Unit 2 CP Chemistry
Modern Atomic Structure
Review of Main Ideas
Key terms:
orbital
orbital diagram
electronegativity
Einstein
energy level
frequency
excited state
Niels Bohr
ionic radius
Hund’s Rule
Name ______________________________
Class ____________ Date _____________
cation
Aufbau Principle
electron affinity
wavelength
Zeff
spectroscope
photon
Max Planck
quantized
s,p,d,f-blocks
valence electron
Pauli exclusion principal
atomic radius
Lewis Dot diagram
shielding eground state
emission spectrum
s,p,d,f
electron configuration
electron spin
electromagnetic radiation
valence eionization energy
anion
-2.18x10-18 J
n2
DEelectron = E photon
DEelectron = E f - Ei
En =
c = l ×n
h×c
E=
l
1.
Calculate the frequency of a photon of light that has a wavelength of 700 nm (don’t forget to convert nm→m)
2.
Calculate the energy of a photon that has a wavelength of 350 nm.
3.
Calculate the frequency of a photon that has 4.50 x 10-19 J of energy.
4.
What is the relationship between the energy of a photon and it’s wavelength?
5.
What did Max Planck theorize?
6.
What happens to an electron as it absorbs a photon?
7.
How do we know that energy levels in a hydrogen atom are not continuous, meaning an electron could exist
anywhere?
8.
What is meant by the ground state of an atom?
9.
What formula determines the # of e- for any specific principal energy level?
10. Complete the chart
Energy
# of
Level
sublevels
1
2
3
4
sublevel
letters
# of orbitals in each
sublevel type
total # of e-
11. How many electrons can there be on any given sublevel or orbital?
Sublevel
Orbital
s
p
d
f
g
1
12. Calculate the energy that is released as a photon when an electron falls from n=4 to n=3.
13. Explain the emission spectrum of hydrogen. What produces it?
14. Identify the elements based on their configurations:
a. [Ar]4s13d10
b. [Kr]5s14d5
c. [Xe]7s25f4
15. How do atoms become ions?
16. What is the difference between an energy level and an orbital?
17. What information does the orbital filling diagram tell you that an electron configuration does not?
18. When filling a p orbital, why do you separate the electrons into different orbitals, before pairing them up. What is
the name of this rule?
19. What information does a Lewis dot diagram tell you?
20. What do the elements in the same period have in common?
21. What do elements in the same group have in common?
22. Which elements might rearrange their valence electrons in order to achieve stable outer energy levels?
23. How does Bohr's "electron orbits" differ from the "electron orbitals"?
24. Why do some properties of elements repeat in periodic patterns?
25. What types of ions do metals form?
26. What types of ions do nonmetals form?
27. For each of the following properties, indicate whether fluorine or iodine has a larger value:
a) electronegativity
b) electron affinity
c) atomic radius
d) ionization energy.
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28. Explain why radius of an atom cannot be measured directly.
29. Which elements are characterized as having their d orbitals fill with electrons as you move left to right across a
period?
30. What is the difference between a core (a.k.a. shielding) electron and a valence electron?
31. Explain why it’s harder to remove a shielding electron than a valence electron from an atom?
32. An element forms an cation when ionized. On what side of the periodic table is the element located? Explain.
33. What is the Periodic Law?
34. What is Hund’s Rule?
35. What is the Aufbau Principle?
36. What is the Pauli Exclusion Principle?
37. Explain why each successive ionization of an electron requires a greater amount of energy.
38. Na+ and Mg2+ ions each have ten electrons surrounding their nuclei. Which ion would you expect to have the larger
radius? Explain.
39. Draw the Lewis dot diagram, and electron configuration for the following elements.
Lewis-dot
Electron configuration (use noble gas shortcut)
Sybmol
diagram
Ne
Cu
S2Pu
Ca2+
40. How does the atomic radius change from left to right across a period. Explain why?
41. How does the atomic radius change from top to bottom for a group. Explain why?
3
42. How does the ionization energy change from left to right across a period. Explain why?
43. How does the ionization energy change from top to bottom for a group. Explain why?
44. Why do large atoms have small ionization energies?
45. Why are metal ions smaller than their metal atoms?
46. Why are nonmetal ions larger than their nonmetal atom?
47. Why does F have a higher electronegativity than Li? Be specific?
48. What electron configuration has the greatest stability?
49. Order the following groups from smallest to largest radii.
a. I-, Ba2+, Cs+, Xe
b.
Kr, Br-, Se2-, Sr2+
c.
Li+, K, Cs, Fr
50. Order the following groups from highest to lowest ionization energy.
a) Na, Mg, Al, S
b)
Br, F, At, I
c)
Cu, Zn, Ga, Se
51. Describe effective nuclear charge (Zeff).
52. Calculate the Zeff for Ca and Br. Can Zeff be used to explain atomic radius trends? Explain.
53. Write the reaction showing how aluminum absorbs 950 kJ of energy and is thus ionized.
54. Why is the 3rd ionization energy of Mg so much larger than it’s 2 nd IE?
55. Why is the ionization energy for Al less than Mg?
56. In which group are the three species isoelectronic (all have same # of e-)?
(A) S2-, K+, Ca2+
(B) Sc, Ti, V2+
(C) O2-, S2-, CI(D) Mg2+, Ca2+, Sr2+
57. The ionization energies for element X are listed in the table below. On the basis of the data, element X is most
likely
Ionization Energies for element X (kJ mol-1)
First
Second
Third
Fourth
Fifth
580
1,815
2,740
11,600
14,800
(A) Na
(B) Mg
(C) Al
(D) Si
58. Which property generally decreases across the periodic table from sodium to chlorine?
(A) 1st ionization energy (B) Atomic mass
(C) Ionic radius
(D) Atomic radius
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5
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