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CHAPTER 3 Atomic Structure: Explaining the Properties of Elements 10/17/2018 CHAPTER 3 Atomic Structure: Explaining the Properties of Elements We are going to learn about the electronic structure of the atom, and will be able to explain many things, including atomic orbitals, oxidation numbers, and periodic trends. 1 10/17/2018 Chapter Outline 3.1 Waves of Light 3.2 Atomic Spectra 3.3 Particles of Light: Quantum Theory 3.4 The Hydrogen Spectrum and the Bohr Model 3.5 Electrons as Waves 3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals 3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions 3.10 The Sizes of Atoms and Ions 3.11 Ionization Energies 3.12 Electron Affinities The Electromagnetic Spectrum Continuous range of radiant energy, (also called electromagnetic radiation). 2 10/17/2018 Electromagnetic Radiation Mutually propagating electric and magnetic fields, at right angles to each other, traveling at the speed of light c a) Electric b) Magnetic Speed of light (c) in vacuum = 2.998 x 108 m/s Properties of Waves - in the examples below, both waves are traveling at the same velocity Long wavelength = low frequency Short wavelength = high frequency 3 10/17/2018 · u = wavelength, = frequency, u = velocity Units: wavelength = meters (m) frequency = cycles per second or Hertz (s-1) wavelength (m) x frequency (s-1) = velocity (m/s) Example: A FM radio station in Portland has a carrier wave frequency of 105.1 MHz. What is the wavelength? 4 10/17/2018 Chapter Outline 3.1 Waves of Light 3.2 Atomic Spectra 3.3 Particles of Light: Quantum Theory 3.4 The Hydrogen Spectrum and the Bohr Model 3.5 Electrons as Waves 3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals 3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions 3.10 The Sizes of Atoms and Ions 3.11 Ionization Energies 3.12 Electron Affinities Atomic Emission (Line) Spectra 5 10/17/2018 Blackbody Radiation Photoelectric Effect • phenomenon of light striking a metal surface and producing an electric current (flow of electrons). • If radiation below threshold energy, no electrons released. 6 10/17/2018 Blackbody Radiation and the Photoelectric Effect Explained by a new theory: Quantum Theory • Radiant energy is “quantized” – Having values restricted to whole-number multiples of a specific base value. • Quantum = smallest discrete quantity of energy. • Photon = a quantum of electromagnetic radiation Quantized States Quantized states: Continuum states: discrete energy smooth transition levels. between levels. 7 10/17/2018 The energy of the photon is given by Planck’s Equation. E = h h = 6.626 × 10−34 J∙s (Planck’s constant) Sample Exercise 7.2 What is the energy of a photon of red light that has a wavelength of 656 nm? The value of Planck’s constant (h) is 6.626 × 10-34 J . s, and the speed of light is 3.00 × 108 m/s. 8 10/17/2018 Chapter Outline 3.1 Waves of Light 3.2 Atomic Spectra 3.3 Particles of Light: Quantum Theory 3.4 The Hydrogen Spectrum and the Bohr Model 3.5 Electrons as Waves 3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals 3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions 3.10 The Sizes of Atoms and Ions 3.11 Ionization Energies 3.12 Electron Affinities 9 10/17/2018 The Hydrogen Spectrum and the Rydberg Equation 1 1 1 = 1.097102 nm1 2 2 λ ni n f Exercise 7.4: using the Rydberg Equation What is the wavelength of the line in the emission spectrum of Hydrogen corresponding from ni = 7 to nf = 2? 1 1 1 = 1.097102 nm1 2 2 λ ni n f 10 10/17/2018 Using the Rydberg Equation for Absorption What is the wavelength of the line in the absorption spectrum of Hydrogen corresponding from ni = 2 to nf = 4? The Bohr Model of Hydrogen Neils Bohr used Planck and Einstein’s ideas of photons and quantization of energy to explain the atomic spectra of hydrogen 1 1 1 photon = 1 .of097 102 nm1 2 2 lightλ (h) ni n f n = 4 n = 3 n = 2 + h h n = 1 n = 2 n = 1 n = 3 absorption emission 11 10/17/2018 Electronic States • Energy Level: • An allowed state that an electron can occupy in an atom. • Ground State: • Lowest energy level available to an electron in an atom. • Excited State: • Any energy state above the ground state. “Solar system” model of the atom where each “orbit” has a fixed, QUANTIZED energy given by - 2.178686 x 10−18 Joules En=− n2 where n = “principle quantum number” = 1, 2, 3…. This energy is exothermic because it is potential energy lost by an unbound electron as it is attracted towards the positive charge of the nucleus. 12 10/17/2018 E = h E = h The Rydberg Equation can be derived from Bohr’s theory - Ephoton = DE = Ef - Ei 13 10/17/2018 Example DE Calculation Calculate the energy of a photon absorbed when an electron is promoted from ni = 2 to nf = 5. Sample Exercise 3.5 How much energy is required to ionize a ground-state hydrogen atom? Put another way, what is the ionization energy of hydrogen? 14 10/17/2018 Strengths and Weaknesses of the Bohr Model • Strengths: • Accurately predicts energy needed to remove an electron from an atom (ionization). • Allowed scientists to begin using quantum theory to explain matter at atomic level. • Limitations: • Applies only to one-electron atoms/ions; does not account for spectra of multielectron atoms. • Movement of electrons in atoms is less clearly defined than Bohr allowed. Chapter Outline 3.1 Waves of Light 3.2 Atomic Spectra 3.3 Particles of Light: Quantum Theory 3.4 The Hydrogen Spectrum and the Bohr Model 3.5 Electrons as Waves 3.6 Quantum Numbers 3.7 The Sizes and Shapes of Atomic Orbitals 3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions 3.10 The Sizes of Atoms and Ions 3.11 Ionization Energies 3.12 Electron Affinities 15 10/17/2018 Light behaves both as a wave and a particle - Classical physics - light as a wave: c = = 2.998 x 108 m/s Quantum Physics - Planck and Einstein: photons (particles) of light, E = h Several blind men were asked to describe an elephant. Each tried to determine what the elephant was like by touching it. The first blind man said the elephant was like a tree trunk; he had felt the elephant's massive leg. The second blind man disagreed, saying that the elephant was like a rope, having grasped the elephant's tail. The third blind man had felt the elephant's ear, and likened the elephant to a palm leaf, while the fourth, holding the beast's trunk, contended that the elephant was more like a snake. Of course each blind man was giving a good description of that one aspect of the elephant that he was observing, but none was entirely correct. In much the same way, we use the wave and particle analogies to describe different manifestations of the phenomenon that we call radiant energy, because as yet we have no single qualitative analogy that will explain all of our observations. http://www.wordinfo.info/words/images/blindmen-elephant.gif 16 10/17/2018 Standing Waves Not only does light behave like a particle sometimes, but particles like the electron behave like waves! WAVE-PARTICLE DUALITY Combined these two equations: E = mc2 and E = h, therefore - 17 10/17/2018 The DeBroglie wavelength explains why only certain orbits are "allowed" - a) stable b) not stable Sample Exercise 3.6 (Modified): calculating the wavelength of a particle in motion. (a)Calculate the deBroglie wavelength of a 142 g baseball thrown at 44 m/s (98 mi/hr) (b)Compare to the wavelength of an electron travelling at 2/3’s the speed of light. Mass electron = 9.11 x 10-31 kg. 18 10/17/2018 Using the wavelike properties of the electron. Close-up of a milkweed bug Atomic arrangement of a Bi-Sr- Ca-Cu-O superconductor 19 10/17/2018 Chapter Outline 3.1 Waves of Light 3.2 Atomic Spectra 3.3 Particles of Light: Quantum Theory 3.4 The Hydrogen Spectrum and the Bohr Model 3.5 Electrons as Waves 3.6 Quantum Numbers NOTE: will do Section 3.7 first 3.7 The Sizes and Shapes of Atomic Orbitals 3.8 The Periodic Table and Filling Orbitals 3.9 Electron Configurations of Ions 3.10 The Sizes of Atoms and Ions 3.11 Ionization Energies 3.12 Electron Affinities E. Schrödinger (1927) The electron as a Standing Wave . mathematical treatment results in a “wave function”, . = the complete description of electron position and energy . electrons are found within 3-D “shells”, not 2-D Bohr orbits . shells contain atomic “orbitals” (s, p, d, f) 20 10/17/2018 E. Schrödinger (1927) The electron as a Standing Wave .Each orbital can hold up to 2 electrons .The probability of finding the electron = 2 .Probabilities are required because of “Heisenberg’s Uncertainty Principle” We can never SIMULTANEOUSLY know with absolute precision both the exact position (x), and momentum (p = mass·velocity or mv), of the electron. Dx·D(mv) h/4 Uncertainty in Uncertainty in position momentum If one uncertainty gets very small, then the other becomes corresponding larger. If we try to pinpoint the electron momentum, it's position becomes "fuzzy".
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