AP Chemistry – Chemistry I Review Assignment
Nomenclature
1. Name these binary compounds of two nonmetals.
IF7 – iodine heptafluoride
N2O5 – dinitrogen pentoxide
XeF2 – xenon difluoride
N2O4 – dinitrogen tetroxide
As4O10 – tetrarsenic decoxide
SF6 – sulfur hexafluoride
PCl3 – phosphorous trichloride
S2Cl2 – disulfur dichloride
2. Name these binary compounds with a fixed charge metal.
AlCl3 – aluminum chloride
MgO – magnesium oxide
BaI2 – barium iodide
KI – potassium iodide
SrBr2 – strontium bromide
Na2S – sodium sulfide
CaF2 – calcium fluoride
Al2O3 – aluminum oxide
3. Name these binary compounds of cations with variable charge.
CuCl2 – copper (II) chloride
Fe2O3 – iron (III) oxide
SnO – tin (II) oxide
PbCl4 – lead (IV) chloride
Cu2S – copper (I) sulfide
HgS – mercury (II) sulfide
AuI3 – gold (III) iodide
CoP – cobalt (III) phosphide
4. Name these compounds with polyatomic ions.
Fe(NO3)3 – iron (III) nitrate
NaOH – sodium hydroxide
Cu2SO4 – copper (I) sulfate
Ca(ClO3)2 – calcium chlorate
KNO2 – potassium nitrite
NaHCO3 – sodium hydrogen carbonate
NH4NO2 – ammonium nitrite
Cu2Cr2O7 – copper (I) dichromate
5. Name these binary acids
HCl(aq) – hydrochloric acid
HI(aq) – hydroiodic acid
6. Name these acids with polyatomic ions.
HClO4 – perchloric acid
H2SO4 – sulfuric acid
HC2H3O2 – acetic acid
H3PO4 – phosphoric acid
HNO2 – nitrous acid
H2CrO4 – chromic acid
H2C2O4 – dichromic acid
H2CO3 – carbonic acid
7. Name these compounds appropriately.
CO – carbon monoxide
NH4CN – ammonium cyanide
HIO3(aq) – iodic acid
NI3 – nitrogen triiodide
AlP – aluminum phosphide
OF2 – oxygen difluoride
LiMnO4 – lithium permanganate
HClO(aq) – hypochlorous acid
HF(aq) – hydrofluoric acid
SO2 – sulfur dioxide
CuCr2O7 – copper(II) dichromate
K2O – potassium oxide
FeF3 – iron(III) fluoride
KC2H3O2 – potassium acetate
MnS – manganese(II) sulfide
8. Write the formulas.
Tin (IV) phosphide – Sn3(PO4)4
Copper (II) cyanide – Cu(CN)2
Magnesium hydroxide – Mg(OH)2
Sodium peroxide – Na2O2
Sulfurous acid – H2SO3
Lithium silicate – Li4SiO4
Potassium nitride – K3N
Chromium (III) carbonate – Cr2(CO3)3
Gallium arsenide - GaAs
Cobalt (II) chromate – CoCrO4
Zinc fluoride – ZnF2
Dichromic acid – H2Cr2O7
Solubility rules
9. Review solubility rules and identify each of the following compounds as soluble or insoluble in water.
Na2CO3 – Sol.
CoCO3 – Insol.
Pb(NO3)2 – Sol.
K2S – Sol.
BaSO4 – Sol.
(NH4)2S – Sol.
AgI – Insol.
Ni(NO3)2 – Sol.
KI – Sol.
FeS – Insol.
PbCl2 – Insol.
CuSO4 – Sol.
Li2O – Sol.
Mn(C2H3O2)2 – Sol.
Cr(OH)3 – Insol
AgClO3 – Insol.
Sn(SO3)4 – Insol.
FeF2 – Insol.
10. Predict whether each of these double replacement reactions will give a precipitate or not based on the solubility of the
products. If yes, identify the precipitate.
silver nitrate and potassium chloride – Yes, AgCl
AgNO3 + KCl  AgCl + KNO3
magnesium nitrate and sodium carbonate – Yes, MgCO3
Mg(NO3)2 + Na2CO3  MgCO3 + 2NaNO3
strontium bromide and potassium sulfate – Yes, SrSO4
SrBr2 + K2SO4  2KBr + SrSO4
cobalt (III) bromide and potassium sulfide – Yes, Co2S3
2CoBr3 + 3K2S  Co2S3 + 6KBr
ammonium hydroxide and copper (II) acetate – Yes, Cu(OH)2
2NH4OH + Cu(C2H3O2)2  Cu(OH)2 + 2NH4C2H3O2
lithium chlorate and chromium (III) fluoride – Yes, Cr(ClO3)3
3LiClO3 + CrF3  Cr(ClO3)3 + 3LiF
Balancing Equations
11. Balance the following equations with the lowest whole number coefficients.
S8 + 12O2  8SO3
C10H16 + 8Cl2 
10C + 16HCl
4Fe + 3O2 
2Fe2O3
2C7H6O2 + 15O2  14CO2 + 6H2O
2KClO3  2KCl + 3O2
2H3AsO4 
As2O5 + 3H2O
V2O5 + 6HCl  2VOCl3 + 3H2O
3Hg(OH)2 + 2H3PO4  Hg3(PO4)2 + 6H2O
Stoichiometry and Limiting Factor
12. Given the equation below, what mass of water would be needed to react with 10.0g of sodium oxide?
Na2O + H2O  2NaOH
 1 mol Na 2 O   1 mol H 2 O   18 g H 2 O 
 
 
  2.9 g H 2 O
10 g Na 2 O 
 62 g Na 2 O   1 mol Na 2 O   1 mol H 2 O 
13. 2NaClO3  2NaCl + 3O2
What mass of sodium chloride is formed along with 45.0 g of oxygen gas?
 1 mol O 2   2 mol NaCl   58.44 g NaCl 
 
 
  54 .8 g NaCl
45.0 g O 2 
 32 g O 2   3 mol O 2   1 mol NaCl 
14. 4NH3 + 5O2  4NO + 6 H2O
What mass of water will be produced when 100.0 g of ammonia is reacted with
excess oxygen?
 1 mol NH 3   6 mol H 2 O   18 g H 2 O 
 
 
  158 g H 2 O
100 .0 g NH 3 
 17 g NH 3   4 mol NH 3   1 mol H 2 O 
15. If the reaction in #14 is done with 25.0g of each reactant, which would be the
limiting factor?
 1 mol NH 3   4 mol NO 
 
  1.47 mol NO
25.0 g NH 3 
 17 g NH 3   4 mol NH 3 
 1 mol O 2   4 mol NO 
 
  0.62 mol NO
25.0 g O 2 
 32 g O 2   5 mol O 2 
Thus, the O2 is the limiting reagent
16. Na2S + 2AgNO3  Ag2S + 2NaNO3
If the above reaction is carried out with 50.0g of sodium sulfide and 35.0g of silver
nitrate, which is the limiting factor?
 1 mol Na 2S   1 mol Ag2S 
 
  0.641 mol Ag2S
50 .0 g Na 2S 
 78.04 g Na 2S   1 mol Na 2S 
 1 mol AgNO3   1 mol Ag2S 
 
  0.103 mol Ag2S
35.0 g AgNO3 
 169 .87 g AgNO3   2 mol AgNO3 
Thus the AgNO3 is the limiting reagent
What mass of the excess reactant remains?
 1 mol Na 2S   78 .04 g Na 2S 
 
  8.04 g Na 2S consumed
0.103 mol Ag2S consumed 
 1 mol Ag2S   1 mol Na 2S 
Thus, (50.0 – 8.04) = 41.96 g Na2S are left
What mass of silver sulfide would precipitate?
0.103 mole Ag2S would precipitate, thus 25.5 g Ag2S would precipitate
17. 6NaOH + 2Al  2Na3AlO3 + 3H2
What volume of hydrogen gas (measured at STP) would result from reacting 75.0g of sodium
hydroxide with 50.0g of aluminum?
 1 mol NaOH   3 mol H 2   22.4 L of H 2 @ STP 
  21 .0 L of H 2 @ STP
 
 
75.0 g NaOH 
1 mol H 2
 40 g NaOH   6 mol NaOH  
