Electrons in Atoms

1. Electrons in Atoms Chapter 4

2. What were the early steps in the development of atomic theory? • John Dalton – Billiard Ball Theory. Atom was indivisible. • J.J. Thomson – Plum Pudding Model. Atom was composed of smaller particles, including negative electrons. • E. Rutherford – Nuclear Model.

3. Cathode Ray Tubes & Discovery of Electrons Memory Jogger source source Movie of cathode ray tube

4. Thomson’s Plum-Pudding Model The positive charge is evenly smeared out. The negative charge is in bits – like chips. source

5. Rutherford’s Experiment - 1911 source

6. Rutherford Model • Nucleus contains all positive charge & most of the mass. • Nucleus is very small in volume - only 1/10,000th of atomic diameter. • Electrons occupy most of the volume of the atom.

7. source Rutherford model was the nuclear atom. Rutherford did not speculate on how the electrons were arranged around the nucleus. source

8. Later Models • N. Bohr – Shell Model • Schrodinger – Wave Mechanics

9. - + - - - - +6 - - source

10. Atomic Emission Spectra • Put some gas in a glass tube and apply a voltage across the ends. Produce light. • The color of the light depends on the gas in the tube. • Every element produces its own unique color.

11. The emission spectrum of an element is the set of frequencies (or wavelengths) emitted.

12. Discharge Tubes (Neon Lights) This lamp contains hydrogen.

13. 1912 – no one could explain these. Called a Line Spectrum.

14. Atomic Emission Spectra: Line spectra

15. Why is the emission spectra useful? • We can use it to determine if a given element is present in a sample. (Identification!) • We can use it to learn fundamental information about atomic structure. • Neon lights

16. Problems with the Rutherford Model • Why doesn’t the electron spiral into the nucleus? • How are the electrons arranged? • Why do different elements exhibit different chemical behavior? • How are the atomic emission spectra produced?

17. Light

18. Electromagnetic Spectrum

19. Electromagnetic Radiation

20. Waves • For any wave: • velocity = wavelength X frequency =  • Wavelength has units of length. • Frequency has units of inverse time (time-1 or 1/time). No symbol for “cycles.”

21. Electromagnetic Radiation • All moves at the same speed – the speed of light! • Which is 3.0 X 108 m/s. Symbol = c.

22. Photoelectric Effect

23. Photoelectric Effect • Shine light on a clean metal surface in a vacuum. • Observe electrons ejected from the metal surface. • Measure how many electrons are ejected from the metal surface as a function of the light frequency. • Also, measure the kinetic energy of the electrons.

24. Results

25. Photoelectric Effect • For a given metal, no electrons were emitted if the frequency of the light was below a certain minimum. • Didn’t matter how intense the light was. • At the minimum frequency, increasing the intensity of the light increased the # of electrons emitted but not their K.E. • Increasing the frequency above the minimum increases the K.E. of the emitted electrons.

26. Stumped! • If light is a wave, any frequency should be able to eject an electron.

27. Einstein: dual nature of light • Sometimes, light acts like a particle. • E = h where h is Planck’s constant and  is the frequency of the light.

28. Quick light review • In the visible spectrum, • Blue light: shorter wavelength, higher frequency, more energy • Red light: longer wavelength, lower frequency, less energy

29. Interlude: Electromagnetic Radiation All electromagnetic radiation has a velocity of 3.00 X 108 m/sec c = 

30. Bohr - 1913 • Link the electrons in the atom to the emission of the photons. • Link the chemical behavior to the arrangement of the electrons in the atom.

31. Bohr Diagram • Shows all the electrons in orbits or shells about the nucleus. E3 n=3 E2 n=2 E1 n=1

32. Bohr vs. Lewis Diagram • Shows valence electrons • Symbol represents kernal • Kernal = nucleus + all inner shells • Shows all electrons • Arranged by energy level • Shows composition of nucleus

33. Bohr’s Model • Electrons travel only in specific orbits. • Each orbit has a definite energy. • The energy of the electron must match the energy of the orbit. So the electron is only allowed to have some energies, not any energy. NEW! Bohr assigned a quantum number, n, to each orbit.

34. Bohr’s Model • The orbit closest to nucleus is the smallest & has the lowest energy. It has n = 1. • Outer orbits hold more electrons than inner orbits. • The larger the orbit, the more energy associated with it. source

35. Bohr’s Model • Atoms emit radiation when an electron jumps from an outer orbit to an inner orbit. • Atoms absorb energy when an electron jumps from an inner orbit to an outer orbit. • Outer orbits determine atom’s chemical properties. source Bohr Applet

36. Bohr’s Model Orbits are NOT evenly spaced – they get closer together the farther they are from the nucleus. Orbits are 3-D not 2-D.

37. Bohr Diagram n=3 E3 = 23* E2 = 15* n=2 E1 = 5* n=1 What is the energy change of the electron if it moves from E1 to E3? It must absorb 18 units of energy. From E2 to E1? It must release 10 units of energy. * Made-up Numbers!!!

38. Bohr Model • Energy is absorbed when the electron moves to a higher orbit, farther from nucleus. Endothermic process. • Energy is released when the electron drops to a lower orbit, closer to nucleus. Exothermic process.

39. Emission & Absorption Spectra of Elements Each line in the bright line spectrum represents 1 electron jump. The blue lines have shorter wavelength and higher frequency, so they are light from “bigger” jumps. The red lines have longer wavelength and lower frequency, so they are light from “smaller” jumps.

40. Blue lines represent a bigger drop than red lines.

41. Hydrogen Atom

42. Flame Tests • Volume 2, CCA • Na • Sr • Cu • Animation

43. The energy levels get closer together away from the nucleus. Larger orbits can hold more electrons.

44. Emitted Light • The energy of the emitted light, E = h, matches thedifferencein energy between 2 levels. • Again, we don’t know the absolute energy of the energy levels, but we can observe how far apart they are from each other.

45. source source Tiger Graphic – potential energy Tiger Graphic – electron orbits

46. A ladder is often used as an analogy for the energy levels of an atom. But it’s a little bit different – How? Potential Energy

47. Max Capacity of Bohr Orbits

48. Ground State • Each electron is in the lowest energy orbit available. Lowest energy state of an atom. • The Bohr configurations in the reference table are ground state configurations.

49. Excited State • Many possible excited states for each atom. One or more electrons excited to a higher energy level. • We can give electron configurations for the excited states as well as the ground state. • You need to recognize excited state configurations.

50. Excited State Configurations • For the smaller elements, it’s easy. • First level holds 2. • Second level holds 8. • If upper levels fill before the first or second is full, it’s an excited state configuration.