Notes: Ch. 3 Atom Models, Electrons and Light

1. Notes: Ch. 3 Atom Models, Electrons and Light p. 90-94

2. Electron Location---Atom Models • atoms are too small to be seen, so scientists have developed models that show their properties and structure • these models have evolved over time

3. Thomson’s “Plum-pudding” Model (see Fig. 7 p. 81) • suggested negatively charged e- (electron(s) are embedded in a positively charged ball of matter • disproved by Rutherford in 1909

4. Rutherford’s “Nuclear” Model (see Fig. 15 p. 90) • suggested that e- orbit around the nucleus • discarded because it couldn’t explain why e- do not crash into the nucleus due to the attraction of the negatively charged e- with the positively charged protons

5. Bohr’s “Energy Level” Model Suggested: • e- can only be certain distances from the nucleus • each distancecorresponds to a certain quantity of energy that the e- possesses

6. Bohr Model (cont.) • if a given e- is as close to the nucleus as it can be, it is in its lowest energy level (state) • different energy levels exist; the farther from the nucleus an e- is, the higher the energy level it is in • the difference in energy between different energy levels is called quantum energy

7. Bohr Model (cont.) • gain or loss of energy by an e- can cause it to change energy levels • e- can be in only one energy level or another, not in-between levels (see worksheet: Electron Cloud Energy Levels)

8. Modern Model • suggest e- are located in orbitals • also called quantum model

9. Electrons act Like Both Particles and Waves---Electrons are Particles • Thomson showed that e- behave like particles that have mass • the extremely small mass of the e- in the cathode tube was still enough to turn the paddle wheel

10. Electrons act Like Both Particles and Waves---Electrons are Waves • Louis De Broglie (1924) noticed the e- in Bohr’s model behaved like waves • like waves, e- could have only certain frequencies • these frequencies correspond to certain levels of energy; for e- this corresponds to the specific energy levels in which the e- are found

11. Electrons are Waves (cont.) • further experiments showed that e- beams can interfere with each other like waves can

12. Significance of Dual-Nature (Particle/Wave) • modern model (quantum atom model) accounts for this dual nature by suggesting e-s are located in orbitals • orbital-(def) a region in the atom where there is a high probability of finding electrons

13. Significance (cont.) • each orbital region corresponds to a specific energy level • orbitals are regions where e- are “likely” to be found • also called electron cloudsbec. they do not have sharply defined boundaries

14. Electrons and Light Light • by 1900, light was considered to be moving waves of energy with specific frequencies, speeds and wavelengths

15. Frequency • frequency-(def) the number of wave cycles completed in one second • answers the question: “how often?”

16. Wavelength • wavelength-(def) the distance between 2 consecutive peaks or troughs of a wave • answers the question: “how long?”

17. Speed • speed-in empty space ( or a vacuum) all light travels at the same speed: 2.998 x 108 m/s • answers the question: “how fast?”

18. Electromagnetic Spectrum • wavelengths of light vary from 105 m (longest) to 10-13 m (shortest) • longer wavelengths (lower frequency) possess less energy; shorter wavelengths (higher frequency) possess greater energy • the broadrange of wavelengths of light make up the electromagneticspectrum

19. EM Spectrum (cont.) • electromagnetic spectrum-(def) all the frequencies or wavelengths of electromagnetic radiation

20. ROY-G-BIV • only a smallportion of the em spectrum is visible; values of 700 nm (red light) to 400 nm (violet light) can be seen

21. Figure 18 p. 92

22. EM Spectrum

23. Photoelectric Effect • proposed by Einstein in 1905 to explain the release of electrons when light strikes metal

24. Einstein’s Observations • he noticed that for each given metal, a specific frequency of light was needed; below the necessary frequency no e- could be removed no matter how long the light was applied

25. Einstein’s Observations • if light was wave-like, then simply increasing exposure time should have been enough to provide the needed energy for e- release; it wasn’t

26. Einstein’s Explanation • Einstein believed light was also particle-like; to remove an e-, a particle of light must possess at least a minimum energy (minimum frequency)

27. Einstein’s Explanation (cont.) • a light particle is called a photon • Einstein proposed light has properties of both waves and particles

28. Light Emission • when a high voltage current is passed through a tube of hydrogen gas at low pressure, lavender light is seen • passing this light through a prism shows this light to be made of a few other colors

29. Light Emission (cont.) • this spectrum of colors is called a line-emission spectrum • these lines are produced by the hydrogen atom’s e- as they move to lower energy levels (producing light)

30. Hydrogen Line-Emission

31. Hydrogen Line-Emission (cont.)

32. Line-Emission (cont.)

33. Figure 20 p. 94 Hydrogen atoms can only move bet. certain energy states (shown as n=1 to n=7). When dropping to a lower state an electron emits a specific wavelength of light; only some are in the visible spectrum.

34. Various Line-Emission Spectra --Experiments with other gaseous elements showed that each element has a unique line-emission spectrum with a different pattern of colors

35. Light provides Info About Electrons • normally an e- is in a state of lowest possible energy or “ground state”; if an e- gainsenergy it moves up to an “excited state” • ground state-(def) the lowest energy state of a quantized system • excited state-(def) state in which an atom has more energy than it does in its ground state

36. Energy Must be Absorbed for an e- to Move to a Higher Energy Level • as the e- moves back to a lower level it releases this energy in the form of a photon • the energy of the photon corresponds to the change in energy the atom experiences as the e- moves from an excited state to a lower state

37. Energy Must be Absorbed for an e- to Move to a Higher Energy Level • as the e- moves back to a lower level it releases this energy in the form of a photon

38. Energy Must be Absorbed for an e- to Move to a Higher Energy Level (cont.) • the energy of the photon corresponds to the change in energy the atom experiences as the e- moves from an excited state to a lower state

39. Identify each Atom Model as: 1) Bohr’s, 2) Rutherford’s, 3) Dalton’s, 4)Thomson’s or 5) The Quantum Model

40. Flame Test