Development of Atomic Models

1. Development of Atomic Models

2. Democritus • Greek philosopher • 400 BC

3. Question • Is there a limit to the number of times matter could be divided?

4. Democritus Theory • Eventually, you would reach a piece that was “indivisible” • Named this smallest piece of matter “atomos,” meaning “not to be cut.”

5. Atomos • Small, hard particles. • Differ in shape and size for each substance

6. Aristotle and Plato • All matter made up of combination of earth,fire, air and water. Aristotle

7. The Four Elements?? • This concept influenced early chemists called alchemists.

8. Buried in History “Atomos”theory was ignored and forgotten for more than 2000 years!

9. John Dalton (early 1800’s) • Performed careful scientific experiments. • Coined the term “atom”.

10. Dalton’s Atomic Theory • Matter is made of tiny indivisible particles called atoms. • Atoms of an element are alike, and different from atoms of other elements.

11. Dalton’s Atomic Theory • Compounds are atoms of different elements combined in fixed proportions. • Chemical reactions involverearrangement of atoms. • Atoms cannot be created or destroyed, but are conserved.

12. Pages from Dalton’s Journal

13. Hard Spheres Dalton’s model is called the “Hard Spheres Model”

14. JJ Thomson (1897)

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

16. Cathode Rays were negatively charged

17. Cathode Rays were particles

18. http://youtu.be/XU8nMKkzbT8 • http://youtu.be/Z61zCaAFky4 • http://youtu.be/IdTxGJjA4Jw

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

20. What about the Positive? • But…matter is neutral. • Must be a positive charge in the atom to balance the negative.

21. Plum Pudding Model • Positively charged sphere with with negatively charged particles scattered throughout.

22. Yummy…

23. Ernest Rutherford (1908) • Physicist who worked with the new field of radioactive emissions.

24. Different Types of Radiation • Used a magnetic field to determine there were three types of radiation. • Alpha (α) • Beta (β) • Gamma (γ)

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

26. Gold Foil Experiment • Shot alpha particles, at a very thin piece of gold foil. • These particles have a positive charge • Fluorescent screen shows where the particles went.

27. Observation: Almost all alpha particles passed straight through the gold foil. Conclusion: Most of the atom’s volume is empty space.

28. Observation: A few alpha particles were deflected at an angle or bounced back. Conclusion: Atoms have a very small, dense positively charged nucleus.

30. Nucleus is extremely small compared to the size of the atom as a whole. • Deflections happened rarely (1/8000).

31. The Nuclear Model Rutherford’s Model is called the “Nuclear Model”

32. Comparison to Thomson • Positively charge contained in nucleus. • Negatively particles scattered outside nucleus. • Not dispursed evenly.

33. http://chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHERFD.html • http://chemmovies.unl.edu/ChemAnime/RUTHERFD/RUTHERFD.html • http://youtu.be/wzALbzTdnc8 • http://youtu.be/XBqHkraf8iE

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

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

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

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

38. 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?

39. 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

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

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

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

43. 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”)

44. Spectral Lines are Different for Each Element

45. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swfhttp://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

46. What Causes Spectral Lines? Jumping Electrons!!

47. 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

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

49. Jumping Electrons • They quickly “fall back down” to the ground state. (stable) • They emit a photon of energy that corresponds to how far they jumped.

50. This photon of energy is seen as a spectral line! • Each spectral line corresponds to a specific photon of energy that is released.