AcademicAcademic ChemistryChemistryUNIT 9 (Chapter 9)UNIT 9 (Chapter 9)

STOICHIOMETRYSTOICHIOMETRY

Name: Class Period: Test Date: 2/11/15

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CalendarMonday Tuesday Wednesday Thursday Friday

JANUARY 26

Investigation: All That Remains (pg. 3)

HW: Complete pg. 3 and pg. 6

pre-lab

27Notes #1 Mole Review

Cu-Fe Lab: Day #1 (pg. 5-6)

HW: Independent Practice Pg. 4

28Notes #1 Mol -> mol conversions

HW: Assessment pg. 9

29Notes #2 Other Unit conversions

HW: Assessment pg. 12

30Cu-Fe Lab: Day #2 (pg. 5-6)

Extra Stoichiometry Practice (pg. 13)

HW: Extra Practice pg. 13

FEBRUARY 2Cu-Fe Lab: Day #3 (pg. 5-6)

Extra Stoichiometry Practice

HW: Lab pg. 6 and Complete pg. 14

3QUIZ

S’MORE’S LABP.15-16

HW: Finish Lab (P.15-16)

4Notes #3 – Limiting Reagent

HW: Assessment pg. 18

5 Limiting Reagent Practice Day

6Notes #4 – Percent Yield

HW: Assessment pg. 21 & Review

9Computer Lab Practice

HW: Finish Test Review (pg. 22-24)

10 Test Review Due (pg. 22-24)

HW: Study for Test

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UNIT 9 TEST

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Begin Gases

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Investigation: All That Remains

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Predict the outcome and then test your prediction: How much solid mass will remain after decomposing 2.00g of sodium hydrogen carbonate by heating? (Hint: 2NaHCO3(s) -> Na2CO3(s) + CO2(g) + H2O(g))

Prediction (include your reasoning):

Design: Plan your investigation with your team. Write your procedure, including any diagrams, in the space below. Be prepared to discuss your plan with the class.

Safety: Discuss safety precautions that must be used and list them below.

Investigate: After class discussion, write in the space below any changes to your original procedure.

Data and Calculations:

Analysis and Conclusions:1. How did your final mass compare with your prediction?

2. What did you observe as you heated the sodium hydrogen carbonate?

3. Does your final mass measurement include either the water or carbon dioxide products? Explain.

4. Does your result match your prediction? Why or why not?

5. Is there a limiting or excess reagent? Explain.

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NOTES 1: MOLE REVIEW

For conversions between mass, volume (of a gas), and number of particles, we use the ___________________________to guide our use of _____________________ _________________________.

With dimensional analysis, you multiply the given by one or more conversion factors in the form of a fraction.

Example:

Practice Problems1. What is the mass of 4.00 moles of oxygen gas?

2. What is the volume of 3.4 x 1025 molecules of CH4 at STP?

Independent Practice:3. What is the mass of the helium inside a balloon with a volume of 2.3 liters?

4. How many moles of water are in a 45 gram sample?

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Chemical Conversion Factors to Know:

1 mole = molar mass (g)1 mole = 6.02 x 1023 particles (atoms, ions, molecules, formula units)1 mole = 22.4 L of a gas at STP

STP = standard temperature and pressure (T = 273K, P = 1 atm)

Cu-Fe Stoichiometry Lab

In this lab exercise, iron nails will react with a solution of copper (II) chloride to produce iron (II) chloride and copper. When the reaction is complete, the mass of iron used as a reactant and the mass of copper metal produced will be determined. Then using the process of stoichiometry, the theoretical yield of copper can be calculated, followed by a percent yield calculation to determine the efficiency of this chemical reaction.

Purpose: To use stoichiometry to predict the amount of copper produced.To calculate the percent yield of copper using the actual yield found by experimentation.

Safety:1) Wear safety goggles and lab aprons at all times.2) All acid spills should be neutralized with baking soda (NaHCO3).

Materials:2 nails 100-mL beaker wash bottle tongsfunnel filter paper electronic balance 25-mL graduated cylinder

Procedure: Day 1 1. Place a clean dry beaker on a balance and reset the balance by pressing "tare". Then add 4.25

grams of copper (II) chloride. 2. Obtain 2 clean, dry nails. Buff the nails with sandpaper to remove the lacquer. Then find the mass

of both of the nails together. Record the mass on the data table. 3. Add about 25 ml of tap water to the copper (II) chloride and stir the mixture until all crystals

dissolve. 4. Drop the nails in the solution and add more water (if needed) until both of the nails are covered by

the solution. Label the beaker with your name and put the beaker at the back of the lab for at least 24 hours.

Day 2 1. Obtain your beaker. Notice how the color has changed! Carefully decant (pour) the iron (II) chloride

solution covering the nails from the beaker into the sink. Do not lose any copper! 2. Using a pair of tongs, pick up each nail (one at a time). Using a slow stream of water from a bottle,

wash all of the copper from the nails back into the beaker. 3. Dry the nails with paper towel. Find their mass together and record in the data table. 4. Mass a piece of filter paper, record its mass, and insert the paper into a funnel using the proper

folding technique. 5. Pour the copper in solution into the funnel. Be careful not to overflow the paper. 6. Using distilled water, wash the copper by filling the funnel. Do this twice. (You must do this to wash

the iron (II) chloride from the copper.) 7. Now wash the copper with 25 ml of HCl. (This will dissolve any remaining iron filings) 8. Wash one more time with water. Then carefully remove the paper with the copper, unfold, and set

on a paper towel to dry overnight. Put your name on the paper towel.

Day 31. Find the mass of your filter paper and copper.

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2. Empty copper into beaker at front of lab. Pre-Lab Questions1. What should you do to clean up an acid spill?

2. Write the balanced chemical equation for this reaction: iron and copper (II) chloride react to produce iron (II) chloride and copper.

3. Label this reaction as one of the 5 general types of chemical reactions (synthesis, decomposition, single-replacement, double-replacement, or combustion).

Data: 1. Mass of iron nails before reaction (Day 1) _______________

2. Mass of iron nails after reaction (Day 2) _______________

3. Mass of filter paper (Day 2) _______________

4. Mass of filter paper and copper (Day 3) _______________

5. Mass of iron that reacted (#1 - #2) _______________

6. Mass of copper produced (#4 - #3) _______________

Post-Lab Questions1. What was the actual yield of copper (data #6)?

2. Using the mass of iron that reacted (data #5) and your balanced chemical equation, determine the theoretical yield of copper...how many grams of copper should have been produced (Hint: Use mole-map).

3. Calculate the % yield (% yield = ) of copper for this reaction.

4. List two credible sources of error in this experiment. You may NOT use human error unless you cite a specific example.

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NOTES #2: MOL TO MOL CONVERSIONS

RECALL: Balancing Chemical EquationsDirections: Using coefficients, balance the following equation -

____ C6H6 + ____O2 → ____CO2 + ____H2O

Cooking Analogy: Grill Master K.T. Tiger has the art of grilled cheese sandwich making down to a science. The Grill

Master's recipe requires 2 pieces of cheese between 2 slices of bread, grilled to perfection. What is the coefficient ratio of the ingredients to the product?

2 + 2 →

____ : ____ : ____

Grill Master Tiger knows that a 20 pack of sliced bread and a 20 pack of sliced cheese will always make the same number of grilled cheese sandwiches with no leftovers. How many?

____________________________________________________________________________________

What happens to the Grill Master's grilled cheese sandwiches if he changes the quantities of ingredients? Will he have enough ingredients? Will there be leftovers?

____________________________________________________________________________________

o Practice grilled cheese sandwich making here:http://phet.colorado.edu/en/simulation/reactants-products-and-leftovers

What Cooking Really Is...STOICHIOMETRY:

_____________________________________________________________________________________ N2 (g) + 3H2 (g) → 2NH3 (g)

o How many molecules of each reactant are required to produce 2 molecules of product? Ratio?

______________________________________________________________________________o How many moles of each reactant are required to produce 2 moles of product? Ratio?

______________________________________________________________________________

Instead of cups, teaspoons, or tablespoons, we have…

1. ______________________________________________ :

___ molecule of nitrogen gas reacts with ___ molecules of hydrogen gas to produce ___ molecules

of ammonia gas.

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It's always in the same coefficient ratio; just like 2 slices of bread plus 2 slices of cheese produce 1

grilled cheese sandwich!

2. ______________________________________________ :

___ mol of nitrogen gas reacts with ___ mol of hydrogen gas to produce ___ mol of ammonia gas.

3. ______________________________________________ :

Law of conservation of mass says mass of reactants must ________________ mass of products.

Mass of nitrogen gas = _________ and mass of hydrogen gas = ________

Their sum equals the mass of the products = ____________

4. ______________________________________________ :

1 mol of gas = __________________ at Standard Temperature and Pressure (STP)

o _________ L of nitrogen gas reacts with __________ L of hydrogen gas to produce ________ L

of ammonia.

Mol to Mol Conversion Calculations: N2 (g) + 3H2 (g) → 2NH3 (g) What is the mole coefficient ratio of the above equation?

____ : ____ : ____

Because we know the ratio, we can calculate to find the number of moles of another substance.

Example: How many mol of NH3 are produced when 0.60 mol of nitrogen gas reacts with hydrogen gas?

Practice ProblemsDirections: Using the balanced chemical equation, calculate the following mol conversions -

MnO2 + 4HCl → MnCl4 + 2H2O1. How many mol of H2O are produced when 3.20 mol of MnO2 reacts with hydrochloric acid?

2. How many mol of HCl are consumed (used) when 1.65 mol of Manganese (IV) chloride are produced?

3. How many mol of water are produced when 4.35 moles of MnCl4 are also produced?

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Assessment: Interpreting Coefficients as MolesDirections: Given the equation, calculate each of the following; balance the equation if necessary -

____CH4 (g) + ____O2 (g) → ____CO2 (g) + ____H2O (g)

1. How many mol of carbon dioxide (CO2) are formed when 40 mol of oxygen (O2) is consumed?

2. How many mol of methane (CH4) are needed to form 200 mol of water?

3. How many mol of oxygen (O2) combine with 0.05 mol of methane (CH4)?

____NO (g) + ____O2 (g) → ____NO2 (g)

4. How many mol of oxygen (O2) combine with 500. mol of NO?

5. How many mol of NO2 are formed from 0.25 mol of NO?

6. a. If you have 80. mol of NO, how many mol of oxygen (O2) would you use?

b. If you had started with 200. mol of oxygen (O2), how many would you have left?

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NOTES #3: MASS TO MASS AND OTHER UNIT CONVERSIONS

RECALL: The Mole Highway...it's been expanded and construction is complete!Like before, the Mole Highway can be used as a map toward setting up an appropriate conversion

HOW IT'S USED:1. Find the starting point; use the value, substance, and unit you are given2. Find the ending point; use the value, substance, and unit are you being required to solve3. You must stay on the highway!!!4. Each road taken represents 1 step in your conversion!5. Once the destination is reached, solve mathematically by multiplying across the top, multiplying

across the bottom, and dividing the top value by the bottom value.

NOTE:1. Before any math can be done, a BALANCED CHEMICAL EQUATION is required.2. If the starting value isn't in the unit "mol," your first step is to convert it there.3. Going from mol of substance A to mol of substance B requires a Mol to Mol conversion; USE YOUR

COEFFICIENTS FROM THE BALANCED CHEMICAL EQUATION!

Mass to Mass Conversion Calculations N2(g) + 3H2 (g) -> 2NH3 (g) Like yesterday, in order to go from the mass of one substance, to the mass of a new substance, a mol to

mol conversion will be necessary.

Example: Calculate the number of grams of NH3 produced by the reaction of 5.40 g of H2 with excess N2 (g).

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Practice Problems - ALL UNITS***Write a balanced formula, and then set up your conversion calculations.***

1. How many grams of O2 (g) are produced when a sample of 29.2 g of water decomposes?

2. Using the same equation, how many liters of hydrogen gas are produced when 1.33 x 1017 molecules of water decompose?

2SO2 (g) + O2 (g) → 2SO3 (g)3. How many liters of O2 are needed to produce 19.8 L of SO3?

4. How many molecules of oxygen are consumed in the formation of 187.4 L of SO3?

5. How many molecules of Sulfur dioxide are consumed in the formation of 4.41 x 1027 molecules of sulfur trioxide?

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Assessment: Mixed Mole Conversions

1. Find the molar mass of the following:

a. 1 mol of CH4 = ____________g of CH4

b. 1 mol of Ag2SO4 = _______________g of Ag2SO4

c. 1 mol of P2O5 = _____________g of P2O5

d. ___________g of NaBr = 1 mol of NaBr

e. ___________g of Al(NO3)3 = 1 mol of Al(NO3)3

f. ___________g of Ba3(PO4)2 = 1 mol of Ba3(PO4)2

CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)

2. How many moles of O2 are needed to produce 5.2 moles of CO2?

3. How many liters of CH4 are needed to produce 0.38 grams of H2O at STP?

4. How many molecules of CO2 are produced from 12.9 grams of O2?

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Mass, Volume, and Particle Stoichiometry Extra PracticeBalance:

___Ba3(PO4)2 + ___HI (g) ___BaI2 + ___H3PO4

1. How many grams of BaI2 are produced from 4.37 x 1027 molecules of HI?

2. How many formula units of BaI2 are produced from 96.2g of HI?

3. How many grams of H3PO4 are produced from 2.0 moles of Ba3(PO4)2?

4. How many liters of HI are needed to produce 72.54g of BaI2 at STP?

5. Practice drawing the expanded Mole Highway from memory. Remember, the only addition is that Mol can now be converted to Mol.

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Mass, Volume, and Particle Stoichiometry Extra Practice 2

1. If you have 1.91 x 1025 molecules of ethane (C2H6), how many molecules of oxygen are needed to react completely? (molecules= particles!)

C2H6 + O2 CO2 + H2O

2. Aluminum and sulfur react to produce aluminum sulfide. How many grams of sulfur are needed to produce 18.62 grams of aluminum sulfide?

3. Ammonia and water combine to produce ammonium hydroxide. At STP, how many liters of water are needed to produce 6.54 liters of ammonium hydroxide?

4. How many atoms of phosphorus are needed to produce 16.0 x 1024 molecules of phosphine (PH3)? (atoms & molecules = particles!)

P + H2 PH3

5. How many grams of oxygen are needed to produce 6.55 x 1024 formula units of potassium oxide?K + O2 K2O?

6. How many liters of hydrogen gas will be produced if 38.5 g of hydrochloric acid reacts with zinc?Zn + HCl ZnCl2 + H2

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Name Date Period

Sweet StoichiometryObjective Students will be able to

Perform stoichiometric calculations using a fictional reaction. Explain the concept of limiting reagent.

Background Information Engineers and scientists get paid thousands of dollars each year just to make sure that their companies are maximizing the profits by minimizing waste. For many of these engineers and scientists, this means maximizing the products of a chemical reaction. In this experiment, you will be the engineer in charge of manufacturing S’mores. Materials

Symbol Substance Molar MassM Marshmallow 11.50 grams/molG Graham Cracker 22.33 grams/molC Piece of Chocolate 31.97 grams/mol

__G + __M + __C __G2MC

Procedure 1. Balance the above equation.2. Count each your raw materials (marshmallows, graham crackers, and pieces of chocolate).

Record your initial counts in the data table provided.3. “Manufacture” your S’mores. Record the total number of S’mores made in the data table

provided.4. Record the number of each raw material left over after you have finished making your

S’mores.5. Enjoy your product as you answer the analysis questions.

Data Initial Counts of Raw Materials Final Counts of Raw Materials

MarshmallowsGraham CrackersPieces of Chocolate

Analysis Questions 1. Which raw material did you have the most of initially? Which did you have the least of

initially?

Most = Least =

2. Which raw material limited the number of S’mores that you could make?

3. Scientists and engineers ideally would like all reactions to run so that every reactant is used in what they call stoichiometric quantities. This means that there would be no waste or left over raw materials. If your goal is to make ten S’mores, how many

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Marshmallows

Graham Crackers

Pieces of Chocolate

S’Mores

marshmallows, graham crackers and pieces of chocolate would you order so that you would have stoichiometric quantities?

4. In your initial experiment, were your raw materials given to you in stoichiometric quantities? Explain why or why not.

5. When reactants are not used in stoichiometric quantities, you end up with limiting reactant(s) and excess reactant(s). Using your book, define the terms: limiting reactant and excess reactant. Your book using reagent instead of reactant, but they mean the same thing.

6. In your initial experiment, identify the limiting reagent(s) and the excess reagent(s).

7. Many students think that the limiting reactant is simply the reactant that you have the smallest quantity of in the beginning. Did your limiting reactant start off with the fewest number?

8. How many molecules of S’mores could you make if you had 100 grams of chocolate?

9. How many grams of Marshmallows are used to produce 16 grams of S’Mores (G2MC)?

10. How many atoms of Chocolate are necessary to produce 6.55 x 1025 molecules of S’Mores?

11.BONUS: If you started with 50 moles of chocolate and 60 moles of graham crackers, how many S’mores could you make? Which reactant would serves as the limiting reagent?

NOTES #4: Limiting Reagents (Reactants)

Back to cooking…1. What is Grill Master Tigers' recipe for grilled cheese sandwiches?_______________________________________2. What happens if the Grill Master receives only half of his order of cheese for the day?____________________3. So, in this case, what is limiting the number of grilled cheese sandwiches he can make? ___________________

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4. What if the Grill Master received 3 times the amount of cheese he needed?_____________________________5. What would be limiting him in this case? ________________________________________________

Limiting Reagent (aka Limiting Reactant) The same thing applies to chemical reactions…

N2 (g) + 3H2 (g) → 2NH3 (g)

___ mol of N2(g) reacts with ___ mol of H2 (g) mol to form __ mol of NH3 (g).

What happens if you only have 0.5 mol of nitrogen gas?

_____________________________________________________________________________________

How much hydrogen gas would you use?

_____________________________________________________________________________________

How much ammonia would you make?

_____________________________________________________________________________________

Limiting Reagent/Reactant :______________________________________________________________

_____________________________________________________________________________________

Excess Reagent/Reactant :__________________________________________________________

_____________________________________________________________________________________

Example: Sodium chloride can be prepared by the reaction of sodium metal with chlorine gas. Suppose that 6.70 mol of Na reacts with 3.20 mol of Cl2. What is the limiting reagent?

Follow these five steps…. Step 1: ______________________________________________________________________________

Step 2: ______________________________________________________________________________ To do this, convert the given info for each reactant (mol, in this problem) to mol of product. If the original substances are given in grams, you’d have an extra step – converting mass to moles –

but the rest of the process would be the same. NOTE: If the original substances are given in grams, what do you need to do before this step?

o Convert to mol.

Step 3: ______________________________________________________________________________ You will have _____ answers for moles of product. Why? _________________________________

________________________________________________________________________________ So which one do I use? _____________________________________________________________ The __________________ of the two answers is the ______________________________ because

you only have enough reactants to make that amount. You will run out of the reactants before the larger amount is made.

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From this, you can figure out your limiting reagent and excess reagent.

Step 4: ______________________________________________________________________________ Once you know your maximum product, look all the way to the ________ of that calculation. The

reactant that produced that smaller amount of product is your _____________________. It “limited” you to making the ______________________________________________. The other reactant – that started the other calculation – is the _________________________.

Step 5: ______________________________________________________________ Convert mol of your limiting reagent to mol of product (or whatever the original question asks for).

Practice Problem2Cu (s) + S (s) → Cu2S (s)

1. What is the limiting reagent when 80.0 g Cu reacts with 25.0 g of S?

2. What is the maximum number of grams of Copper (I) sulfide produced?

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Assessment: Limiting Reagent/ReactantDirections: Reference Notes #3 to complete the following Limiting Reagent word problems. Find a balanced chemical equation if necessary and follow the four steps.

1. What mass of water can be produced from 2.0 mol of H2 and 4.0 mol of O2?

2. What mass of water can be produced from 16 g of H2 and 8 g of O2?

3. What mass of water can be produced from 0.49 g of H2 and 1.3 g of O2?

4. What mass of water can be produced from 3.5 mol of H2 and 4.5 mol of O2?

____Cu + ____AgNO3 --> ____Cu(NO3)2 + ____Ag

5. How many mol of Ag can be produced from 6.3 mol of Cu and 4.2 moles of AgNO3?

6. What mass of Cu(NO3)2 can be produced from 5.5 g Cu and 1.95 g AgNO3?

7. How many mol of Ag can be produced from 1.5 g Cu and 7.2 g of AgNO3?

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NOTES #5: Percent Yield

Cooking, yet again… Grill Master Tiger's recipe yields 10 grilled cheese sandwiches when 20 slices of bread and 20 slices of

cheese are used

So, the recipe tells you the theoretical yield: ________________________________________________

_____________________________________________________________________________________

What you actually make is the actual yield: _________________________________________________

_____________________________________________________________________________________

Percent yield :

_____________________________________________________________________________________

_____________________________________________________________________________________

Formula: ____________________________ x 100Food for Thought: Could the percent yield normally be larger than 100%? Smaller? What might cause the % yield to vary?

Example: Using the following equation, what is the theoretical yield of CaO if 24.8g CaCO3 is heated? What is the percent yield if 13.1g of CaO is produced?

Follow the steps below… o Step 1: Convert grams of the given substance to grams of the questioned substance.

Which means…

24.8 g of CaCO3 = _____g of CaO.

o Step 2: Use the formula to find percent yield.

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Assessment: Percent Yield ApplicationsDirections: Reference Notes #4 to complete the following percent yield word problems. All work must be shown.

1. A baker has a recipe for cookies that says he can bake a yield of 250 cookies. However, when he does bake them, he finds that he can only make 125. What is the percent yield? (Assume that all cookies are the exact same size as directed by the recipe)

2. A chemist runs an experiment and produces 100 g of sulfuric acid. The All-Knowing Chemistry Book states he should have made 123 g of sulfuric acid. What is the percent yield?

3. A farmer out checking his fields looks in on his Hawaiian flower tree. He notices only 15 blooming flowers. This is 73% of what it should be. What is the Hawaiian flower tree’s theoretical yield?

4. A man shelling pecans notices that each pecan is only producing 39% of the meat expected. In the Pecan Man’s Handbook, he reads that each pecan should have a mass of 200 g. What is the actual yield of each pecan?

5. Genny collected 1.05g of sodium carbonate by decomposing 2.00 g of sodium hydrogen carbonate. According to the chemical equation

She should have been able to collect more. Use stoichiometry and the percent yield equation to find the percent yield of Genny’s experiment.

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Academic Chemistry Unit 9 Test Review (Due on Feb 10th)STOICHIOMETRY TEST on Wednesday, February 11th

1. Define the following:

a. Stoichiometry

b. Limiting reagents

c. Excess reagents

d. Percent Yield

e. Actual Yield

f. Theoretical Yield

2. Find the molar mass of the following:

a. __________g of CS2 c. _________g of P2O5

b. __________g of Ag3PO4

Directions: Balance the following equations and perform the appropriate conversions and calculations.

____HCl ____H2 + ____Cl23. How many grams of HCl are needed to produce 5.2 moles of Cl2?

4. How many moles of H2 are produced from 6.3 L of HCl at STP?

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5. How many moles of Cl2 are produced from 12.9 grams of HCl? (see equation on previous page)

____BaSO4 + ____HI (g) ____BaI2 + ____H2SO4

6. How many moles of BaI2 are produced from 7.6 moles of HI?

7. How many grams of BaI2 are produced from 9.2L of HI at STP?

8. How many grams of H2SO4 are produced from 2.0 moles of BaSO4?

9. How many molecules of HI are needed to produce 72.54g of BaI2?

10. What is the volume of H2SO4 produced from .035 grams of HI at STP?

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11. Balance the following equation and answer the questions that follow:

____Fe + ____H2O ____Fe3O4 + ____H2

a. What is the molar mass in grams of iron IV oxide?

b. What is the limiting reagent if 36.0g of water reacts with 167.0g of Fe?

c. Using the information in part b, If only 105 grams of Fe3O4 were produced, what is the percent

yield of the reaction?

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