1. I. Waves & Particles Ch. 6 – Electronic Structure of Atoms

2. Properties of Waves z Many of the properties of light may be described in terms of waves even though light also has particle-like characteristics. z Waves are repetitive in nature

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

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

5. Electromagnetic Radiation z Electromagnetic radiation: (def) form of energy that exhibits wavelike behavior as it travels through space zTypes of electromagnetic radiation: y visible light, x-rays, ultraviolet (UV), infrared (IR), radiowaves, microwaves, gamma rays

6. Electromagnetic Spectrum z All forms of electromagnetic radiation move at a speed of about 3.0 x 10 8 m/s through a vacuum (speed of light) zElectromagnetic spectrum: made of all the forms of electromagnetic radiation

7. B. EM Spectrum LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY

8. LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY ROYG.BIV redorangeyellowgreenblueindigoviolet

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

10. B. EM Spectrum GIVEN: = ? = 434 nm = 4.34  10 -7 m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  10 -7 m = 6.91  10 14 Hz zEX: Find the frequency of a photon with a wavelength of 434 nm.

11. C. Quantum Theory z Photoelectric effect: emission of electrons from a metal when light shines on the metal z Hmm… (For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum – why did light have to be of a minimum frequency?)

12. C. 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

13. C. Quantum Theory zPlanck (1900) vs. Classical TheoryQuantum Theory

14. C. Quantum Theory zEinstein (1905) yObserved - photoelectric effect

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

16. C. Quantum Theory GIVEN: E = ? = 4.57  10 14 Hz h = 6.6262  10 -34 J·s WORK : E = h E = ( 6.6262  10 -34 J·s ) ( 4.57  10 14 Hz ) E = 3.03  10 -19 J zEX: Find the energy of a red photon with a frequency of 4.57  10 14 Hz.

17. C. 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

18. 6.3. Bohr Model of the Atom Ch.6-

19. Excited and Ground State zGround state: lowest energy state of an atom zExcited state: an atom has a higher potential energy than it had in its ground state zWhen an excited atom returns to its ground state, it gives off the energy it gained as EM radiation

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

21. B. Bohr Model z 2) e - exist only in orbits with specific amounts of energy called energy levels y When e- are in these orbitals, they have fixed energy y Energy of e- are higher when they are further from the nucleus

22. B. Bohr Model zTherefore…Bohr model leads us to conclude that: ye - can only gain or lose certain amounts of energy yonly certain photons are produced

23. B. 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, visible, and UV lines for the H atom

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

25. III. Wave Behavior of Matter Ch. 6 - Electrons in Atoms

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

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

28. A. Electrons as Waves zDiffraction: (def) bending of a wave as it passes by the edge of an object z Interference: (def) when waves overlap (causes reduction and increase in energy in some areas of waves)

29. 6.5: Quantum Model Chapter 6

30. A. Quantum Mechanics zHeisenberg Uncertainty Principle yImpossible to know both the velocity and position of an electron

31. A. Quantum Mechanics zSchrödinger Wave Equation (1926) yfinite # of solutions  quantized energy levels ydefines probability of finding an e -

32. B. Quantum Mechanics z Schrodinger wave equation and Heisenberg Uncertainty Principle laid foundation for modern quantum theory zQuantum theory: (def) describes mathematically the wave properties of e- and other very small particles

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

34. C. Quantum Numbers UPPER LEVEL zFour Quantum Numbers: ySpecify the “address” of each electron in an atom

35. C. Quantum Numbers 1. Principal Quantum Number ( n ) y Main energy level ySize of the orbital yn 2 = # of orbitals in the energy level

36. C. Quantum Numbers s p d f 2. Angular Momentum Quantum # ( l ) yEnergy sublevel yShape of the orbital (# of possible shapes equal to n) y values from 0 to n-1

37. C. Quantum Numbers If l equals…Then orbital shape is… 0s 1p 2d 3f Principle quantum # followed by letter of sublevel designates an atomic orbital

38. C. Quantum Numbers 3. Magnetic Quantum Number ( m l ) yOrientation of orbital  Specifies the exact orbital within each sublevel

39. C. Quantum Numbers z Values for m l : m = - l … 0… + l

40. C. Quantum Numbers pxpx pypy pzpz

41. zOrbitals combine to form a spherical shape. 2s 2p z 2p y 2p x

42. C. Quantum Numbers 4. Spin Quantum Number ( m s ) yElectron spin  +½ or -½ yAn orbital can hold 2 electrons that spin in opposite directions.

43. C. Quantum Numbers 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron zPauli Exclusion Principle yNo two electrons in an atom can have the same 4 quantum numbers. yEach e - has a unique “address”:

44. C. Quantum Numbers zn=# of sublevels per level zn 2 =# of orbitals per level zSublevel sets: 1 s, 3 p, 5 d, 7 f

45. Wrap-Up Quantum #SymbolWhat it describes Possible values Principle quantum # n main E level, size of orbital n = positive whole integers Angular Momentum Quantum # l sublevels and their shapes 0 to (n-1) Magnetic Quantum # m l orientation of orbital - l … 0 … + l Spin Quantum # m s electron spin+1/2 or -1/2

46. Feeling overwhelmed? Read Section 4-2!