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I. Waves & Particles (p. 91 - 94)

Ch. 4 - Electrons in Atoms. I. Waves & Particles (p. 91 - 94). A. Waves. Wavelength ( ) - length of one complete wave Frequency ( ) - # of waves that pass a point during a certain time period hertz (Hz) = 1/s Amplitude (A) - distance from the origin to the trough or crest. . crest.

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I. Waves & Particles (p. 91 - 94)

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  1. Ch. 4 - Electrons in Atoms I. Waves & Particles(p. 91 - 94)

  2. A. Waves • Wavelength () - length of one complete wave • Frequency () - # of waves that pass a point during a certain time period • hertz (Hz) = 1/s • Amplitude (A) - distance from the origin to the trough or crest

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

  4. B. EM Spectrum HIGH ENERGY LOW ENERGY

  5. R O Y G. B I V red orange yellow green blue indigo violet B. EM Spectrum HIGH ENERGY LOW ENERGY

  6. B. EM Spectrum • Frequency & wavelength are inversely proportional c =  c: speed of light (3.00  108 m/s) : wavelength (m) : frequency (Hz)

  7. C. Quantum Theory • Planck (1900) • Observed - emission of light from hot objects • Concluded - energy is emitted in small, specific amounts (quanta) • Quantum - minimum amount of energy change

  8. Classical Theory Quantum Theory C. Quantum Theory • Planck (1900) vs.

  9. C. Quantum Theory • Einstein (1905) • Observed - photoelectric effect

  10. C. Quantum Theory • Einstein (1905) • Concluded - light has properties of both waves and particles “wave-particle duality” • Photon - particle of light that carries a quantum of energy

  11. C. Quantum Theory • The energy of a photon is proportional to its frequency. E: energy (J, joules) h: Planck’s constant (6.6262  10-34 J·s) : frequency (Hz) E = h

  12. Ch. 4 - Electrons in Atoms II. Bohr Model of the Atom(p. 94 - 97)

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

  14. B. Bohr Model • e- exist only in orbits with specific amounts of energy called energy levels • Therefore… • e- can only gain or lose certain amounts of energy • only certain photons are produced

  15. Energy of photon depends on the difference in energy levels Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom B. Bohr Model 6 5 4 3 2 1

  16. C. Other Elements • Each element has a unique bright-line emission spectrum. • “Atomic Fingerprint” Helium • Bohr’s calculations only worked for hydrogen! 

  17. ConcepTest According to the energy diagram below for the Bohr model of the hydrogen atom, if an electron jumps from E1 to E2, energy is... absorbed emitted not involved absorbed

  18. Ch. 4 - Electrons in Atoms III. Quantum Model of the Atom(p. 98 - 104)

  19. A. Electrons as Waves • Louis de Broglie (1924) • Applied wave-particle theory to e- • e- exhibit wave properties QUANTIZED WAVELENGTHS

  20. A. Electrons as Waves QUANTIZED WAVELENGTHS

  21. VISIBLE LIGHT ELECTRONS A. Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS

  22. B. Quantum Mechanics • Heisenberg Uncertainty Principle • Impossible to know both the velocity and position of an electron at the same time

  23. B. Quantum Mechanics • SchrödingerWave Equation (1926) • finite # of solutions  quantized energy levels • defines probability of finding an e-

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

  25. UPPER LEVEL C. Quantum Numbers • Specify the “address” of each electron in an atom

  26. C. Quantum Numbers 1. Principal Quantum Number ( n ) • Energy level • Size of the orbital • n2 = # of orbitals in the energy level

  27. s p d f C. Quantum Numbers 2. Angular Momentum Quantum # ( l ) • Energy sublevel • Shape of the orbital

  28. C. Quantum Numbers • n = # of sublevels per level • n2 = # of orbitals per level • Sublevel sets: 1 s, 3 p, 5 d, 7 f

  29. C. Quantum Numbers 3. Magnetic Quantum Number ( ml) • Orientation of orbital • Specifies the exact orbitalwithin each sublevel

  30. C. Quantum Numbers px py pz

  31. 2s 2px 2py 2pz C. Quantum Numbers • Orbitals combine to form a spherical shape.

  32. C. Quantum Numbers 4. Spin Quantum Number ( ms) • Electron spin  +½ or -½ • An orbital can hold 2 electrons that spin in opposite directions.

  33. C. Quantum Numbers • Pauli Exclusion Principle • No two electrons in an atom can have the same 4 quantum numbers. • Each e- has a unique “address”: 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin # energy level sublevel (s,p,d,f) orbital electron

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