Download Unit 8 Homework Packet

Name:
Per:
Date:
Unit 8 - Chemical Quantities (Stoichiometry)
Chemistry 1
Information Given by Chemical Equations
1. Explain why, in the balanced chemical equation C + O2 → CO2, we know that 1 g of C will not react exactly
with 1 g of O2.
2. For each of the following reactions, give the balanced equation for the reaction and state the meaning of the
equation in terms of numbers of individual molecules and in terms of moles of molecules. The first one is
done for you.
a. 2NO(g) +
O2(g) →
2NO2(g)
- 2 molecules of nitrogen monoxide react with 1 molecule of oxygen to produce 2 molecules of
nitrogen dioxide.
- 2 moles of nitrogen monoxide react with 1 mole of oxygen to produce 2 moles of nitrogen dioxide.
b. AgC2H3O2 (aq) +
CuSO4(aq) →
Ag2SO4(s) +
c. PCl3(l) +
H2O(l) →
H3PO3(l) +
d. C2H6(g) +
Cl2(g) →
C2H5Cl(g) +
Cu(C2H3O2)2(aq)
HCl(g)
HCl(g)
Mole-Mole Relationships
3. True or false? For the reaction represented by the chemical equation: 2H2O2(aq) → 2H2O(/) + O2(g)
if 2.0 g of hydrogen peroxide decomposes, then 2.0 g of water and 1.0 g of oxygen gas will be produced.
Unit8-StudHand1213.doc
1
4. Consider the balanced equation CH4(g) + 2O2(g) → CO2(g) + 2H2O(g).
a. What is the mole ratio that would enable you to calculate the number of moles of oxygen needed to
react exactly with a given number of moles of CH4(g)?
b. What mole ratios would you use to calculate how many moles of each product form from a given
number of moles of CH4?
5. Balance the following equation:
Ag(s) +
H2S(g) →
Ag2S(s) +
H2(g)
a. Identify the mole ratios that you would use to calculate the number of moles of each product that
would form from 5 moles of silver reacting.
6. For each of the following unbalanced chemical equations, first balance them. Then calculate how many
moles of each product would be produced by the complete conversion of 0.125 mol of the reactant indicated
in boldface. State clearly the mole ratio used for the conversion. The first one is done for you.
a.
2 FeO(s) +
C(s) →.
2 Fe(l) +
CO2(g)
0.125 mol FeO × 2 mol Fe = 0.125 mol Fe
0.125 mol FeO ×
2 mol FeO
b.
Cl2(g) +
c.
NaI(aq) +
Unit8-StudHand1213.doc
KI(aq) →
KCl(aq) +
Pb(NO3)2(aq) →
1 mol CO 2
= 0.0625 mol CO2
2 mol FeO
I2(s)
PbI2(s) +
NaNO3(aq)
2
7. For each of the following unbalanced equations, first balance them. Then indicate how many moles of the
first product listed are produced if 0.625 mol of the second product forms. State clearly the mole ratio used
for each conversion.
a.
KO2(s) +
b.
SeO2(g) +
H2Se(g) →
c.
Fe2O3(s) +
Al(s) →
H2O(/) →
O2(g) +
KOH(s)
Se(s) +
Fe(/) +
H2O(g)
Al2O3(s)
Mass Calculations
8. What quantity serves as the conversion factor between the mass of a sample and the number of moles the
sample contains?
9. For each of the following unbalanced equations, first balance them. Then, calculate how many moles of the
second reactant would be required to react completely with exactly 25.0 g of the first reactant listed.
Indicate clearly the mole ratio used for each conversion.
a.
Mg(s) +
b.
AgNO3(aq) +
NiCl2(aq) →
AgCl(s) +
c.
KHCO3(aq) +
HCl(aq) →
KCl(aq) +
Unit8-StudHand1213.doc
CuCl2(aq) →
MgCl2(aq) +
Cu(s)
Ni(NO3)2(aq)
H2O(l) +
CO2(g)
3
10. For each of the following unbalanced equations, first balance them. Then, calculate how many milligrams
of each product would be produced by complete reaction of 10.0 mg of the reactant indicated in boldface.
Indicate clearly the mole ratio used for the conversion.
a.
Cr(s) +
SnCl4(l) →
b.
Fe(s) +
S8(s) →
c.
Ag(s) +
HNO3(aq) →
CrCl3(s) +
Sn(s)
Fe2S3(s)
AgNO3(aq) +
H2O(l) +
NO(g)
11. Although we usually think of substances as "burning" only in oxygen gas, the process of rapid oxidation to
produce a flame may also take place in other strongly oxidizing gases. For example, when iron is heated and
placed in pure chlorine gas, the iron "burns" according to the following (unbalanced) reaction:
Fe(s) +
Cl2(g) →
FeCl3(s)
How many milligrams of iron (III) chloride result when 15.5 mg of iron is reacted with an excess of chlorine
gas?
Unit8-StudHand1213.doc
4
Mass Calculations Using Scientific Notation
12. If chlorine gas is bubbled through a potassium iodide solution, elemental iodine is produced. The
unbalanced equation is:
Cl2(g) +
KI(aq) →
I2(s) +
KCl(aq)
Calculate the mass of iodine produced when 4.50 x 103 g of chlorine gas is bubbled through an excess amount of
potassium iodide solution.
13. Elemental fluorine and chlorine gases are very reactive. For example, they react with each other to form
chlorine monofluoride.
Cl2(g) + F2(g) → 2ClF(g)
Calculate the mass of chlorine gas required to produce 5.00 × 10-3 g of chlorine monofluoride given an excess
of fluorine gas.
The Concept of Limiting Reactants
14. What is the limiting reactant for a process? Why does a reaction stop when the limiting reactant is
consumed, even though there may be plenty of the other reactants present?
15. Nitrogen (N2) and hydrogen (H2) react to form ammonia (NH3). Consider the mixture of N2(g) (
)
and H2(g) (
) in a closed container as illustrated below. Assuming the reaction goes to completion,
draw a representation of the product mixture. Explain how you arrived at this representation.
Unit8-StudHand1213.doc
5
Calculations Involving a Limiting Reactant
16. Explain how one determines which reactant in a process is the limiting reactant.
17. What does it mean to say a reactant is present "in excess" in a process? Can the limiting reactant be present
in excess? Does the presence of an excess of a reactant affect the mass of products expected for reaction?
18. For each of the following unbalanced chemical equations, suppose that exactly 15.0 g of each reactant taken.
Determine which reactant is limiting, and calculate the mass of the underlined product produced. (Assume
that the limiting reactant is completely consumed.) Be sure to balance the reaction first!
a.
2 Al(s) + 6 HCl(aq) →
15 g Al ×
1 mol Al
27g Al
15 g HCl ×
×
2 AlCl3(aq) + 3 H2(g)
2 mol AlCl 3
2 mol Al
1 mol HCl
36.5 g HCll
×
×
133.5 g AlCl 3
mol AlCl 3
2 mol AlCl 3
6 mol HCl
×
= 74.2 g AlCl3
133.5 g AlCl 3
mol AlCl 3
= 18.3 g AlCl3
HCl is limiting; since it makes less AlCl3
b.
NaOH(aq) +
Unit8-StudHand1213.doc
CO2(g) →
Na2CO3(aq) +
H2O(l)
6
Pb(NO3)2(aq) +
c.
K(s) +
d.
HCI(aq) →
I2(s) →
PbCl2(s) +
HNO3(aq)
KI(s)
19. For each of the following unbalanced chemical equations, suppose that exactly 50.0 g of each reactant is
taken. Determine which reactant is limiting, and calculate what mass of the product in boldface is expected.
(Assume that the limiting reactant is completely consumed.)
a.
NH3(g) +
b.
BaCl2(aq) +
Unit8-StudHand1213.doc
Na(s) →
NaNH2(s) +
Na2SO4(aq) →
H2(g)
BaSO4(s) +
NaCl(aq)
7
20. The more reactive halogen elements are able to replace the less reactive halogens from their compounds.
a. Cl2(g) +
b. Br2(l) +
NaI(aq) →
NaI(aq) →
NaCl(aq) +
NaBr(aq) +
I2(s)
I2(s)
Suppose separate solutions, each containing 25.0 g of NaI, are available. If 5.00 g of Cl2 gas is bubbled into one
NaI solution, and 5.00 g of liquid bromine is added to the other, calculate the number of grams of elemental
iodine produced in each case.
a.
b.
21. If steel wool (iron) is heated until it glows and is placed in a bottle containing pure oxygen, the iron reacts
spectacularly to produce iron(III) oxide. Fe(s) O2(g) → Fe2O3(s) If 1.25 g of iron is heated and placed in a
bottle containing 0.0204 mol of oxygen gas, what mass of on(Ill) oxide is produced?
Unit8-StudHand1213.doc
8
Percent Yield
22. What is the actual yield of a reaction? What is the percent yield of reaction? How do the actual yield and the
percent yield differ from the theoretical yield?
23. The text explains that one reason why the actual yield for a reaction may be less than the theoretical yield is
side reactions. Suggest some other reasons why the percent yield for a reaction might not be 100%.
24. According to his pre-laboratory theoretical yield calculations, a student's experiment should have produced
1.44 g of magnesium oxide. When he weighed his product after reaction, only 1.23 g of magnesium oxide
was present. What is the student's percent yield?
25. The compound sodium thiosulfate pentahydrate, Na2S2O3⋅5H2O, is important commercially to the
photography business as "hypo," because it has the ability to dissolve unreacted silver salts from photographic
film during development. Sodium thiosulfate pentahydrate can be produced by boiling elemental sulfur in an
aqueous solution of sodium sulfite.
S8(s) +
8Na2SO3(aq) +
5H2O(l) →
8Na2S2O3⋅5H2O(s) (Balanced)
a. What is the theoretical yield of sodium thiosulfate pentahydrate when 3.25 g of sulfur is boiled
with 13.1 g of sodium sulfite?
b. Sodium thiosulfate pentahydrate is very soluble in water. What is the percent yield of the synthesis if
a student doing this experiment is able to isolate (collect) only 5.26 g of the product?
Unit8-StudHand1213.doc
9
26. Although they were formerly called the inert gases, at least the heavier elements of Group 8 do form
relatively stable compounds. For example, xenon combines directly with elemental fluorine at elevated
temperatures in the presence of a nickel catalyst.
Xe(g) + 2F2(g) → XeF4(s)
a. What is the theoretical mass of xenon tetrafluoride that should form when 130.g of xenon is reacted
with 100.g of F2?
b. What is the percent yield if only 145 g of XeF4 is actually isolated?
27. One process for the commercial production of baking soda (sodium hydrogen carbonate) involves the
following reaction, in which the carbon dioxide is used in its solid form ("dry ice") both to serve as a source
of reactant and to cool the reaction system to a temperature low enough for the sodium hydrogen carbonate to
precipitate:
NaCl(aq) + NH3(aq) + H2O(l) + CO2(s) → NH4Cl(aq) + NaHCO3(s)
Because they are relatively cheap, sodium chloride and water are typically present in excess. What is the
expected yield of NaHCO3 when one performs such a synthesis using 10.0 g of ammonia and 15.0 g of dry ice,
with an excess of NaCl and water?
Unit8-StudHand1213.doc
10
28. Before going to lab, a student read in her lab manual that the percent yield for a difficult reaction to be
studied was likely to be only 40 % of the theoretical yield. The student's pre-lab stoichiometric calculations
predict that the theoretical yield should be 12.5 g. What is the student's actual yield likely to be?
Challenge
29. The traditional method of analyzing the amount of chloride ion present in a sample was to dissolve the
sample in water and then slowly add a solution of silver nitrate. Silver chloride is very insoluble in water, and
by adding a slight excess of silver nitrate, it is possible effectively to remove all chloride ion from the sample.
Ag+(aq) + Cl-(aq) → AgCl(s)
Suppose a 1.054-g sample is known to contain 10.3% chloride ion by mass. What mass of silver nitrate must be
used to completely precipitate the chloride ion from the sample? What mass of silver chloride will be obtained?
30. When a certain element whose formula is X4 combines with HCl, the results are XCl3 and hydrogen gas.
Write a balanced equation for the reaction. When 24.0 g of hydrogen gas results from such a reaction, it is
noted that 248 g of X4 are consumed. Identify element X.
Unit8-StudHand1213.doc
11
31. When a 5.00-g sample of element X reacts completely a 15.0-g sample of element Y, compound XY is
formed. When a 3.00-g sample of element X reacts with an 18.0-g sample of element 2, compound XZ3 is
formed. The molar mass of Y is 60.0 g/mol. Find molar masses of X and Z.
Unit8-StudHand1213.doc
12