Development of Atomic Models

1. Development of Atomic Models

2. Democritus • Greek philosopher • 400 BC • “Atomos” concept

3. Can matter can be divided forever? • Eventually, a piece would be “indivisible” • “Atomos,” meaning “not to be cut,” is smallest piece of matter

4. John Dalton (early 1800’s) • Coined the term “atom”.

5. Dalton’s Atomic Theory • Matter made of tiny indivisible particles called “atoms”. • Atoms of one element are alike, and different from atoms of other elements.

6. Page from Dalton’s Journal

7. Compounds form when different atoms combine in fixed proportions. • Chemical reactions involverearrangement of atoms. • Atoms can’t be created or destroyed, but are conservedin chemical reactions.

8. Dalton’s Atomic Theory called “Hard Spheres Model”

9. JJ Thomson (1897)

10. Thomson’ Experiments • Studied “cathode rays” (electric current) in a “Crooke’s Tube”. • Fluorescent screen, shows how cathode ray behaved in a magnetic field. Lets draw a typical Crooke’s Tube in our notes.

11. Cathode Rays were negatively charged Cathode Ray Tube and Magnet http://youtu.be/XU8nMKkzbT8 They bent toward (+) plate

12. Cathode Rays were particles They couldn’t pass through matter.

13. JJ is Awesome • Concluded the negative “cathode ray” particles came from within atoms. • Discovered first subatomic particle (electron).

14. What about the Positive? • But…matter is neutral. • Therefore: • A positive charge must exist to balance the negative.

15. Plum Pudding Model Atoms are positively charged spheres with negatively charged particles scattered throughout.

16. Yummy… Brian Cox: Thompson and Discovery of Electron http://youtu.be/IdTxGJjA4Jw

17. Ernest Rutherford (1908) • Physicist who worked in new field of radioactivity.

18. Found 3 Different Types of Radiation • Used magnetic field to isolate three types of radiation. • Alpha (α) • Beta (β) • Gamma (γ)

19. Charges of Radiation • The radiation had different charges. Identify the charge each type of radiation has.

20. Gold Foil Experiment • Shot alpha particles, at very thin piece of gold foil. • Alpha particles have a positive charge, and a mass of 4 amu • Fluorescent screen shows where the particles went. Rutherford Gold Foil http://www.youtube.com/watch?v=5pZj0u_XMbc

21. Observation: Most alpha particles passed straight through gold foil. Conclusion: Atom’s volume is mostly empty space.

22. Observation: A few alpha particles deflected at an angle or bounced back. Conclusion: Atoms have a very small, dense positively charged nucleus. http://www.kentchemistry.com/moviesfiles/Units/AtomicStructure/Rutherford3.htm

23. Nucleus is extremely small compared to the size of the atom as a whole. Deflections happened rarely (1/8000). Modern Example of Gold Foil Experiment in Action http://youtu.be/XBqHkraf8iE

24. The Nuclear Model Rutherford’s Model is called the “Nuclear Model” Brian Cox: Rutherford and the Nucleus http://youtu.be/wzALbzTdnc8

25. Comparison to Thomson • Positively charge only contained in nucleus. • Negatively particles scattered outside nucleus. • Charge is not disbursed evenly.

26. Niels Bohr (1913) • Came up with the “Planetary Model”

27. Bohr’s Theory • Electrons circle nucleus in specific energy levels or “shells”. • The higher the “energy level,” the higher the electron’s energy.

28. Energy Levels • Different energy levels can contain different numbers of electrons.

29. How many per level? • n = the number of the energy level 2n2 = maximum number of electrons an energy level can hold. Ex: Level 3 can hold 2(3)2 = 18 electrons

30. Draw a Bohr Atom • Ex: The Fluorine Atom (F) • Protons = 9 • Neutrons = 10 • Electrons = 9 • How many energy levels do you draw? • How many electrons in each level? Human Bohr Model http://www.youtube.com/watch?v=PLpZfJ4rGts

31. Draw a Bohr Ion • They only difference is that one or more electrons gets added or taken out of the outer energy level. • Ex: The Magnesium Ion (Mg+2) • Protons = 12 • Neutrons = 12 • Electrons = 10

32. (+) Ions (cations) (+) ions are smaller Lost electron(s)

33. (-) Ions (anions) (-) ions are larger Gained electron(s)

34. How Did Bohr Come Up With His Model? • Studied the spectral lines emitted by various elements (especially Hydrogen)

35. What are Spectral Lines? • Energy gets absorbed by an atom causing it to emit a unique set of colored lines. • Used to identify what elements are present in a sample. (elemental “Fingerprint”)

36. Spectral Lines are Different for Each Element

37. Answer: 1

38. What Causes Spectral Lines? Jumping Electrons!! Video of Line Spectra of Hydrogen http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

39. Jumping Electrons Electrons normally exist in the lowest energy level possible called the “ground state”. (stable) “Ground state” e- configurations are written on the periodic table for each element. Ex: Aluminum is 2-8-3 Calcium is 2-8-8-2

40. An Electron Gets “Excited” Electrons can absorb a photon (or “quanta”) of energy and “jump up” to a higher energy level farther from the nucleus. This is called the “excited state”. (unstable)

41. Jumping Electrons • They quickly “fall back down” to the ground state. (stable) • They emit a photon (or “quanta”) of energy that corresponds to how far they jumped. Spectral Lines http://www.youtube.com/watch?v=QI50GBUJ48s

42. This photon of energy is seen as a spectral line! • Each spectral line corresponds to a specific photon of energy that is released. Model Of Hydrogen Atom and Electrons Jumping http://www.upscale.utoronto.ca/PVB/Harrison/BohrModel/Flash/BohrModel.html

43. REMEMBER Absorb Energy Jump Up Emit Energy Fall Down

44. Excited vs Ground State • Periodic table lists ground state electron configurations for neutral atoms. • To recognize an “excited state” configuration, count the electrons and see if the configuration matches the one on the table. • Ex: 2-8-7-3 = 20 electrons • Calcium (atomic # 20) is 2-8-8-2 • So this must be showing one of the ways calcium could be in the excited state.

45. Valence Electrons • Electrons in highest occupied energy level. • Involved in forming bonds with other atoms. • Atoms are most stable when they obtain a “stable octet” of 8 valence electrons • Noble Gases: (Group 18) • Have stable octet already and are “inert” and unreactive • Ex: Argon 2-8-8, Neon 2-8

46. Valence Electrons • Look at the last number in the atom’s electron configuration to determine the number of valence electrons. • Ex: • Al 2-8-3 3 valence • Ca 2-8-8-2 2 valence • F 2-7 7 valence

47. Lewis Dot Diagrams • Shows the number of valence electrons an atom has as “dots” around the atom’s symbol. Phosphorus is 2-8-5

48. Kernel • Nucleus and non-valence electrons • Inner part of atom not involved directly in reactions • Ex: • Al 2-8-3 has 10 kernel electrons and 3 valence electrons

49. The Nature of Light • Study of light has provided important information about the structure of atoms. • Dual Nature of Light: • behaves as both waves and as particles (depending on what type of experiment is being performed.) • Speed of Light: all light waves travel at the same velocity • C = 3.0 x 108 meters/sec

50. Electromagnetic Spectrum • Spectral lines can come from all areas of the EM Spectrum. • Lines of visible colors make up only a small part of the spectrum.