2. Light Periodic Trends Quantum Numbers Electron Configuration Mixed Bag $100 $200 $300 $400 $500

3. Mixed Bag

4. Which element has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 3 ?

5. If you add up the exponents, there are 15 electrons which corresponds to phosphorous.

6. What is the difference between the ground state and the excited state for an electron?

7. The ground state is the lowest energy level for an electron; the excited state is when an electron is at an energy level above its normal state (the excited state is temporary).

8. How many valence electrons does carbon have in it’s outer shell? (Hint: first write out the e - configuration)

9. C = 1s 2 2s 2 2p 2 The second energy level is the outside shell in this case. There are 4 electrons in the second energy level (valence shell) for carbon.

10. How many unpaired electrons does germanium have in it’s valence shell? (Hint: draw the orbital diagram first!)

11. After drawing the orbital diagram, there are 2 unpaired electrons in the germanium atom.

12. How many unpaired electrons are in there in a bromine atom? How does this relate to the charge on a bromide ion?

13. Bromine has 1 unpaired electron in it’s orbital diagram which makes sense since bromide ions have a charge of -1 (they have gained 1 e - to complete their valence shell).

14. Periodic Trends

15. Which atom would have the largest ionization energy – Mg or Cl?

16. Chlorine would have the largest I.E. because it is farther right on the periodic table (e- are held closer so it will be harder to rip the e- off)

17. Define the term “electron affinity”

18. The energy involved when an atom gains an electron. (The opposite of ionization energy)

19. Rank the following elements in order of increasing atomic radius: In, Al, B, Ga

20. B < Al < Ga < In

21. Rank the elements in order of increasing metallic character: N, Zr, Cs, Co

22. N < Co < Zr < Cs

23. Describe the comparative size of a cation compared to the atom from which it forms.

24. Cations are smaller than the atoms from which they form due to the increasing proton/electron attraction due to the loss of electrons.

25. Electron Configuration

26. What is the electron configuration for cobalt?

27. Co = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7

28. What is the noble gas configuration for zirconium?

29. Zr = [Kr]5s 2 4d 2

30. What is the orbital configuration for sulfur?

31. 1s 2s 2p 3s 3p

32. What do the coefficient, letter, and exponent signify when writing electron configurations?

33. Coefficient = principal energy level Letter = sublevel Exponent = # of electrons in sublevel

34. What is the electron configuration for iridium?

35. Ir = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 7

36. Quantum #’s

37. What is the maximum number of electrons that can be contained in an: s sublevel? p sublevel? d sublevel? f sublevel?

38. s sublevel – 2 e - p sublevel – 6 e - d sublevel – 10 e - f sublevel – 14 e -

39. Which sublevels (s, p, d, or f) are located in the fourth principal energy level?

40. The third energy level has s, p, d, and f sublevels present.

41. How many total electrons can fit into the third principal energy level? (Take into account all sublevels present in the 3 rd energy level)

42. The third energy level contains s, p, and d sublevels so 2 + 6 + 10 = 18 total electrons!

43. How do you draw arrows in an orbital diagram and why are they drawn that way?

44. Arrows are drawn with the first pointing up and the second pointing down due to the Pauli Exclusion Principal (electrons must have opposite spins in the same orbital).

45. Define the Aufbau Principle and how it relates to the electrons in an atom.

46. The Aufbau Principle says that electrons completely fill from the lowest energy level to the highest energy level.

47. Light

48. Describe the dual nature of light

49. Light behaves as both a particle and a wave.

50. What is the wavelength of a portion of light having a frequency of 4.73 x 10 12 Hz?

51. Use c = λν… λ = 6.34 x 10 -5 m

52. Calculate the energy of a photon of light having a frequency of 3.11 x 10 13 kHz.

53. Use E = hv… E = 2.06 x 10 -17 J

54. What is the wavelength of a portion of light having an energy of 1.44 x 10 -12 J?

55. Use both light equations! First calculate the frequency… v = 2.17 x 10 21 Hz Then use v to find λ… λ = 1.38 x 10 -13 m

56. Describe why you saw colors in the flame during the flame test lab (think about what is going on with the electrons)

57. As the element samples are placed in the flame, the electrons are given energy so they jump from the ground state to the excited state. When the electrons go back down to the ground state, they emit this extra energy which we see as light (the color depends on the λ and v.

58. Final Jeopardy What are the general shapes of the four orbitals (s, p, d, and f)?

59. Final Jeopardy s = spherical p = peanut d = dumbell (clover) f = flower