Development of Atomic Models

1. Development of Atomic Models

2. Democritus • Greek philosopher (400 BC) • Matter can’t be divided forever • Eventually, a piece is “indivisible” • “Atomos,” =“not to be cut”

3. Dalton’s Atomic Theory (early 1800’s) • Matter made of tiny indivisible particles called “atoms”. • Each element’s atoms are different.

4. Compounds form when atoms combine in fixed proportions. • Chemical reactions involverearrangement of atoms. • Atoms are conservedin chemical reactions.

5. Page from Dalton’s Journal

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

7. JJ Thomson (1897) • 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.

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

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

10. Thomson’s Conclusions • Negative “Cathode Ray” particles must come from within atoms. • Discovered first subatomic particle • Negative electron However, matter is neutral so a positive charge must exist to balance the negative!!

11. Thomson’s Theory: • Atoms are positively charged spheres with negatively charged particles scattered throughout. Called the “Plum Pudding Model” Brian Cox: Thompson and Discovery of Electron http://youtu.be/IdTxGJjA4Jw

12. Ernest Rutherford (1908) • Physicist working in new field of radioactivity.

13. Found 3 Types of Radiation • Used magnetic field to isolate three types of radiation by their charge. • Alpha (α) • Beta (β) • Gamma (γ)

14. Gold Foil Experiment • Shot positive alpha particles, at thin piece of gold foil. • Fluorescent screen shows where particles went.

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

16. Observation: Some 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

17. 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

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

19. Comparison to Thomson • Positive charge only in nucleus. • Negative particles scattered outside nucleus. • Charge not disbursed evenly. Simulator: https://phet.colorado.edu/sims/html/rutherford-scattering/latest/rutherford-scattering_en.html

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

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

22. Different energy levels can contain different numbers of electrons.

23. 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

24. Energy Level configurations of neutral atoms are below its symbol. • Shows # of electrons in each occupied energy level Ex: Bromine 2-8-18-7

25. Draw a Bohr Atom Ex: Fluorine Atom (F) • Protons = 9 • Neutrons = 10 • Electrons = 9 • How many energy levels do you draw? • How many electrons in each level?

26. Outermost Electrons Valence Electrons: • Electrons found in highest occupied energy level (furthest from nucleus) • Involved in forming bonds with other atoms • “Stable Octet” of valence is desirable • What Group on the Table is stable as atoms and doesn’t need to form bonds?

28. Noble Gases: (Group 18) • “Inert” and unreactive (have stable octet) • Ex: Argon 2-8-8, Neon 2-8 Note: Helium is a Noble Gas too! It only has 2 electrons but its one energy level is complete and stable

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

30. Looking for Trends • Where do we see a similar valence electron pattern on the Periodic Table? • Down a Group? (column) • Across a Period? (row)

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

33. 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

34. Draw a Bohr Ion • One or more electrons gets added or taken out of the outer energy level. Ex: Magnesium Ion (Mg+2) • Protons = 12 • Neutrons = 12 • Electrons = 10

35. Pause for Thought • What do you think happens to the size of an atom when it gains or loses electrons and becomes an ion?

36. Positive Ions (cations) are smaller Lose electron(s) Ca atom Ca+2 ion 2-8-8-2 2-8-8

37. Negative Ions (anions) are larger Gain electron(s) P atom P-3 ion 2-8-5 2-8-8

38. How Did Bohr Come Up With His Model? • Studied spectral lines emitted by various elements

39. What are Spectral Lines? • Energy absorbed by an atom causes it to emit a unique set of colored lines.

40. Spectral lines are different for each element • Used to identify elements present in a sample. (“Fingerprint”)

42. What Causes Spectral Lines? • Electrons jumping between energy levels • Electrons absorb energy then emit it https://phet.colorado.edu/en/simulation/legacy/discharge-lamps Video of Line Spectra of Hydrogen http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

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

44. Excited State Electrons absorb a photon (or “quanta”) of energy and “jump up” to a higher energy level. This is called “excited state”. (unstable)

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

46. 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

47. REMEMBER Absorb Energy Jump Up Emit Energy Fall Down

48. Recognizing Excited State • Count total electrons given 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

49. Modern Atomic Model • Sometimes called: • Charge Cloud Model • Wave Mechanical Model • Quantum Theory • Theory created to help describe spectral line signatures of multi-electron elements Crash Course: History of Atomic Theory (9:45) https://www.youtube.com/watch?v=thnDxFdkzZs

50. How is it Different from the Planetary Model? • We can’t tell exactly where an electron is!! Heisenberg’s Uncertainty Principle: We can’t know exact location and momentum of an electron at the same time.