Chapter 7: Atomic Structure & Periodicity

1. Chapter 7: Atomic Structure & Periodicity • What is the origin of color in matter? • Demo: flame tests • What does this have to do withthe atom? • Why are atomic properties periodic? fireworks: 07m01an1 CHMVID14, 07z01an1

2. Development of Modern Atomic Theory • Theelectron

3. Development of Modern Atomic Theory • Atomicmodel g

4. Development of Modern Atomic Theory • Quantum theory

5. The Wave Nature of Light • visible light radio waves Figure 6.2microwavesX-raysinfraredultravioletg rays • What do all these types of radiation have in common? 07m03an1

6. 7.1 Electromagnetic Radiation • Radiation carries energy through space. • All forms of electromagnetic radiation exhibit wave properties. • Velocity, Wavelength, Amplitude, and Frequency of Light Velocity, c = 2.998 x 108 m/s in vacuum l, wavelength in units of m u, frequency in units of s-1 or Hz amplitude relates to light intensity

7. Wavelength and frequency Figure 6.3are relatedinversely l u = c low-frequency wave high-frequency wave Types of Electromagnetic Radiation

9. Electromagnetic Radiation • Practice Problems: • What is the wavelength of light with a frequency 5.89 x 105 Hz? • What is the frequency of blue light with a wavelength of 484 nm?

10. Interaction of Matter and Light • Some radiation sources may emit only a single wavelength (e.g., a laser); called monochromatic • Radiation that spans a whole array of wavelengths is called continuous • Continuous radiation, such as visible light, can be separated into a spectrum, or sequence of colors by a prism or a diffraction grating. • Demo: colors on CD

11. Visible Spectrum • continuous spectrum chm3d18s

12. Visible Spectrum • Not all radiation sources produce a continuous spectrum. Line spectra are observed for elements, such as hydrogen • Emission Figure 6.7spectrum • Demo: gas discharge tubes • Why lines? chm3d19s

13. Visible Spectrum • Absorptionspectrum Figure 6.9

14. Visible Spectrum • Comparisonof hydrogen absorption and emission spectra • http://astro.u-strasbg.fr/~koppen/discharge/

15. Spectra Discussion Questions • Describe the relationship between the color change of the filament with temperature, and the colors in the observed spectrum of white light. • Order the colors of the rainbow according to the amount of energy needed to produce them, from lowest to highest

16. Spectra Discussion Questions • Based on the spectrum of the fluorescent lamp, what is the main difference between the radiation emitted by a heated solid (filament) and the radiation emitted by an excited atomic gas (hydrogen gas in the fluorescent tube)? • What can you say about the energy spread of the radiation emitted by an atomic gas?

17. What this means • Only certain energy states are allowed for the hydrogen atom. • Can only give off or absorb certain energies. • Use DE = hn = hc / l • Energy in the in the atom is quantized.

18. 7.2 Quantized Energy and Photons • Planck - studied radiation from heated matter and proposed that vibrating atoms can only have certain energies. • Planck found the energy came in chunks with size hn DE = h u or DE = nhn where n is an integer. and h is Planck’s constant h = 6.626 x 10-34 J s • these packets of hn are called quantum chm3d20s

19. 7.2 Quantized Energy and Photons • Einstein - studied the emission of electrons from metals struck by light (photoelectric effect) and proposed that light must have some particle properties and that these “particles”, or photons, are quantized (can have only certain energy values: E = h u)

20. Practice • The laser light of a CD is 7.80 x 102 m. What is the frequency of this light? • What is the energy of a photon of this light? • What is the energy of a mole of these photons?

21. What is the wavelength? • of an electron with a mass of 9.11 x 10-31kg traveling at 1.0 x 107m/s? • of a softball with a mass of 0.10 kg moving at 125 mi/hr?

22. 7.4 Bohr's Model • Rutherford showed that the atom contains a positively charged nucleus surrounded by electrons • Bohr wondered about what the electrons are doing in the atom. If they are stationary, they should fall into the nucleus, so they must be moving at tremendous speeds. But then why isn’t light emitted continuously and then why doesn’t the electron “run down” or lose its energy?

23. Niels Bohr • Niels Bohr developed the quantum model of the hydrogen atom. • He said the atom was like a solar system. • The electrons were attracted to the nucleus because of opposite charges. • Didn’t fall in to the nucleus because it was moving around.

24. The Bohr Model • Evidence shows only certain energies were emitted. • Only specific energy levels are allowed. • Transitions between energy levels absorb or emit light

25. The Bohr Model • For each energy level the energy can be found using: E = -2.178 x 10-18J (Z2 / n2 ) • n is the energy level • Z is the nuclear charge, which is +1 for hydrogen. • n = 1 is called the ground state • when the electron is removed, n = ¥ • E = 0

26. We are worried about the change • When the electron moves from one energy level to another. DE = Efinal - Einitial DE = -2.178 x 10-18J Z2 (1/ nf2 - 1/ ni2)

28. Examples 1.Calculate the energy need to move an electron from its first to the third energy level. 2.Calculate the energy released when an electron moves from n= 4 to n=2 in a hydrogen atom. 3.Calculate the energy released when an electron moves from n= 5 to n=3 in a He+1 ion

29. Hydrogen Spectrum • Other series of lines are observed in other ranges of wavelength (UV, IR) • All are summarized by the Rydberg equation, l-1 = R(n1-2 - n2-2), where the n1 and n2 are called quantum numbers • Other elements gave other sets of lines • How can we explain this behavior of atoms when they are heated or radiated?

30. Bohr's Model • In summary, Bohr introduced the idea that the energies of electrons in atoms were quantized -- allowed only certain values, characterized by the principal quantum number, n • Invented idea of orbits • Unfortunately, this idea only worked for atoms and ions that had but a single electron

31. 7.2-5 Wave nature of light & matter • Demo: wave interferenceLook at a bright light through a slit formed by holding your thumbs close together. This phenomenon is used to study light and matter.

32. Diffraction of Light (7.2) • When light passes through, or reflects off, a series of thinly spaced line, it creates a rainbow effect • because the waves interfere with each other.

33. Exploring Interference Behavior • http://www.colorado.edu/physics/2000/schroedinger/index.html

34. A wave moves toward a slit.

35. Comes out as a curve

36. with two holes

37. Two Curves with two holes

38. Two Curves with two holes Interfere with each other

39. Two Curves with two holes Interfere with each other crests add up

40. Several waves

41. Several waves Several Curves

42. Several waves Several waves Several Curves Interference Pattern

43. What will an electron do? • It has mass, so it is matter. • A particle can only go through one hole. • A wave through both holes. • An electron does go though both, and makes an interference pattern. • It behaves like a wave. • Other matter has wavelengths, but too short to notice.

44. Diffraction of Electrons and X-Rays • Davisson and Germer - found that electrons can be diffracted by solids just like X-rays • wave-particleduality

45. Diffraction of Electrons by Ice Just as light is diffracted by the grooves on a CD, a beam of electrons is diffracted by the planes of atoms in a crystal of ice. Debbie Leedy

46. Double slit e- diffraction patterns • http://www.xys.org/sci_forum/db/8/156.html

47. Meaning for Bohr Model? • Doesn’t really work. • Only works for hydrogen atoms. • Electrons don’t move in circles. • The quantization (shelf concept) of energy seems right, but not because electrons are circling the nucleus like planets.

48. 7.5 The Wave Behavior of Matter – Quantum Mechanical Model of the Atom • Does an electron behave like a particle or like a wave? Radiation appeared to behave sometimes like a wave and sometimes like a particle. What about electrons? • de Broglie - suggested that particles can have wavelike properties • wavelength is associated with particles; important only for atoms and subatomic particles

49. 7.5 The Wave Behavior of Matter – Quantum Mechanical Model of the Atom • A totally new approach. • De Broglie said matter could be like a wave. • De Broglie said they were like standing waves. • (similar to the vibrations of a stringed instrument.)

50. See figure 7.10