Transcript Lecture 6
Chemistry 103 Lecture 6 Outline I. Electronic Structure Classical vs. Quantum Mechanics Orbitals/Quantum Numbers Electron Configurations Learning Check 1. Which of the following pairs are isotopes of the same element? 2. In which of the following pairs do both atoms have 8 neutrons? A. B. C. 15 15 8 7 X X 12 14 6 6 X X 15 16 7 8 X X 3 Electronic Structure Quantum Mechanics Nature of Electrons in Atoms Modern Periodic Table Periodic Law of the Elements – when elements are arranged in a particular order (increasing atomic number), elements of similar properties occur at periodic intervals The Theoretical basis for the periodic law lies in electronic theory (Nature of the electrons in an atom) http://www.woodrow.org/teachers/chemistry/institutes/1992/MENDELEEV.GIF Classical Physics In the late 1800’s, Classical (Newtonian) Physics worked so well for the macroscopic world, most academics thought we had discovered all there was to know about the subject Classical Mechanics There is no limit to the number of observables we can measure simultaneously These observables are continuous Quantum Mechanics Unfortunately, extremely small particles (electrons) do not follow the laws of classical (Newtonian) physics. The new physics that mathematically treats small particles is called Quantum Mechanics. Periodicity of Periodic Table Objective: Placing Electrons about the Nucleus of an Atom for a Particular Element. MODEL DEVELOPED (Quantum Numbers) APPLICATION (Electron Configurations) Quantum Model Attempts to explain certain properties of light’s interaction with matter. One early experimental observation that could not be explained classically: Emission spectra of heated gases. Light To embark upon the topic of quantum mechanics, we need to understand the nature of light. Electromagnetic Spectrum The electromagnetic spectrum Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Light Is a Form of Energy Light is a form of electromagnetic radiation Light travels as waves Light carries radiant energy through space Properties of Waves All waves have: Wavelength: (l) horizontal distance between two corresponding points on a wave (units are usually m) Frequency: (n) the number of complete wavelengths that pass a stationary point in a second (units are usually Hz, s-1) Question! Which wave has a higher frequency? Light Energy and Photons Light is a stream of small particles called photons that have • Energy related to their frequency. Using Plank’s constant (h) High energy with a high frequency Low energy with a low frequency Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Electromagnetic Spectrum The electromagnetic spectrum • Arranges forms of energy from lower to higher • Arranges energy from longer to shorter wavelengths • Shows visible light with wavelengths from 700-400 nm Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Light Equations All electromagnetic radiation moves through a vacuum at a specific speed. c = 3.00 x 108 m/s Wave Equations for light: c = ln Energy Equation for light: E = hn h = Planck’s constant h = 6.626 x 10-34 J s Light Equations Wave Equations for light: c = ln Energy Equation for light: E = hn Substituting frequency (n = c/l) E = hc l Light Calculations What is the frequency of green light if it has a wavelength of 500. nm? Light Calculations A wavelength of 850. nm is used for fiber-optic transmission. What is the energy of this wavelength? Spectrum White light that passes through a prism • Is separated into all colors called a continuous spectrum • Gives the colors of a rainbow Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Atomic Emission Spectrum An atomic spectrum consists of lines • Of different colors formed when light from a heated element containing an element passes through a prism spectrum Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Atomic Emission Spectra When gases are heated, they give off light at certain frequencies. In other words, atoms absorb or emit energy only at specific wavelengths (specific energies) (Bodner and Pardue, 1995) Quantum Mechanics Discrete values for observables - quantized values. Probability - does not tell us where an electron is exactly, tells us the probability of an electron being at some point in space Quantum Mechanics Summarized A Mathematical Model that described what was being observed experimentally. The Math defined the Theory. The creators of Quantum Mechanics had difficulties when it came to interpretations made in a classical world. In reference to “Quantum Mechanics” Schrodinger (1887-1961) “I don’t like it, and I’m sorry I ever had anything to do with it.” Bohr (1885 - 1961) “Anyone who is not shocked by Quantum Theory has not understood it. Levine “We cannot hope to obtain a proper understanding of microscopic particles based on models taken from our experiences in the macroscopic world.” Development of Theory - early steps Bohr’s Model of the Atom His Goal was to create a model that would explain the atomic spectrum of Hydrogen Electrostatic Quandary Postulates that there are only certain positions about the nucleus an electron can reside STATIONARY STATES OF MOTION Bohr’s Model of the Atom •When electrons absorb energy (in the form of light), they move from a lower energy level to a higher one. •When electrons move from a higher energy level to a lower energy level, they give off energy in the form of light (“emission”) Electron Energy Levels Electrons are arranged in specific energy levels that • Are labeled n = 1, n = 2, n = 3, and so on • Increase in energy as n increases • Have the electrons with the lowest energy in the first energy level (n=1)closest to the nucleus Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Energy Level Changes • An electron absorbs energy to “jump” to a higher energy level. When an electron falls to a lower energy level, energy is emitted. In the visible range, the emitted energy appears as a color. Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings Adding Probability The very nature of locating an electron’s position causes uncertainty it another property (kinetic energy). Can’t know both simultaneously UNCERTAINTY Quantum Numbers An introduction into the math behind the theory. Quantum Numbers - The Model Shell (n) QUANTUM NUMBERS Primary or Shell Quantum number (n = 1, 2, 3…) (Bohr’s Model interpretation & notation) Electron Shell A region of space about a nucleus that contains electrons that have approximately the same energy and that spend most of their time approximately the same distance from the nucleus Quantum Numbers - The Model Shell (n) l=0. Subshell (l) l=1 l=2 l=3 Electron Subshells/Sublevels Every shell (n) has subshell(s) (l) l = 0, 1, … n - 1 Discribes the spatial distribution for an electron (PROBABILITY) l = 0 (s), l = 1(p), l = 2, (d), l = 3 (f), … Electron Orbitals (n & l ) Orbital Region of space where two electrons are likely to be found (90% probability) Have different shapes depending on which subshell (l quantum number) they are in Quantum Numbers & Orbitals n=1,l=0 First shell or energy level, you have one orbital: 1s An “s” orbital as a spherical shape around the nucleus Quantum Numbers and Orbitals n=2 l=0 l=1 Orbital Name 2s Orbital Name 2p s Orbitals An s orbital • Has a spherical shape around the nucleus • Increases in size around the nucleus as the energy level n value increases Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings 42 Quantum Numbers & Orbitals n=3 l = 0(s), l = 1(p), l = 2 (d) n=4 l = 0(s), l=1(p), l=2(d), l=3(f)