1. I. Waves & Particles Ch. 4 - Electrons in Atoms

2. Light and Electrons zBecause light and electrons have common properties, understanding one helps to understand the other.

3. Electromagnetic radiation zEnergy that exhibits wave-like behavior as it travels zIncludes: gamma rays, X-rays, infrared, visible spectrum, microwaves, ultraviolet rays, radio and TV waves

4. EM Spectrum LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY

5. LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY ROYG.BIV redorangeyellowgreenblueindigoviolet

6. Waves zWavelength ( ) - length of one complete wave (measured in m, cm, nm) zFrequency ( ) - # of waves that pass a point during a certain time period yhertz (Hz) = 1/s (s -1 ) zAmplitude (A) - distance from the origin to the trough or crest

7. Waves A greater amplitude (intensity) greater frequency (color) crest origin trough A

8. EM Spectrum zFrequency & wavelength are inversely proportional c = c:speed of light (3.00  10 8 m/s) :wavelength (m, nm, etc.) :frequency (Hz)

9. EM Spectrum GIVEN: = 7.50 x !0 12 Hz = ? c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 7.50  10 12 Hz = 4.00  10 -5 m zEX: Calculate the wavelength of radiation whose frequency is 7.50 x !0 12 Hz.

10. Light as Particles zA property which could not be explained in terms of waves was a phenomenon known as the photoelectric effect – refers to the emission of electrons from a metal when heated or lit.

11. Quantum Theory zPlanck (1900) yObserved - emission of light from hot objects yConcluded - energy is emitted in small, specific amounts (quanta) yQuantum - minimum amount of energy change

12. Quantum Theory zPlanck (1900) vs. Classical TheoryQuantum Theory

13. E:energy (J, joules) h:Planck’s constant (6.6262  10 -34 J·s) :frequency (Hz) E = h zThe energy of a photon is proportional to its frequency.

14. Quantum Theory GIVEN: E = ? = 3.55  10 17 Hz h = 6.6262  10 -34 J·s WORK : E = h E = ( 6.6262  10 -34 J·s ) ( 3.55  10 17 Hz ) E = 2.35  10 -16 J zEX: Find the energy of a photon with a frequency of 3.55  10 17 Hz.

15. Quantum Theory zEinstein (1905) yObserved - photoelectric effect

16. Quantum Theory zEinstein (1905) yConcluded - light has properties of both waves and particles “wave-particle duality” yPhoton - particle of light, having zero mass, carrying a quantum of energy

17. Quantum Theory zRadiation is emitted and absorbed only in whole numbers of photons

18. II. Bohr Model of the Atom Ch. 4 - Electrons in Atoms

19. A. Line-Emission Spectrum ground state excited state ENERGY IN PHOTON OUT

20. B. Bohr Model zLinked the atom’s electron with photon emission ze - exist only in paths, or orbits, with specific amounts of energy called energy levels zTherefore… ye - can only gain or lose certain amounts of energy yonly certain photons are produced

21. B. Bohr Model 1 2 3 4 5 6 zEnergy of photon depends on the difference in energy levels ze - jumps up when energy is absorbed- gives off light when is falls back down

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

23. zBohr’s model of the atom explained electrons as particles.

24. A. Electrons as Waves zLouis de Broglie (1924) yApplied wave-particle theory to e - ye - exhibit wave properties

25. B. Quantum Mechanics zHeisenberg Uncertainty Principle yImpossible to know both the velocity and position of an electron at the same time

26. B. Quantum Mechanics zSchrödinger Wave Equation (1926) ytreated e- moving around the nucleus as waves ydefines probability of finding an e - ydefines mathematically the wave properties of electrons

27. B. Quantum Mechanics Radial Distribution Curve Orbital zOrbital (“electron cloud”) yRegion in space where there is 90% probability of finding an e -