Electromagnetic Spectrum in Atoms

1. Chapter 4 Arrangement of electrons in atoms

2. We are all familiar with light but what is “visible” is just a very, very small portion of the electromagnetic spectrum What colors make up the rainbow? Red, Orange, Yellow, Green, Blue, Indigo, Violet (ROYGBIV) Visible Light

3. The E-M Spectrum X-rays(Cancerous in large doses; small doses medical scanning) Ultraviolet ( not as harmful - sunburn; black lights) Radio Rays(TV, Radio, other communications) Microwaves(Cooking, communications) Gamma Rays(Very Harmful / Cancerous) Infrared(Heat, communication)

4. The Development of a New Atomic Model • 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

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

6. Gamma Rays AM radio Television channels FM radio Short wave radio Radar Microwave Radio Waves X- Rays infrared The Electromagnetic Spectrum Decreasing wavelength Increasing frequency Increasing photon energy V i s i b l e L i g h t UV Rays R O Y G B I V Red Orange Yellow Green Blue Indigo Violet

7. Electromagnetic Spectrum • Frequency & wavelength are inversely proportional c =  c: speed of light (3.00  108 m/s) : wavelength (m, nm, etc.) : frequency (Hz)

8. WORK:  = c   = 3.00  108 m/s 4.34  10-7 m Electromagnetic Spectrum • EX: Find the frequency of a photon with a wavelength of 434 nm. GIVEN:  = ?  = 434 nm = 4.34  10-7 m c = 3.00  108 m/s = 6.91  1014 Hz

9. So why is the electromagnetic spectrum so important to chemistry? • Why is the steel emitting light when it is heated? • We take it for granted that when things get hot they turn red then orange and finally white; but that isn’t good enough any more

10. Black Body Radiation Colors

11. Black Body Radiation Colors

12. So why is the electromagnetic spectrum so important to chemistry? • Incandescence is heat made visible – the process of turning heat energy into light energy. • Our usage of terms like "red hot," "white hot," and so on, is part of the color sequence black, red, orange, yellow, white, and bluish white, seen as an object is heated to successively higher temperatures.

13. So why is the electromagnetic spectrum so important to chemistry? • The light produced consists of photons emitted when atoms and molecules release part of their thermal vibration energy. • For increasing temperatures, the sequence of radiated colors is: black, red, orange, yellow-white, bluish-white.

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

15. Energy and Light • 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

16. Energy and Light • EX: Find the energy of a red photon with a frequency of 4.57  1014 Hz. GIVEN: E = ?  = 4.57  1014 Hz h =6.6262  10-34 J·s WORK: E = h E = (6.6262  10-34 J·s) (4.57  1014 Hz) E = 3.03  10-19 J

17. Bohr hypothesized that instead of haphazardly orbiting the nucleus, electrons had clearly defined orbits – very similar to the planetary orbits circling our sun His model is (cleverly) named the Planetary Model Niels Bohr and the Bohr model of the atom

18. Niels Bohr Bohr Model (1913)

19. Bohr’s Proof • Bohr said this: If you assume that the electrons have clearly defined orbits that are congruent to the energy levels…

20. Bohr’s Proof • … then when an electron gets “excited” it jumps to a higher energy level. When it “relaxes” it emits a certain wavelength of light. • Bohr showed the energy of an electron in an atom is quantized, which means it has a particular numerical value, not some arbitrary number.

21. Excitation of Hydrogen Atoms Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 328

22. Return to Ground State

23. Bohr’s Proof 1 1 1 - = 1.097373 x 107 m-1 nr2 ne2 λ n=7 n=6 n=5 n=4 Paschen Series (ir) n=3 Balmer Series (vis and uv) n=2 Lyman Series (uv) n=1

24. 410 nm 434 nm 486 nm 656 nm 1 nm = 1 x 10-9 m = “a billionth of a meter” Emission Spectrum of an Element 1 nm = 1 x 10-9 m = “a billionth of a meter”

25. Continuous and Line Spectra

26. Hydrogen to Steel • If Hydrogen emits 4 distinct wavelengths of light when its one electron is excited what can we extrapolate to that of steel which is made mostly of iron? • http://jersey.uoregon.edu/vlab/elements/Elements.html

27. Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. methane gas wooden splint sodium ion copper ion strontium ion calcium ion Include link to web page http://www.unit5.org/christjs/flame%20tests.htm

28. Fireworks

29. Composition of Fireworks • Gunpowder • Sulfur, charcoal, potassium nitrate (saltpeter) • Salts (to give color) • Red = lithium • Green = copper

30. Good News Bad News • Good News • Bohr’s Model works and moves us along in the development of the Atomic Theory • End of this little unit • Bad (Frustrating) News • Lots of Math • Everything I taught you only works for Hydrogen and therefore is completely wrong and obsolete.

31. Check for Understanding c= λν E=hν c=3.0 x108 m/s h=6.626 x 10-34 J s • What is the frequency of a radar photon with an energy of 7.2 x 10-24 J? • What is the frequency of light having a wavelength of 6.20x10-7m?

32. - e + - e + e - e + + + + e - + e e - e + e + e Models of the Atom Greek model (400 B.C.) Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125

33. Your Current View of the Atom nucleus electrons

34. Again… so why is it so important to chemistry? • Einstein (1905) • Observed - photoelectric effect

35. Again… so why is it so important to chemistry? • Einstein (1905) • Concluded - light has properties of both waves and particles “wave-particle duality” • Photon - particle of light that carries a quantum of energy

36. Quantum Mechanical Model Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

37. Modern View • The atom is mostly empty space • Two regions • Nucleus • protons and neutrons • Electron cloud • region where you might find an electron • Also called the electron cloud model

38. Excited state e- e- Modern View of Atom Electrons can only be at specific energy levels, NOT between levels. Ground state

39. - e + - e + e - e + + + + e - + e e - e + e + e Models of the Atom Greek model (400 B.C.) Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Charge-cloud model (present) Bohr’s model (1913) Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125

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

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

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

43. Quantum Theory • quantum theory- • Describes mathematically the wave properties of electrons and other small particles • orbital- a region of an atom in which there is a high probability of finding electrons • Today’s atomic model predicts quantized, or particular energy levels for electrons. • does not describe the exact path or location electrons take or can be found around the nucleus • concerned with the probability, or likelihood, of finding an electron in a certain position • Two electrons can occupy each orbital, also called an electron cloud.

44. UPPER LEVEL Quantum Numbers • Four Quantum Numbers: • Specify the “address” or “seat” of each electron in an atom Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

45. Quantum Numbers 1. Principal Quantum Number( n) • Energy level (ladder rungs) • Size of the orbital • Positive integer 1s 2s 3s

46. Quantum Numbers 1. Principal Quantum Number • > number, further away from the nucleus • 1- right next to the nucleus • 3- further away from nucleus • > number, higher the energy level • n = 2 greater energy level than n = 1 • these electrons have more energy than electrons in the n = 1 level 1s 2s 3s

47. s p d f Quantum Numbers 2. Angular Momentum Quantum #( l) • Energy sublevel (orbital) • Shape of the orbital • Often represented by letters than numbers Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

48. Quantum Numbers y y y z z z x x x px pz py

49. d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 336

50. 2s 2px 2py 2pz Quantum Numbers • Orbitals combine to form a spherical shape. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem