04 - Atoms and Molecules - My Full Notes PDF

Module structure  Fundamentals of light  Propagation of light in waveguides  Light interaction with matter  Lasers  Photobiology basics  Biophotonics applications  Bioimaging  Tissue engineering  Topics to be covered  Atoms and Molecules  Molecular level...

Module structure  Fundamentals of light  Propagation of light in waveguides  Light interaction with matter  Lasers  Photobiology basics  Biophotonics applications  Bioimaging  Tissue engineering  Topics to be covered  Atoms and Molecules  Molecular level interactions  Absorption, spontaneous emission, stimulated emission  Fate of excited molecules  Fluorescence  Light scattering  Rayleigh, Mie, Raman, spectroscopy  At the start  ‘Indivisible’ particle – the atom  Making up different types of elements  ‘Plum pudding’ idea – Joseph J. Thomson (1897)  Discovered negative charges coming from atoms Democritus Wikipedia  Failures of Classical Mechanics  At the end of the 19th century, a number of observations related to tiny particles simply could not be explained by Classical Mechanics  These included areas such as:  Blackbody radiation  Photoelectric effect  Discrete emission bands  A different system for describing small objects was developed – Quantum Mechanics  Quantization of energy  Max Planck suggested that energy could only be absorbed or emitted in whole multiples of  EM radiation from a source can only have energies and so on  Using this, his equations fit the observed measurements of blackbody radiation perfectly  Planck’s ideas were revolutionary and were not widely accepted until Einstein used them to explain another classical mechanics failure – the photoelectric effect Planck's constant = 6.626068 × 10-34 m2 kg / s  Rutherford model  ‘Planetary model’ – Ernest Rutherford (1909)  Positively charged heavy nucleus with orbiting negative charges https://www.britannica.com/science/Rutherford-model/images- videos  Bohr model (1913)  Electrons orbit at specific distances from the nucleus  Have an associated energy  Larger orbits have higher energy and vice versa  Radiation is absorbed and https://physicalsciencepaula.weebly.com/shellslevels.html emitted by electrons moving between orbits.  Electron transitions  An excited electron can transfer to a lower (more stable) energy level, giving off its excess energy as a photon (radiative decay)  Higher the energy difference, higher the energy of the photon  Since E=hf, a high energy transition will emit a high frequency photon  As energy levels are discrete (not continuous), only certain frequency photons can be emitted.  Identify elements through spectroscopy ν (Nu) = frequency P. Atkins and J. de Paula, Physical Chemistry 8th ed., Oxford University Press, 2006  Quantum Mechanics model  De Broglie postulated electrons can behave like waves (1924)  Solved problems with Bohr’s model  Orbitals are described by Schrödinger's wave equations (1926)  Orbitals are defined as areas around the nucleus where electrons with a particular energy are most likely to be found.  Higher the electron energy, the more complex the orbital shape https://courses.lumenlearning.com/chemistryatomsfirst/chapter/ development-of-quantum-theory-2/  Wave-particle duality of matter Matter Particle like behaviour Wave-like behaviour 1 Kinetic Energy  mv 2 h 2 Wavelength   mv Momentum mv Translational energy Quantized energy described by obtained through classical quantum mechanics mechanics Photons can behave like particles and electrons can behave like waves  Parts of the Atom  Atoms have electrons, protons and neutrons  Electrons (-) and Protons (+) have the same but opposite charge. Neutrons are neutral  Protons and neutrons have roughly the same mass  Electrons are ~1836 times lighter (rest mass)  (Hydrogen does not have neutrons)  # electrons = # protons  When it is not, it is known as an ‘ion’  Changing # neutrons produces ‘isotopes’  Filling orbitals with electrons Na Mg Al Si P S Cl Ar n=3 Li Be B C N O F Ne n=2 H He n=1 P. Atkins and J. de Paula, Physical Chemistry 8th ed., Oxford University Press, 2006  Orbital shapes 1s 2s 3s 3p 2p https://winter.group.shef.ac.uk/orbitron/ 3d  Allowed transitions in the Hydrogen atom  Spectral emission bands from atomic Hydrogen have been identified by a number of scientists  They represent the permitted radiative electron Numbers are given by 1/λ transitions between quantized orbits in the H atom P. Atkins and J. de Paula, Physical Chemistry 8th ed., Oxford University Press, 2006  Molecular energies  A molecule is able to exhibit four types of energies:  Electronic  Vibrational  Rotational  Translational  Only the first three show quantized effects  Rotational movements are only pronounced in the gaseous state and so will not be considered here  Electronic and vibrational energies play key roles in spectroscopy  Molecular potential energy curve  The molecular potential energy curve (opposite) shows the molecule is the most stable (has the lowest energy) when the nuclei are separated by the equilibrium bond length Re  -De is the depth of the minimum point of the energy curve below the potential energy of infinitely separated stationary atoms P. Atkins and J. de Paula, Physical Chemistry 8th ed., Oxford University Press, 2006  Molecular vibrational states  From analysis of a harmonic oscillator, a diatomic molecule can be shown using Schrödinger’s equations to have quantized vibrational states  Lowest vibrational energy is called the bond disassociation energy Do which cannot be zero, even at 0 Kelvin  The bottom graph shows that a molecule has a finite number of quantized vibrational energy states and that they become closer as the nuclear separation distance increases P. Atkins and J. de Paula, Physical Chemistry 8th ed., Oxford University Press, 2006  Typical normal vibrational modes CO2 H2O  The vibrational displacement patterns of a molecule are called normal modes  A linear molecule (like CO2) made up of N atoms has 3N-5 vibrational modes  A non-linear molecule (like H2O) made up of N atoms has 3N-6 vibrational modes  How many normal vibrational modes does a diatomic (e.g. H2) molecule have?  In a polyatomic molecule, one v1 = symmetric stretch normal mode can be excited v2 = bending without exciting another. Normal v3 = asymmetric stretch modes are generally P. Atkins and J. de Paula, Physical Chemistry 8th ed., Oxford University Press, 2006 independent of each other  Summary  Energy is absorbed and emitted in discrete quanta; E = hf  Quantum mechanics is used to describe the electron orbital ‘shape’ in an atom  Electrons can transition between certain energy levels – absorbing or emitting a photon of equal energy in the process  A molecule’s bonds can bend and stretch as determined by its vibrational energy