1. Quantum Theory & Bohr’s Model of the Atom Ch. 4 - Electrons in Atoms

2. A. Quantum Theory zPlanck (1900) yObserved - emission of specific colors of light from hot objects yConcluded - energy is emitted in small, specific amounts (quanta) yQuantum - minimum amount of energy that can be lost or gained by an atom.

3. A. Quantum Theory zPlanck (1900) vs. Classical TheoryQuantum Theory

4. A. Quantum Theory zEinstein (1905) yObserved - photoelectric effectphotoelectric effect

5. A. Quantum Theory zEinstein (1905) yConcluded - light has properties of both waves and particles “wave-particle duality” yPhoton - particle of light that carries a quantum of energy

6. A. Quantum Theory E:energy (J, joules) h:Planck’s constant (6.626  10 -34 J/Hz) :frequency (Hz) E = h zThe energy of a photon is proportional to its frequency of the light.

7. A. Examples GIVEN: E = ? = 4.57  10 14 Hz h = 6.626  10 -34 J/Hz WORK : E = h E = ( 6.626  10 -34 J/Hz ) ( 4.57  10 14 Hz ) E = 3.03  10 -19 J zEX 1: Find the energy of a red photon with a frequency of 4.57  10 14 Hz.

8. A. Examples GIVEN: E = ? = 1.0  10 -9 m c = 3.00  10 8 m/s h = 6.626  10 -34 J/Hz WORK : E = h = hc/ λ = (6.626 ·10 -34 J/Hz)(3.00 ·10 8 m/s) 1.0 x 10 -9 m E = 2.0  10 -16 J zEX 2: Find the energy of a photon whose wavelength is 1.0  10 -9 m.

9. B. Bohr’s Model zLinked photon emission with an atom’s electrons (e - ) zIn Bohr’s model, e - exist only in orbits with specific amounts of energy called energy levels

10. B. Bohr’s Model zTherefore… ye - can only gain or lose certain amounts of energy (quanta) yonly certain photons are produced yeach photon has a unique frequency, and therefore a unique color of light is seen

11. B. Bohr’s Model ground state (low energy orbit; closer to nucleus) excited state (high energy orbit; further from nucleus) ENERGY IN PHOTON OUT e-e-e-e- electron “relaxes” Line Emission Spectrum produced

12. D. Bohr Model 1 2 3 4 5 6 zEnergy of photon depends on the difference in energy levels zBohr’s calculated energies matched the IR (Paschen), visible (Balmer), and UV (Lyman) lines for the H atom Hydrogen Emission Spectrum

13. D. Other Elements zEach element has a unique bright-line emission spectrum. y“Atomic Fingerprint” Helium zBohr’s calculations only worked for hydrogen! 