1 / 30

Electronic Structure of Atoms (i.e., Quantum Mechanics)

Electronic Structure of Atoms (i.e., Quantum Mechanics). Brown, LeMay Ch 6 AP Chemistry Monta Vista High School. What does light have to do with the atomic model?.

nicolebailey
Download Presentation

Electronic Structure of Atoms (i.e., Quantum Mechanics)

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Electronic Structure of Atoms(i.e., Quantum Mechanics) Brown, LeMay Ch 6 AP Chemistry Monta Vista High School

  2. 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.

  3. 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 Hz- Hertz, hz= cycles/s or 1/s) c = l n

  4. 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). • where Planck’s constant (h) = 6.63 x 10-34 J-s Animation Link Max Planck(1858-1947)

  5. The Photoelectric Effect • Spontaneous emission of e- from metal struck by light; first explained by Einstein in 1905A 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)

  6. 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 absorptionspectra.

  7. Hydrogen: contains 1 red, 1 green, 1 blue and 1 violet. Atomic Emission Spectra of C and H Carbon: Contains many more emission lines as compared to H. Why?

  8. 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: principal quantum number (n) = 1, 2, 3, 4, … Rydberg’s constant (RH) = 2.178 x 10-18 J Niels Bohr(1888-1962) Johannes Rydberg(1854-1919)

  9. 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.

  10. 5 4 3 2 1 E5 E4 E3 E2 E1 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 Increasing Energy, E Principal Quantum Number, n

  11. Phased out!! ni = initial orbital of e- nf = final orbital of e- in its transitionMovie on e transition

  12. Theodore Lyman (1874 - 1954) 5 4 3 2 1 FriedrichPaschen(1865 - 1947) n ? Phased out! JohannBalmer(1825 – 1898) FrederickBrackett(1896 – 1988)

  13. 6.4: Matter is a Wave Planck said: E = h c / l Einstein said: E = m c2 Louis DeBroglie said (1924): h c / l = m c2 h / l = m c Therefore: Louisde Broglie(1892 - 1987)

  14. 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!!

  15. 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 within small confining structures.

  16. Heisenberg’s Uncertainty Principle (1927) It is impossible to determine the exact position and exact momentum (p) of an electron. p = m v • 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 may change its motion (its momentum). WernerHeisenberg(1901 – 1976)

  17. Electron density distribution in H atom

  18. 6.5: Quantum Mechanics & Atomic Orbitals Schrödinger’s wave function: • Relates probability (Y2) of predicting position of e- to its energy. ErwinSchrödinger(1887 – 1961) Where: U = potential energy x = position t = time m = mass i =√(-1) http://daugerresearch.com/orbitals/index.shtml

  19. Probability plots of 1s, 2s, and 3s orbitals

  20. 6.6: Representations of Orbitalswww.orbitals.com; animation 1, Draft of a letter from Bohr to Heisenberg (never sent) s orbital p orbitals

  21. d orbitals f orbitals: very complicated

  22. 6.7: Filling Order of Orbitals 7p 6d 7s 5f x 7 6p 5d 6s 4f x 7 5p Increasing Energy 4d 5s 4p 3d 4s 3p 3s 2p 2s 1s • Aufbau principle: e- enter orbitals of lowest energy first • Relative stability & average distance of e- from nucleus

  23. Animation for filling of Orbitals 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Use the “diagonal rule” (some exceptions do occur). Sub-level maxima: s = 2 e- p = 6 e- d = 10 e- f = 14 e- …

  24. Pauli exclusion principle (1925): no two e- can have the same four quantum numbers; e- in same orbital have opposite spins (up and down) • Hund’s rule: e- are added singly to each equivalent (degenerate) orbital before pairing • Ex: Phosphorus (15 e-) has unpaired e- inthe valence (outer) shell. • 1s 2s 2p 3s 3p WolfgangPauli(1900 – 1958) FriedrichHund(1896 - 1997)

  25. p1 p2 p3 p4 p5 p6 f1 d2 f2 f3 d3 f4 d5 f5 d5 f6 d6 d7 f7 f8 d8 f9 d10 f10 d10 f11 f12 f13 f14 6.9: Periodic Table & Electronic Configurations s block f block d block p block s2 s1 s2 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p d1 3d 4d 5d 6d 3d 4d 5d 6d 4f 5f Notable Exceptions: Cr & Mo: [Ar] 4s1 3d5 not [Ar] 4s2 3d4 Cu, Ag, & Au: [Ar] 4s13d10 not [Ar] 4s23d9

  26. Electronic Configurations [Ar] 4s23d1 [Ar] 4s23d104p1 1s22s22p63s23p64s23d104p1

  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 • Expanded Electron Configuration • Condensed Electron Configurations • Orbital Notation • 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

More Related