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Chemistry Honors Quarterly 2 Review 2013-2014 School Year
Name: _______________________________ Teacher: _____________________
The Electron:
1.
Light is observed during a flame test of an element when electrons move from
__________________________ energy levels to __________________________ energy levels.
2.
The _________________________________ Principle states that electrons always fill
The lowest energy orbitals of an atom first.
3.
The _______________________________ Principle states that an orbital can hold a
maximum of 2 electrons which must be spinning in opposite directions.
4.
What is Hund’s rule for electron configurations?
5.
How is the ground state of an atom different from the excited state?
6.
Why is it not possible for the 2f orbital to exist?
7. Illustrate the shape and energy of the four electron orbitals.
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8. For all of the following elements:
 Write the electron configuration
 Write the noble gas configuration
 Draw the Lewis dot notation for each
 draw the orbital diagram for the highest energy sublevel
 Determine how many valence electrons each element has.
 Determine the charge each element would have as an ion
a. Oxygen
b. Fluorine
c. Sodium
d. Aluminum
e. Phosphorus
f. Cadmium
g. Argon
9. What are the four sublevel blocks of the periodic table?
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10. Give an example of an electron transition that would absorb energy into the atom.
How is this transition different from one that would release energy from the
atom?
11. The concept of electrons existing in specific orbits around the nucleus was the
contribution of
12. The number of orbitals in a d sublevel is
13. The number of electrons in the third principal energy level in an atom having the
electron configuration 1s22s22p63s23p2
14. The total number of orbitals that contain at least one electron in an atom having the
configuration 1s22s22p63s23p2 is
15. Which of these elements has two s and six p electrons in its outer energy level?
a.
He
b.
O
c.
Si
d.
Ar
16. Which element is not a noble gas?
a.
Ra
c.
He
b.
d.
Xe
Ar
17. Which element has the largest number of unpaired electrons?
a.
F
b.
S
c.
Cu
d.
N
18. How many unpaired electrons are in the electron configuration [Ar]4s13d5?
19. Groups 13-18 form the area on the periodic table where the electron sublevels being
filled are
The Periodic Table
1. Why did Mendeleev leave blank spaces on his periodic table?
3. What is the difference between families/groups and periods/series?
4. Label the family of elements that each of the following elements belongs to.
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a. sodium _____________________________
b. magnesium __________________________
c. iron ______________________________
d. chlorine ______________________________
e. neon ______________________________
5. Which period do each of the following belong to?
a. strontium _________________
b. oxygen __________________
c. the element whose electron configuration ends in 4s 2 _______________
6. Predict what happens to each of these periodic properties as you move down a group of elements on the
periodic table and explain why.
a. atomic radius _________________________
b. electronegativity _______________________
7. Predict what happens to each of these periodic properties as you move from left to right across a series on
the periodic table and explain why.
a. atomic radius _________________________
b. electronegativity (electron affinity) _______________________
9. What is the defining characteristic of the noble gas family of elements?
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10. What family of elements is defined by the “d” electron sublevel?
12.Define electronegativity.
13. What do elements in the same vertical column all have in common?
14. Where are the metals located on the PT?
15. Describe the physical properties of metals.
16. Where are the nonmetals located on the PT?
17. Describe the physical properties of nonmetals.
18. What do we call the elements that lie between the metals and the nonmetals,
having some properties of each?
19. Give two examples of those types of elements.
20. What do we call an atom that has either gained or lost electrons?
21. What is the name given to a positive ion?
22. How does a positive ion form?
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23. What is the name given to a negative ion?
24. How does a negative ion form?
25. The lanthanide and actinide series of elements are
a.
all artificially made
b.
filling in d level electrons
c.
transition elements
d.
filling in f level electrons
26. The element having the structure 1s22s22p63s23p2 is in Group
27. In Group 15, the element having the smallest atomic radius is
28. In Group 14, the most metallic element is
29. Which group in the periodic table contains the least reactive elements?
30. Which group on the periodic table contains the Alkali Metals?
31. An atom of fluorine is smaller than an atom of oxygen. One possible explanation is
that, compared to oxygen, fluorine has
a.
a larger mass number
b.
a greater nuclear charge
c.
a smaller atomic number
d.
more unpaired electrons
32. If the size of the fluorine atom is compared to the size of the fluoride ion,
a. they would both be the same size
b. the ion is larger than the atom
c. the atom is larger than the ion
d. the sizes depend on the reaction
33. Sodium is a very active metal because
a.
it has a low ionization energy
b.
it has a relatively small atomic mass
c.
it has only one outermost electron
d.
all of the above
34. Which of the following formulas is incorrect?
a.
Na+
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b.
c.
d.
SAl3+
F-
35. Give an element that belongs to each of the following categories:
a.
Halogens
b.
Lanthanide Series
c.
Alkali Metals
d.
Transition Metals
e.
Noble Gases
f.
Alkali Earth Metals
g.
Period 5
i.
p-block elements
36. Give the element that fits the following criteria:
a.
largest in Group 14
b.
highest electronegativity in period 2
c.
lowest 1st ionization energy in group 8
d.
smallest in period 5
e.
least reactive of the halogens
f.
lowest electronegativity in the oxygen family
g.
greatest 1st ionization energy in period 3
37. What are the only two elements that exist as liquids at normal atmospheric pressure
and 25oC?
Bonding & Compounds
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1.
When metals form compounds they tend to _________________ their valence electrons
and become _________________.
2.
When nonmetals form compounds they tend to _____________________ their valence
electrons and become ________________________.
3.
As the electronegativity difference between two elements that are bonded increases, the
percent _____________________________ character increases.
4.
5.
The octet rule states that all stable atoms and ions have ______________ valence electrons.
____________________________ compounds are formed by the transfer of electrons from
one atom to another. _____________________________ compounds, on the other
hand, are formed when elements share electrons.
6.
For each of the following compounds, determine if it is ionic or covalent and name it.
a. NaF ____________________________________________
b. SiF4 ______________________________________________
c. K2SO4 ____________________________________________
d. LiNO3 ______________________________________________
e. N2O4 _______________________________________________
f. CO _________________________________________________
g. FePO4 _____________________________________________
7.
Write formulas for the following compounds.
a. boron trihydride __________________________________
b. potassium oxide ___________________________________
c. aluminum sulfate _________________________________
d. iron III hydroxide _________________________________
e. sulfur dioxide _____________________________________
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8.
Draw Lewis Structures for the following.
a. NH3
b. CO2
c. CF4
d. H2O
9. Which of the following is an incorrect formula?
a.
NaCl
b.
K2O
c.
AlO
d.
BaO
10. Which of the following has bonding that is ionic?
a.
H2
b.
MgF2
c.
H2O
d.
CH4
11. Which of the following is a correct Lewis structure?
12. Which of the following is an incorrect Lewis structure?
13. When a magnesium atom participates in a chemical reaction, it is most likely to
gain/lose
electrons.
14. If X represents an element of Group 13, what is the general formula for its oxide?
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15. As the difference in electronegativity between two elements decreases, the tendency
for the elements to form a covalent bond
16. The number of electrons in a triple bond is
17. The number of lone (unshared) pairs of electrons in H2O is
18. In which of the following is the formula correct for the name given?
a.
c.
e.
g.
i.
k.
m.
o.
q.
s.
u.
w.
y.
aa.
ac.
ae.
ag.
ai.
ak.
copper (II) sulfate, CuSO4
mercury (I) carbonate, HgCO3
calcium acetate, Ca(CH3COO)2
dichlorine heptoxide, Cl2O7
sulfurous acid, H2SO3
lead (II) chromate, PbCrO4
mercury (II) sulfate, HgSO4
sodium hypochlorite, NaClO
cadmium cyanide, Cd(CN)2
carbonic acid, H2CO3
iron (III) iodide, FeI2
carbon monoxide, CO
sodium bromide, Na2Br
potassium hydroxide, POH
zinc sulfate, ZnSO4
tin (IV) nitrate, Sn(NO3)4
chloric acid, HCl
cobalt (II) chloride, CoCl2
zinc oxide, ZnO2
j.
l.
n.
p.
r.
t.
v.
x.
z.
ab.
ad.
af.
ah.
aj.
al.
b.
ammonium hydroxide, NH4OH
d.
phosphorous triiodide, PI3
f.
hypochlorous acid, HClO
h.
magnesium iodide, MgI
potassium permanganate, KMnO4
iron (II) phosphate, FePO4
dinitrogen pentoxide, N2O5
sodium dichromate, Na2Cr2O7
bismuth (III) oxide, Bi3O2
silver oxide, Ag2O
tin (II) fluoride, TiF2
phosphoric acid, H3PO3
hydrosulfuric acid, H2S
sodium carbonate, Na2CO3
sulfur trioxide, SO3
iron (II) sulfate, FeSO4
aluminum sulfide, Al2S3
acetic acid, CH3COOH
tin (IV) nitrite, Sn(NO3)4
19. What is a diatomic element and which elements are diatomic?
20. What is a binary compound?
21. Determine the oxidation states for each atom in the following molecules.
CO2
CF4
H2O
PI3
N2O5