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Electronic Structure of Atoms (i.e., Quantum Mechanics) Brown, LeMay Ch 6 AP Chemistry Monta Vista High School 1 What does light have to do with the atomic model? Scientists knew the nature of light but knew little about the nature of matter. To understand the nature of matter, scientists studied the changes caused in light by interaction of matter. From these studies, scientist tried to extrapolate information about the nature of matter. 2 6.1: Light is a Wave Electromagnetic spectrum: A form of radiant energy (can travel without matter) Both electrical and magnetic (properties are perpendicular to each other) http://imagine.gsfc.nasa.gov/Videos/general/spectrum.mov Speed of Light: c = 3.0 x 108 m/s (in a vacuum) http://www.astronomynotes.com/light/s3.htm Wavelength (l): distance between wave peaks (determines “color” of light), measured in nm, m etc. Frequency (n): # cycles/sec (measured in HzHertz, hz= cycles/s cor= 1/s) ln 3 6.2: Light is a Particle (Quantum Theory) Blackbody radiation: * Blackbody: object that absorbs all EM radiation that strikes it; it can radiate all possible wavelengths of EM; below 700 K, very little visible EM is produced; above 700 K visible E is produced starting at red, orange, yellow, and white before ending up at blue as the temperature increases ◦ discovery that light intensity (energy emitted per unit of time) is proportional to T4; hotter = shorter wavelengths “Red hot” < “white hot” < “blue hot” Interactive Link • Planck’s Theory: (explained blackbody radiation by quantization of energy transfer) Blackbody radiation can be explained if energy can be released or absorbed in packets of a standard size he called quanta (singular: quantum). E hn hc l where Planck’s constant (h) = 6.63 x 10-34 J-s E hn Animation Link hc Max Planck (1858-1947) l 4 The Photoelectric Effect Spontaneous emission of e- from metal struck by light; first explained by Einstein in 1905 A quantum strikes a metal atom and the energy is absorbed by an e-. If the energy is sufficient, e- will leave its orbital, causing a current to flow throughout the metal. To explain photoelectric effect, quantization of light was put forth by Einstein. Animation Albert Einstein (1879-1955) 5 6.3: Bohr’s Model of the H Atom (and only H!) Applied quantization of energy transfer to the atomic model Studied atomic spectrum of H to come up with atomic model. Atomic emission spectra: Most sources produce light that contains many wavelengths at once. Animation However, light emitted from pure substances may contain only a few specific wavelengths of light called a line spectrum (as opposed to a continuous spectrum). Animation Atomic emission spectra are inverses of atomic absorption spectra. 6 Atomic Emission Spectra of C and H Hydrogen: contains 1 red, 1 green, 1 blue and 1 violet. Carbon: Contains many more emission lines as compared to H. Why? 7 Niels Bohr theorized that e-: ◦ Travel in certain “orbits” around the nucleus, or, are only stable at certain distances from the nucleus ◦ If not, e- should emit energy, slow down, and crash into the nucleus. Allowed orbital energies are defined by: RH 2.178 10 En 2 n n2 18 principal quantum number (n) = 1, 2, 3, 4, … Rydberg’s constant (RH) = 2.178 x 10-18 J Johannes Rydberg (1854-1919) Niels Bohr (1888-1962) 8 Think, Pair, Share Activity With your elbow partner, describe Electromagnetic radiation, blackbody radiation, Plank’s theory and Photoelectric effect. Address each of the above in the following terms: 1. What is it? 2. Why was it important? 3. What existing theory or concept, it approved/disapproved. • 9 5 4 E3 3 E2 2 E1 1 Principal Quantum Number, n Increasing Energy, E E5 E4 As n approaches ∞, the e- is essentially removed from the atom, and E∞ = 0. • • ground state: lowest energy level in which an e- is stable excited state: any energy level higher than an e-’s ground state 1 1 E R H 2 2 ni nf E R H 1 1 2 2 n h h n f ni Phased out!! E R H 1 1 2 2 n h h n i nf ni = initial orbital of enf = final orbital of e- in its transition Movie on e transition 11 5 4 3 2 Friedrich Paschen (1865 - 1947) n Theodore Lyman (1874 - 1954) Johann Balmer (1825 – 1898) Phased out! 1 Frederick Brackett (1896 – 1988) ? Figure 1: Line series are transitions from one level to another. Series Transition down to (emitted) or up from (absorbed)… Type of EMR Lyman 1 UV Balmer Paschen Brackett 2 3 4 Visible IR Far IR 6.4: Matter is a Wave Planck said: E=hc/l Einstein said: E = m c2 Louis DeBroglie said (1924): h c / l m c2 h/lmc Louis de Broglie Therefore: (1892 - 1987) Particles (with mass) have an m = h / cl associated wavelength Waves (with a wavelength) have an l h / mc associated mass and velocity 13 Neils Bohr Model: Partner Activity On a sheet of paper, take turns with your partner drawing Bohr’s model of atom. Draw the following in context of Bohr’s Model: 1.nucleus 2.energy levels (1,2,3,4) 3.an electron in energy level 2 4. Show an electron transition from energy level 2 to 3 5. Write formula for calculating this energy change and calculate energy. 6. Give each other high fives!! 14 IBM – Almaden: “Stadium Corral” This image shows a ring of 76 iron atoms on a copper (111) surface. Electrons on this surface form a two-dimensional electron gas and scatter from the iron atoms but are confined by boundary or "corral." The wave pattern in the interior is due to the density distribution of the trapped electrons.Their energies and spatial distribution can be quite accurately calculated by solving the classic problem of a quantum mechanical particle in a hard-walled box. Quantum corrals provide us with a unique opportunity to study and visualize the quantum behavior of electrons Heisenberg’s Uncertainty Principle (1927) It is impossible to determine the exact position and exact momentum (p) of an electron. p=mv To determine the position of an e-, you have to detect how light reflects off it. But light means photons, which means energy. When photons strike an e-, they Werner may change its motionHeisenberg (its momentum). (1901 – 1976) 16 Electron density distribution in H atom 17 6.5: Quantum Mechanics & Atomic Orbitals Schrödinger’s wave function: Relates probability (Y2) of predicting position of e- to its energy. h2 d 2Y dY E UY ih 2 2m dx dt Erwin Schrödinger (1887 – 1961) Where: U = potential energy x = position t = time m = mass i =√(-1) http://daugerresearch.com/orbitals/index.shtml 18 Probability plots of 1s, 2s, and 3s orbitals 19 6.6: Representations of Orbitals www.orbitals.com; animation 1, Draft of a letter from Bohr to Heisenberg (never sent) s orbital p orbitals 20 d orbitals f orbitals: very complicated 6.7: Filling Order of Orbitals 1. Aufbau principle: e- enter orbitals of lowest energy first 7p 7s 6s 5s 6p 5p 4p 6d 5d 5f x7 4f x7 4d 3d 4s 3p 3s 2p 2s 1s • Relative stability & average distance of e- from nucleus 22 Animation for filling of Orbitals Use the “diagonal rule” (some exceptions do occur). Sub-level maxima: s = 2 ep = 6 ed = 10 ef = 14 e… 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 23 2. Pauli exclusion principle (1925): no two e- can have the same four quantum numbers; e- in same orbital have opposite spins (up and down) Wolfgang Pauli (1900 – 1958) 3. Hund’s rule: e- are added singly to each equivalent (degenerate) orbital before pairing Ex: Phosphorus (15 e-) has unpaired e- in the valence (outer) shell. 1s 2s 2p 3s 3p Friedrich Hund (1896 - 1997) 24 6.9: Periodic Table & Electronic Configurations s block f block d block p block s2 s1 s2 1s 2s 3s 4s 5s 6s 7s p1 p2 p3 p4 p5 p 6 d1 3d 4d f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 5d 6d 4f 5f 2p d2 d3 d5 d5 d6 d7 d8 d10d103p 4p 3d 5p 11 12 13 14 4d f f f f 6p 5d 7p 6d Notable Exceptions: Cr & Mo: [Ar] 4s1 3d5 not [Ar] 4s2 3d4 Cu, Ag, & Au: [Ar] 4s13d10 not [Ar] 4s23d9 Electronic Configurations Element Standard Configuration Nitrogen 1s22s22p3 Scandium 1s22s22p63s23p64s23d1 Gallium 1s22s22p63s23p64s23d104p1 Noble Gas Shorthand [He] 2s22p3 [Ar] 4s23d1 [Ar] 4s23d104p1 26 Noble Gas Shorthand Element Standard Configuration Lanthanum 1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s25d1 [Xe] 6s25d1 Cerium 1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s25d14f1 [Xe] 6s25d14f1` 1s2 2s22p6 3s23p6 4s23d104p6 Praseodymium 5s24d105p6 6s24f3 [Xe] 6s24f3 27 Electron Configuration for Ions Valence Electrons: Only s and p e are valence electrons. The maximum number of valence e that an atom can have is 8. WHY? Write the electron configurations for the following ions: Cr + Cr3+ Ground State Electron Config. V. Excited State Electron Configuration 28 Ways to Represent Electron Configuration 1.Expanded Electron Configuration 2.Condensed Electron Configurations 3.Orbital Notation 4.Electron Dot Structure Write the above four electron configurations for Zinc, Zinc ion and Cu ion. Paramagnetic Diamagnetic Why are some ions colored and some aren’t? 29 Electron Configuration and Para- and Diamagnetism demo + activity 30