Balancing Chemical Equations PART I: Background Information: Word and Formula Equations The first step in writing a chemical equation is to identify the facts to be represented. It is helpful to write a word equation, an equation in which the reactants (left) and products (right) in a chemical reaction are represented by words. A word equation has only qualitative (descriptive) meaning. It does not give the whole story because it does not give the quantities of reactants used or products formed. A formula equation represents the reactants (left) and products (right) of a chemical reaction by their symbols or formulas. Chemical formulas must be written correctly according to the rules of nomenclature. Word Equation (A) An aqueous solution of hydrochloric acid reacts with an aqueous solution of sodium hydroxide to produce liquid water and aqueous sodium chloride. Formula Equation (A) HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) Word Equation (B) Solid calcium carbonate decomposes when heated to yield solid calcium oxide and carbon dioxide gas. Formula Equation (B) CaCO 3 (s) CaO (s) + CO 2 (g) Word Equation (C) Gaseous ethene and gaseous hydrogen will react in the presence of a platinum catalyst to form ethane gas. Formula Equation (C) C 2 H 4 (g) + H 2 (g) Pt C 2 H 6 (g) Group Activity: 1. What does the arrow represent in a chemical equation?
2. What do the following symbols in the equations above represent? (s) (g) (l) (aq) Pt 3. In the table below, indicate the reactant(s) and product(s) for each of the equations above. EQUATION REACTANTS PRODUCTS Name Formula Name Formula A B C 4. In chemical equation (C), is platinum a reactant, a product, or neither? Explain. Model: Atoms Are Conserved in Chemical Reactions. Chemical reactions are described by chemical equations. According to the Law of Conservation of Mass (Matter), atoms are neither created nor destroyed when chemical reactions take place. Therefore, the number of atoms of each element must be identical on the reactant (left) and the product (right) sides of a balanced chemical reaction.
Reaction (D) CuO (s) + H 2 (g) Cu(s) + H 2 O(g) + + Reaction (E) CO(g) + O 2 (g) CO 2 (g) + Group Activity: 5. How many of each type of atom are represented on each side of the equation in reaction (D)? Reactant Side Product Side Cu O H 6. Does this equation (D) obey the Law of Conservation of Mass (Matter)? Explain. 7. How many of each type of atom are represented on each side of the equation in reaction (E)? Reactant Side Product Side C O 8. Does this equation (E) obey the Law of Conservation of Mass (Matter)? Explain. 9. If either Reaction D or E does not obey the Law of Conservation of Mass (Matter), draw additional molecules on the reactant and/or product side in order to make the number of all atoms equal on both sides of the equation. (Remember that once written correctly, chemical formulas cannot be changed. In other words, molecules must remain intact.)
10. In reaction (D), how many H 2 O molecules are produced for every H 2 molecule that is consumed? 11. In reaction (E), how many CO 2 molecules are produced for every O 2 molecule consumed? 12. In reaction (E), how many CO molecules react with every O 2 molecule? Independent Activity: 1. Read pages 243 through 248. 2. Write word and formula equations for the chemical reaction that occurs when solid sodium oxide is added to water at room temperature and forms sodium hydroxide (dissolved in the water). Include symbols for physical states in the formula equation. 3. Translate the following chemical equation into a sentence: PbCl 2 (aq) + Na 2 CrO 4 (aq) PbCrO 4 (s) + NaCl(aq) 4. Work problems 18 and 19 on page 270.
Balancing Chemical Equations PART II: Background Information: Balanced Chemical Equations A balanced chemical equation represents a chemical reaction that obeys the Law of Conservation of Mass (Matter). It can be thought of as describing how many atoms or molecules of reactants are consumed in order to produce a certain number of atoms or molecules of products. A coefficient is a small whole number that appears in front of a formula in a chemical equation. The coefficient represents the relative number of atoms or molecules. Recall from Part I of this activity the following reactions: Reaction (D) CuO (s) + H 2 (g) Cu(s) + H 2 O(g) + + Reaction (E) CO(g) + O 2 (g) CO 2 (g) + In order to answer question 9, Reaction D was already drawn correctly, but Reaction E should have been redrawn as pictured below: + This picture shows 2 carbon atoms and 4 oxygen atoms on each side of the equation. These atoms are represented as 2 CO molecules and 1 O 2 molecule on the reactant side and 2 CO 2 molecules on the product side of the equation. Group Activity: 1. Based on the correct drawings write a formula equation that will obey the Law of Conservation of Mass (Matter) for both Reaction D and E. In order to do this, coefficients may be required.
2. Is the number of molecules identical on the reactant and product sides of both of these balanced equations? 3. Explain how your answer to # 2 can be consistent with the idea that atoms are neither created nor destroyed when chemical reactions take place. 4. Why must we balance chemical equations? Model: Balanced Chemical Equations a) Zn (s) + 2 HCl (aq) H 2 (g) + ZnCl 2 (aq) b) 2 H 2 (g) + O 2 (g) 2 H 2 O (g) c) BaCl 2 (aq) + 2 AgNO 3 (aq) Ba(NO 3 ) 2 (aq) + 2 AgCl (s) d) 2 C 6 H 14 (g) + 19 O 2 (g) 12 CO 2 (g) + 14 H 2 O (g) Group Activity: 5. What is the coefficient for hydrochloric acid in equation a? 6. What is the coefficient for barium nitrate in equation c? 7. How many barium nitrates are on the product side of equation c? Does your answer agree with your answer for # 6 above? 8. How many barium atoms are on each side of equation c? Explain how you arrived at this answer. 9. How many carbon atoms are on each side of equation d? Explain how you arrived at this answer. 10. How many oxygen atoms are on each side of equation d? Explain how you arrived at this answer.
11. Balance the following chemical equations: a) Cr(s) + S 8 (s) Cr 2 S 3 (s) b) NaHCO 3 (s) Na 2 CO 3 (s) + CO 2 (g) + H 2 O (g) c) Fe 2 S 3 (s) + HCl (g) FeCl 3 (s) + H 2 S (g) d) CS 2 (l) + NH 3 (g) H 2 S (g) + NH 4 SCN (s) e) Fe 2 (SO 4 ) 3 (aq) + KOH (aq) K 2 SO 4 (aq) + Fe(OH) 3 (s) f) C 25 H 52 (l) + O 2 (g) CO 2 (g) + H 2 O (g) 12. The trial-and-error method for balancing equations can be successful; however it can be made easier by following some simple guidelines. Develop a few simple guidelines to achieve a balanced chemical equation. 13. Use your guidelines to balance the following chemical equations: a) H 3 PO 4 H 4 P 2 O 7 + H 2 O b) C 6 H 6 + O 2 CO 2 + H 2 O c) Al(OH) 3 + H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 O 14. What adjustments, if any, need to be made to your guidelines in order to balance the equations above? Independent Activity: 15. Read pages 250 through 254. Carefully study the sample problems 8.3, 8.4, and 8.5 on pages 252 254. 16. Work practice problem on page 254. 17. Write the balanced chemical equation for the gaseous reaction of methane (CH 4 ) with oxygen (O 2 ) to form carbon dioxide (CO 2 ) and water (H 2 O). 18. Chapter 8: page 270, questions #20 and #21. 19. Chapter 8: page 272, questions #41 and #47.
Balancing Chemical Equations Teacher Notes Time Required: Estimated 20-30 minutes for Part I and 50 minutes for Part II. It is suggested that Part I and Part II be completed on consecutive days. Prerequisites: Chemical formulas and nomenclature Learning Objectives: 1. Write word equations correctly. 2. Write formula equations correctly, using proper symbolic notation. 3. Correctly balance chemical equations. 4. Explain how a balanced chemical equation illustrates the Law of Conservation of Mass (Matter). Considerations: This activity is a wonderful way to introduce the concept of writing and balancing chemical equations. Be careful not to answer the students questions for them. If a group asks a question concerning content, rather than supplying the answer, ask a question in return that will steer them toward discovering the answer for themselves. Another way to intervene with the group may be to have the Spokesperson or Presenter for the group to restate their reasoning so that they have the opportunity to pinpoint misconceptions or faulty reasoning on their own. Closure: It is always good to have the students reflect on what they have learned before leaving class. One way to do this is to stop the activity about 5 minutes before the end of class and have students take out a sheet of paper. Have them record one thing that they learned as a result of the lesson that day. Then have them to ask a question concerning the lesson. The question could be something that is still unclear to them or it could be something they are curious about or a subject that they think may be related to the lesson.
Answers to Questions in Group Activity: Part I: 1. What does the arrow represent in a chemical equation? Yields or produces 2. What do the following symbols in the equations above represent? (s) solid (l) liquid (g) gas (aq) aqueous heat Pt platinum catalyst 3. In the table below, indicate the reactant(s) and product(s) for each of the equations above. EQUATION REACTANTS PRODUCTS Name Formula Name Formula A hydrochloric acid HCl water H 2 O sodium hydroxide NaOH sodium chloride NaCl B calcium carbonate CaCO 3 calcium oxide CaO carbon dioxide CO 2 C ethene C 2 H 4 ethane C 2 H 6 hydrogen H 2 4. In chemical equation (C), is platinum a reactant, a product, or neither? Explain. Neither, because it is not on either side of the arrow.
5. How many of each type of atom are represented on each side of the equation in reaction (D)? Reactant Side Product Side Cu O H 1 1 2 1 1 2 6. Does this equation (D) obey the Law of Conservation of Mass (Matter)? Explain. Yes, because there are the same number of each atom on both sides of the equation. 7. How many of each type of atom are represented on each side of the equation in reaction (E)? Reactant Side Product Side C O 1 3 1 2 8. Does this equation (E) obey the Law of Conservation of Mass (Matter)? Explain. No, because there are NOT the same number of oxygen atoms on each side of the equation. 9. If either Reaction D or E does not obey the Law of Conservation of Mass (Matter), draw additional molecules on the reactant and/or product side in order to make the number of all atoms equal on both sides of the equation. (Remember that once written correctly, chemical formulas cannot be changed. In other words, molecules must remain intact.) + 10. In reaction (D), how many H 2 O molecules are produced for every H 2 molecule that is consumed? one
11. In reaction (E), how many CO 2 molecules are produced for every O 2 molecule consumed? two 12. In reaction (E), how many CO molecules react with every O 2 molecule? two Part II 1. Based on the correct drawings write a formula equation that will obey the Law of Conservation of Mass (Matter) for both Reaction D and E. In order to do this, coefficients may be required. (D) CuO (s) + H 2 (g) Cu(s) + H 2 O(g) already balanced (E) 2 CO(g) + O 2 (g) 2 CO 2 (g) 2. Is the number of molecules identical on the reactant and product sides of both of these balanced equations? No!!!!!!! NOTE: This question asks about molecules, not atoms. In reaction (D), there are 2 molecules on the reactant side and 1 atom and 1 molecule on the product side. In reaction (E), there are 3 molecules on the reactant side and 2 molecules on the product side. If your students do not see this, hint that they look back at their picture from part I. 3. Explain how your answer to # 2 can be consistent with the idea that atoms are neither created nor destroyed when chemical reactions take place. Because the atoms are balanced, even though the molecules are not. 4. Why must we balance chemical equations? In order to obey the Law of Conservation of Mass (Matter). 5. What is the coefficient for hydrochloric acid in equation a? 2 6. What is the coefficient for barium nitrate in equation c? 1 7. How many barium nitrates are on the product side of equation c? Does your answer agree with your answer for # 6 above? 1, yes 8. How many barium atoms are on each side of equation c? Explain how you arrived at this answer. 1, because there is 1 Ba in BaCl 2 and 1 Ba in Ba(NO 3 ) 2
9. How many carbon atoms are on each side of equation d? Explain how you arrived at this answer. 12, because there are 2 molecules of C 6 H 14 and 12 molecules of CO 2 10. How many oxygen atoms are on each side of equation d? Explain how you arrived at this answer. 38, because there are 19 molecules of O 2, 12 molecules of CO 2 = 24 atoms of O, and 14 molecules of H 2 O = 14 atoms of O 11. Balance the following chemical equations: a) 16 Cr(s) + 3 S 8 (s) 8 Cr 2 S 3 (s) b) 2 NaHCO 3 (s) Na 2 CO 3 (s) + CO 2 (g) + H 2 O (g) c) Fe 2 S 3 (s) + 6 HCl (g) 2 FeCl 3 (s) + 3 H 2 S (g) d) CS 2 (l) + 2 NH 3 (g) H 2 S (g) + NH 4 SCN (s) e) Fe 2 (SO 4 ) 3 (aq) + 6 KOH (aq) 3 K 2 SO 4 (aq) + 2 Fe(OH) 3 (s) f) C 25 H 52 (l) + 38 O 2 (g) 25 CO 2 (g) + 26 H 2 O (g) 12. The trial-and-error method for balancing equations can be successful; however it can be made easier by following some simple guidelines. Develop a few simple guidelines to achieve a balanced chemical equation. Student answers may vary. 13. Use your guidelines to balance the following chemical equations: a) 2 H 3 PO 4 H 4 P 2 O 7 + H 2 O b) 2 C 6 H 6 + 15 O 2 12 CO 2 + 6 H 2 O c) 2 Al(OH) 3 + 3 H 2 SO 4 Al 2 (SO 4 ) 3 + 6 H 2 O 14. What adjustments, if any, need to be made to your guidelines in order to balance the equations above? Student answers may vary. However, the purpose of this question is to show students that there are a few helpful guidelines to use when balancing equations: Save H and O for last. Keep polyatomic ions together when possible (treat them as a unit). Combustion reactions may require doubling to remove fractions from the coefficients.
Can A Molecule Be Identified by its Percent Composition? Background Information: Percent Composition The percent composition (by mass) of an element in a molecule is the mass of the element in the molecule divided by the mass of the entire molecule times 100. Or, because the number of atoms (molecules) is proportional to the number of moles of atoms (molecules), Percent composition of element x = mass of x in one mole of the compound x 100 Mass of one mole of the compound % Ba in BaCl 2 = mass of Ba in 1 mole of BaCl 2 x 100 mass of 1 mole of BaCl 2 % Ba in BaCl 2 = 137.3 grams Ba x 100 = 65.91% Ba in BaCl 2 208.3 grams BaCl 2
Group Activity 1. Fill in the missing molecular formulas using the structural formulas in Table 1. 2. Find the % composition of benzene, cyclobutane, and 1-octene. 3. Verify that the % composition given for ethyne in Table in Table 1 is correct. 4. Is it possible, given the original data in Table 1, to determine the % composition by mass of H for 2-butene without using the equation given in the model above? If so, how? 5. Based on the data in Table 1, is it possible to determine the molecular formula of a compound solely from its percent composition? Why or why not? 6. What is similar about the molecular formulas of cyclobutane and 1-octene? 7. Compare the molecular formulas of ethyne and benzene? 8. What mathematical relationship do all compounds with the same % composition have? Background Information: Empirical Formula The empirical formula of a compound describes the relative number of each type of atom in the compound. It is given in terms of the smallest-possible-whole-number ratios (as subscripts). For example, the empirical formula of ethane is CH 3. (Note that the subscript 1 is omitted.) Group Activity 9. Write the empirical formulas for all of the compounds in Table 1. Independent Activity 1. The molecule 2-hexene has the molecular formula C 6 H 12. Refer to Table 1 and determine the percent composition of H in the molecule. 2. Determine the percent composition of each element in acetic acid, CH 3 COOH. 3. A molecule containing only nitrogen and oxygen, contains 30.4% N by mass. (a) How many grams of N would be found in a 100.0 g sample of the compound? (b) How many grams of O would be found in the same sample of the compound? (c) How many moles of N would be found in a 100.0 g sample? (d) How many moles of O would be found in the same sample?
(e) What is the ratio of number of moles of O to number of moles of N? (f) What is the empirical formula of the compound? 4. An unknown liquid contains 38.7% C and 51.6% O by mass. The remainder of the compound is H. What is the empirical formula? 5. Holt, Rinehart, Winston 2002 Text. Chapter 7,p. 229-231 practice problems 1-3, p. 231; p. 232-233, practice problems 1 and 2 and section review problems 1-5, p. 233; p. 237 problems 36-39, 47, 50, and 52.
Limiting Reagent (Reactant) Model 1: The S more A delicious treat known as a S more is constructed with the following ingredients and amounts: 2 graham crackers 1 chocolate bar 2 marshmallows At a particular store, these items can be obtained only in full boxes, each of which contains one gross (144) of items. A gross is a specific number of items, analogous (but not equal) to one dozen. The boxes of items have the following net weights ( the weight of the material inside the box): Box of graham crackers Box of chocolate bars Box of marshmallows 9.0 pounds 36.0 pounds 3.0 pounds Group Activity 1. Based on the information given, which of the three ingredients (a graham cracker, a chocolate bar, or a marshmallow) weighs the most? Which weighs the least? 2. If you have a collection of 100 graham crackers, how many chocolate bars and how many marshmallows do you need to make S mores with all of the graham crackers? 3. If you have a collection of 1000 graham crackers, 1000 chocolate bars, and 1000 marshmallows, how many S mores can you make? What will be left over? 4. If you have 36.0 pounds of graham crackers, 36.0 pounds of chocolate bars, and 36.0 pounds of marshmallows, which item do you have the most of? The least? Explain your reasoning 5. a) If you attempt to make S mores from the material described in question #4, what item will you run out of first? (This item is known to chemists as the limiting reagent because it is the reactant which limits the amount of product which can be made.) b) How many gross of S mores will you have made? c) How many gross of each of the two left over items will you have? d) How many pounds of each of the left over items will you have? e) How many pounds of S mores will you have?
6. Using G as the symbol for graham crackers and Ch as the symbol for the chocolate bars and M for marshmallows, write a balanced chemical equation for the production of S mores. 7. Why is it not correct to state that if we start with 36 pounds each of G, Ch, and M, then we should end up with 3 x 36 = 108 pounds of S mores? Model 2: Water 8. 25.0 grams of hydrogen and 25.0 grams of oxygen react to form water. a) Write the balanced chemical equation and show all steps. b) Assuming that the reaction goes to completion, which element will be totally consumed in the formation of water? c) What is the limiting reagent? d) How many grams water can be produced? e) How many grams of the excess element remain unreacted? Independent Activity #1 Read p. 288-292 Modern Chemistry Text 1. Given the balanced chemical equation: 2 NO(g) + O 2 (g) 2 NO 2 (g) calculate the mass of nitrogen dioxide that can be made from 30 grams of NO and 30 grams of O 2. 2. Zinc and iodine, I 2, react to form zinc iodide, ZnI 2 (the reactants and the product are all solids at room temperature). a) Write a balanced chemical equation for this reaction. b) Suppose that 50.0 g of zinc and 50.0 g of iodine are used to form zinc iodide. 1) Assuming that the reaction goes to completion, which element will be totally consumed in the formation of the zinc iodide? 2) What is the limiting reagent? 3) How many grams of zinc iodine can be produced? 4) How many grams of the excess element remain unreacted?
3. Acetylene gas, HCCH, is commonly used in high temperature torches. a) Write a balanced chemical equation for the reaction of acetylene with hydrogen gas (H 2 ) to form ethane (C 2 H 6 ). b) How many grams of ethane can be produced from a mixture of 30.3 grams of HCCH and 4.14 grams of H 2? 4. Titanium (Ti) is a strong, lightweight metal that is used in the construction of rockets, jet engines, and bicycles. It can be prepared by reacting TiCl 4 with Mg metal at very high temperatures. The products are Ti(s) and MgCl 2 (s). a) Provide a balanced chemical reaction for the reaction described above. b) How many grams of Ti metal can be produced from a reaction involving 3.54 x 10 4 g of TiCl 4 and 6.53 x 10 3 g of Mg? Independent Activity #2 1. Modern Chemistry Text - Page 297, problems 22 27
LABORATORY SAFETY (How can we all survive chemistry labs?) Why: The chemistry laboratory is a potentially hazardous place to work. There are dangerous substances and equipment that must be handled with care. This exercise is designed to construct student knowledge of common lab safety issues. and to focus on the importance of laboratory safety and the procedures to execute before, during, and after laboratory activities. Group Members Group Activity 1. Brainstorm as many lab safety rules as you can within your group (time limit 5 minutes). 2. Determine the top five safety rules from the list generated by your group and rank them in order of importance. 3. Your instructor will supply you with a random list of lab safety rules. Rank these safety rules in order of decreasing importance (most important to least important). 4. Place an asterisk beside the top five rules identified by your group in question 2 that appear on the instructor s list. 5. Compare your group s brainstormed list to the instructor s list. Account for any differences that you find. Developed by Alabama Science in Motion teachers Summer, 2003. Page 1 Revised 8/2/2005
6. Design an emergency procedure to use in the event of each of the situations listed below in which the instructor is incapacitated (unconscious) and students have to react responsibly to the situation to reduce injury and panic. A chemical table or floor fire An individual s clothing catches on fire A cut or a burn injury An eye injury An acid spill A base spill A gas or noxious fumes Developed by Alabama Science in Motion teachers Summer, 2003. Page 2 Revised 8/2/2005
7. Identify these safety symbols and determine their proper location in your school lab. Floor Plan: In the space below draw a sketch of the floor plan of your school s chemistry laboratory. There should be 2 exits from your lab. (NOTE: Sometimes a window is considered an alternative exit.) Locate both exits on your floor plan and label where the safety goggles, safety shower, eye wash station, and fire extinguisher are located. Developed by Alabama Science in Motion teachers Summer, 2003. Page 3 Revised 8/2/2005
Review Activity Each group will choose one of the safety procedures and create a laboratory safety poster for display in the classroom. Developed by Alabama Science in Motion teachers Summer, 2003. Page 4 Revised 8/2/2005
Instructor s Safety List Acids should be transferred to water only (never add water to acid). Aprons must be worn at all times. Bunsen burners should be monitored at all times. Chemicals must be transferred from large stock bottles to small manageable containers. Fire blankets should be used on individuals in case of fire. Fire extinguishers should not be used on individuals. Broken glassware should be disposed of in its proper container. Goggles must be worn at all times. Long hair should be tied back. Loose clothing should not be worn. No eating or drinking in the laboratory. No horseplay in the laboratory. Open-toed shoes should not be worn. Solid substances should not be disposed of in the sink unless approved by your instructor. Unlabeled glassware with unknown substances should not be used. Developed by Alabama Science in Motion teachers Summer, 2003. Page 5 Revised 8/2/2005
The Mole Concept Why: Samples of matter are composed of numerous particles (atoms, ions, or molecules). Because the numbers of particles are so huge, chemists have established a unit of measure called the mole (abbreviated mol) for use in counting them. In this activity you will discover the relationship between the term mole as it relates to the number of particles and the mass of the particles. Model 1: Below are some equivalent quantities with which you may or may not be familiar: 1 pair of shoes = 2 shoes 1 dozen donuts = 12 donuts 1 ream of copy paper = 500 sheets of copy paper 1 gross of pencils = 144 pencils One elephant has one trunk and four legs. Equivalent quantities may be used as conversion factors. A conversion factor is written as a fraction where the numerator and denominator are expressed in different units but are equal quantities. For example: 1 ream of copy paper OR 500 sheets of copy paper 500 sheets of copy paper 1 ream of copy paper Because the numerator and denominator represent the same thing, only in different units, conversion factors are always = 1. Equivalent quantities (conversion factors) used in the study of chemistry: 1 mole of items = 6.022 x 10 23 items (Avogadro s Number) (items = atoms, molecules, ions or formula units) One methane molecule, CH 4, contains one carbon atom and four hydrogen atoms. Key Questions: If your answer is extremely large or small, use scientific notation form. Include units for all answers. Show your work. 1. How many trunks are found in one dozen elephants? Page 1
2. How many legs are found in one dozen elephants? 3. How many trunks are found in a ream of elephants? 4. How many legs are found in a pair of elephants? 5. How many carbon atoms are found in one dozen methane (CH 4 ) molecules? 6. How many hydrogen atoms are found in one dozen methane molecules? 7. How many carbon atoms are found in a pair of methane molecules? 8. How many hydrogen atoms are found in a ream of methane molecules? 9. How many elephants are there in one mole of elephants? 10. How many trunks are found in one-half mole of elephants? 11. How many legs are found in one mole of elephants? 12. How many methane molecules are there in one mole of methane? 13. How many carbon atoms are found in one mole of methane molecules? 14. How many hydrogen atoms are found in one-half mole of methane molecules? Page 2
Model 2: The mass of 1 mole of any pure substance = average atomic mass expressed in grams of that substance. For example: 1.008 amu = average atomic mass for H 1.008 g H = 1 mol H 15. Where would you find the average atomic mass for carbon? 16. What is the average mass (in amu) of one carbon atom? 17. What is the mass (in grams) of one mole of carbon atoms? 18. What is the average mass (in amu) of one methane molecule? 19. What is the mass (in grams) of one mole of methane molecules? 20. How is the mass in amu of one atom of an element related to the mass in grams of one mole of that element? 21. Refer to questions #12 and 19. If you measured out the mass of methane in your response to question 19, how many methane molecules would you have? 22. If you measure out 58.69 g of Ni, how many nickel atoms would you have? 23. If you want to measure out 3.011 x 10 23 atoms of gold, how many grams would you need to measure out? Page 3
Complete the Table below by supplying the necessary information. Substance Number of Particles Mass in grams Al 6.022 x 10 23 atoms H 2 O 18.02 g O 2 1.505 x 10 23 molecules Ne 10.09g Skill Development Exercises: 1. Indicate whether each of the following statements is true or false and explain your reasoning. A. One mole of NH 3 weighs more than one mole of H 2 O. B. There are more carbon atoms in 48 grams of CO 2 than in 12 grams of diamond (a form of pure carbon). C. There are equal numbers of nitrogen atoms in one mole of NH 3 and one mole of N 2. D. The number of Cu atoms in 100 grams of Cu(s) is the same as the number of Cu atoms in 100 grams of copper (II) oxide, CuO. E. The number of Ni atoms in 100 moles of Ni(s) is the same as the number of Ni atoms in 100 moles of nickel (II) chloride, NiCl 2. F. There are more hydrogen atoms in 2 moles of NH 3 than in 2 moles of CH 4. 2. Use grammatically correct sentences to describe how to calculate the number of H atoms in any number of moles of NH 3. Page 4
The Mole Concept Teacher Notes Time Required: One 50 minute class period. Prerequisites: Knowledge of average atomic mass, the use of the Periodic Table to find an element s average atomic mass, and the calculation of a formula mass Learning Objectives: 1. Understand the relationship between the unit mole and Avogadro s number. 2. Define the gram formula mass of a substance. 3. Understand the relationship between the mole and the gram formula mass. Performance Criteria: 1. Perform conversions between mass, moles, and numbers of particles. Considerations: This activity is a wonderful way to introduce the concept of moles to an introductory chemistry class. No previous knowledge of moles is required. Be careful not to answer the students questions for them. If a group asks a question concerning content, rather than supplying the answer, ask a question in return that will steer them toward discovering the answer for themselves. Another way to intervene with the group may be to have the Spokesperson or Presenter for the group to restate their reasoning so that they have the opportunity to pinpoint misconceptions or faulty reasoning on their own. Closure: It is always good to have the students reflect on what they have learned before leaving class. One way to do this is to stop the activity about 5 minutes before the end of class and have students take out a sheet of paper. Have them record one thing that they learned as a result of the lesson that day. Then have them to ask a question concerning the lesson. The question could be something that is still unclear to them or it could be something they are curious about or a subject that they think may be related to the lesson. Page 5
Answers to Questions: Model 1 Key Questions: 1. 12 trunks 2. 48 legs 3. 500 trunks 4. 8 legs 5. 12 carbon atoms 6. 48 hydrogen atoms 7. 2 carbon atoms 8. 2000 hydrogen atoms 9. 6.022 x 10 23 elephants 10. 3.011 x 10 23 trunks 11. 2.409 x 10 24 legs 12. 6.022 x 10 23 molecules 13. 6.022 x 10 23 carbon atoms 14. 1.204 x 10 24 hydrogen atoms Model 2 15. periodic table 16. 12.01 amu 17. 12.01 g 18. 16.05 amu 19. 16.05 g 20. One atomic mass unit (amu) equals the mass in grams of one mole of an element. 21. 6.022 x 10 23 molecules 22. 6.022 x 10 23 atoms 23. 98.99 g Substance Number of Particles Mass in grams Al 6.022 x 10 23 atoms 26.98 g H 2 O 6.022 x 10 23 molecules 18.02 g O 2 1.505 x 10 23 molecules 8.00 g Ne 3.011 x 10 23 atoms 10.09g Page 6
Skill Development Exercises: 1. a. False, NH 3 1 mole = 17.03 g, H 2 O 1 mole = 18.02 g b. True, there are 6.57 x 10 23 atoms of carbon in 48 g of CO 2 and only 6.016 x 10 23 atoms of carbon in 12 g of diamond c. False, 1 mole of NH 3 = 6.022 x 10 23 atoms of nitrogen, 1 mole of N 2 = 1.2044 x 10 24 atoms of nitrogen d. False, 100 g of Cu = 7.57 x 1023 atoms of copper, 100 g of CuO = 9.48 x 10 23 atoms of copper e. True, Since 1 molecule of NiCl 2 contains 1 atom of nickel then 1 mole of NiCl 2 would contain 6.022 x 10 23 atoms of nickel. One mole of nickel also contains 6.022 10 23 atoms of nickel. f. False, In 2 moles of NH 3 there are 3.61 x 10 24 hydrogen atoms but in 2 moles of CH 4 there are 4.82 x 10 24 hydrogen atoms. 2. In one molecule of NH 3 there are three atoms of hydrogen and in mole of NH 3 there are 6.022 x10 23 molecules of NH 3. To determine the number of hydrogen atoms in any number of moles, multiply the number of moles by 6.022 x 10 23 to convert to molecules. Then convert molecules to atoms by multiplying by three. Page 7
1 What is an Atom? Background Information: An atom is defined as the smallest particle of an element that retains the chemical properties of that element. All atoms are composed of the same basic particles, but all atoms are not the same. The Periodic Table is a useful tool for determining the number of particles in different atoms. Here is the Periodic Table entry for the element carbon: 6 C 12.011 Atomic number Atomic symbol Atomic mass Model: Schematic Diagrams for Various Atoms and Ions
2 Group Members: Group Activity: 1. How many protons are found in 12 C? 13 C? 13 C -? 2. How many neutrons are found in 12 C? 13 C? 13 C -? 3. How many electrons are found in 12 C? 13 C? 13 C -? 4. Based on the background information, a) what do all carbon atoms (and ions) have in common? b) what do all hydrogen atoms (and ions) have in common? 5. Look at the Periodic Table provided. The atomic number (normally referred to as Z) is the number given above each element s symbol. What is the significance of this atomic number? 6. What do all nickel (Ni) atoms have in common? (refer to the periodic table if needed) 7. How is the mass number (normally referred to as A) of an atom (or ion) determined? (NOTE: The mass number is the superscript number to the left of the atomic symbol.) 8. What structural feature is different in isotopes of a particular element?
3 Model: Schematic Diagram of Various Atomic Changes Loss of electrons Neutral Atom Gain of electrons Positive Ion Change in atomic mass (or neutrons) Isotope Negative Ion 9. a) What distinguishes a neutral atom from an ion? b) How is the charge on an ion determined? 10. Where is most of the mass of a neutral atom, ion and isotope found? Within the nucleus or outside of the nucleus? Explain your reasoning. 11. What is the mass (in grams) of one 1 H atom? Of one 12 C atom? Review Questions/Problems (Individual Effort): 1. Read pages 75-82 in your textbook. 2. Study Sample Problem 3.1 on page 77. Review Practice Problems on page 78.
4 3. Complete the following table: ISOTOPE ATOMIC NUMBER 31 P 15 ATOMIC MASS NUMBER OF ELECTRONS 18 O 8 19 39 18 58 Ni 2+ 58 27 59 25 14 N 58 Zn 2+ 3 7 3 3 6 3 19 F - 4. How many electrons, protons and neutrons are found in each of the following? a) 24 Mg b) 23 Na + c) 35 Cl d) 35 Cl - e) 56 Fe 3+ f) 15 N g) 16 O 2+
5 h) 27 Al 3+ 5. What is the mass (in grams) of 4.35 x 10 6 atoms of 12 C? 6. a) Define atomic mass (A). b) Define atomic number (Z). Sources: 1. Davis, Raymond E., et al, Modern Chemistry, Holt, Rinehart and Winston, 2002, pages 75-82. 2. Farrell, John J. and Moog, Richard S., Chemistry: A Guided Inquiry, John Wiley & Sons, 1999, pages 1-4.
6 What is an Atom? Teacher Notes Time Required: One 50 minute class period. Prerequisites: Knowledge of basic dimensional analysis, calculation of unit conversions Learning Objectives: 1. Identify the basic subatomic particles. 2. Understand the relationship between subatomic particles, atomic number, and atomic mass. 3. Identify the nucleus of an atom and the subatomic particles contained in it. Performance Criteria: 1. Recognize the relationship between the atomic number of an element and its placement on the Periodic Table. Considerations: This activity is a wonderful way to introduce the concept of basic atomic structure. No previous knowledge of atoms is required. Be careful not to answer the students questions for them. If a group asks a question concerning content, rather than supplying the answer, ask a question in return that will steer them toward discovering the answer for themselves. Another way to intervene with the group may be to have the Spokesperson or Presenter for the group to restate their reasoning so that they have the opportunity to pinpoint misconceptions or faulty reasoning on their own. Closure: It is always good to have the students reflect on what they have learned before leaving class. One way to do this is to stop the activity about 5 minutes before the end of class and have students take out a sheet of paper. Have them record one thing that they learned as a result of the lesson that day. Then have them to ask a question concerning the lesson. The question could be something that is still unclear to them or it could be something they are curious about or a subject that they think may be related to the lesson.
7 Answers to Questions in Group Activity: 1. 6, 6, 6 2. 6, 7, 7 3. 6, 6, 7 4. a) all have 6 protons b) all have 1 proton 5. It is the same as the number of protons. 6. They must all have 28 protons. 7. It is the total of the protons and neutrons. 8. The number of neutrons is different (which makes the mass different). 9. a) an ion is charged b) adding the protons (pluses) and electrons (minuses) together equals the charge 10. inside the nucleus because that is where the protons and neutrons are located. The atomic masses don t agree if you add in the electrons that are outside of the nucleus. 11. 1.0078 amu x 1.6606 x 10-24 g = 1.6736 x 10-24 g for the 1 H atom 1 amu 12 amu x 1.6606 x 10-24 g = 1.9927 x 10-23 g for the 12 C atom 1 amu