a million times more energetic than the typical radio photon. The energies of X-ray photons range from hundreds to thousands of times higher than that of optical photons. • Very low temperatures (hundreds of degrees below zero Celsius) produce low energy radio and microwave photons, whereas cool bodies like ours (about 30 degrees Celsius) produce infrared radiation. Very high temperatures (millions of degrees Celsius) produce X-rays. The absorption of UV or visible radiation corresponds to the excitation of outer electrons. There are three types of electronic transition which can be considered; Transitions involving p, s, and n electrons Transitions involving charge-transfer electrons Transitions involving d and f electrons When an atom or molecule absorbs energy, electrons are promoted from their ground state to an excited state. In a molecule, the atoms can rotate and vibrate with respect to each other. These vibrations and rotations also have discrete energy levels, which can be considered as being packed on top of each electronic level. . Absorbing species containing p, s, and n electrons Absorption of ultraviolet and visible radiation in organic molecules is restricted to certain functional groups (chromophores) that contain valence electrons of low excitation energy. The spectrum of a molecule containing these chromophores is complex. This is because the superposition of rotational and vibrational transitions on the electronic transitions gives a combination of overlapping lines. This appears as a continuous absorption band. Possible electronic transitions of p, s, and n electrons are; Sigma to sigma* Transitions An electron in a bonding sigma orbital is excited to the corresponding antibonding orbital. The energy required is large. For example, methane (which has only C-H bonds, and can only undergo sigma to sigma* transitions) shows an absorbance maximum at 125 nm. Absorption maxima due to sigma to sigma* transitions are not seen in typical UV-Vis. spectra (200 - 700 nm) n to sigma* Transitions Saturated compounds containing atoms with lone pairs (non-bonding electrons) are capable of n to sigma* transitions. These transitions usually need less energy than sigma to sigma * transitions. They can be initiated by light whose wavelength is in the range 150 - 250 nm. The number of organic functional groups with n to sigma* peaks in the UV region is small. N to pi* and pi to pi* Transitions Most absorption spectroscopy of organic compounds is based on transitions of n or pi electrons to the pi* excited state. This is because the absorption peaks for these transitions fall in an experimentally convenient region of the spectrum (200 - 700 nm). These transitions need an unsaturated group in the molecule to provide the pi electrons. Molar absorbtivities from n ® pi* transitions are relatively low, and range from 10 to100 L mol-1 cm-1 . pi ® pi* transitions normally give molar absorbtivities between 1000 and 10,000 L mol-1 cm-1 . The solvent in which the absorbing species is dissolved also has an effect on the spectrum of the species. Peaks resulting from n to pi* transitions are shifted to shorter wavelengths (blue shift) with increasing solvent polarity. This arises from increased solvation of the lone pair, which lowers the energy of the n orbital. Often (but not always), the reverse (i.e. red shift) is seen for pi to pi* transitions. This is caused by attractive polarisation forces between the solvent and the absorber, which lower the energy levels of both the excited and unexcited states. This effect is greater for the excited state, and so the energy difference between the excited and unexcited states is slightly reduced - resulting in a small red shift. This effect also influences n to pi* transitions but is overshadowed by the blue shift resulting from solvation of lone pairs. Charge - Transfer Absorption Many inorganic species show charge-transfer absorption and are called charge-transfer complexes. For a complex to demonstrate charge-transfer behaviour, one of its components must have electron donating properties and another component must be able to accept electrons. Absorption of radiation then involves the transfer of an electron from the donor to an orbital associated with the acceptor. Molar absorbtivities from charge-transfer absorption are large (greater that 10,000 L mol-1 cm-1).