NAME____________________________________ PER____________ DATE DUE____________

ACTIVE LEARNING IN CHEMISTRY EDUCATION

"ALICE"

CHAPTER 9

THE MOLECONCEPT

(Part 2)

Unit ConversionsStoichiometry

Mass-Mass & Mass-Volume9-1 ©1997, A.J. Girondi

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© 1997 A.J. Girondi, Ph.D.505 Latshmere DriveHarrisburg, PA 17109

alicechem@geocities.com

Website: www.geocities.com/Athens/Oracle/2041

9-2 ©1997, A.J. Girondi

SECTION 9.1 The Law of Conservation of Matter

In an earlier chapter you experimented with the law of conservation of matter (or mass). The Lawstates that "matter is neither created nor destroyed in a chemical reaction." Therefore, the total mass ofthe reactants must equal the total mass of the products in any chemical reaction. In addition to provingthis law experimentally using real masses, it can also be demonstrated using atomic masses from theperiodic table. Problem 1 will illustrate this.

Problem 1. Follow the directions below.

a. Finish balancing the equation shown below.

3 NO2 + _____H2O -----> _____HNO3 + _____NO

b. Now, add the atomic masses of all the atoms on the reactant (left) side of the equation:

Total mass of reactants:__________

c. Next, add the atomic masses of all the atoms on the product (right) side of the equation:

Total mass of products:__________

Do your answers support the law of conservation of mass?__________ If not, check to see if you

balanced the equation properly. For the reaction 2 H2 + O2 ----> 2 H2O, if 2.80 g of H2 combine with

22.2g of O2, how many grams of water should be produced if the law of conservation of mass is true?

{1}__________

SECTION 9.2 Converting Grams to Moles and Moles to Grams

Using your knowledge of unit analysis, determine which equation below is correct for changingmoles to grams. Circle the correct equation:

moles X molesgrams

= grams OR moles X gramsmoles

= grams

Sample Problem: Change 7.65 moles of sulfur to grams of sulfur.

Solution:

7.65 moles S X 32.07 g S1 mole S

= 245.34 g S (rounds to 245 g)

9-3 ©1997, A.J. Girondi

Sample Problem: Change 1.43 moles of NO2 to grams of NO2.

Solution:

1.43 moles NO2 X 46.01 g NO2 1 mole NO2

= 65.8 g NO2

Problem 2. Use the equation you just circled to calculate the mass in grams of each of the following.Use the periodic table, and round atomic masses to the hundredths column. Show your work.

a. 3.00 moles Na X =

b. 3.50 moles CaCO3 X =

c. 1.50 moles Ba(OH)2 X =

d. 0.800 mole Fe3O4 X =

e. 0.00200 mole Pb(C2H3O2 )2 X =

Once again use your knowledge of dimensional analysis to determine which equation below iscorrect for calculating the number of moles. Circle the correct equation.

grams X

gramsmole

= moles OR grams X mole

grams = moles

Sample Problem: Convert 345 grams of Ca to moles of Ca.

Solution:

345 g Ca X 1 mole Ca 40.08 g Ca

= 8.61 moles Ca

Sample Problem: Convert 288 grams of NaNO3 to moles of NaNO3 .

Solution:

288 g NaNO3 X 1 mole NaNO3 85.01 g NaNO3

= 3.39 moles NaNO3

9-4 ©1997, A.J. Girondi

Problem 3. Use the equation you just circled to calculate the number of moles for each of the following.Use the periodic table, and round atomic masses to the hundredths column. Show your work.

a. 200. g F2 X =

b. 180. g Ca X =

c. 10.0 g Na2S X =

d. 68.6 g H2SO4 X =

e. 216 g P2O5 X =

Problem 4. This problem will give you additional practice working with mole problems. Use atomicmasses rounded to the hundredths column, use dimensional analysis, show your work, and assume allsituations involving gases are at STP. Show complete set-up for each problem.

a. What is the mass in gramsof 4.2 moles of Na2CO3

b. Calculate the number ofmoles in 14.5 g of C4H10.

c. Calculate the mass in grams of 0.00400moles of KMnO4.

d. How many moles are presentin 24.5 g of K2Cr2O7?

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e. Calculate the number of molesin 15.5 g of Ca3(PO4)2.

ACTIVITY 9.3 Determination of the Gram–Atomic Mass (GAM) of Silver

There are several methods for determining the gram atomic mass (gam) of an element. You willrecall that the gram-atomic mass of an element is the mass in grams of one mole of that element. In thisactivity, the gram-atomic mass of silver will be determined using a compound (silver oxide) of knowncomposition, Ag2O. Gather the following equipment: crucible and cover, ring stand, iron ring, claytriangle, crucible tongs, spatula, burner, balance, goggles or safety glasses, apron.

1. Clean a crucible and cover. Place the crucible in the clay triangle as shown in Figure 9.1 below. Heatthe crucible and cover for about 2 minutes. Be sure to tilt the cover as illustrated. Balance it carefully toavoid breakage. Put out the flame and use your crucible tongs to put the crucible and lid on the base ofthe ring stand to cool.

2. Measure the mass of the crucible + cover.Record this mass as (a) in Table 9.1.

3. Obtain about 1.75 g of silver oxide powder.Add this compound to the crucible. With thecover on the crucible, measure the mass of thecrucible and its contents. Record this mass as(b).

4. To remove oxygen from the oxide compound,tilt the cover as before and strongly heat thecrucible, cover and contents in the hottest part ofthe flame for 10 minutes. Allow the crucible tocool. The crucible should now contain only silvermetal. It may look white, but it is silver metal.Measure and record the mass of the crucible,cover, and silver metal (d).

Figure 9.1 Heating a Crucible

crucible

pipestemtriangle

lid

5. If time permits, reheat strongly for an additional 5 minutes. After cooling, again measure the mass of thecrucible, cover, and metal to check for constancy of mass.

9-6 ©1997, A.J. Girondi

Table 9.1The Atomic Mass of Silver Metal

a. mass of crucible + cover __________ g

b. mass of crucible + cover + Ag2O __________ g (this is the total mass before heating)

c. mass of oxide compound (Ag2O) used (b–a) __________ g

d. mass of crucible + cover + Ag metal __________ g (this is the total mass after heating)

e. mass of silver metal. (d–a) __________ g

6. Do not discard the silver in the crucible. Your instructor will have a container available in which you canput it. Try scratching the silver with a spatula or other metal object. This should reveal the shiny, silveryluster of the metal. If there is a magnifying glass or binocular scope in your lab, use it to observe the metalmore closely. If you have trouble removing the silver from the crucible, give it to your instructor. The silvercan be recycled for use in other experiments.

1. Find the mass of oxygen in your sample of Ag2O. (b–d)

__________ g O

2. Convert the number of grams of oxygen in your sample (from calculation 1 above) to moles of oxygen.(Oxygen was not in its diatomic form (O2) here, since it was combined with silver.) 16.00 g O = 1 mole O.

__________ mole O

3. Using the result of calculation 2 above, find the number of moles of silver metal in your sample ofcompound (Ag2O). Since the formula is known to be Ag2O, this means that the compound contains 2moles of Ag for every one mole of O. To find moles of Ag in your sample, just multiply the result ofcalculation 2 above, by two.

__________ mole Ag

4. Find the molar mass (g/mol) of the silver metal. (Divide the mass in grams of silver metal in Table 9.1 (e)above by the number of moles of Ag found in calculation 3 above.)

__________ g/mole

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5. Obtain the accepted value for the atomic mass of silver from the periodic table. Compare it to yourexperimentally determined value and determine your percentage of error. Check your reference materialsif you forget the formula for percent error. Show calculations below.

Accepted Value from Periodic Table = ____________ g/mole Ag = 1 gram-atomic mass (gam) of Ag

Observed (experimental) Value = ____________ g/mole Ag = 1 gram-atomic mass (gam) of Ag

Calculation of % Error. Show Work.

% Error = _____________

This experiment usually yields results with a higher percentage error than most of the activities you will

perform in this course. Nevertheless, this procedure should produce less than 15% error. Does your %

error fall within this range? __________ If not, how do you account for the excess error? ____________

______________________________________________________________________________

ACTIVITY 9.4 Balancing An Equation By Experiment

In this activity you will observe the reaction of iron nails with a solution of copper (II) chloride anddetermine the number of moles involved in the reaction. By converting grams of iron metal consumed andgrams of copper metal produced to moles, you will see how experimental data can be used to balance anequation.

Procedure:

1. Find the mass of a clean, empty, dry 250 mL beaker which you have marked with your initials. Recordthe mass to the nearest 0.01 g in Table 9.2 below.

2. Add about 8.00 g of copper (II) chloride crystals to the beaker (note that when a word like about is used,it means that you do not have to use exactly that amount. However, it is important to know precisely howmuch you are using.) Record the mass of CuCl2 in Table 9.1. Add about 50 mL of distilled water to thebeaker, and swirl or stir the solution to dissolve all of the solid.

3. Obtain two nails. Clean the nails using a piece of fine sandpaper. Find and record the mass of the twonails (together). If you have at least 25 minutes left in your class, place the nails into the solution and allowthe reaction continue for at least 20 minutes. Occasionally swirl the solution in the beaker. (If your classtime is short, you can allow the reaction to run longer, but do not allow it to continue for more than a fewhours.) During that time, you will see the formation of copper in the beaker. At the same time the iron nailswill be partially consumed. The solution will slowly change color from blue to green as some of the blueCuCl2(aq) is consumed and replaced by pale yellow FeCl3(aq). This is a single replacement reaction, andthe balanced equation is:

Fe(s) + CuCl2(aq) ----> Cu(s) + FeCl2(aq)

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4. After about 20 minutes, remove any remaining pieces of nail using your forceps or crucible tongs. Ifany copper is sticking to the nails, use your wash bottle to rinse as much of the copper as you can backinto the beaker. Dry the nails on a paper towel. When the nails are dry, find their mass and record in Table9.1. The nails can then be discarded.

5. Decant means to pour off only the liquid from a container that contains both solid and liquid. Carefullydecant the liquid from the solid. (No filtering is necessary.) Pour the liquid into another beaker so that incase you lose any solid, you can recover it. (You will always lose a little - don't worry about that.) Afterdecanting, rinse the solid copper in your beaker with some distilled water from your wash bottle. Decantthe rinse water. Repeat this rinsing three more times.

Table 9.2The Iron - Copper (II) Chloride Reaction

1. Mass of empty, dry, labeled beaker __________ g (Before the Reaction)

2. Mass of beaker + CuCl2 __________ g

3. Mass of two iron nails(before reaction) __________ g

4. Mass of two iron nails(after reaction) __________ g

5. Mass of beaker + dry copper __________ g

6. Mass of copper produced (5 – 1) __________ g

6. Next, wash the solid copper with about 25 mL of 1.0 M hydrochloric acid (HCl). Decant again, and rinseonce more with distilled water. Place the beaker containing the copper in an oven to dry until the next day.

7. When the copper is completely dry, find and record the mass of the beaker plus the copper. Scrapethe copper out of the beaker and discard it in the waste can. Clean and return all equipment. Perform thefollowing calculations, rounding atomic masses to the hundredths column. Show your work.

Calculations:

1. Referring to items 3 and 4 in Table 9.2, calculate the mass in grams of iron used in the reaction.

2. Using your answer to calculation 1 above, calculate the number of moles of iron used in the reaction.

3. Using items 1 and 5 in Table 9.2, calculate the number of grams of copper metal produced and enterthe result in the data table.

9-9 ©1997, A.J. Girondi

4. Using your answer to calculation 3 above, determine the number of moles of copper metal produced.

5. Use your results from calculations 2 and 4 above to determine a simple whole–number ratio of moles ofiron reacted to moles of copper produced. (Hint: divide each quantity of moles by the smallest of the twoand round to the closest whole number.) If this confuses you , consider the following. Suppose youhave a ratio such as 0.64 to 0.16. To change this to an equivalent whole-number ratio we will divide eachnumber the the smallest of the two which is 0.16. Since 0.64 ÷ 0.16 = 4, and since 0.16 ÷ 0.16 = 1, theequivalent whole-number ratio is 4 to 1.

Experimental whole–number ratio of Fe to Cu is: _______ to _______

6. The balanced equation for this reaction was shown previously in the activity. What Fe to Cu mole ratio ispredicted by the balanced equation for this reaction? ______ to ______

7. The equation for this reaction is Fe(s) + CuCl2(aq) ----> Cu(s) + FeCl2(aq) How does your experimentalratio compare to the Fe – Cu mole ratio predicted by the equation?

______________________________________________________________________________

SECTION 9.5 Introduction To Mass-Mass Problems

You have seen that the coefficient numbers in a balanced chemical equation may represent therelative number of moles or the relative number of molecules involved in a reaction. These reacting ratiosdo not change. The reaction below is a balanced chemical equation. Two moles of Li2O decompose toform four moles of Li and one mole of O2. Note the 2:4:1 mole ratio in the following equation:

2 Li2O ----> 4 Li + O2

The ratio of moles of reactant (Li2O) to moles of Li produced is 2:4 in this balanced equation. Knowingthis relationship can help to solve a wide variety of problems such as the next sample problem.

Sample Problem: Referring to the equation above, determine the number of moles of Li producedwhen 3.00 moles of Li2O decompose.

Begin this problem by writing the information given and the information you are hoping to find as follows:

3.00 moles Li2O X ? ? ?

? ? ? = ???? moles Li

Once you have the basic framework of your equation, you simply build the rest of the equation using dataoffered in the problem or from the balanced equation. Enter the data in the correct positions, so that yourunits cancel properly and leave you with only the units you want in the answer:

9-10 ©1997, A.J. Girondi

4:2 ratio from balanced equation

3.00 moles Li2O X 4 moles Li

2 moles Li2O = 6.00 moles Li

Because these problems begin with a mass (grams or moles) and end with a mass (grams or moles), theyare sometimes called "mass-mass" problems. Notice that complete units are very, very important. Wedidn't just use moles, but we used moles Li and moles Li2O. In chemistry, when you do calculations whichrelate quantities of two substances used or produced in a chemical reaction, you are doing what is knownas stoichiometry. Stoichiometric problems give you information about one substance, and ask forinformation about a different substance in a chemical reaction. For this reason, you need to refer to thechemical equation for the reaction. Most of the problems in the rest of this chapter involve stoichiometry.Your calculations should always include complete units. Your teacher will be looking for them! Try thisnext problem, yourself. Show your work.

Problem 5. Using the equation from the last sample problem, determine the number of moles of O2

produced when 3.00 moles of Li2O decompose.

3.00 moles Li2O X

=

Other varieties of mass-mass problems are shown below. Examine the examples and then try theproblems which follow.

Sample Problem: In the reaction 2 H2 + O2 ----> 2 H2O, how many grams of H2O will be formed if we use4.5 grams of O2?

Note that you are asked to change information about O2 into information about H2O. (Change 4.5grams O2 into grams of H2O) Hmmmm. This is stoichiometry. In order to change information about one

substance into information about another substance, you must first know a relationship between thosetwo substances. Where do we find such a relationship? In the equation for the reaction! Whatrelationship do we already know between O2 and H2O? Well, thanks to the balanced equation, we knowthat 1 mole of O2 is needed to produce 2 moles of H2O. The relationship that we already know is a moleratio. Therefore, in order to change information about O2 into information about H2O, we must do it in theterms of the relationship that we already know. In other words, we must do it in terms of moles in this case.

Since we cannot change grams of O2 directly into grams of H2O, we will first change the 4.5 g of O2

into moles of O2. Then, we can change moles of O2 into moles of H2O. Finally, we will change moles ofH2O into grams of H2O. This will all be done using unit analysis. Note that the part of the solution wherewe change from one substance to the other (O2 to H2O) occurs, is the mole-to- mole ratio:

Relating Two Different Substances

Relationship From Balanced Equation unit conversionunit conversion

4.5 g O2 X 1 mole O232.00 g O2

X 2 moles H2O1 mole O2

X 18.02 g H2O1 mole H2O

= 5.1 g H2O

9-11 ©1997, A.J. Girondi

Remember, in order to establish a relationship between two substances, you must first know a relationshipbetween them. Balanced chemical equations provide such relationships in the form of mole-to-moleratios. Thus, problems which change information about one substance to information about anothersubstance must make use of the mole-to-mole ratio in the balanced equation. If the information given inthe problem is not in the form of moles, it has to be changed to moles inorder to then use the mole to moleratio. Thus, the mole to mole ratio is the "heart of" a stoichiometry problem. When the problem is set up,there may be any number of "unit conversions" before or after the mole-to-mole ratio, depending on theinformation given and requested in a particular problem. Note the two unit conversions used in theillustration above. The diagram in Figure 9.2 outlines the process. There are many other unit conversionsbesides the ones listed in Figure 9.2.

The problem below is simpler than the last sample problem since the information is given in moles. Fill inthe missing parts of the set-up so that all units divide out except grams of Al2O3 and solve the problem.

Problem 6. In the reaction: 4 Al + 3 O2 ---> 2 Al2O3 how many grams of Al2O3 will be formed if we startwith 5.0 moles of Al?

5.0 moles Al X moles Al

X 1 mole Al2O3

= __________ g Al2O3

Use as many unit conversions as needed to change the starting information to moles.

Use mole-to-mole ratio in the balanced equation to convert to moles of the substance in the answer.

Is the information you are going to start with given in moles?

NO

YES

Do you want your answer to be expressed in moles?

YES

You are Done!

NO Use as many unit conversions as needed to change moles to the units requested in the answer.

SOME UNIT CONVERSIONS1 g-atomic mass = 1 mole1 g-formula mass = 1 mole22.4 L = 1 mole gas @ STP6.02 X 1023 atoms = 1 mole6.02 X 1023 molecules = 1 mole

START

Stoichiometry?

Figure 9.2 Flow Diagram for Solving Stoichiometry Problems

9-12 ©1997, A.J. Girondi

Solve the following problems using dimensional anaylsis. All measurements must include units!

Problem 7. In the reaction: C + 2 Cl2 ----> CCl4, how many grams of CCl4 can be prepared starting with8.6 grams of Cl2?

Problem 8. In the reaction: FeCl3 + 3 NaOH ---> Fe(OH)3(s) + 3 NaCl, how many grams of Fe(OH)3 canbe produced if we start with 0.50 moles of NaOH?

Problem 9. In the reaction 2 Al + 3 Pb(NO3)2 ---> 2 Al(NO3)3 + 3 Pb, how many grams of Al are neededto produce 68 grams of Pb?

Problem 10. In the reaction in problem 9 above, how many moles of Pb will be formed if 23.2 grams ofAl(NO3)3 are formed?

Problem 11. In the reaction: BaCO3 ----> BaO + CO2, how many grams of CO2 can be produced fromthe decomposition of 45 grams of BaCO3?

9-13 ©1997, A.J. Girondi

ACTIVITY 9.6 The "Silver Tree" Reaction

The purpose of this activity is to use mass-mass calculations to predict the outcome of anexperiment. In order to accomplish this, you will be measuring the masses of both the reactants and theproducts. A single replacement reaction is involved in which copper replaces the silver in silver nitrate:

Cu(s) + 2 AgNO3(aq) ----> Cu(NO3)2(aq) + 2 Ag(s)

Notice that one of the products will be pure silver metal. This activity requires that you perform a number oflaboratory procedures over a period of 3 days.

Procedure for Day 1: (About 15 minutes required)

1. Obtain a clean, dry 100 or 150 mL beaker and use a magic marker to label it as "Beaker A" and includeyour name(s).

Figure 9.3 Copper Coil in Beaker

2. Weigh the beaker and record the mass in the Table9.3. (Note: all masses in this activity should bemeasured to the nearest 0.01 or 0.001 g dependingon the balances available. Use the most precisebalance available, and use balances of the sameprecision throughout the activity.)

3. Measure about 25 to 30 centimeters of copper wireand coil it around a test tube. Measure and record itsmass.

4. Leave some wire uncoiled to make a hook which willfit over the edge of the beaker. (See Figure 9.3)

5. Measure between 2 and 3 grams of silver nitratecrystals. Record the exact mass of the crystals.Caution: silver nitrate crystals can create dark stains orspots on your skin and clothing. The stained area willnot darken for several hours. Use care in handlingAgNO3. Clean up any spilled crystals with a damptowel.

6. Fill beaker A about two-thirds full of distilled water. Add the AgNO3 crystals. Stir with a glass stirring roduntil most of the crystals have dissolved. The coils of your copper wire should not be too close together.When the coil is immersed in the solution, most of the coils should be submerged. (See Figure 9.3)

7. Place the copper coil into the solution and observe the beginning of the reaction. Store the beaker inyour drawer until the next day.

Procedure for Day 2: (About 10 minutes required)

8. Lift the copper wire from beaker A and shake off any silver residue which has formed. Rinse the wireusing a wash bottle with distilled water, allowing the rinse water to fall into beaker A. Set the wire aside todry until the next day. DO NOT discard the wire until you have weighed it and recorded its mass!

9. Obtain a second clean, dry 100 or 150 mL beaker. Decant most of the solution in beaker A into thesecond beaker. Try not to lose too many of the solid silver crystals as you decant the solution. You do not

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need to decant all of it. Discard the solution in the second beaker rinse it with water.10. Squirt a little distilled water over the silver metal in beaker A. Decant most of the rinse water into thesink, retaining the silver metal.

11. Place beaker A containing the wet silver metal into a warm drying oven until the next day. Any waterleft in this beaker will evaporate.

Procedure for Day 3: (About 10 minutes required)

12. Weigh the dry copper wire, which can then be discarded in the waste can. Record the mass in the datatable.

13. Weigh beaker A (with contents) and record. Empty the silver metal in beaker A into the containerprovided by your instructor. It can be purified and used in other experiments. Remove any markings fromthe beaker with alcohol and a towel.

Table 9.3The “Silver Tree” Reaction

Item(s) Mass in Grams

1. Beaker A __________

2. Copper Wire (day 1) __________

3. Copper Wire (day 3) __________

4. Copper Metal Reacted (2-3) __________

5. Silver Nitrate (day 1) __________

6. Beaker A + Silver Metal (day 3) __________

7. Silver Metal Formed (6-1) __________

Calculations:

1. Here is where your expertise in mass-mass problems will come in handy! Based on the number ofgrams of copper that were consumed (item 4 in Table 9.3), predict how many grams of silver should havebeen produced. Show work.

__________ g Ag should be produced

2. The actual amount of silver produced is item 7 in Table 9.3. Is it close to your predicted value incalculation 1 above? __________ Calculate your percentage error using the formula for %E. You canfind it in your reference notebook. The accepted value (A) is your prediction from calculation 1. Theobserved value is item 7 in Table 9.3. Show work.

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3. Calculate the number of moles of copper which reacted. Be sure to use as many significant figures asyou are entitled to. Show work.

4. Calculate the number of moles of silver produced. Be sure to use as many significant figures as you areentitled to.

5. Using the answers to calculations 3 and 4 above, determine the whole number ratio of moles of Cuused to moles of Ag produced. Do this by dividing moles of copper (calculation 3 above) and moles ofsilver (calculation 4 above) by the smallest of the two values. Show work.

According to the balanced equation for this reaction, what should the mole to mole ratio be for Cu to Ag?

______ Cu(s) + 2 AgNO3(aq) ----> Cu(NO3)2(aq) + ______ Ag(s)

How does your experimentally–determined Cu:Ag ratio in calculation 5 compare to that in the balanced

equation?______________________________________________________________________

SECTION 9.7 Introduction to Mass–Volume Problems

The important unit conversion which you know that relates mass to volume (moles to liters) is:1 mole of a gas = 22.4 L @ STP. In mass-volume problems, the same basic procedure is followed nomatter how complex a problem may become. The key to success is to always include all units when youset up the problem. In that way, you can be sure to set it up so that all units cancel properly.

Problem 12. Determine the volume in liters (at STP) of nitrogen gas, N2, needed to produce 92.0

grams of product, NO2, in the reaction: N2(g) + 2 O2(g) ----> 2 NO2(g)

The two relationships which you need to use which are not given in the problem are:

22.4 L = 1 mole gas and NO2 = 46.0 g/mole.

Furthermore, the ratio of moles of N2 to moles of NO2 = 1:2. The basic framework has been drawn for youbelow. (We can call this framework a "fencepost" because it begins to look a bit like one. In that event,doing dimensional analysis is called "fenceposting.") The problem has been partially completed for you

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below. You should now fill in the blanks and calculate the answer. Make sure to use complete units and tobe certain that all units divide out except those you want in the answer. (Problems which attempt toestablish a relationship between a mass and a volume are called mass-volume problems.)

92.0 g NO2 X g NO2

X moles N2 X

moles N2 = _______ L N2

The next four problems refer to the following equation. Use the same approach as you did to solve theprevious problems. Show all of your work, including complete units. Assume STP for gases. 2 KClO3(s) ----> 2 KCl(s) + 3 O2(g)

Problem 13. How many moles of KClO3 are required to form 2.5 liters of O2?

Problem 14. How many grams of KClO3 are needed to produce 11.2 liters of O2?

Problem 15. How many liters of O2 are formed when 367.8 g of KClO3 decompose?

Problem 16. How many moles of KCl are produced when 2.3 L of O2 are produced?

SECTION 9.8 A Variety of Stoichiometry Problems To Solve

Use the equation below to solve problems 17 through 20. Assume STP conditions for any gases. Showyour work, and include units with all measurements.

3 H2(g) + N2(g) ----> 2 NH3(g)

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Problem 17. How many moles of ammonia (NH3) would be produced by reacting 6.00 moles ofhydrogen (H2) with an adequate amount of nitrogen (N2)?

Problem 18. If 56.0 grams of N2 are reacted with a sufficient amount of H2, how many moles of product(NH3) are formed?

Problem 19. How many grams of NH3 are formed when 5.60 liters of N2 react with a sufficient amount ofH2?

Problem 20. How many liters of NH3 will be formed when 3.01 X1023 molecules of H2 react with asufficient amount of N2?

The following problems involve all of the information you have gathered thus far about moles. Use

unit analysis to set up these problems and to be sure your units cancel properly. Show all your work, andassume STP conditions for any gases.

Problem 21. What is the volume in liters occupied by 10.0 grams of CO2 gas at STP? (Hint: you mustfirst find the molecular mass of CO2.)

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Problem 22. If 1.00 mole of a gas has a mass of 18.0 grams, and you have 15.0 grams of the gas, whatwould its volume be at STP?

Problem 23. If the molecular mass of a gas is 34.0 g/ 1 mole, what would be the mass of 30.0 L at STP?

Problem 24. What is the volume in liters of 6.40 g of O2 gas at STP?

Problem 25. A gas sample with a mass of 50.0 grams has a volume of 40.0 L. What is the gram-molecular mass (GMM) of the gas? (GMM is epxressed in units of grams/mole.) Hint: This problem is a bitdifferent form the others. Notice that you are given grams and liters but asked for grams/mole. Why notstart your set-up with 50.0 grams/40.0 liters and then just convert liters to moles?

Problem 26. Calculate the gram-molecular mass, GMM, (in g/mole) of a gas if 11.5 g of the gas has avolume of 5.60 L at STP? (You want to end the problem with units of grams/moles.)

Problem 27. Is the gas mentioned in problem 26 above H2O, SO2, or NO2? Show how you know.

9-19 ©1997, A.J. Girondi

SECTION 9.9 Learning Outcomes

By now you must realize that the mole is a central concept in chemistry and is very useful whenworking with chemical equations. When you are certain you understand each one, place a check mark inthe space preceding each learning outcome below. Then, move on to chapter 10.

_____1. Use dimensional (unit) analysis to make conversions between moles, grams, liters, and molecules.

_____2. Calculate mole–mole relations from balanced chemical equations.

_____3. Calculate mass–mass relations from balanced chemical equations.

_____4. Calculate mass–volume relations from balanced chemical equations.

_____5. Use the molar volume concept in problems involving gases.

9-20 ©1997, A.J. Girondi

SECTION 9.10 Answers to Questions and Problems

Questions: {1} 25.0 g

Problems:

1. a. 3, 1, 2, 1; b. 156.05 amu; c. 156.05 amu2. a. 68.97 g (69.0 rounded); b. 350.32 g (350. rounded); c. 257.03 g (257 rounded)

d. 185.24 g (185 rounded); e. 0.651 g3. a. 5.26 moles; b. 4.5 moles; c. 0.128 mole; d. 0.699 mole; e. 1.52 moles4. a. 445.16 g rounds to 450 g; b. 0.249 mole; c. 0.632 g; d. 0.0833 mole; e. 0.0500 mole5. 1.50 moles O2

6. 254.9 g Al2O3 (250 g rounded)7. 9.3 g CCl48. 18 g Fe(OH)39. 5.9 g Al10. 0.163 mole Pb11. 10. g CO2

12. 22.4 L N2

13. 0.074 mole KClO3

14. 40.9 g KClO3

15. 100.8 L O2

16. 0.068 mole KCl17. 4.00 moles NH3

18. 4.00 moles NH3

19. 8.52 g NH3

20. 7.47 L NH3

21. 5.09 L CO2

22. 18.7L23. 45.5 g24. 4.48 L O2

25. 28.0 g/mole26. 46.0 g/mole27. NO2

9-21 ©1997, A.J. Girondi

SECTION 9.11 Student Notes

9-22 ©1997, A.J. Girondi