2nd quarter review 
Chemistry
Name: _KEY_____________________________________
Date: _________________________
Write the electron configurations for the following elements.
1. O: 1s22s22p4
2. Ca:1s22s22p63s23p64s2
3. Cu: 1s22s22p63s23p64s23d9
4. Sn: 1s22s22p63s23p64s23d104p65s24d105p2
5. Rn: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6
Name the groups of the periodic table.
1. __Alkali Metals________________
2. ___Alkaline Earth Metals_________
17. __Halogens_______________
18. ___Noble Gases_____________
D-Block: ___Transition Metals________
4f: __Lanthanide___________
5f: ___Actinide_________________
Circle the element that fits the statement
Li
N
K
S
Al
Ga
V
Te
Si
Li
As
H
Hg
Na
Pb
B
Ca
Si
P
Ca
Cl
Si
Al
Nb
I
Ge
Be
Se
Li
Tl
Mg
Bi
C
Sc
S
As
Sc
Ar
P
Si
Ta
Xe
Sn
B
Br
Na
Pb
Al
Po
N
Ti
metal
smallest ionization energy
largest atomic mass
member of the halogen family
greatest electron affinity
largest atomic radius
largest atomic number
member of noble gases
4 energy levels
member of alkali metals
6 valence electrons
nonmetal
member of transition metals
electron distribution ending in s2p1
metalloid
gas at room temperature
electron distribution ending in s2d2
Chemistry
2nd quarter review 
Trend Review
1. Rank the following elements by increasing atomic radius: carbon, aluminum,
oxygen, potassium.
O, C, Al, K
2. Rank the following elements by increasing electronegativity: sulfur, oxygen,
neon, aluminum.
Ne, Al, S, O
3. Why does fluorine have a higher ionization energy than iodine?
Effective nuclear charge is higher – the + charge of the nucleus is stronger for F
than I – the nucleus of F is able to attract electrons to it better than I.
4. Why do elements in the same family generally have similar properties?
They have the same number of valence electrons.
5. Indicate whether the following properties increase or decrease from left to right
across the periodic table.
a. atomic radius (excluding noble gases)
Decreases
b. first ionization energy
Increases
c. electronegativity
Increases
6. What trend in atomic radius occurs down a group on the periodic table? What
causes this trend?
It increases because of electron shielding – the core electrons block the valence
electrons from feeling the pull of the + nucleus and it allows the valence electrons
to get further and further from the nucleus thus expanding the size of the atom.
7. What trend in ionization energy occurs across a period on the periodic table?
What causes this trend?
It increases because of effective nuclear charge – the atoms are more able to
remove an electron from an atom because the nucleus is gaining + protons and
the atom is getting smaller. As you move across the table, the atoms want to gain
electrons more and more (until you get to the noble gases).
8. Circle the atom in each pair that has the largest atomic radius.
a. Al or B
c. S or O
e. Br or Cl
b. Na or Al
d. O or F
f. Mg or Ca
9.
Circle the atom in each pair that has the greater ionization energy.
a. Li or Be
c. Na or K
e. Cl or Si
b. Ca or Ba
d. P or Ar
f. Li or K
Chemistry
2nd quarter review 
10. Define electronegativity.
Electronegativity is the ability of a bonded atom (one that is bonded to
something else already) to attract electrons to it.
11. Circle the atom in each pair that has the greater electronegativity.
a. Ca or Ga
c. Li or O
e. Cl or S
b. Br or As
d. Ba or Sr
f. O or S
Bonding Review
12.
List three differences between ionic and covalent compounds:
Ionic bonds transfer electrons and covalent bonds share electrons.
Ionic bonds are stronger.
Ionic bonds have a metal bonded to a nonmetal (they are far apart on the
PT) and covalent bonds are between 2 nonmetals (close together on the PT).
13. Explain why ionic compounds are formed when a metal bonds with a nonmetal but
covalent compounds are formed when two nonmetals bond.
The metal and nonmetal in an ionic bond have a large electronegativity
difference so 1 atom takes the other atoms electrons. The 2 nonmetals in a
covalent bond doesn’t have a large electronegativity difference so 1 atom
doesn’t pull the other atom’s electrons off of it – they still share the electron.
Naming review
1) Na2CO3 ____sodium carbonate____________________
2) NaOH _____sodium hydroxide_____________________________
3) MgBr2 ____magnesium bromide_____________________________
4) KCl ____potassium chloride______________________________
5) FeCl2 ____iron (II) chloride_________________________________
6) FeCl3 ____iron (III) chloride_________________________________
7) Zn(OH)2 zinc (II) hydroxide__ the (II) could be omitted b/c zinc is a constant charge
8) BeSO4 __beryllium sulfate_____________________________
9) CrF2 ____chromium (II) fluoride_______________________________
10) Al2S3 ___aluminum sulfide_____________________________
2nd quarter review 
Chemistry
11) sodium phosphide ___Na3P_________________
12) magnesium nitrate _____Mg(NO3)2___________________________
13) lead (II) sulfite ____PbSO3______________________________
14) calcium phosphate ___Ca3(PO4)2_____________________________
15) ammonium sulfate __(NH4)2SO4____________________________
16) silver cyanide _____AgCN
Ag is a constant charge of +1 ___
17) aluminum sulfide __Al2S3_______________________________
18) beryllium chloride __BeCl2________________________________
19) copper (I) arsenide _Cu3As______________________________
20) iron (III) oxide ____Fe2O3__________________________________