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Chapter 6: Electronic Structure of Atoms

Chapter 6: Electronic Structure of Atoms. Pages 207-247. Brief Explanation of the Chapter. Explore Quantum Theory Nature of Light Arrangement of electrons in atoms Exploration of reactivity History. Wave Nature of Light. Knowledge of electrons come from light

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Chapter 6: Electronic Structure of Atoms

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  1. Chapter 6: Electronic Structure of Atoms Pages 207-247

  2. Brief Explanation of the Chapter • Explore Quantum Theory • Nature of Light • Arrangement of electrons in atoms • Exploration of reactivity • History

  3. Wave Nature of Light • Knowledge of electrons come from light • How can we numerically define the speed of light? • 3.00 x 108 m/s • How do we explain the term periodic in reference to waves? • Define Wavelength • Define Frequency

  4. Wavelength of Light • How are the wavelength and frequency of electromagnetic radiation related? • Why do different types of electromagnetic radiation have different properties? • How do the wavelength and frequency of an X-ray compare with those of the red light from a neon sign? • How do we express frequency?

  5. Practice • How did you determine which of the waves represented blue, and which represented red?

  6. Practice • A laser used in eye surgery to fuse detached retinas produces radiation with a wavelength of 640.0 nm. Calculate the frequency of this radiation. • An FM radio station broadcasts electromagnetic radiation at a frequency of 103.4 MHz. Calculate the wavelength of this radiation?

  7. Quantized Energy and Photons • Wave model of light does not explain three very important phenomena. • The emission of light from hot objects (blackbody radiation) • The emission of electrons from metal surfaces on which light shines (photoelectric effect) • The emission of light from electronically excited gas atoms (emission spectra)

  8. Hot Objects and Quantization of Energy (blackbody radiation) • 1900- Max Planck • Assumed energy either released or absorbed by atoms in “chunks” of some minimum size. • termed QUANTUM (fixed amount) to smallest quantity of energy that can be emitted or absorbed • Proposed that energy (E) of a single quantum equals constant times the frequency of radiation E=hv

  9. Planck’s Constant • h= 6.626 x 10-34 joule-second (J-s) • Can be emitted or absorbed in whole number multiples of hv (hv, 2hv, 3hv) • If 3hv has been released, we say: 3 quanta of energy have been emitted • Stairs Vs. Ramp

  10. Photoelectric Effect and Photons • Albert Einstein • Light shining on a clean metal surface causes surface to emit electrons • Minimum frequency of light (different for different metals) is required for emission. • Light with frequency of 4.60 x 1014 s-1 or greater causes cesium metal to emit electrons. • Radiant energy striking surfaces behaves like stream of energy particles (PHOTONS) • Does light act more like a wave or a particle?

  11. Practice Problems • A laser emits light that a frequency of 4.69 x 10^14/s. What is the energy of one photon of this radiation? • If the laser emits a pulse containing 5.0 x 10^17 photons of this radiation, what is the total energy of that pulse? • If the laser emits 1.3 x 10^-2 J of energy during a pulse, how many photons are emitted?

  12. Bohr and Line Spectra • Neils Bohr (1913) • Theoretical explanation of line spectra

  13. Line Spectra • Continuous Vs. Line Spectra, how do they differ? • How do we produce a bright line spectra? • How is that energy emitted? • We know there is a definite change in energy, this corresponds to: • Definite Frequency • Definite Wavelength

  14. Bohr’s Model • Assumed electrons move in circular orbits around nucleus • Goes against physics!! • Charged particle moving in circular orbit should continually lose energy and spiral into positively charged nucleus • Based model on 3 postulates: • Only orbits of certain radii (with specific energies) permitted for electrons in H atom. • Electron in permitted orbit is in “allowed” energy state and therefore does not radiate energy • Energy is emitted or absorbed as a photon only when moving energy levels

  15. Energy States of Hydrogen Atom • Bohr began calculating possible energy levels and found they fit into a specific equation.

  16. December 13th, 2012 • Do Now: 1.Einstein’s 1905 paper on the photoelectric effect was the first important application of Planck’s quantum hypothesis. Describe his original hypothesis, and explain how Einstein made use of it in his theory of the photoelectric effect. 2. A laser emits a wavelength of 987 nm. In what portion of the spectrum is this found? Its output energy is absorbed in a detector that measures a total energy of 0.52 J over 32 seconds. How many photons per second are being emitted?

  17. Energy States of a Hydrogen Atom • With three postulates and classical equations for motion and interacting charges, Bohr calculated energy corresponding to orbits. E= (-hcRH) (1/n2) = (-2.18 x 10-18 J) (1/n2) h= plancks constant c= speed of light RH= Rydbergs Constant n= principal quantum number (range from 1- infinity)

  18. Energy of Hydrogen cont. • Each orbit = different n value • Energies of electrons of a H atom given by this equation are negative for all values of n (most negative being closest to 1) • Lower the energy, more stable the atom • Lowest energy state = ground state • When in higher = excited state

  19. Yikes! • How is the radius effected as n becomes infinitely larger? • How is the energy effected as n becomes infinitely larger? • The state in which the electron is removed from the nucleus is the reference, or zero energy, state. • ∆E = Ef – Ei = Ephoton = hv

  20. Hydrogen Atom Cont. • Bohr’s model states that specific frequencies of light satisfy the previous equation. Therefor we can state: • If nf is smaller than ni, the electron is moving closer to the nucleus and change in energy is negative *indicates a release of energy* • This equation can be used to calculate frequency or wavelength as well.

  21. Limitations to Bohr • Bohr model can not explain line spectra beyond the Hydrogen atom (except in crude manner) • As we have seen, electrons also have wavelike properties • Things kept from Bohr Model: • Electrons exist in certain discrete energy levels described by quantum numbers • Energy is involved in moving an electron from one level to another

  22. Wave Behavior of Matter • Louis de Broglie: • Studied Dual Nature of Particles (wave and radiant energy) • Suggested that the electron is associated with a particular wavelength. • Wavelength of electron is dependent upon mass and velocity ( m and v) • λ= h/ mv • Quantity mv = momentum • Termed “matter waves” to describe wave characteristics of particles

  23. Practice Problem • Calculate the velocity of a neutron whose de Broglie wavelength is 500 pm. The mass of a neutron is 1.67492716 x 10-24 g.

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