1. 1 Modern Atomic Theory

2. 2 In the Rutherford model electrons traveled about the nucleus in an orbit. The Problem with Rutherford Scientists know that just like a orbiting satellite eventually crashes to ground, an electron will slow down and crash into the nucleus. Something else must be going on.

3. 3 #1 Electromagnetic Spectrum increasing frequency increasing wavelength High frequency Higher energy (more wavelengths) Low frequency Low energy (fewer wavelengths) Longer wavelengths (orange and red light) ShortWavelength (Purple and blue light) To understand the next model we must review what we know about light.

4. 4 #2 Continuous spectrum of White Light All the colors (frequencies) of light - infinite White light has all the different colors of light waves

5. 5 #3 Bright line spectrum of Excited Hydrogen Gas Only four colors (frequencies) of light This spectrum has holes! When electricity is forced through hydrogen gas (think “neon” light) only certain colors of light are emitted.

6. 6 #4 Line Spectra of Other Elements Each element has its own distinct lines – Like a fingerprint – can be used to identify the elements in an unknown sample!

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8. 8 Excited Gases in neon lights And Aurora – the northern lights These are examples of excited state gases!

9. 9 Flame tests: Different ions emit unique colors of light when samples are held in the flame of a Bunsen burner Useful for identification SodiumLithium

10. 10 How can the bright line spectrum be explained? Click here For video

11. 11 Excited electron states Electrons orbit in "shells” close to the nucleus: “ground” state Electrons absorb energy and move to higher shells become “excited” How can the bright line spectrum be explained?

12. 12 Excited = unstable electrons fall back to ground state Electrons release energy as light Color of light depends on magnitude of the drop Ex: Small drop = red light Larger drop = blue How can the bright line spectrum be explained?

13. 13 Many different drops back to ground state creates the atomic spectra. The colors corresponds to the difference in energy between the two levels But why not all the colors? Why only certain colors? How can the bright line spectrum be explained? SimulationSimulation: How atoms emit light Simulation

14. 14 “Quantum” Theory: Max Planck (1900) Energy is “quantized” - Light energy absorbed and emitted as discrete packets of energy called “photons” Energy of electrons is not continuous, it is gained and lost in discrete units called quanta (sing. Quantum)

15. 15 Niels Bohr model of the atom (1913) Electron is restricted to QUANTIZED energy state Electrons can only orbit in certain discrete Energy Levels (shells) depending on the amount of energy they contain Higher energy levels = electrons with higher energy Normally electrons sit in lowest energy position available, close to nucleus = “Ground state”

16. 16 If Rutherford was right electrons can exist in any orbit: All colors would be emitted If Bohr is right electrons can only sit on energy levels: Only specific colors emitted

17. 17 Learning Check 1. 1.What is the source of the bright line spectrum of an element? 2. 2.How does the ground state and excited state of an atom compare in terms of energy? 3. 3.What determines the color of the light in an individual bright line? 4. 4.In what way is an atomic spectrum (bright line spectrum) like a fingerprint? How is it useful?

18. 18 Review:

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20. 20 Bohr Diagrams May show nucleus particles Shows energy levels and electrons – paired or unpaired in orbitals Ex: Al electron configuration  2-8-3 2 electrons in 1 st level 8 electrons in 2 nd level 3 electrons in 3 rd level

21. 21 Shows energy levels and electrons in orbitals Total electrons = 8 First 2 electrons go in as orbital pair in 1st Energy level Next two electrons 2nd level as a pair The next three enter separate orbitals The remaining 1 electron combines with one of the Three unpaired electrons Bohr Diagram for oxygen Electron configuration : 2-6 Orbital pair 8P +

22. 22 Try these: Draw Bohr diagrams for the following atoms: 1.Be2. F 3. Na4. P 2-2 2-7 2-8-12-8-5

23. 23 Arrangement of electrons in the atom = Electron configuration Rules: Each energy level can hold a certain maximum number of electrons 1st holds 2 e- 2nd holds 8 e- 3rd holds 18 [ 2n 2 ] Ex: Carbon: 2 – 4 2 electrons in 1 st energy level 4 electrons in 2 nd energy level 2(1) 2(1+3) 2(1+3+5)

24. 24 P 15 30.973 2 - 8 - 5 F 9 18.998 2 - 7 Electron configurations on the periodic table How many electrons occupy the 2 nd energy level of a phosphorus atom in the ground state?

25. 25 Electrons are divided between Kernel and valence electrons Aluminum Kernel = [Ne]: 2 - 8 + 3 valence electrons Chlorine 2 – 8 - 3 2 - 8 - 7 3 valence electrons 7 valence electrons Properties are connected to valence electrons: Elements in the same group have same # valence electrons

26. 26 An easier way: Lewis (Electron Dot) Structures For Elements and Ions – dots or x’s used to represent valence electrons Atoms hydrogen: H Magnesium: Mg e- configuration: 1 2-8-2 Oxygen: O Chlorine: Cl 2-6 2-8-7

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29. 29 Learning check: 1.Write the electron configuration for an atom of Ca in the ground state: 2.How many kernel and valence electrons does calcium have? 3.Which noble gas has a kernel like calcium? 4.Draw the Bohr diagram and Lewis electron dot symbols for Lithium, nitrogen and fluorine

30. 30Review

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32. 32 Excited state electron configurations Atom in ground state – electrons in lowest E levels = lowest energy state ex: LiCF 2-12-62-7 Become Excited: 1-22-5-11-7-1

33. 33 44 Which is an electron configuration for an atom of chlorine in the excited state? HDYK? (1) 2–8–7 (3) 2–8–6–1 (2) 2–8–8 (4) 2–8–7–1 1 As an electron in an atom moves from the ground state to the excited state, the electron (1) gains energy as it moves to a higher energy level (2) gains energy as it moves to a lower energy level (3) loses energy as it moves to a higher energy level (4) loses energy as it moves to a lower energy level Still has 17 electrons

34. 34 12 2-1 2-22-42-32-5 2-62-7 2-8 2-8-12-8-2 Periodic table is organized by electron arrangement 2-8-32-8-4 2-8-52-8-62-8-72-8-8 2-8-8-1Etc. Groups with same # valence electrons 1e 2e 3e4e5e 6e7e 8e Periods = # energy levels 1 energy level 2 energy levels 3 energy levels Etc.

35. 35 Why can’t the outer shell hold more than 8 electrons? K (its hard to pack more electrons into an already crowded space)

36. 36 Learning check 1. A neutral Sulfur atom is in the ground state Without looking directly at the electron configuration determine: a. the number of valence electrons b. the number of occupied energy levels c. the electron configuration 2. A neutral atom has the ground state electron configuration of 2-8-8. a. What group and period is it located in? HDYK? b. What properties will it have? Group 16 = 6 e - Period 3 = 3 levels 2-8-6 8 valence e - = group 18 3 levels = period 3 Noble gas = chemically inert gas

37. 37 Modern view of the atom Wave Mechanical model Electrons as standing waves

38. 38 Only standing waves are allowed

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40. 40 Arrangement of Electrons in Atoms Electrons in atoms are arranged as PRINCIPLE ENERGY LEVEL (shell): n SUBLEVEL (shape): SUBLEVEL (shape): ORBITAL (region): Orbital = Region where electron is most likely to be found Region of “highest probability”

41. 41 3 In the wave-mechanical model, an orbital is a region of space in an atom where there is (1) a high probability of finding an electron (2) a high probability of finding a neutron (3) a circular path in which electrons are found (4) a circular path in which neutrons are found

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43. 43 Periodic Table II: Electron arrangement and the table

44. 44 Honors: So why the weird shape? s p d f The sublevels don’t fill in order! s “Diagonal Rule”

45. 45 1s 1 1s 2 2s 1 2s 2 2p 1 p2p2 p3p3 p4p4 p6p6 p5p5 3s 1 s2s2 3p 1------------------------  p 6 4s 1 s 2 Shape of the table follows Filling of sublevels

46. 46 1s 1 1s 2 2s 1 2s 2 2p 1 p2p2 p3p3 p4p4 p6p6 p5p5 3s 1 s2s2 3p 1------------------------  p 6 4s 1 s 2 3d 1----electrons in lower E levels---------- 3d 10 4p 1—back outside---- P 6 5s 1----2 4d 1--------------------etc.--------------------------- 4d 10 5p 1---------------------------6 6s 1----2 4f 1---------electrons going two E levels inside---------------------- 4f 14 5d 1-------------------------------------------------- 5d 10 6p 1---------------------etc. 7s 1----2 5f 1---------------------etc.-------------------------------------------------------- 5f 14 6d 1-------------------------------------------------- 6d 10 etc.

47. 47 Why are d and f orbitals always in lower energy levels? d and f orbitals require LARGE amounts of energy due to electron repulsion (hard to pack more electrons into an already crowded space) It’s better (lower in energy) to skip a sublevel that requires a large amount of energy for higher level but lower energy K

48. 48 HHeNeArHg

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51. 51 Lithium barium strontium copper potassium sodium calcium