1. The Electromagnetic Spectrum and the Model of the Atom Part I Chemistry – Mrs. Cameron

2. The Purpose of Science The purpose of science to make models that explain natural phenomena.

3. Scientific Models A model is the best possible explanation which accounts for all observed phenomenon and has predictability. 1) An unanswered question means change the model. 2) Predictability is the test of a good or “true” model.

4. Radioactivity and Light have been tools to discover the structure of atoms. Do you Remember? –Thompson’s Cathode Ray Experiment –Rutherford’s Gold foil Experiment

6. Thomson’s Plum-Pudding Model

7. Ernest Rutherford’s Gold Foil Experiment

8. Results of the Rutherford experiment (a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct (b) Actual results

9. Rutherford’s Nuclear Model of the atom Small, dense, positively charged nucleus Contains protons (+1 charge) Contains neutrons (no charge) Remainder of the atom is mostly empty space Contains electrons (-1 charge) in the “empty space”

10. Let’s Talk about LIGHT New Evidence: Continuous and Line Spectra

11. White Light Given off by objects heated to a very high temperature. When an object is heated, it first gives off a red glow then, as more energy is added it begins to glow “white hot.”

12. When white light is passed through a prism, the light refracts, or bends to display all of its component colors

13. The Wave Nature of Light All waves exhibit similar characteristics and properties: Crest – the top of a wave Trough – the bottom of the wave (on next slide) Origin – the center line through which the wave oscillates Wave crest or peak

14. The Wave Nature of Light Origin  Amplitude – distance from the origin to the crest or the origin to the trough of a wave. (Indicates intensity) Frequency ( ) – the number of waves that pass a given point in a given amount of time Wavelength (λ) - distance from crest to crest or trough to trough Speed (Velocity) – the amount of distance covered in a specified amount of time. ←trough Crest 

15. The Wave Nature of Light

16. Wave Behaviors Reflection – a wave strikes an object and bounces off Refraction – the bending of waves Diffraction – the bending of waves around an opening or around the edge of an object Interference – the ability of two or more waves to add together forming regions of large or small amplitude.

18. Wave Interference patterns

19. 19 Light Exhibits Interference Constructive interference: waves “in- phase” create waves of greater amplitude ( they add) Destructive interference: waves “out- of-phase” create waves of lower amplitude (they cancel out)

20. The Electromagnetic Spectrum Visible light is one type of electromagnetic radiation. The arrangement of electromagnetic waves by their wavelength is called the electromagnetic spectrum The spectrum ranges from high energy, shorter wavelengths of radiation to long wavelength, lower energy radiation. The visible portion of the spectrum is very small.

21. The Wave Nature of Light

22. Ionizing Radiation  Ionizing Radiation  onizing Radiation  Nonionizing Radiation 

24. X-rays

25. Units for Wave Characteristics Wavelength (λ “lambda”) = meters with a Greek prefix (nanometers, nm) Frequency ( “nu”) –cycles per second or Hertz (1/s, sec -1, Hz) Speed – meters per second (m/s)

26. Wave Calculations –All electromagnetic radiation travels at the same speed through the vacuum of space: –3.00 x 10 8 m/s (c) –Wavelength and frequency are inversely proportional: λ = c/ –Energy and frequency are directly proportional: E = h E = energy in joules, h = Plank’s constant 6.626  10 -34 J·s

27. The Dual Nature of Light Light behaves as both a wave and as a particle. Evidence: 1) The glowing of heated metals (first infra red, then visible as Tº  ) 2) The photoelectric effect

28. The Work of Max Planck (1858-1947) Able to predict the wavelengths of light changes with temperature Energy must be emitted in a “Quantized” way, or restricted to certain quantities. “Quantum” (singular) or “Quanta” (Plural) A quantum is a packet of energy VERY SMALL Related through Planck’s Constant h= 6.626  10 -34 J·s, E = h

29. The Photoelectric Effect When light is shined on a piece of metal, electrons are ejected from the metal Only light containing enough energy (of a certain wavelength and frequency) works

30. The Photoelectric Effect

31. Einstein and the Photoelectric effect Albert Einstein (1879-1955) proposes that light has a particle nature too. He relates Planck’s idea of “quantized energy” to light. Light energy quanta = photons Photons transfer energy to electrons when light strikes the metal. Photons must be of sufficient energy for this to occur.

32. What do you see? Depending on how you look at this it can be an old lady or a young lady turning her head. The picture has a dual nature

33. Line Spectra of Elements A line spectrum is a spectrum that contains only certain colors, or wavelengths of light. The rainbow is a continuous spectrum Elements emit line spectra when they are vaporized in an intense flame or with electricity.

34. Continuous spectrum of white light Line spectra of elements

35. Line Spectra and the Quantization of Energy

36. Neils Bohr (1885-1962) Planetary Model of the Atom Combines Rutherford’s nuclear model and Planck’s quanta of energy. Electrons must have certain energy levels in which they travel. (Energy Level Postulate) Energy of Electrons must be quantized to explain line spectra. (Transitions Between Energy Levels)

37. Bohr Model Ground state = energy level closest to the nucleus Excited state = energy levels farther away from the nucleus Energy levels represented by quantum numbers ex. n = 1, n= 2, n = 3 etc.

39. Emission Spectrum of Hydrogen http://upload.wikimedia.org/wikipedia/com mons/4/4c/Emission_spectrum-H.png In general, the line spectrum of an element is rather complicated The line spectrum of hydrogen, with a single electron, is the simplest

40. Atomic line spectra tell us that when an excited atom loses energy, not just any arbitrary amount can be lost This is possible if the electron is restricted to certain energy levels The energy of the electron is said to be quantized The Evidence of Line Sprectra:

41. Radiation absorbed – electron moves from the ground state to an excited state. Can’t maintain this higher energy level Electron falls back down to the ground state Radiation is emitted as it returns to the ground state The energy emitted or absorbed =  energy levels

42. Emission of Light During Movement of Electrons http://iws.collin.edu/biopage/faculty/m cculloch/1406/outlines/chapter%2010 /Ma10-8b.JPG

44. J.J. Balmer equation for visible spectrum of Hydrogen v = 1/ = 1.097 x 10 7 /m (1/n 1 2 - 1/n 2 2 ) Which led to….. The Rydberg equation can be used to calculate all the spectral lines of hydrogen

45. Rydberg Equation Used to calculate the wavelengths of all the spectral lines of hydrogen. Atomic spectra indicate that when an excited atom loses energy, the energy is in discrete amounts - or “quantized.” n 1 and n 2 are positive integers = Rydberg constant = 109,678 cm -1

46. Bohr proposed that the electrons moved around the nucleus is fixed paths or orbits much like the planets move around the sun The orbits, labeled with the integer n, have energy This equation allows the calculation of the energy of any orbit

47. Bohr was able to use this model to calculate the energies of the light given off by the hydrogen. E = h Unfortunately, the model became too mathematically complicated for any element with more electrons than hydrogen. Also, electrons not completely explained as particles

48. Limitations of the Bohr Model It cannot explain the spectra of atoms other than hydrogen. Electrons do not move about the nucleus in circular orbits. However - the model introduces two important ideas: The energy of an electron is quantized: electrons exist only in certain energy levels described by quantum numbers. Energy gain or loss is involved in moving an electron from one energy level to another.

49. Matter Waves Louis DeBroglie (1892-1897) –Electrons (matter) have a dual nature just like light. –Proved it mathematically –Evidence: diffraction patterns produced by beams of electrons

50. http://www.microscopy.ethz.ch/TEM_ED_e xamples.htm Electron Diffraction Patterns

51. E = mc 2 and E = h λ = c/ (solve for ) = c/ λ (Replace with c/ λ) E = h c/ λ (Both sides of the equation equal “E” so you can equate one to the other) hc/ λ = mc 2 (solve for λ) λ = hc = h mc 2 mc (replace the speed of light with the speed of the particle) λ = h ms You can calculate the wavelength of any object! Matter Wave Equations:

52. λ = h ms Why do we not see these waves? Because for relatively large objects their mass is too large and their speed too slow for the wavelengths to be observed. Remember, h = 6.626 x 10 -34 J · s Electrons: - small mass (9.109 x 10 -28 g) - travel REALLY fast so we can observe the wave behaviors

53. To be Continued………