Stoichiometry Mark Denniston December 6, 2001
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Table of Contents Unit Plan I
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Introduction Misconceptions Text Analysis Concept Map
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Unit Plan II
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Objectives Discussion Standards Addressed Lesson Plans 1. Chemical Measurements 2. Mole Conversions 3. Empirical and Molecular Formulas 4. Review: Lessons 1-3 5. Stoichiometry 6. Stoichiometry Lab Prep 7. Stoichiometry Lab 8. Solving Stoichiometry Problems 9. Stoichiometry Web Site Review 10. Limiting Reactants and Percent Yield Unit Plan III
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Unit Test Unit Test Answer Key Correlation Between Unit Objectives and Test Alternate Assessment Strategies Conclusion
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39 47 55 57 58
Introduction The question I have chosen for my unit plan is “What is stoichiometry and why is it important to everyday life?” Stoichiometry is the study of the quantitative, or measurable, relationships that exist in chemical formulas and chemical reactions. In simpler terms, stoichiometry refers to the ratio of substances in molecules or chemical reactions. Possibly the most frequently asked question of a teacher is “Why do I need to know this?” In the case of stoichiometry it is important that students understand the concepts for several reasons, including: •
Understanding how energy is produced. Based on their stoichiometry, certain chemicals produce energy more efficiently than others.
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Understanding how the human body uses food. Certain foods are more efficient at providing nutrition and energy than others, based on their chemical composition.
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Chemistry principles such as stoichiometry are necessary in other disciplines. For example, in biology a knowledge of stoichiometry is necessary to understand how humans convert glucose into ATP for energy.
Understanding the principles of stoichiometry contributes to scientific literacy as an adult. When stoichiometric principles are considered, educated decisions can be made about what type of heating system would be the most efficient to install in your home. When the issue of pollution arises, you will be informed as to what causes it. You can predict what the resulting products will be if you mix things together. On a simpler level, you will appreciate why it is important to follow a cooking recipe closely to produce the desired results.
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Misconceptions Students often have misconceptions about chemical reactions and stoichiometry. In talking to an experienced chemistry teacher, I found that a common misconception is that substances always react in a 1:1 ratio. I can remember that this is a misconception I had early in my chemistry career. Along the same lines, students often assume that substances react by equal weight (i.e., 1 lb of HCl reacts with 1 lb of CaO to form 1 lb of H2O and 1 lb of CaCl). The text I chose also lists common misconceptions, such as a mole always contains 6.02E23 atoms. In the case of one mole of Cl2, there are actually two moles of Cl atoms. Text Analysis The text I chose is Chemistry: Connections to Our Changing World, 2002, written by LeMay, Beall, Robblee, and Brower, and published by Prentice Hall. When studying any area of science there are common terms that texts often use. To help understand and distinguish between the terms, I have prepared definitions for them: Science fact: A science fact refers to a piece of knowledge that is believed to be true all of the time. For example, a fact stated in the text is “A mole contains 6.02E23 particles.” (p.319). A description of the text is given later. Science generalization: A science generalization is something that is true most of the time or applies to most of something. A science generalization is often an opinion. An example in the text is, “The gram is the desired unit of mass.” (p. 313) Concept: A concept is a broad, big-picture type of idea that includes many individual, specific pieces of knowledge. The text discusses the mole concept, which
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includes more specific detail such as Avogadro’s number, atomic mass, etc. (p. 310339) Law: A law is much like a fact in that it is a widely accepted belief. A law is a theory that has stood the test of time. An example of a law in the text is the Law of Conservation of Matter (p.349-350) Theory: A theory is a belief that does a good job of explaining the way things work. Theories can’t always be proved, but also may not be able to be disproved. There were no specific examples of theories in the section of text I chose. Empirical Entity: An empirical entity is something that was derived through experimentation or hands-on activity. For example, experiments determined a mole to contain 6.02E23 particles (p. 315). Theoretical Entity: A theoretical entity is something that was derived through thought or calculation, such as combining several equations to form another. I did not find any solid examples of this in the section of text I chose.
The section of text dealing with stoichiometry gives examples of how to calculate percent yield. An experiment could easily be set up in which students use stoichiometric principles to calculate the expected yield of a product in a reaction. The students could then actually perform the experiment, measuring the amount of product produced (actual yield) to calculate the percent yield. Another experiment could be set up to introduce the concept of titration to determine the molar ratios of a reaction. Overall I found this text to be very good at presenting the concepts of stoichiometry. The text clearly explains the principles involved in making calculations and contains step-by-
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step examples of how to work through problems. In addition, the text also includes in each chapter sections on “Problem Solving: Chemistry at Work,” “Chemistry in Action,” and “Exploration” that help students relate the concepts they are learning to real-life situations. These sections also contain experiments or activities that students can perform to better understand stoichiometry. The text does a good job of offering opportunities for all different types of learners, which is a key to being a good educator. I would primarily use the text as a reinforcement of lecture material. After presenting the material in class, students would have the opportunity to review the concepts using the textbook in study hall or at home. The text would also allow students to broaden their knowledge using some of the tools mentioned in the previous paragraph. I might give students extra credit points for completing activities outlined in the textbook that we don’t have time for in regular class. For the stoichiometry unit I would use the text as a good source of practice problems, which are a necessary part of learning stoichiometry. In general, I would use the text as a supplement to lecture and classroom activity.
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Objectives Discussion At the beginning of the lesson plans in the next section the learning objectives for that lesson are listed. After each objective the following code is used to denote the cognitive or affective level for that objective: K C AP AN S E
Cognitive Knowledge Comprehension Application Analysis Synthesis Evaluation
RC RS V O C
Affective Receiving Responding Valuing Organizing Characterizing
For example, (C-AP) would indicate a cognitive objective at the application level.
Academic Standards The following Wisconsin State Academic Standards are addressed in the lesson plans: A.12.2 Show* how conflicting assumptions about science themes* lead to different opinions and decisions about evolution*, health, population, longevity, education, and use of resources, and show* how these opinions and decisions have diverse effects on an individual, a community, and a country, both now and in the future. A.12.3 Give examples that show* how partial systems*, models*, and explanations* are used to give quick and reasonable solutions that are accurate enough for basic needs A.12.5 Show* how the ideas and themes* of science can be used to make real-life decisions about careers, work places, life-styles, and use of resources. A.12.6 Identify* and, using evidence* learned or discovered, replace inaccurate personal models* and explanations* of science-related events. C.12.3 Evaluate* the data collected during an investigation*, critique the data-collection procedures and results, and suggest ways to make any needed improvements D.12.4 Explain* how substances, both simple and complex, interact* with one another to produce new substances 14
D.12.5 Identify* patterns in chemical and physical properties and use them to predict* likely chemical and physical changes and interactions D.12.6 Through investigations*, identify* the types of chemical interactions*, including endothermic, exothermic, oxidation, photosynthesis, and acid/base reactions G.12.2 Design, build, evaluate, and revise models* and explanations related to the earth and space, life and environmental, and physical sciences
Lesson Plan 1 15
Unit:
Stoichiometry
Subject: Chemical Measurements Rationale: The scale for expressing atomic masses is based on the mass of carbon-12. The sum of the atomic masses of all the atoms in a compound is called the formula mass of the compound and is measured in atomic mass units. In the laboratory, however, the gram is the desired unit of measurement. A unit called the mole establishes a relationship between the atomic mass unit and the gram. A mole is equivalent to Avogadro’s number of particles, 6.02 x 1023, and the mass in grams of one mole of a substance is called its molar mass. The molar mass of a compound in grams is equal to the compound’s formula mass in amu. Learning Objectives: Students should be able to: 1. 2. 3.
Define a mole and describe its importance. (C-K) Identify and use Avogadro’s number. (C-AP) Define molar mass and use it to relate the mass of a substance to the number of particles in that substance. (C-C)
Important Terms: atomic mass, formula mass, mole, Avogadro’s number, molar mass WI State Standards: A.12.6, D.12.4, D.12.5 Materials Needed: 12 g graphite, 55.8 g iron filings, balance Procedure: Students will be taking notes during lecture. 5-10 minutes 1. Introduce atomic mass and formula mass. • atomic mass unit – based on C-12 mass • formula mass = sum of atomic masses • Q: What is formula mass of SO2? A: 64 amu (32+16+16) 15-20 minutes 2. Introduce mole concept. • mole – relates amu to grams 16
• Avogadro’s number = 6.02 x 1023 - number of particles in 12 g of C-12 - one mole has Avogadro’s number of particles • molar mass – mass in grams of one mole of a substance - numerically equal to atomic mass or formula mass 5 minutes 3. Demonstration Place one mole (12 g) of graphite, or carbon, on balance with one mole (55.8g) of iron filings. Explain to students that there is one mole of each substance, but do not tell them masses. Students should observe: -
there is an equal number of atoms of each substance mole of iron has more mass than one mole of carbon, indicating that an iron atom has more mass than a carbon atom use periodic table to find molar atomic mass of Fe and C
10-15 minutes 4. Discuss the forms moles can take: • atoms in an element • molecules in a diatomic element or compound • formula units in ionic substances Have class answer these questions out loud: Q: How many formula units are in 2 moles of Cl-? A: 12.04 x 1023 Q: How many moles of H+ ions and O- ions are in one mole of H2O? A: 2 moles H+, one mole OAssociated reading: p. 310-322 Assessment: -
Gauge student participation and understanding while working sample problems with class. Section Review p. 322 (1-4), due next class
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Conclusion: -
Review important terms and their meanings Explain how moles will be the basis for converting between measurements (next lesson)
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Lesson Plan 2 Unit:
Stoichiometry
Subject: Mole Conversions Rationale: The mole is the central unit for converting the amount of a substance from one type of measurement to another. The number of moles of a substance can be calculated once the mass of the substance is known. Knowing the number of moles allows for a direct conversion to the number of particles. Two conversion factors are thus needed to convert mass to number of particles. The mole is also used to convert between the number of particles of a gas and the volume of a gas at STP. This molar volume of the gas is identical for all gases and has the value of 22.4 L/mol. Learning Objectives: Students should be able to: 1. Convert among the number of particles, moles, and mass of a substance. (C-AP) 2. Define molar volume and use it to solve problems. (C-AP) Important terms: molar volume WI State Standards: D.12.4, D.12.5 Materials: Mole Conversions worksheet Procedure: Students will be taking notes during lecture. 10-15 minutes 1. Discuss how to make the following conversions, using the line method: • mass ↔ moles • moles ↔ particles • mass ↔ particles Work examples from the book in front of the class. Use sample problems from teacher’s edition pages 326-327. Ask students to lead you through the problem step by step.
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5-10 minutes 2. Explain the concept of molar volume. • molar volume – volume of one mole of gas at STP = 22.4 L/mol - STP = 0°C, 1 atm 20-25 minutes 3. Allow students the remainder of class to work on homework problems and ask questions Associated reading: p. 323-331 Assessment: -
Gauge student participation and understanding while working sample problems with class. Homework problems (due next class)
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Review important terms and the 3 types of conversions Preview next lesson (formulas tell us about the composition of materials)
Conclusion:
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Chemistry
Name: Period:
Moles Conversion Worksheet 1. How many atoms are in 0.62 moles of Ag?
2. How many moles of BaNO3 contain 102.3 g?
3. How many atoms are in 56.3 g of copper?
4. A room with a volume of 4000 L contains how many moles of air at STP?
5. A glass of milk contains 5 g calcium. How many atoms of calcium is that?
6. If you burned 4.0 x 1024 molecules of natural gas (methane CH4), what mass of methane did you burn?
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Lesson Plan 3 Unit:
Stoichiometry
Subject: Empirical and Molecular Formulas Rationale: Each chemical compound has a unique combination of different types of atoms. When the formula of a compound is known, the percent composition of each element in the compound is determined by comparing its mass to the total mass of the compound. Similarly, if the percent composition is known, the empirical formula of the compound can be determined. Empirical formulas give the simplest whole-number ratio of the atoms of the elements in the compound. Molecular formulas give the actual number of atoms of each element. The molecular formula is always a whole-number multiple of the empirical formula. Learning Objectives: Students should be able to: 1. Calculate the percent composition of elements in a given formula. (C-AN) 2. Describe how percent composition can be used to make intelligent consumer decisions. (A-V) 3. Use percent composition to find the formula of an unknown sample. (C-AN) 4. Determine empirical and molecular formulas. (C-AN) Important terms: percent composition, empirical formula, molecular formula WI State Standards: D.12.4, D.12.5 Materials: Empirical and Molecular Formulas worksheet Procedure: Students will take notes during lecture. 10-15 minutes 1. Introduce percent composition - review concepts of atomic and molecular weight - ask class what the atomic weight of C is (12 g) - ask class what the molecular weight of HCl is (36.5 g) Work sample problem on the board with class. Go step by step so class can take notes.
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Q: What is the percent hydrogen in water? Atomic weights: H = 1.0 g, O = 16.0 g Water = 2H (1.0 x 2) + 1O (16.0) = 18.0 g Percent hydrogen = (2.0 g / 18.0 g)*100 = 11% Ask class if everyone understands how we arrived at this answer. Work more problems if necessary. Relate percent composition to dietary guidelines on the labels of food products. 10-15 minutes 3. Introduce empirical formula - simplest whole number ratio of atoms - review reducing to lowest common denominator Work sample problem on the board for class. Go slow enough so they can take notes. Q: Given a compound that is 80% carbon and 20% hydrogen, find the empirical formula. Step 1. Work on a 100 g basis. 80 g C, 20 g H. Step 2. Convert to moles.
80 g C / (12.0 g C/mol) = 6.67 mol C 20 g H / (1.01 g H/mol) = 19.8 mol H
Step 3. Divide each mole value by the smallest value. C 6.67 / 6.67 = 1.00 H 19.8/6.67 = 2.97 (round to 3.00) Step 4. Empirical formula = CH3 10-15 minutes 3. Introduce molecular formula. - multiple of empirical formula, shows actual number of atoms of each element in compound - Ask class question: Which formula do you think is more useful, empirical formula or molecular formula? Why? A: molecular tell us more Work sample problem on the board for class. Go slow enough so they can take notes. Be aware that students may confuse molar mass with working on a 100 g basis. Q: Ribose has molar mass = 150 g/mol, and is composed of 40% C, 6.67% H, and 53.3% O. Find the molecular formula.
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Step 1. Find empirical formula. See above. C 40 g / (12.0 g/mol) = 3.33 H 6.67 g / (1.01 g/mol) = 6.60 O 53.3 g / (16.0 g/ml) = 3.33 Empirical formula = CH2O Step 2. Find formula mass of empirical formula. C = 12.0 , H = 1.0 (x2), O = 16.0 = 30.0 Step 3. Divide molar mass by empirical formula mass. 150.0 / 30.0 = 5 Step 4. Multiply empirical formula by that number. C5H10O5 Associated Reading: p. 332-339 Assessment: -
Observe students while working sample problems to gauge understanding. Consider responses to questions posed during discussion. Homework problems due next class.
Conclusion: Review important terms and what they mean. Explain to students that they will be given the opportunity during next class to put their knowledge to the test in a game.
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Chemistry Empirical and Molecular Formulas Worksheet
Name: Period:
1. What is the percent composition of carbon, oxygen, and hydrogen in C6H12O6?
2. Determine the empirical formula of a compound containing 5.75 g Na, 3.5 g N, and 12.0 g O.
3. Find the molecular formula of a compound that contains 42.56 g of palladium and 0.80 g of hydrogen. The molar mass of the compound is 216.8 g/mol.
4. Octane, a compound of hydrogen and carbon, has a molar mass of 114.26 g/mol. If the compound contains 18.17g/mol of hydrogen, what is its molecular formula?
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Lesson Plan 4 Unit:
Stoichiometry
Subject: Review – Sections 10.1, 10.2, 10.3 (see last 3 lessons) Rationale: This lesson will review chemical measurements, mole conversions, and empirical and molecular formulas, while tying them together at the same time. Learning Objectives: Students will be able to explain how chemical measurements, mole conversions, and empirical and molecular formulas relate to one another. (C-C) Important terms: see last 3 lessons WI State Standards: D.12.4, D.12.5 Materials: none Procedure: Entire class period 1. Instructor will divide students into two evenly matched teams. 2. In a “Family Feud” style game, the students will review sections 10.1, 10.2, and 10.3. a. One student (take turns) from each team will come to the front of the room. b. Instructor will read questions (from text or made-up). Student to raise his/her hand first gets first chance at answering. c. Each correct answer is worth 2 points, while an incorrect answer results in subtraction of 1 point. 3. At the end of the game each time will be allowed to wager their points on the final question, “Jeopardy” style. Assessment: Visually check to make sure all students are participating during game. Conclusion -
Congratulate students on a great game. Explain how the previous 3 lessons will lead into the next chapter on Stoichiometry.
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Lesson Plan 5 Unit:
Stoichiometry
Subject: Stoichiometry Rationale: Stoichiometry describes the quantitative relationships between reactants and products in chemical reactions. Stoichiometric calculations depend upon balanced chemical equations. The coefficients of the balanced equation indicate the ratio of reactants and products taking part in the reaction. . Learning Objectives: Students should be able to: 1. Define stoichiometry. (C-K) 2. Relate and compare stoichiometry to balancing chemical equations. (C-AP) 3. Describe, citing examples, how stoichiometry affects personal and public-policy decision-making. (A-V) Important terms: stoichiometry, mole-mole problem WI State Standards: A.12.2, A.12.5, D.12.4, D.12.5 Materials: none Procedure: 5 minutes 1. Review rules for balancing equations a. Write a formula equation with correct symbols and formulas. b. Count the number of atoms of each element on each side of the equation. c. Balance atoms by using coefficients. d. Check your work by counting atoms of each element. 5 minutes Students should take notes during this portion of lecture. 2. Explain what stoichiometry is. - study of quantitative, or measurable, relationships that exist in chemical formulas and chemical reactions
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Give examples, then ask students to think of some (i.e., how much of A reacts with how much of B, etc.) Connect importance of balanced equations with stoichiometry
Q: How might this be important in the real world? A: manufacturing, medicine, other examples Write chemical reactions of CH4 vs. C3H8 reacting with O2. From a purely environmental standpoint, which fuel produces more CO2 per mole? 5-10 minutes 3. Introduce mole-mole problems Work sample problem on board for class. Go slow enough so they can take notes. Q: How many moles of HCl are needed to react with 5.70 moles of Zn? Step 1. Write balanced equation. 2 HCl + Zn → ZnCl2 + H2 Step 2. Determine molar ratio. (2 HCl for every 1 Zn) Step 3. Cross multiply with given number of moles. HCl / Zn = 2 / 1 = X / 5.70 X = 11.40 moles HCl 30 minutes 4. Allow students the rest of the class period to work on homework assignment. Associated Reading: p. 348-353 Assessment: -
Students will complete Section Review (1-4), p. 353. Due next class period.
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Review how stoichiometry and balanced equations are related. Review mole-mole problems. Explain how today’s lesson will allow us to solve stoichiometric problems.
Conclusion:
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Lesson Plan 6 Unit:
Stoichiometry
Subject: Stoichiometry Lab Prep Rationale: Stoichiometry describes the quantitative relationships between reactants and products in chemical reactions. Stoichiometric calculations depend upon balanced chemical equations. The coefficients of the balanced equation indicate the ratio of reactants and products taking part in the reaction. Learning Objectives: Using their knowledge of stoichiometry, students will be able to design a lab experiment to determine molar ratios in a reaction. Important terms: stoichiometry, mole-mole problem WI State Standards: A.12.3, D.12.4, D.12.5, D.12.6, G.12.2 Materials: none Procedure: Entire class period Divide students into groups of 4. Given the following items: 1 M NaOH 1 M HCL 1 M H2SO4 2 burets beaker pH indicator Students will design an experiment to determine how they can calculate the molar ratio of sodium hydroxide to hydrochloric acid, and sodium hydroxide to sulfuric acid. Note: Do not provide students the balanced equations! By next class, students should have an introduction and procedure written in their lab notebooks. Students should also select LAMP duties as done in previous labs.
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Note: Be sure to stress SAFETY. Students should include SAFETY precautions in their procedure. Assessment: -
Students will hand in a copy of the procedure at the beginning of next class period.
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Review the parts of the lab write-up students need to have done for next class. Remind students that next class period they will be performing the lab.
Conclusion:
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Lesson Plan 7 Unit:
Stoichiometry
Subject: Stoichiometry Lab Rationale: Stoichiometry describes the quantitative relationships between reactants and products in chemical reactions. Stoichiometric calculations depend upon balanced chemical equations. The coefficients of the balanced equation indicate the ratio of reactants and products taking part in the reaction. Learning Objectives: Students should be able to: 1. Use titration techniques to calculate molar ratios in a chemical equation. (C-S) 2. Evaluate lab results and summarize in a lab write-up. (C-E) Important terms: stoichiometry, mole-mole problem WI State Standards: A.12.3, C.12.3, D.12.4, D.12.5, D.12.6, G.12.2 Materials: For each group: 1 M NaOH 1 M HCL 1 M H2SO4 2 burets beaker pH indicator Procedure: Entire class period Student groups will perform the lab experiment they designed. Instructor will be available to answer questions and steer groups in the right direction. Make sure students are wearing proper personal protective equipment. Final lab write-up should include: Introduction Conclusions (with balanced chemical equations) 32
Procedure Discussion (students should include what they might do differently next time, what went well, etc.) Assessment: -
Observe students during lab to make sure all are participating. Lab notebooks with lab write-up will be handed in before next class period.
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Have students summarize what they liked/disliked about lab. Explain how knowledge learned in last 3 lessons will allow them to solve stoichiometric problems (next lesson).
Conclusion:
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Lesson Plan 8 Unit:
Stoichiometry
Subject: Solving Stoichiometry Problems Rationale: The major categories of stoichiometry problems are mass-mass, mass-volume, and volumevolume problems. The molar ratios in balanced equations can be used to determine the mass of one product or reactant if the mass of another product or reactant is known. Similarly, the volume of a gas in a reaction can be calculated from the mass or volume of another substance in the reaction. Learning Objectives: Students should be able to: Identify and solve different types of stoichiometric problems. (C-AN) Important terms: mass-mass problem, mass-volume problem, volume-volume problem WI State Standards: D.12.4, D.12.5 Materials: Stoichiometry Problems worksheet Procedure: Students will take notes during this portion of the lecture. 5 minutes 1. Explain that moles are the common quantitative unit for these problems. 10-15 minutes 3. Work sample mass-mass problems on the board. Have students help lead you through process. 10-15 minutes 4. Work sample mass-volume problems on the board. Have students help lead you through process. Review molar volume of gas at STP = 22.4 L/mol.
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10-15 minutes 5. Work sample mass-volume problems on the board. Have students help lead you through the process. Associated Reading: p. 354-364 Assessment: -
Students will complete homework assignment for next class period.
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Review 3 types of problems. Remind students to meet in computer lab for next lesson.
Conclusion:
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Chemistry
Name: Period:
Stoichiometry Problems Worksheet
1. Determine the mass of lithium hydroxide produced when 0.38 g of lithium nitride reacts with water according to the following equation: Li3N + 3 H2O → NH3 + 3 LiOH
2. Determine the mass of sodium chloride produced when chlorine reacts with 0.29 g of sodium iodide.
3. How many liters of carbon dioxide are produced at STP when 400 g of CaCO3 react with HCl according to the equation: CaCO3 + 2 HCl → CaCl2 + CO2 + H2O
4. How many liters of oxygen are necessary for the combustion of 425 g sulfur, assuming the reaction takes place at STP?
5. When 0.75 L of hydrogen gas reacts with bromine gas, what volume of hydrogen bromide gas will be produced?
6. Propane (C3H8) burns in oxygen to produce carbon dioxide and water vapor. What volume of CO2 is produced when 2.8 L O2 is consumed?
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Lesson Plan 9 Unit:
Stoichiometry
Subject: Stoichiometry Web Site Review Rationale: The major categories of stoichiometry problems are mass-mass, mass-volume, and volumevolume problems. The molar ratios in balanced equations can be used to determine the mass of one product or reactant if the mass of another product or reactant is known. Similarly, the volume of a gas in a reaction can be calculated from the mass or volume of another substance in the reaction. Learning Objectives: Students will be able to: 1. Identify and solve different types of stoichiometric problems. (C-AN) 2. Evaluate a chemistry website on the internet. (C-E) Important terms: see previous stoichiometry lessons WI State Standards: D.12.4, D.12.5 Materials: none Procedure: Class will meet in computer lab. Have students pair up and go to the website http://www.quia.com/jg/4059.html. Students can challenge each other to games or can take turns playing. If taking turns, students should help each other review. Assessment: Observe students to make sure they are all participating and understanding the concepts of the games. Conclusion: -
Q: What happens if there is an excess amount of one reactant? A: Next lesson.
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Lesson Plan 10 Unit:
Stoichiometry
Subject: Limiting Reactants and Percent Yield Rationale: When reactants are not present in stoichiometric proportions, the amount of product produced is determined by the quantity of the limiting reactant. The limiting reactant is completely used up in the reaction, and any reactants present in excess will be left over. The percent yield of a reaction is calculated by dividing the amount of product actually produced by the amount of product predicted by the balanced equation for the reaction. Percent yield is usually less than 100 percent. Learning Objectives: 1. Be able to determine the limiting reactant of a chemical reaction. (C-AN) 2. Be able to calculate the amount of product formed in a chemical reaction when reactants are present in nonstoichiometric proportions. (C-AN) Important terms: limiting reactant, expected yield, actual yield, percent yield WI State Standards: D.12.4, D.12.5 Materials: Limiting Reactants and Percent Yield worksheet Procedure: Students will take notes during this portion of the lecture. 10-15 minutes 1. Explain to students how to identify the limiting reactant in a reaction. 10-15 minutes 2. Work sample limiting reactant problems on the board for the class. 10-15 minutes 3. Explain concepts of expected yield, actual yield, and how to calculate percent yield. Work a sample problem on the board.
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Assessment: -
Students will complete homework assignment. Due next class period.
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Review concepts from lecture. Preview next lesson’s lab.
Conclusion:
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Chemistry
Name: Period:
Percent Yield/Limiting Reactants Worksheet
1. What is the limiting reactant if 1.22 g of O2 react with 1.05 g of H2 to produce water?
2. Identify the limiting reactant when 5.1 g of lithium react with 1.5 L of fluorine gas at STP to produce lithium fluoride.
3. When 4.1 g of Cr is heated with 9.3 g of Cl2, what mass of CrCl3 will be produced?
4. What mass of SO2 is produced from the reaction between 31.5 g of S and 8.65 g of O2?
5. Determine the percent yield for the reaction between 3.74 g of Na and excess O2 if 5.34 g of Na2O2 is recovered.
6. Determine the percent yield for the reaction between 6.92 g of potassium and 4.28 g of oxygen if 7.36 g of KO2 is produced.
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Stoichiometry Unit Test
Name _____________________ Date_______________________ Period_____________________
TOTAL SCORE:
Directions: Place your name, the date, and the period in the upper right hand corner. This test is worth a total of 100 points. Read all of the directions carefully for each section of the test. Calculators may be used. Use the periodic table on the wall for reference when needed. Scrap paper will be provided and will be collected with your test. A. Multiple Choice (20 points): Select the best answer to each question. Please print your answer using capital letters only. 1. Which of the following has the greatest mass? A. 4.2 mol of carbon B. 8.34 x 1024 atoms of lead C. 9500 formula units of calcium carbonate (CaCO3) D. 12.6 g aluminum nitrate (Al2(NO3)2) 2. The mass of one mole of lithium hydroxide (LiOH) is: A. 24 g B. 31 g C. 25 g D. 47 g 3. Percent yield is the quantity of product actually produced compared with the quantity: A. of product expected. B. of the limiting reactant. C. usually produced on average. D. of the reactant in excess. 4. The equation for the synthesis of ammonia is: N2 + 3 H2 → 2 NH3. How may moles of H2 are needed to produce 6 mol NH3? A. 4 B. 6 C. 8 D. 9 5. If 16.4 g Na2CrO4 are combined with 26.2 g AgNO3 in a double replacement reaction, what will be the limiting reactant? A. Na2CrO4 B. AgNO3 C. NaNO3 D. Ag2CrO4
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6. The empirical formula and molecular formula for a compound: A. are the same B. are different C. can be the same or different 7. The percent composition of Mg in Mg(OH)2 is: A. 58.3 % B. 61.4 % C. 41.7 % D. 24.8 % 8. In 60.0 g of N2O there are: A. 3.61 x 1025 atoms B. 1.00 x 1023 molecules C. 8.21 x 1023 molecules D. 2.77 x 1025 formula units 9. The coefficients in a balanced chemical equation represent ratios of all of the following except: A. mass B. moles C. volumes D. particles 10. The empirical formula for a compound that contains 36 g carbon, 8 g oxygen, and 6 g hydrogen is: A. C4H6O B. C6H12O C. C9H2O4 D. C6H12O6 B. Fill in the Blank (20 points): Write the correct answer in the space provided. 1. One mole of gas occupies _____________ at STP. 2. The amount of product recovered from a chemical reaction is the ______________ yield. 3. One mole of calcium carbonate contains _______________ particles. 4. The molar mass of HCl is ______________. 5. When 3 moles of calcium chlorate are decomposed in the reaction Ca(ClO3)2 → CaCl2 + 3 O2, _______ moles of oxygen gas are produced.
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6. When 5.0 g of aluminum combine with 3.0 g of sulfur to produce aluminum sulfide (Al2S3), the limiting reactant is ______________. 7. While a molecule is the smallest representation of a molecular compound, a formula unit is the smallest representation of a(n) ______________ compound. 8. All stoichiometry problems are based upon the _____________ ratios that exist between the substances in a reaction. 9. _____________________ is the study of the quantitative, or measurable, relationships that exist in chemical formulas and chemical reactions. 10. In the reaction of H2SO4 with NaOH, the molar ratio of H2SO4 to NaOH is _________.
C. Matching (16 points): Match the descriptions below with the correct answer from the list of words on the right. Print, and use capital letters only. No words will be used twice. 1. percent by mass of each element in a compound
A. molar volume
2. mass in grams of one mole of a substance
B. atomic mass
3. mass of an atom expressed relative to the mass of one carbon-12 atom
C. Avogadro’s number
4. sum of the atomic masses of all the atoms in a compound
D. expected yield
5. the number of particles in a mole is known as this
E. formula mass
6. the amount of product that should be produced in a chemical reaction
F. molar mass
7. the volume occupied by one mole of gas at STP is known as
G. limiting reactant
8. determines the quantity of products in a chemical reaction
H. percent composition
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D. Short Answer. Answer each question below as completely as possible. Point values are given for each question. Partial credit will be given. 1. List the 3 main types of stoichiometric problems. (3 points) a. b. c. 2. What is the difference between an empirical formula and a molecular formula? (3 points)
3. List the 4 steps involved in solving a stoichiometric problem. (4 points) a. b. c. d. 4. Cite two examples of how stoichiometry is important in everyday life. (2 points) a. b. 5. Moles allow for easy conversion among these three quantities. (3 points) a. b. c.
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6. Name the two acids used in the titration lab. (2 points) a. b. E. Essay. (5 points) Answer the following question as completely as possible. Partial credit will be given. Briefly describe how titration techniques can be used to determine molar ratios in a chemical reaction.
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F. Problem-Solving. Complete the following problems. Show all your work. Point values for each problem are given. Partial credit will be awarded. 1.
The explosive, TNT, is composed of 37.0% carbon, 2.20% hydrogen, 18.5% nitrogen, and 42.3% oxygen. (5 points) a. Determine the empirical formula for TNT.
b. The molar mass for TNT is 227 g/mol. What is the molecular formula for TNT?
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2.
What mass of barium chloride (BaCl2) is needed to react completely with 46.8 g of sodium phosphate according to the given equation. (4 points) 3 BaCl2 + 2 Na3PO4 → Ba3(PO4)2 + 6 NaCl
3.
When octane (C8H18) is burned in oxygen, carbon dioxide and water are produced. If 320 g of octane are burned and 392 g of water are recovered, what is the percent yield of the experiment? (5 points) C8H18 + 25 O2 → 16 CO2 + 18 H2O
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4. Determine the percent composition of sucrose (C12H22O11). (3 points)
5. What mass of lead(II) iodide will be produced when 16.4 g of lead(II) nitrate is added to 28.5 g of potassium iodide? What is the limiting reactant? (5 points) Pb(NO3)2 + 2 KI → PbI2 + 2 KNO3
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Stoichiometry Unit Test TOTAL SCORE:
ANSWER KEY
Name _____________________ Date_______________________ Period_____________________
Directions: Place your name, the date, and the period in the upper right hand corner. This test is worth a total of 100 points. Read all of the directions carefully for each section of the test. Calculators may be used. Use the periodic table on the wall for reference when needed. Scrap paper will be provided and will be collected with your test. A. Multiple Choice (20 points): Select the best answer to each question. Please print your answer using capital letters only. B
1. Which of the following has the greatest mass? A. 4.2 mol of carbon B. 8.34 x 1024 atoms of lead C. 9500 formula units of calcium carbonate (CaCO3) D. 12.6 g aluminum nitrate (Al2(NO3)2)
A
2. The mass of one mole of lithium hydroxide (LiOH) is: A. 24 g B. 31 g C. 25 g D. 47 g
A
3. Percent yield is the quantity of product actually produced compared with the quantity: A. of product expected. B. of the limiting reactant. C. usually produced on average. D. of the reactant in excess.
D
4. The equation for the synthesis of ammonia is: N2 + 3 H2 → 2 NH3. How may moles of H2 are needed to produce 6 mol NH3? A. 4 B. 6 C. 8 D. 9
A
5. If 16.4 g Na2CrO4 are combined with 26.2 g AgNO3 in a double replacement reaction, what will be the limiting reactant? A. Na2CrO4 B. AgNO3 C. NaNO3 D. Ag2CrO4
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C
6. The empirical formula and molecular formula for a compound: A. are the same B. are different C. can be the same or different
C
7. The percent composition of Mg in Mg(OH)2 is: A. 58.3 % B. 61.4 % C. 41.7 % D. 24.8 %
C
8. In 60.0 g of N2O there are: A. 3.61 x 1025 atoms B. 1.00 x 1023 molecules C. 8.21 x 1023 molecules D. 2.77 x 1025 formula units
A
9. The coefficients in a balanced chemical equation represent ratios of all of the following except: A. mass B. moles C. volumes D. particles
B
10. The empirical formula for a compound that contains 36 g carbon, 8 g oxygen, and 6 g hydrogen is: A. C4H6O B. C6H12O C. C9H2O4 D. C6H12O6
B. Fill in the Blank (20 points): Write the correct answer in the space provided. 1. One mole of gas occupies _22.4 L________ at STP. 2. The amount of product recovered from a chemical reaction is the __actual_______ yield. 3. One mole of calcium carbonate contains ___6.02 x 1023________ particles. 4. The molar mass of HCl is ___36.5 g/mol______. 5. When 3 moles of calcium chlorate are decomposed in the reaction Ca(ClO3)2 → CaCl2 + 3 O2, ___9____ moles of oxygen gas are produced.
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6. When 5.0 g of aluminum combine with 3.0 g of sulfur to produce aluminum sulfide (Al2S3), the limiting reactant is _sulfur__________. 7. While a molecule is the smallest representation of a molecular compound, a formula unit is the smallest representation of a(n) __ionic_______ compound. 8. All stoichiometry problems are based upon the __molar______ ratios that exist between the substances in a reaction. 9. _Stoichiometry_____________ is the study of the quantitative, or measurable, relationships that exist in chemical formulas and chemical reactions. 10. In the reaction of H2SO4 with NaOH, the molar ratio of H2SO4 to NaOH is _1:2_____.
C. Matching (16 points): Match the descriptions below with the correct answer from the list of words on the right. Print, and use capital letters only. No words will be used twice. H
1. percent by mass of each element in a compound
A. molar volume
F
2. mass in grams of one mole of a substance
B. atomic mass
B
3. mass of an atom expressed relative to the mass of one carbon-12 atom
C. Avogadro’s number
E
4. sum of the atomic masses of all the atoms in a compound
D. expected yield
C
5. the number of particles in a mole is known as this
E. formula mass
D
6. the amount of product that should be produced in a chemical reaction
F. molar mass
A
7. the volume occupied by one mole of gas at STP is known as
G. limiting reactant
G
8. determines the quantity of products in a chemical reaction
H. percent composition
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D. Short Answer. Answer each question below as completely as possible. Point values are given for each question. Partial credit will be given. 7. List the 3 main types of stoichiometric problems. (3 points) a. mass - mass b. mass - volume c. volume - volume 8. What is the difference between an empirical formula and a molecular formula? (3 points) An empirical formula gives the simplest whole number ratio of atoms of elements in a compound. A molecular formula gives the actual number of atoms of elements in a compound. 9. List the 4 steps involved in solving a stoichiometric problem. (4 points) a. Write a balanced chemical equation. b. Convert given information to moles. c. Determine molar ratio between given and unknown. d. Convert moles of unknown to units you are seeking. 10. Cite two examples of how stoichiometry is important in everyday life. (2 points) a. Look for specific examples of safety, economics, public policy decisions, consumer decisions b. 11. Moles allow for easy conversion among these three quantities. (3 points) a. mass b. particles c. volume
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12. Name the two acids used in the titration lab. (2 points) a. HCl b. H2SO4 E. Essay. (5 points) Answer the following question as completely as possible. Partial credit will be given. Briefly describe how titration techniques can be used to determine molar ratios in a chemical reaction. If you use solutions of equal concentration, one substance can be added to another until an indicator turns color, indicating the reaction is complete. The volume of each solution used can be measured and calculated as a ratio. This ratio gives you the molar ratio of the chemical reaction.
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F. Problem-Solving. Complete the following problems. Show all your work. Point values for each problem are given. Partial credit will be awarded. 1.
The explosive, TNT, is composed of 37.0% carbon, 2.20% hydrogen, 18.5% nitrogen, and 42.3% oxygen. (5 points) a. Determine the empirical formula for TNT. C
37.0 g = 3.08 mol
/1.32 = 2.33
x3 = 7
H
2.2 g = 2.18 mol
/1.32 = 1.66
x3 = 5
N
18.5 g = 1.32 mol
/1.32 = 1.00
x3 = 3
O
42.3 g = 2.64 mol
/1.32 = 2.00
x3 = 6
C7H5N3O6 or C7H5(NO2)3
b. The molar mass for TNT is 227 g/mol. What is the molecular formula for TNT? Formula weight of C7H5(NO2)3 = 227 g/mol. Therefore the molecular formula is C7H5(NO2)3.
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2.
What mass of barium chloride (BaCl2) is needed to react completely with 46.8 g of sodium phosphate according to the given equation. (4 points) 3 BaCl2 + 2 Na3PO4 → Ba3(PO4)2 + 6 NaCl molecular weight of Na3PO4 = 164.0 g/mol. molecular weight of BaCl2 = 208.2 g/mol. 46.8 g/ 164 g/mol = 0.285 moles Na3PO4 molar ratio = 2/3 = 0.285/x
x = 0.428 moles BaCl2
0.428 moles BaCl2 x 208.2 g/mol = 89.0 g BaCl2
3.
When octane (C8H18) is burned in oxygen, carbon dioxide and water are produced. If 320 g of octane are burned and 392 g of water are recovered, what is the percent yield of the experiment? (5 points) C8H18 + 25 O2 → 16 CO2 + 18 H2O Molecular weight of octane = 114 g/mol 320 g octane / 114 g/mol octane = 2.8 mol octane molar ratio = 2 / 18 = 2.8 / x
x = 25.3 mol H2O
25.3 mol H2O x 18.0 g/mol = 454.7 g H2O = expected yield 392 / 454.7 = 86.2 %
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4. Determine the percent composition of sucrose (C12H22O11). (3 points) molecular weight of sucrose = 342 g/mol C
12.0 g/mol x 12 = 144
/342
42.1%
H
1.0 g/mol x 22 = 22
/342
6.4%
O
16.0 g/mol x 11 = 176
/342
51.5%
5. What mass of lead(II) iodide will be produced when 16.4 g of lead(II) nitrate is added to 28.5 g of potassium iodide? What is the limiting reactant? (5 points) Pb(NO3)2 + 2 KI → PbI2 + 2 KNO3 (16.4 g Pb(NO3)2 x 331 g Pb(NO3)2 / mol) x (1 mol PbI2 / 1 mol Pb(NO3)2 ) x (461 g PbI2 / mol PbI2 ) = 22.8 g PbI2
limiting reactant
28.5 g KI x(1 mol KI / 166 g KI) x (1 mol PbI2 / 2 mol KI) x 461 g PbI2 / mol PbI2 = 39.6 g PbI2
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CORRELATION BETWEEN UNIT OBJECTIVES AND TEST QUESTIONS Unit Objectives
Test Question
Lesson 1 Chemical Measurements 1. Define a mole and describe its importance. C-K)
A2, homework
2. Identify and use Avogadro’s number. (C-AP)
B3, C5, homework
3. Define molar mass and use it to relate the mass of a substance to the number of particles in that substance. (C-C)
B4, C2, homework
LESSON 2 Mole Conversions 1. Convert among the number of particles, moles, and mass of a substance. (C-AP) 2. Define molar volume and use it to solve problems. (C-AP)
A1, A9, D5, homework B1, C7
LESSON 3 Empirical and Molecular Formulas 1. Calculate the percent composition of elements in a given formula. (C-AN) 2. Describe how percent composition can be used to make intelligent consumer decisions. (A-V) 3. Use percent composition to find the formula of an unknown sample. (C-AN) 4. Determine empirical and molecular formulas. (C-AN)
A7, C1, F4, homework D4, class discussion F1 A6, A10, D2, F1, homework
LESSON 4 Review Sections 10.1, 10.2, 10.3 1. Students will be able to explain how chemical measurements, mole conversions, and empirical and molecular formulas relate to one another. (C-C)
Class game
LESSON 5 Stoichiometry 1. Define stoichiometry. (C-K)
B9, homework
2. Relate and compare stoichiometry to balancing chemical equations. (C-AP)
A3, A9, B5, homework
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3. Describe, citing examples, how stoichiometry affects personal and public-policy decision-making. (A-V)
D4
LESSON 6 Stoichiometry Lab Prep 1. Using their knowledge of stoichiometry, students will Lab write-up be able to design a lab experiment to determine molar ratios in a reaction. LESSON 7 Stoichiometry Lab 1. Use titration techniques to calculate molar ratios in a chemical equation. (C-S) 2. Evaluate lab results and summarize in a lab write-up. (C-E)
B10, E1, lab write-up Lab write-up
LESSON 8 Solving Stoichiometry Problems 1. Identify and solve different types of stoichiometric problems. (C-AN)
D1, D3, F2, homework
LESSON 9 Stoichiometry Web Site Review 1. Identify and solve different types of stoichiometric problems. (C-AN) 2. Evaluate a chemistry website on the internet. (C-E)
D1, D3, F2, homework homework
LESSON 10 Limiting Reactants and Percent Yield 1. Be able to determine the limiting reactant of a chemical reaction. (C-AN)
A5, B6, F5, homework
2. Be able to calculate the amount of product formed in a chemical reaction when reactants are present in nonstoichiometric proportions. (C-AN)
B2, C8, F3, F5, homework
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ALTERNATE ASSESSMENT STRATEGIES Because some objectives can’t be addressed appropriately through a standard test, alternate assessment strategies must be used. The alternate assessment strategies I chose and the objectives they apply to are discussed below. Alternate Strategy #1 – Writing a letter to a public official I would use this strategy with objective 5-3. Students would take a stand on an issue and back it up using evidence they have obtained through research, class lectures, etc. In addition to giving them real-life examples of how stoichiometry is important, it would hopefully help them to become good citizens as well. As evidence of the student’s success, I would review the letter before they send it out, and also have them bring me the response they receive from the public official. In an assignment like this, I would most likely give the points in an all or none fashion. If the student made an effort and wrote a decent letter, they would receive the points. I may consider giving extra credit points for bringing the response from the public official. Alternate Strategy #2 – Designing a game or website I would use this strategy with objective 9-2. Students would be allowed to use their creativity, while also integrating computers into the project. Students would also learn what makes a good website or game. As evidence of the student’s success, I would have the student give me a demonstration of the game or take me on a “tour” of their website. As far as grading considerations, I would provide the students with a rubric outlining the criteria, which would include such topics as creativity, organization, level of difficulty, etc. Alternate Strategy #3 – Designing a lab experiment I would employ this strategy with objective 6-1. Like strategy #2, this strategy focuses on the affective rather than cognitive abilities of the student. Students must use prior knowledge to predict what will happen in their design. As evidence of the student’s success, I would have the student include their procedure in the written lab report they turn in at the completion of the lab. To make sure they are on the right track, I would preview their procedures before the lab. For grading, I would have a rubric for the students showing them what a complete lab report should include. For the procedure they design, I would have specific items that I would be looking for and would assign points based on that criteria.
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CONCLUSION Although writing this unit plan was difficult and time-consuming, it was probably one of the most worthwhile tasks I have performed while in college. It is unlikely that I will write this detailed of a unit plan when I am teaching, mainly for lack of time, but writing this unit plan forced me to consider what goes into a good lesson. Even if I don’t write it down in future lesson plans I will at least be thinking of what I should include to make the lesson meaningful and effective.
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