CHM 130 HW2 Ch. 5-9
Spring 2017
Instructor: _____________ Name: ______KEY_______
Due Friday March 24 by noon, no later.
1. As wavelength increases, frequency __decreases_____ and energy _____decreases______.
2. Fill in the blanks in this electromagnetic spectrum:
Gamma
X Ray
___UV_____ visible ____IR_____ microwave ____radio_____
3. Which scientist decided that electrons are quantized in their orbits? ____Bohr_____
4. How does a neon sign work in your own words?
A gas is put in a glass discharge tube and plugged into electricity. The energy excites the
electrons and so they jump to an outer orbital absorbing the energy. Then they jump back to
where they started and they must release the energy. It is released as a specific wave of light.
Each atoms has their own specific electron jumps so their own specific colors.
5. How many electrons can fit into the third energy level? _____18____
6. How many electrons can fit into a d sublevel? ______10__________
7. How many orbitals are in a p sublevel? ______3________
8. How many orbitals are on the second energy level? ______4______
9. How many electrons can fit on the fourth energy level? ______32_______
10. How many electrons can fit into an orbital? _____2________
11. How many sublevels are on the fourth level? ____4________
12. Write electron configurations for the following atoms and ions:
a. Ca2+ _______1s22s22p63s23p6___________________
b. F- ______1s22s22p6_________________________
c. O _________1s22s22p4____________________________
d. K ________1s22s22p63s23p64s1___________________________
e. P3- ________1s22s22p63s23p6___________________________
f.
Li ____________1s22s1___________________________________
13. How many valence electrons do the following atoms have?
a. P ____5______
b. Al _____3___
c. C _____4______
d. Na ___1____
14. What does a sodium atom want to do in order to become stable like a noble gas?
Na wants to lose its one outer electron to be s2p6 like the noble gas neon and get charged +1.
15. What does a phosphorus atom want to do in order to become stable like a noble gas?
P wants to gain 3 electrons to become -3 charged and be s2p6 like argon.
16. What are the charges for the following when they become ions?
a. Na ___+1___
b. N ____-3___
c. Al __+3__
d. I __-1___
e. Rb ___+1___
3-
17. P is isoelectronic with what atom? ___argon______
18. Ca2+ is isoelectronic with what atom? ____argon______
19. Fill in this table for these atoms and ions:
Cl -
25
Mg2+
Atoms or Ion
37
37
25
# protons
17
17
12
12
# electrons
17
18
12
10
# neutrons
20
20
13
13
mass
37
37
25
25
Cl
Mg
20. In which group is calcium? ___alkaline earth metals___________
21. In which group is fluorine? ______halogens__________
22. Why is a chlorine atom smaller than a silicon atom when chlorine’s mass is larger than silicon’s? (Do
not quote the trend – explain!)
Cl has 17 protons and they pull the outer electrons inwards more than Si’s 14 protons can. The
are both filling in the third level, and so more protons pulls electrons tighter.
23. Which atom is largest?
a. In
b. I
c. Sr
d. Mg
e. Cl
24. For each PAIR, circle which is larger:
a. Ca or Ca2+
b. Al or Al3+
c. S or S2-
25. Define ionization energy:
IE is the energy required to remove an outer electron from an atom
26. True or False? IE is generally very low for small atoms. __false_____
27. True or False? IE is generally very high for metals. ___False_____
28. Explain why IE is generally high for the noble gases. (Do not quote the trend – explain!)
The noble gases are already s2p6 in their outer configuration, and are stable and happy. They
do NOT want to gain or lose any electrons. So removing an electron from them is very bad and
takes a lot of energy. They hold onto their electrons tightly and don’t readily let go.
29. Define electronegativity:
Ability to pull bonded electrons close.
30. What is the most electronegative element? _____F______
31. Why don’t noble gases have electronegativity values?
To have an electronegativity value you must bond. You can’t pull bonded electrons close
without a bond to another atom. And noble gases don’t bond.
32. Noble gases are very stable and inert. Why? What is so special about them?
Noble gases, except He which is 1s2, are all s2p6 which is 2+6=8 outer electrons. Thus the octet
rule! Atoms are very stable and happy with 8 outer electrons!
33. True or False? When atoms gain and lose electrons an ionic bond can form. ___true___
34. True or False? When atoms share electrons an ionic bond can form. ___false_____
35. True or False? Cations are larger than their corresponding atom. ___false____
36. Circle the covalent molecules:
NaCl
CO
H2
CaO
NH3
37. True or False? Breaking a bond is exothermic. ____false______
38. True or False? Sharing 4 electrons is called a double covalent bond. ___true____
39. Draw Lewis dot structures for the following:
SF2
# val e- = ___20__
Polar bonds? Yes or No
Polar molecule? Yes or No
Shape: __bent_____________
SiBr4
# val e- = _32__
Polar bonds? Yes or No
Polar molecule? Yes or No
Shape: _____tetrahedral___________
40. Draw lewis dot structures for the following ions:
NO2 -1
# val e- = __18___
Shape: ____linear______
Shape: ______bent________
41. Circle the bond if it is polar:
# val e- = __14___
OCl -
C-H
- +
+ -
F-N
S-F
+ O-O
42. Add the delta notation above.
43. Write the correct formulas for the following compounds;
a. Barium acetate ______ Ba(C2H3O2)2_______
b. Calcium cyanide ______ Ca(CN)2_________
c. Sodium nitride ________Na3N__________
d. Iron(III) phosphide ______FeP_______
e. Phosphorus pentachloride ___PCl5________
f.
Copper(II) sulfite ____CuSO3________
g. Aluminum sulfate _____Al2(SO4)3_______
h. Silicon tetrafluoride _____SiF4________
i.
Carbon dioxide _____CO2_______
j.
Cobalt(III) iodide ______CoI3_________
k. Strontium nitrite _____Sr(NO2)2________
l.
Potassium sulfide ______K2S__________
m. Nickel(III) nitrate _______Ni(NO3)3__________
44. Write the names for the following compounds:
a. AuBr3 _____gold(III) bromide__________
b. SF2 ___________sulfur diflouride__________________
c. P2O5 _________diphosphorus pentoxide__________________
d. Mn(C2H3O2)2 _____manganese(II) acetate____________
e. ZnS _________zinc sulfide______________
f.
PbO2 ________lead(IV) oxide______________________
g. SnCl2 _________tin(II) chloride__________________________
Mg-O
Br-O
h. KI _____________potassium iodide_________________________
i.
Al2(SO3)3 ________aluminum sulfite__________________________
j.
SO3 __________sulfur trioxide________________________________
k. MgCl2 ___________magnesium chloride__________________________
l.
Na2CrO4 ________sodium chromate_______________________
m. FeSO4 __________iron(II) sulfate________________________________
45. What ion does an Arrhenius acid produce in water? __H+____
46. What ion does an Arrhenius base produce in water? ___OH -_____
47. Circle the Arrhenius acid in this reaction: HBr(aq) + LiOH(aq)  LiBr(aq) + H2O(l)
HI(aq) + NH3(aq)  NH4+(aq) + I- (aq)
48. Circle the Bronsted Lowry base in this reaction:
49. Which of the following is a strong base? C
a.
b.
c.
d.
50. Give the names for three strong acids: ___hydrochloric acid______
________sulfuric acid
and nitric acid_________________________
51. Give the names for two weak acids: hydrofluoric acid, carbonic acid, acetic acid, phosphoric acid
52. If the pH is 5, the solutions is:
53. Circle the soluble compounds:
a. acidic
Na2S
b. neutral
MgS
LiOH
c. basic
Al(OH)3
CaCO3
BaSO4
54. Draw the following electrolytes in beakers of water:
NaNO3
KBr
AlPO4
MgS
55. If [H+] is 10-5 then the pH = ___5_____.
56. If [H+] is 10-12 then the pH = ____12____.
57. How is a strong electrolyte different from a weak electrolyte and from a non-electrolyte?
Strong electrolytes break apart 100% into ions (all ions). Weak electrolytes break apart into few
ions, like 1-5% ions, the rest remains bonded together. Non-electrolytes do not break apart at
all, zero ions.