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Western University
Chemistry 1027A/B
Tom Haffie

1 UNIT 2 – STRUCTURE & PROPERTIES Origins of Quantum Theory  Kirchhoff’s work on “blackbodies” was start of quantum theory but Planck was credited with starting quantum theory  “blackbody” used to describe an ideal perfectly black object i.e. does not reflect any light & emits various forms of light as a result of temperature Planck’s Quantum Hypothesis  when an object is heated it was observed that it will glow  initially it will glow red, then white, these changes do not depend on the composition of the solid  if we look at the intensity of the different colours we get a bell shaped curve 2  scientists couldn’t explain why the bell curve didn’t follow what was expected “the classical theory” until Planck came up with an equation to explain this (1900)  Planck’s theory was that the energies of the atoms in the solids were multiples of small quantities of energy i.e. energy was not continuous  Einstein continued this work by pointing out that Planck’s hypothesis led to the conclusion that light emitted by a hot solid is also “quantized” i.e. sent out as bursts of energy  Each of these bursts were referred to as a quantum of energy (read money analogy in text on page 169)  lead to thinking that as temperature increased, the proportion of each larger quantum becomes greater 3 The Photoelectric Effect  Maxwell proposed that light was an electromagnetic wave composed of electric and magnetic fields that could exert forces on charged particles (classical theory of light)  Hertz discovered photoelectric effect which involved the effect of electromagnetic radiation or light on substances  classical theory said that the intensity (brightness) shining on the metal would determine the kinetic energy of the freed electrons i.e. the brighter the light, the greater the energy of the emitted electrons 4  further work showed that this was incorrect and that it was the frequency (colour/energy) of the light that was the most important characteristic of the light to produce this effect  Einstein won the Nobel Prize (1905) for using Planck’s ideas to explain this effect  Einstein proposed the idea that light consisted of a stream of packets of energy or quanta which were later called photons  a photon of red light contained less energy that a photon of UV light  Einstein used idea that energy of photon could be transferred to the electron and that some of the energy was used to by the electron to break free with the rest being left over as kinetic energy in the ejected electron  The electron could not break free of the atom unless a certain quantity of energy is absorbed from a single photon (see marble in bowl explanation) 5 The Bohr Atomic Theory 6 BOHR MODEL (approx 1913) The Bohr model of the atom was based on the line spectra of the Hydrogen atom. The model also incorporated the concept developed by Einstein regarding the particle behaviour of light during emission or absorption (photon or quanta of energy). Postulates: 1. Energy Levels  an electron can only have specific energy values in an atom  the path followed by the electrons is a circular orbit (spherical in 3-D)  these energies are called energy levels given the name Principal Quantum Number, n  the orbit closest to the nucleus is given n=1, with the lowest energy  as one moves outward the energies get larger and “n” increases (n=1,2,3,4,...)  an electron can only circle in one of the allowed orbits WITHOUT a loss of energy 2. Transition Between Energy Levels  an electron in an atom can only change energy by going from one energy level to another level, i.e. NOT in-between  light is emitted when an electron falls from a higher energy level to a lower energy level  an electron generally remains in its ground state – the lowest energy level possible (n=1 for hydrogen)  when an electron is given energy it can be bumped to a higher energy level called the excited state  as the electron drops from the excited state down to lower energy levels it emits light  Bohr was able to show that his model was able to match the Balmer series (Visible light, drop to n=2) and predict the 7 Paschen series ( IR light , drop to n=3) and the Lyman series (UV light, drop to n=1) Although the Bohr model could explain the behaviour of the hydrogen atom, it could not fully explain the behaviour of other atoms QUANTUM NUMBERS During and subsequent to Bohr’s work, others continued to probe the behaviour of light emitted or absorbed by the atom. As a result, additional components were added to Bohr’s model to improve its ability to explain the behaviours observed. This led to the introduction of quantum numbers, in addition to Bohr’s Principle Quantum Number. Principal Quantum Number, n  Goes back to Bohr model which labeled the shells  Today called Principal Quantum Number, i.e. n=1,2,3……etc.  n=1 is closest to the nucleus with the lowest energy  Relates primarily to the main energy of an electron e.g. Fig. 1 “energy staircase” where energy levels are like unequal steps Secondary Quantum Number, l  Michelson found that the spectrum of H was composed of more than one line (experimental observation using spectrometry indicated that main lines of hydrogen spectra composed of more than one line – line splitting (Michelson, 1891)  Sommerfeld (1915) used elliptical orbits to extend the knowledge of the time by explaining Michelson’s work 8  He introduced the concept of there being additional electron subshells (or sublevels) that formed part of the energy levels and used the concept of the secondary quantum number to describe this concept  I.e. each energy “step” was a group of several little steps  “l” relates to the shape of the electron orbit and the number of values for l equals the volume of n  i.e. if n=3, then l=0,1,2 as letters: l = s, p, d, f, g Magnetic Quantum Number, m l  it was observed that if a gas discharge tube was placed near a strong magnet, some single lines split into new lines  called normal Zeeman effect after Zeeman who 1 observed st this (1897)  this was explained by Sommerfeld & Debye (1916) who thought that orbits may exist at varying angles and that the energies may be different when near strong magnets  for each value of l, m cln vary from –l to +l (each value represents a different orientation)  i.e. if l=1 then m cal be –1,0 or +1 if l=2 then m cln be –2,-1,0,+1, or +2 Summary: The magnetic quantum number m , relates tolthe direction of the orbit of the electrons. The number of values of m l represents the number of orientations of the orbits that we can have. Spin Quantum Number, m s  needed to explain additional spectral line-splitting & different kinds of magnetism  ferromagnetism-associated with substances containing Fe, Co & Ni  paramagnetism-weak attraction to strong magnets (individual atoms vs. collection of atoms)  paramagnetism couldn’t be explained until Wolfgang Pauli suggested that electrons spin on their axis (1925) 9  could spin only 2 ways (clockwise vs. counterclockwise) and he used only 2 numbers to describe this: m s +1/2 (clockwise) or –1/2 (
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