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Unit 3 Packet: The Mole NameWPHS Chemistry

Unit 3

The Mole

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Chemistry: Unit 3 Outline: The MoleAssignment WB Page Number Score Out ofPodcast 3.1 (CB 1-5) Online Lab: Pennies pg 5-6 100Worksheet A pg 13 Podcast 3.2 (CB 7-9) Online Lab: Mole Lab 100Worksheet B (long) pg 14-17 Mole Video (Dr. Don) pg 12 Demo: Measure out 1 mole of NaCl, H2O, NaHCO3 and show to your teacher

In class

Podcast 3.3 (CB 11-13) Online Worksheet C Pg 17-18 Worksheet D Pg 19-20 Technology Lab Teacher Handout 100Lab: Magnesium and Oxygen Formula

Pr 8-11 100

THL 2: amount of Oxygen in Air—do at home for parents

Pg 4 100

Review Pg 22-23Unit 2 Exam (You must score 85/100 to move to the next unit)

In Class 100

(You must score 85/100 on all assignments with a number to move to the next unit. For those assignments with a check, you need to do it to the satisfaction of your teacher)

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Visualizing Chemistry Notebook Set-Up: Chapter 7: Chemical Quantities7-1: The Mole: A Measurement of Matter?1 What is the mole (pg 171-173)

Why was the mole unit created? What is Avagadro’s number and why is it so big?

7-2: Mole Mass and Mole Volume Relationships3 Molar Mass—The Mass of one Mole (page 176-183)

Define molar mass Read the sections on the mass of one mole and fill in this chart

Name Formula Show Work Molar mass (g/mol)

Iron II hydroxideAluminum sulfateSodium bicarbonateDinitrogen pentoxide

5 The Volume of a mole of gas Write out the conditions for standard temperature and pressure

7 Converting between moles-grams-particles-litersLeave this page blank for teacher directed notes

9 Converting between moles-grams-particles-litersLeave this page blank for teacher directed notes

7.3: Percent Composition and Chemical Formulas11 Percent composition (Continued from previous page) (188-190)

Define Percentage composition Use sample problem 7-10 on page 189 to complete the table below

Chemical Molar Mass Percent 1st Element Percent 2nd Element Percent 3rd Element

FeSO4

NaClO3

K3PO4

C6H12O6

Use sample 7-12 on page 191 to do practice problem 34 on page 19213 Derive a compound’s empirical formula from its percentage composition (page 192-193)

Define empirical formula Copy down the following rhyme:

o Percent to masso Mass to moleo Divide by smallo Times ‘til whole

15 Calculate the molecular formula of a compound from its empirical formula and from its formula mass (pg 194-195)

Define molecular formula Finish the following statement: Molecular formulas are _______ of empirical formulas Copy down Table 7.2 on page 194

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Take Home LabParent/Student ExperimentTitle: Amount of Oxygen in Air

Subject/Concept: Chemistry - Combustion and Percent Composition

Purpose: The purpose of this activity is to observe the amount of oxygen required for a candle to burn in a closed environment and then to calculate the amount of oxygen originally present.

Materials: • small candle (emergency candles work well)• food coloring• shallow pan filled two-thirds with water• glass jar

Procedure:1. Attach the candle to the bottom of the shallow pan. A few drops of hot wax work nicely.2. Place a piece of masking tape on the side of the jar from the top to the bottom. This will be used to

mark water levels during the experiment.3. Fill the dish two-thirds with water and color the water. 4. Place the jar over the UNLIT candle, and mark the starting height of the water. 5. Light the candle and cover it quickly with the jar. Observe.6. Mark the ending height of the water on the masking tape, and repeat the experiment several more

times. 7. Measure the original height of the air column in the glass and the ending heights of the air column.

Average the ending heights.8. Calculate the amount of air (oxygen) used by the candle. This will be the difference between your

starting and ending heights of air.9. Calculate the percent of air that is oxygen. This will be the amount of oxygen used divided by the

amount of total air to start x 100.

Questions:1. What was the height of air in the jar at the beginning of the experiment?2. What was the average height of air in the jar after the burn?3. What was the difference in heights?4. What was the percent of oxygen that you obtained? Compare this to the actual value of 21%.5. Why does the water rise in the jar?

For Credit:To receive credit, your parent or guardian must write a short note confirming that you performed

the experiment for them and explained the results to their satisfaction using the concept of combustion. Attach your note to the back of this sheet.

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PENNY COUNTING BY WEIGHING

PURPOSE: To make a model of counting by weighing.

MATERIALS: A handful of pennies, a balance

PROCEDURE:1. Determine the average mass of a penny by weighing 25 pennies and dividing the total mass by 25.2. Repeat step 1 two more times with different pennies, and take the average of your three results.3. Weigh about three-fourths of you total number of pennies. 4. Calculate how many pennies you weighed.5. Count the number of pennies in your sample and compare that to the number you calculated in steps 3 & 4.6. Repeat steps 3,4 and 5 with a different sample size.

DATA TABLE:

QUESTIONS:1. Did the number of pennies you counted in the sample (step three) equal the number you calculated by weighing (step two)? If there was not agreement, propose an explanation.

2. Explain how you would use the balance to “count out” 185 pennies.

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3. What is the advantage of using a larger sample size in step 1? What is a disadvantage?

4. How are the pennies like atoms in this experiment?

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Mole Lab: Making a precipitate of BaSO4

Background: The unit of measure for the amount of a substance is the mole. A mole is 6.02 x 1023 particles (such as atoms or molecules). The mass of one more of any substance is found using the periodic table (calculate molar mass).

Purpose: The purpose of this lab is to measure chemical quantities.

Materials: small test tubes, balance, weigh boat, BaCl2, Na2SO4, H2O, centrifuge, graduated cylinder

Calculations:Convert moles of barium chloride to grams

Convert moles of sodium sulfate to grams

Procedure:1. Weigh 0.0025 mole of barium chloride and place in one test tube2. Weigh 0.0035 mole of sodium sulfate and place in another test tube3. Add 3 mL of water to each test tube 4. Agitate (shake or tap) each mix until the chemical inside it dissolves5. After both substances have completely dissolved combine them into one of the

test tubes (this will create a milky solution)6. Place the test tube in the centrifuge, making sure that another group’s test tube is

directly across from your test tube and allow the centrifuge to run for 1 minute7. Bring your test tube to your teacher to check the amount of precipitate and sign

your group’s papers

Questions:1. What is the evidence that a chemical reaction has occurred?

2. When you mixed the two solutions, you created BaSO4. What is the name of this chemical?

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MAGNESIUM OXIDE PRODUCTION

Objectives: In this investigation, you will prepare magnesium oxide, calculate the percent composition of your product, and determine the product’s empirical and molecular formulas.

Equipment:crucible and cover tongsBunsen burner clay trianglering stand iron ringwash bottle glass stirring rod15 cm Mg ribbon analytic balancedistilled water GOGGLES

Procedure:1. Wash and dry your hands (moisture on your hands will react with the magnesium

ribbon).2. Record the mass of a clean, dry crucible and cover. 3. Obtain a piece of magnesium ribbon approximately 15 cm long from your

instructor and scrape both sides of the magnesium with the scissor blade to remove corrosion (when the corrosion is removed the Mg will appear shiny). Coil the Mg loosely around a pencil. Remove the pencil, place the magnesium in the crucible and record the mass of the magnesium, crucible and lid.

4. Place the crucible, cover and magnesium on a clay triangle as shown in the figure.

5. Adjust the crucible and cover on the clay triangle so that the lid is ajar. This position will allow a steady flow of air into the crucible. Heat the crucible gently

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for 3 minutes then strongly for 5 minutes. CAUTION: Do not look directly at the burning magnesium. Allow the crucible to cool for 5 minutes.

6. Crush the contents of the crucible into powder using a glass stirring rod. Using about 15 drops of water from a wash bottle, rinse powder from the the stirring rod into the crucible. Note any odor as you add the water. Odor:___________________

7. Heat the crucible strongly, without the cover, for 5 minutes to dry the residue. Cool the crucible and contents for 5 to 10 minutes. It should be cool enough to NOT burn your hand. Record the mass of the crucible, cover, and contents.

ANALYSIS: Show all measurements and calculated numbers in the spaces provided in the data table. (Show units and substance symbol or formula for each measurement taken or number calculated.)

Measurement

mass of crucible, cover, and Mg before heating

mass of empty crucible and cover

mass of magnesium

mass of crucible, cover and residue after heating

mass of residue (magnesium oxide produced)

moles of oxygen in the magnesium oxide residue

moles of magnesium in the magnesium oxide residue

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Calculations:Calculate the percent composition of magnesium oxide (write the formula, balancing the charges of the Mg ion and the O ion, then find the percent composition by mass).

From your measured value of the mass of residue in your data table, calculate the moles of magnesium and of oxygen in your sample of residue.

Using the moles calculated above; calculate the empirical formula for your sample of magnesium oxide.

Using the empirical formula found above, what is its molecular formula if the molar mass is 40.3 g/mol?

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Questions

1. Does the magnesium in your crucible gain or lose mass? Explain.

2. How would your final ratio change if not all of the magnesium had reacted?

3. How would your final ratio change if there were still some water in the crucible after you stopped heating it?

4. Does your calculated formula for magnesium match the predicted formula (questions #6)?

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The World of Chemistry: Episode 11 - The Mole

1. Why is it important to use the correct amount of materials in a chemical reaction?

2. What names are given to the materials at the beginning and end of a chemical reaction?

3. Atoms and molecules are extremely small. How do chemists "count" them? Can you think of an everyday application of this?

4. a. What did early chemists discover about reactions involving the combination of gases?

b. How did Avogadro explain this?

5. How may a chemical equation such as H2 + Cl2 2 HCl be interpreted?

6. What is true about the mass of a compound?

7. What is the numerical value for Avogadro's Number?

8. When the I V solutions were prepared, quality control was involved. What is quality control?

9. Why did using twice as much magnesium not produce twice as much hydrogen in the demonstration?

10. What ratio of starting materials was found to produce the best epoxy resin?

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Mole WorksheetMolar RelationshipsPart A:Find the molar mass of the following compounds: Show work and include units

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Name ______________Period _______

1. CO2

2. Fe2O3

3. AgCl

4. Ca3(PO4)2

5. W3(PO3)5

6. Fe(C2H3O2)2

7. Calcium Carbonate

8. Lead IV Sulfate

9. Lead IV Sulfite

10. Lead IV Sulfide

11. Lead II sulfate

12. Lead II Sulfite

13. Lead II Sulfide

14. Copper I Sulfide

15. Copper II Sulfite

WPHS Chemistry: Unit 3 Packet: The Mole Name________________

Worksheet B:Directions: Answer the following questions. Set-up all problems using the factor-label method of dimensional analysis and show all your work and units.

1. What is the mass of 7.50 moles of sulfur dioxide (SO2)?

2. How many moles are there in 250.0 grams of sodium phosphate (Na3PO4)?

3. How many grams of potassium sulfate (K2SO4) are there in 25.3 moles?

4. How many atoms are in 1.5 moles of neon?

5. How many moles of SF6 are there in 4,595,000,000,000,000,000 molecules of SF6?

6. What is the volume of 0.38 moles of any gas at STP?

7. Calculate the number of moles in 32.2-L of NH3

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8. Calculate the number of moles in 5.45x1025 atoms of Zn

9. Calculate the number of grams in 3.25-mol of AgNO3

10. What is the mass of 51 liters of oxygen gas?

11. What volume would be occupied by 9.45 x 1024 molecules of CO2 gas at STP?

12. How many calcium atoms would be in a 100 gram sample of calcium metal?

13. How many grams are in 5.6 x 1023 atoms of Zinc?

14. Calculate the number of molecules in 4.56-g of Pb(NO3)2

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15. Calculate the number of liters in 3.25-g of NH3

16. Calculate the number of liters in 5.43x1025 molecules of H2

17. Calculate the number of grams in 3.54-L of CO2

18. Calculate the number of grams in 9.7x1022 molecules of CH3CH2OH

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Part C: Percentage Composition1. Calculate the % composition of Li2O.

2. What is the percentage composition of a carbon-oxygen compound, given that a 95.2 g sample of the compound contains 40.8 g of carbon and 54.4 g of oxygen?

3. What is the percentage composition of N2O4?

4. What is the percentage composition of a compound made from 28 grams of nitrogen and 32 grams of oxygen?

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5. What is the percentage composition of a carbon-hydrogen-fluorine compound which contains 7.2 grams of carbon, 11.4 grams of fluorine, and 1.8 grams of hydrogen?

6. Find the percentage composition of Na2SO4?

7. If a compound is formed from 60.0 liters of nitrogen gas, N2, (at STP) and 180 liters of hydrogen gas, H2, (at STP), what is its percentage composition?

8. Find the percentage composition of a compound formed when 0.4 moles of potassium are reacted with 8.96 liters of O2 gas and 2.41 x 1022 atoms of S.

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Part D: Empirical and Molecular Formulas1. Determine the empirical formula of a compound with 72.4% Fe and 27.6%

Oxygen.

2. Determine the empirical formula of a compound with 65.2% Sc and 34.8% O

3. Determine the empirical formula of a compound with 52.8% Sn, 12.4% Fe, 16% C and 18.8% N.

4. Determine the empirical formula of a compound that contains 2.61-g of carbon, 0.65-g of hydrogen, and 1.74-g of oxygen

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5. Determine the molecular formula for a compound that contains 12.2-g Nitrogen, 27.8-g Oxygen, and a molecular mass of 92.0 g/mol.

6. Determine the molecular formula for a compound that contains 94.1% oxygen and 5.9% hydrogen and a molecular mass of 34 g/mol.

7. Determine the molecular formula for a compound that contains 22.5% Na, 30.4% P and 47.1% O and a molar mass of 306 g/mol

8. Determine the molecular formula of a compound that contains 76% iodine and 24% oxygen and has a molar mass of 334g/mol.

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9. Determine the molecular formula of a compound that contains 48.6% carbon, 8.1% hydrogen, and 43.2% oxygen and has a molar mass of 296-g/mol.

10. Determine the molecular formula of a compound that contains 0.993-g nitrogen, 1.27-g carbon, 0.213-g hydrogen, 2.52-g chlorine and has a molar mass of 423-g/mol.

11. A sample of TNT, a common explosive is analyzed and found to contain 1.03-g of nitrogen, 0.220-g hydrogen, and 1.76-g of carbon. The molar mass is 123 g/mol. What is the molecular formula?

12. Azobenzene is an important intermediate in the manufacture of dyes. It contains 79.1% carbon, 5.55% hydrogen, and 15.4% nitrogen. It has a molar mass of 182-g/mol. What is the molecular formula?

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Unit 3 ReviewBe sure to show all of your work. Make sure that you box your answer and that you have units!1) Determine the molar mass of the following compounds.

a) Nitrogen dioxide

b) NH4NO3

2) Convert the following:a) 4.53 moles of carbon monoxide to grams

b) 0.0067 L of chlorine gas (Cl2) at STP to moles

c) 2.41 1024 molecules of (NH4)SO3 to moles

3) Convert the following:a) 20.6 L of SO2 to grams

b) 4.44 g of iron (II) oxide molecules

c) 8.322 x 1024 molecules of N2 to L

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4) Determine the percent composition of CuSO4

5) Determine the empirical formula of a sample that has: 21.6% sodium (Na), 33.3% chlorine (Cl), and 45.1% oxygen (O)

6) Determine the molecular formula of the following compound:Nitrogen = 30.4% Oxygen=69.6%Molecular weight = 92 g/mol

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Empirical Formula

Empirical Formula

Molecular Formula


Mole Worksheet KEY Name____________________Molar Relationships Period____Directions: Answer the following questions on a separate sheet of paper. Set-up all problems using the factor-label method of dimensional analysis and show all your work and units.

1. What is the mass of 7.50 moles of sulfur dioxide (SO2)?480g

2. How many moles are there in 21.4 grams of nitrogen gas (N2)? 0.764mol

3. How many moles are there in 250.0 grams of sodium phosphate (Na3PO4)? 1.52mol

4. How many grams of potassium sulfate (K2SO4) are there in 25.3 moles? 4402g

5. How many atoms are in 1.5 moles of neon? 9.0x1023atoms

6. How many moles of SF6 are there in 4,595,000,000,000,000,000 molecules of SF6? 7.633x10-6mol

7. How many molecules are there in 7.50 moles of sulfur dioxide (SO2)? 4.52x1024molec

8. What volume is occupied by 7.50 moles of sulfur dioxide gas (SO2) at STP? 168L

9. 49.28 L of oxygen gas is how many moles of gas?2.20mol

10. What is the volume of 0.38 moles of any gas at STP? 8.5L

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11. What is the mass of 51 liters of oxygen gas? 73g

12. What volume would be occupied by 9.45 x 1024 molecules of CO2 gas at STP? 351L

13. How many calcium atoms would be in a 100 gram sample of calcium metal? 1.5x1024atoms

14. How many grams are in 5.6 x 1023 atoms of Zinc? 60.g

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Mole WS 2Two Step Problems

1. Calculate the number of moles in 45.5-g of NH4Cl

2. Calculate the number of moles in 32.2-L of NH3

3. Calculate the number of moles in 5.45x1025 atoms of Zn

4. Calculate the number of grams in 3.25-mol of AgNO3

5. Calculate the number of liters in 0.0045-mol of N2

6. Calculate the number of molecules in 0.00325-mol of O2

Three Step Problems7. Calculate the number of molecules in 4.56-g of Pb(NO3)2

8. Calculate the number of liters in 3.25-g of NH3

9. Calculate the number of liters in 5.43x1025 molecules of H2

10. Calculate the number of grams in 3.54-L of CO2

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11. Calculate the number of grams in 9.7x1022 molecules of CH3CH2OH

12. Calculate the number of molecules in 5.42-L of O2

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Part II1. Calculate the % composition of Li2O.

2. What is the percentage composition of a carbon-oxygen compound, given that a 95.2 g sample of the compound contains 40.8 g of carbon and 54.4 g of oxygen?

3. What is the percentage composition of N2O4?

4. What is the percentage composition of a compound made from 28 grams of nitrogen and 32 grams of oxygen?

5. What is the percentage composition of a carbon-hydrogen-fluorine compound which contains7.2 grams of carbon, 11.4 grams of fluorine, and 1.8 grams of hydrogen?

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6. Find the percentage composition of Na2SO4?

BONUS PROBLEMS:7. If a compound is formed from 60.0 liters of nitrogen gas, N2, (at STP) and 180 liters of hydrogen gas, H2, (at STP), what is its percentage composition?

8. Find the percentage composition of a compound formed when 0.4 moles of potassium are reacted with 8.96 liters of O2 gas and 2.41 x 1022 atoms of S.

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Part E: Empirical and Molecular Formulas1. Determine the empirical formula of a compound with 72.4% Fe and 27.6% Oxygen.

+6

2. Determine the empirical formula of a compound with 65.2% Sc and 34.8% O

+6

3. Determine the empirical formula of a compound with 52.8% Sn, 12.4% Fe, 16% C and 18.8% N.

+8

4. Determine the molecular formula for a compound that contains 12.2-g Nitrogen, 27.8-g Oxygen, and a molecular mass of 92.0 g/mol.

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+8

5. Determine the molecular formula for a compound that contains 94.1% oxygen and 5.9% hydrogen and a molecular mass of 34 g/mol.

+8

6. Determine the molecular formula for a compound that contains 22.5% Na, 30.4% P and 47.1% O and a molar mass of 306 g/mol

+8

Total=44

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8. Determine the molecular formula of a compound that contains 76% iodine and 24% oxygen and has a molar mass of 334g/mol.

9. Determine the molecular formula of a compound that contains 48.6% carbon, 8.1% hydrogen, and 43.2% oxygen and has a molar mass of 296-g/mol.

10. Determine the molecular formula of a compound that contains 0.993-g nitrogen, 1.27-g carbon, 0.213-g hydrogen, 2.52-g chlorine and has a molar mass of 423-g/mol.

11. A sample of TNT, a common explosive is analyzed and found to contain 1.03-g of nitrogen, 0.220-g hydrogen, and 1.76-g of carbon. The molar mass is 123 g/mol. What is the molecular formula?

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12. Azobenzene is an important intermediate in the manufacture of dyes. It contains 79.1% carbon, 5.95% hydrogen, and 15.4% nitrogen. It has a molar mass of 182-g/mol. What is the molecular formula?

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Answer Key

1. Why is it important to use the correct amount of materials in a chemical reaction?

If too little is used the reaction may not proceed very far. The use of too much chemical may result in waste.

2. What names are given to the materials at the beginning and end of a chemical reaction?

Reactants andproducts.

3. Atoms and molecules are extremely small. How do chemists "count" them? Can you think of an everyday application of this?

They weigh them. If you know the number of nails or screws in a given mass, it is quic*er to weigh them instead of counting them individualy.

4. a. What did early chemists discover about reactions involving the combination of gases?

They combined in small whole number ratios.

b. How did Avogadro explain this?

Equal volumes of gases funder the same conditions) contain equal numbers of particles.

5. How may a chemical equation such as H2 + Cl2 2 HCl be interpreted?

It may be interpreted at the molecular basis or in terms of moles.

6. What is true about the mass of a compound?

It is equal to the sum of the masses of the individual atoms in the compound.

7. What is the numerical value for Avogadro's Number?

6.02 X 1023

8. When the I V solutions were prepared, quality control was involved. What is quality control?

The testing of a manufactured product to determine if it contains what it is supposed to contain.

9. Why did using twice as much magnesium not produce twice as much hydrogen in

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the demonstration?

The hydrochloric acid was used up. The magnesium will be totally consumed only if twice as many moles of acid are present. The hydrochloric acid became the limiting reagent.

10. What ratio of starting materials was found to produce the best epoxy resin?

A one - to - one ratio.

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