2. Honors Chemistry Chapter 5 Electrons

3. “The more success the quantum theory has, the sillier it looks.” ~Albert Einstein, Nobel Prize in Physics, 1921. “Quantum mechanics: the dreams that stuff is made of.” ~unknown “If quantum mechanics has not yet profoundly shocked you, you have not yet understood it.” ~Niels Bohr, Nobel Prize in Physics, 1922.

4. Where are electrons located? Outside the nucleus How are they arranged?

6. Bohr model of electron placement  Called the “planetary” model  Electrons closer to the nucleus – lower “energy level”  Electrons farther away from the nucleus – higher “energy level”

7. CCCCalled the “shells” K, L, M, N 2222, 8, 18, 32 electrons SSSStudied the emission spectrum of Hydrogen SSSSpecific colors that are emitted (given off) when an atom releases energy

8. Quantum Staircase An electron in a stable orbit will have a specific, restricted (quantitized) energy: Niels Bohr

9. Max Planck Stated that the object (metal) emits energy in small, specific amounts called Quanta. Quantum is the minimum quantity of energy that can be lost or gained by an atom. Step ladder analogy

10. Albert Einstein  Took Planck’s idea a little further.  He introduced that electromagnetic radiation has a dual wave-particle nature. Light exhibits many wavelike properties Can also be thought of as a stream of particles

11. Two important concepts from Bohr: Electrons exist only in certain discrete energy levels Energy is involved in moving an electron from one level to another IN REALITY, ELECTRONS DO NOT ORBIT THE NUCLEUS LIKE PLANETS ORBITING A STAR!!!! Bohr’s Model

12. Electrons not really in “planetary” orbits Are really in areas of “probability” called “electron clouds” Quantum Model of electron placement

13. Quantum Mechanical Model This model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus

14. Heisenburg’s Uncertainty Principle –it is not possible to know both the velocity and position of a particle at the same time –velocity = speed and direction

15. Orbital - 3 dimensional region around the nucleus where a particular electron can be located  “clouds” - that show a region of high probability of finding an electron  size and shape of “cloud” depends on energies of electrons that occupy them

16. Principal Energy Levels  Indicates main energy level of an electron in an atom  called “shells”  1 = lowest 7 = highest  can be any positive integer

17. Sublevel Indicates the shape of an orbital labeled s, p, d, f

18. s = sphere s = sphere p = dumbbell or figure-eight p = dumbbell or figure-eight d = 4 lobes d = 4 lobes f = complicated f = complicated

22. Orbital shapes for Sc funky orbitals

23. Principal Energy Levels  Are divided into sublevels  the number of sublevels allowed is equal to the principal energy level (n) (up to n=4)  PEL = 11 sublevel  PEL = 22 sublevels  PEL = 33 sublevels  PEL = 44 sublevels

24. Wish you were here? Well, you’re not, so pay attention!

25. “s” sublevel – 1 orbital allowed “s” sublevel – 1 orbital allowed “p” sublevel – 3 orbitals allowed “p” sublevel – 3 orbitals allowed “d” sublevel – 5 orbitals allowed “d” sublevel – 5 orbitals allowed “f” sublevel – 7 orbitals allowed “f” sublevel – 7 orbitals allowed

26. PELsublevels allowed PELsublevels allowed 1s1s1s1s 2s, p 3s, p, d 4s, p, d, f

27. Each sublevel has a certain number of orbitals allowed Sublevelorbitals allowed s 1 p 3 d 5 f 7

28. Maximum of 2 e- in any orbital ! They “spin” in opposite directions

29. WAKE UP!!!!!!!!

30. Don’t give up!

31. You can’t escape!

32. I know you’d rather be here, but it gets better, I promise!

33. Are you ready????????

34. Chart that follows this slide: Principal energy level type of sublevel #orbitals per type #orbitals per level Max. # electrons

35. Principal Energy Level (pel) Type of sublevel # orbitals per type of sublevel # orbitals per pel n 2 Maximum number of electrons per pel 2n 2 1 s 1 1 2 2 s p 1 3 4 8 3 s p d 1 3 5 9 18 4 s p d f 1 3 5 7 16 32

36. Rules for writing electron configurations Add one electron at a time according to these rules: 1. each added electron is placed in a sublevel of lowest energy available (Aufbau Process) 2. No more than 2 electrons can be placed in any orbital (Pauli Exclusion Principle)

37. 3. Before a second electron can be placed in any orbital, all the orbitals of the sublevel must contain at least one electron (Hund’s Rule) 4. No more than 4 orbitals are occupied in the outermost principal energy level of any atom. (next electron must enter the next principal energy level)

38. NOW!!!! We will start writing electron configurations “Regular” and “Exceptions”

39. Stanford explanation

40.  Atoms absorb a SPECIFIC amount of energy – quanta  Electrons “jump” up into energy levels where they really don’t belong  Immediately drop back and release that specific amount of energy in the form of light of specific wavelength and frequency (color)

41. Spectroscopy  Used to study structure of atoms  substances heated –e - move to higher energy levels –“fall back” - release photons of energy of specific wavelength  produce a series of “spectral lines” –are characteristic to specific substances –used as an identifying tool

42.  Ground state - atom where the electrons are in the lowest available energy levels  excited state - atom has electrons that have “jumped” to higher energy levels

43. Can identify elements by the colors they produce FIREWORKS!!! Flame tests – lab we will do

54. flame test lab!