Chapter 3 The Mole and Stoichiometry

Transcription

1 Chapter 3 The Mole and Stoichiometry Chemistry, 7 th Edition International Student Version Brady/Jespersen/Hyslop Brady/Jespersen/Hyslop Chemistry7E, Copyright 015 John Wiley & Sons, Inc. All Rights Reserved

2 Chapter in Context Use the mole and Avogadro s number as conversions between the molecular and laboratory scales of matter Perform calculations involving moles Calculate percent composition of a compound Determine empirical and molecular formulas Perform calculations involving mole of reactants and products in a reaction Determine limiting reactants and amounts of unreacted substance remaining Calculate percent yields for reactions

3 Using Mass to Count 3

4 Defining The Mole Mole: A number equal to the number of atoms in exactly 1 grams of 1 C atoms How many in 1 mole of 1 C? Based on experimental evidence 1 mole of 1 C = atoms 1 C Avogadro s number = N A Number of atoms, molecules or particles in one mole 1 mole of X = units of X 1 mole Xe = Xe atoms 1 mole NO = NO molecules 4

5 Molecular to Laboratory Scale So far, we have looked at chemical formulas and reactions at a molecular scale It is known from experiments that: Electrons, neutrons and protons have set masses Atoms must also have characteristic masses Atoms and molecules are extremely small Need a way to scale up chemical formulas and reactions to carry out experiments in laboratory Mole is our conversion factor 5

6 Atoms Moles of Compounds Atomic Mass Mass of atom (from periodic table) 1 mole of atoms = gram atomic mass = atoms Molecules Molecular Mass Sum of atomic masses of all atoms in compound s formula 1 mole of molecule X = gram molecular mass of X = molecules 6

7 Moles of Compounds Ionic compounds Formula Mass Sum of atomic masses of all atoms in ionic compound s formula 1 mole ionic compound X = gram formula mass of X General Molar mass (MM) = formula units Mass of 1 mole of substance (element, molecule, or ionic compound) under consideration 1 mol of X = gram molar mass of X = formula units 7

8 SI Unit for Amount = Mole 1 mole of substance X = gram molar mass of X 1 mole S = 3.06 g S 1 mole NO = g NO Molar mass is our conversion factor between g and moles 1 mole of X = units of X N A is our conversion factor between moles and molecules 1 mole H O = molecules H O 1 mole NaCl = formula units NaCl 8

9 Your Turn! What is the molar mass of C H 5 COOH? A. 148 g/mol B. 6.5 g/mol C g/mol D. 11 g/mol E g/mol 3 mol C 1.0g C mol C 6 mol H 1.00g H mol H 16.0g O mol O mol O 74.0 g CH5COOH mol C H COOH 5 9

10 Your Turn! What is the formula mass of magnesium phosphate? A. 6.9 g/mol B g/mol C g/mol D g/mol E g/mol 3 mol Mg Formula is Mg 3 (PO 4 ) 4.3g Mg mol Mg mol P 31.00g P mol P 16.0g O 8 mol O mol O 6.9 g Mg3(PO4 ) mol Mg (PO )

11 Learning Check: Using Molar Mass Example: How many moles of iron (Fe) are in g Fe? What do we want to determine? Start g Fe =? mol Fe End What do we know? 1 mol Fe = g Fe Set up ratio so that what you want is on top and what you start with is on the bottom 1 mol Fe g Fe = mol Fe g Fe 11

12 Learning Check: Using Molar Mass Example: If we need mole Ca 3 (PO 4 ) for an experiment, how many grams do we need to weigh out? What do we want to determine? mole Ca 3 (PO 4 ) =? g Ca 3 (PO 4 ) Start End Calculate MM of Ca 3 (PO 4 ) 3 mass Ca = g = 10.4 g mass P = g = g 8 mass O = g = g 1 mole Ca 3 (PO 4 ) = g Ca 3 (PO 4 ) 1

13 Learning Check: Using Molar Mass Set up ratio so that what you want is on the top and what you start with is on the bottom mol Ca 3 (PO 4 ) g Ca3(PO 1 mol Ca (PO ) ) = 5.11 g Ca 3 (PO 4 ) 13

14 Your Turn! How many moles of CO are there in 10.0 g? A mol B mol C mol D mol E. 0.7 mol Molar mass of CO g = 1.01 g C g = 3.00 g O 1 mol CO = g CO 10.0 g CO 1 mol CO g CO = 0.7 mol CO 14

15 Your Turn! How many grams of platinum (Pt) are in mole Pt? A. 195 g B g C g D g E. 9.7 g Molar mass of Pt = g/mol g Pt mol Pt 1 mol Pt = 9.7 g Pt 15

16 Using Moles in Calculations Start with either Grams (Macroscopic) Elementary units (Microscopic) Use molar mass to convert from grams to mole Use Avogadro s number to convert from moles to elementary units 16

17 Macroscopic to Microscopic How many silver atoms are in a 85.0 g silver bracelet? What do we want to determine? 85.0 g silver =? atoms silver What do we know? g Ag = 1 mol Ag 1 mol Ag = Ag atoms g Ag mol Ag atoms Ag 85.0 g Ag 1 mol Ag g Ag mol atoms Ag Ag = Ag atoms 17

18 Using Avogadro s Number What is the mass, in grams, of one molecule of octane, C 8 H 18? Molecules octane mol octane g octane 1. Calculate molar mass of octane Mass C = g = g Mass H = g = g 1 mol octane = 114. g octane. Convert 1 molecule of octane to grams 114. g octane 1 mol octane 1 mol octane molecules octane = g octane 18

19 Learning Check: Mole Conversions Calculate the number of formula units of Na CO 3 in 1.9 moles of Na CO 3. æ formula units Na 1.9 mol Na CO CO ö 3 3 ç è 1 mol Na CO 3 ø = particles Na CO 3 How many moles of Na CO 3 are there in formula units of Na CO 3? æ 1 mol Na formula units Na CO CO ö 3 3 ç è formula units Na CO 3 ø = mol Na CO 3 19

20 Your Turn! How many atoms are in g of U? Molar mass U = g/mole. A atoms B atoms C atoms D atoms E atoms æ ( g U) 1 mol U ö æ ç atoms Uö è38.03 g U ç ø è 1 mol U ø = atoms U 0

21 Your Turn! How much, in grams, do atoms of zinc weigh? A g B. 961 g C g D g E g atoms = 961 g Zn 1mol atoms 65.41g Zn 1mol 1

22 Your Turn! Calculate the mass in grams of FeCl 3 in formula units. (molar mass = g/mol) A. 16. g B g C g D. 41. g E units FeCl 3 1 mol FeCl 3 units FeCl g FeCl 1 mol FeCl 3 3 = 41. g FeCl 3

23 Mole-to-Mole Conversion Factors Can use chemical formula to relate amount of each atom to amount of compound In H O there are three relationships: mol H 1 mol H O 1 mol O 1 mol H O mol H 1 mol O Can also use these on atomic scale: atom H 1 molecule H O 1 atom O 1 molecule H O atom H 1 molecule O 3

24 Stoichiometric Equivalencies Within chemical compounds, moles of atoms always combine in the same ratio as the individual atoms themselves Ratios of atoms in chemical formulas must be whole numbers These ratios allow us to convert between moles of each quantity e.g., N O 5 mol N 1 mol N O 5 5 mol O 1 mol N O 5 mol N 5 mol O 4

25 Stoichiometric Equivalencies Equivalency Mole Ratio Mole Ratio mol N 1 mol N O 5 mol N 1 mol N O 5 1 mol N O 5 mol N 5 mol O 1 mol N O 5 5 mol O 1 mol N O 5 1 mol N O 5 5 mol O mol N 5 mol O 5 mol O mol N mol N 5 mol O 5

26 Calculating the Amount of a Compound by Analyzing One Element Calcium phosphate is widely found in natural minerals, bones, and some kidney stones. A sample is found to contain moles of phosphorus. How many moles of Ca 3 (PO 4 ) are in that sample? What do we want to find? mol P =? mol Ca 3 (PO 4 ) What do we know? Solution mol P 1 mol Ca 3 (PO 4 ) 1 mol Ca3(PO4 ) mol P mol P = 0.43 mol Ca 3 (PO 4 ) 6

27 Your Turn! Calculate the number of moles of calcium in.53 moles of Ca 3 (PO 4 ) A..53 mol Ca B mol Ca C mol Ca D mol Ca E mol Ca.53 moles of Ca 3 (PO 4 ) =? mol Ca 3 mol Ca 1 mol Ca 3 (PO 4 ) 3 mol Ca.53 mol Ca3(PO 4 ) 1 mol Ca3(PO 4 ) = 7.59 mol Ca 7

28 Your Turn! A sample of sodium carbonate, Na CO 3, is found to contain 10.8 moles of sodium. How many moles of oxygen atoms (O) are present in the sample? A mol O B. 7.0 mol O C mol O D. 3.4 mol O E. 16. mol O 10.8 moles of Na =? mol O mol Na 3 mol O 3 mol O 10.8 mol Na = 16. mol O mol Na 8

29 Mass-to-Mass Calculations Common laboratory calculation Need to know what mass of reagent B is necessary to completely react given mass of reagent A to form a compound Stoichiometry comes from chemical formula of compounds Use the subscripts Summary of steps mass A moles A moles B mass B 9

30 Mass-to-Mass Calculations Chlorophyll, the green pigment in leaves, has the formula C 55 H 7 MgN 4 O 5. If g of Mg is available to a plant for chlorophyll synthesis, how many grams of carbon will be required to completely use up the magnesium? Analysis g Mg? g C g Mg mol Mg mol C g C Assembling the Tools g Mg = 1 mol Mg 1 mol Mg 55 mol C 1 mol C = g C 30

31 Ex. Mass-to-Mass Conversion g Mg 1 mol Mg 1 mol C g C g Mg mol Mg mol C g C 1 mol Mg 55 mol C g Mg 1 mol Mg g Mg 55 mol C 1 mol Mg g C 1 mol C = g C 31

32 Your Turn! How many g of iron are required to use up all of 5.6 g of oxygen atoms (O) to form Fe O 3? A g B. 9.8 g C g D. 134 g E. 5.4 g 5.6 g O = 59.6 g Fe mass O mol O mol Fe mass Fe 5.6 g O? g Fe 3 mol O mol Fe 1 mol O 16.0 g O mol Fe 3 mol O g Fe 1 mol Fe 3

33 Your Turn! Silver is often found in nature as the ore, argentite (Ag S). How many grams of pure silver can be obtained from a 836 g rock of argentite? A g B. 78 g C. 364 g D. 775 g E. 418 g 836 g Ag mass Ag S mol Ag S mol Ag mass Ag 836 g Ag S? g Ag 1 mol Ag S mol Ag 1mol Ag S mol Ag g Ag S 47.8g Ag S 1mol Ag S 1mol Ag = 78 g Ag 33

34 Percentage Composition Way to specify relative masses of each element in a compound List of percentage by mass of each element Percentage by Mass mass of element % by mass of element 100% mass of sample Example: Na CO 3 is 43.38% Na 11.33% C 45.9% O What is sum of the percentages? % 34

35 Example - Percent Composition Determine percentage composition based on chemical analysis of substance Example: A sample of a liquid with a mass of g was decomposed into its elements and gave 5.17 g of carbon, g of hydrogen, and.478 g of oxygen. What is the percentage composition of this compound? Analysis: Calculate percentage by mass of each element in sample Tools: Equation for percentage by mass Total mass = g Mass of each element 35

36 Ex. % Composition of Compound For C: For H: For O: æ g C ö ç è g total 100% 5.17 g C 100% = 60.6% C ø g æ g H ö ç è g total 100% g H 100%= 11.11% H ø g æ g O ö ç è g total 100%.478 g O 100% = 8.6% O ø g Sum of percentages: 99.99% Percentage composition tells us mass of each element in g of substance In g of our liquid 60.6 g C, g H, and 8.6 g O 36

37 Your Turn! A sample was analyzed and found to contain g nitrogen and g oxygen. What is the percentage composition of this compound? 1. Calculate total mass of sample Total sample mass = g g = g. Calculate % Composition of N g N g N 100% 100% g total g 3. Calculate % Composition of O g O g total 100% g O g 100% = 5.94% N = 74.06% O 37

38 Your Turn! In a previous You Turn! slide, you found that an 836 g rock of argentite (Ag S) contained 78 g of silver. What are the percentages of silver and sulfur, by mass, in argentite? A. 50.0% Ag, 50.0% S B. 66.6% Ag, 33.4% S C. 1.9% Ag, 87.1% S D. 115% Ag, 108% S E. 87.1% Ag, 1.9% S Calculate % comp. of Ag: g Ag g total 78 g Ag 836 g 100% 100% = 87.1% Ag Because the rest must be sulfur, % S = = 1.9% S 38

39 Percent Compositions and Chemical Identity Theoretical or Calculated Percentage Composition Calculated from molecular or ionic formula. Lets you distinguish between multiple compounds formed from the same two elements If experimental percent composition is known Calculate theoretical percentage composition from proposed chemical formula Compare with experimental composition e.g., N and O form multiple compounds N O, NO, NO, N O 3, N O 4, and N O 5 39

40 Example: Using Percent Composition Are the mass percentages 30.54% N and 69.46% O consistent with the formula N O 4? Procedure: 1. Assume one mole of compound. Subscripts tell how many moles of each element are present mol N and 4 mol O 3. Use molar masses of elements to determine mass of each element in 1 mole Molar Mass of N O 4 = 9.14 g N O 4 / 1 mol 4. Calculate % by mass of each element 40

41 Ex. Using Percent Composition (cont) mol N 4 mol O g N 1 mol N g O 1 mol O = 8.14 g N = g O 8.14 g N %N 100% = 30.54% N in N 9.14 g N O O g O %O 100% = 69.46% O in N 9.14 g N O O 4 4 The experimental values match the theoretical percentages for the formula N O 4. 41

42 Your Turn! If a sample containing only phosphorous and oxygen has percent composition 56.34% P and 43.66% O, is this P 4 O 10? A. Yes B. No 4 mol P 1 mol P 4 O mol O 1 mol P 4 O 10 4 mol P = g/mol P = 13.9 g P 10 mol O = g/mol O = g O 1 mol P 4 O 10 = 83.9 g P 4 O 10 % P % O 13.9 g P 83.9 g P O g O 83.9 g P O % 100% = 43.64% P = 56.36% O 4

43 Determining Empirical and Molecular Formulas When making or isolating new compounds one must characterize them to determine structure and Empirical Formula Simplest ratio of atoms of each element in compound Obtained from experimental analysis of compound Molecular Formula Exact composition of one molecule Exact whole number ratio of atoms of each element in molecule glucose Empirical formula CH O Molecular formula C 6 H 1 O 6 43

44 Three Ways to Calculate Empirical Formulas 1. From Masses of Elements e.g.,.448 g sample of which g is Fe and g is O.. From Percentage Composition e.g., 43.64% P and 56.36% O 3. From Combustion Data Given masses of combustion products e.g., The combustion of a 5.17 g sample of a compound of C, H, and O in pure oxygen gave g CO and 4.51 g of H O 44

45 Strategy for Determining Empirical Formulas 1. Determine mass in g of each element. Convert mass in g to moles 3. Divide all quantities by smallest number of moles to get smallest ratio of moles 4. Convert any non-integers into integer numbers. If number ends in decimal equivalent of fraction, multiply all quantities by the denominator of the fraction Otherwise, round numbers to nearest integers 45

46 1. Empirical Formula from Mass Data When a g sample of a compound was analyzed, it was found to contain g of C, g of H, and g of N. Calculate the empirical formula of this compound. Step 1: Calculate moles of each substance g C g H g N 1 mol C g C 1 mol H g H 1 mol N g N mol C mol H mol N 46

47 1. Empirical Formula from Mass Data Step : Select the smallest number of moles Smallest is mole C = H = mol C mol C mol H mol C Mole ratio Integer ratio = = mol N N = = mol C Step 3: Divide all number of moles by the smallest one Empirical formula = CH 5 N 47

48 Empirical Formula from Mass Data One of the compounds of iron and oxygen, black iron oxide, occurs naturally in the mineral magnetite. When a.448 g sample was analyzed it was found to have g of Fe and g of O. Calculate the empirical formula of this compound. Assembling the tools: 1 mol Fe = g Fe 1 mol O = g O 1. Find the moles of each element, then find the ratio of the moles of the elements.. Calculate moles of each substance 1 mol Fe g Fe g Fe 1 mol O g O g O mol Fe mol O 48

49 1. Empirical Formula from Mass Data. Divide both by smallest number mol to get smallest whole number ratio. Or mol Fe mol Fe mol O mol Fe =1.000 Fe =1.33 O 3 = Fe 3 = 3.99 O Fe 1 O Fe O (1.003) O(1.333) Fe3O3.99 Fe Empirical Formula = Fe 3 O 4 49

50 . Empirical Formula from Percentage Composition New compounds are characterized by elemental analysis, from which the percentage composition can be obtained Use percentage composition data to calculate empirical formula Must convert percentage composition to grams Assume g sample Convenient Sum of percentage composition = 100% Sum of masses of each element = 100 g 50

51 . Empirical Formula from Percentage Composition Calculate the empirical formula of a compound whose percentage composition data is 43.64% P and 56.36% O. If the molar mass is determined to be 83.9 g/mol, what is the molecular formula? Step 1: Assume 100 g of compound g P g O 1 mol P = g 1 mol O = g g P 1 mol P g P = mol P g O 1 mol O g O = 3.53 mol P 51

52 . Empirical Formula from Percentage Composition Step : Divide by smallest number of moles mol P mol P = 3.53 mol mol O P.500 = 5 Step 3: Multiple to get integers =.500 = 5 Empirical formula = P O 5 5

53 3. Empirical Formulas from Indirect Analysis: In practice, compounds are not broken down into elements, but are changed into other compounds whose formula is known Combustion Analysis Compounds containing carbon, hydrogen, and oxygen, can be burned completely in pure oxygen gas Only carbon dioxide and water are produced e.g., Combustion of methanol (CH 3 OH) CH 3 OH + 3O CO + 4H O 53

54 Combustion Analysis Classic Modern CHN analysis 54

55 3. Empirical Formulas from Indirect Analysis: Carbon dioxide and water are separated and weighed separately All C ends up as CO All H ends up as H O Mass of C can be derived from amount of CO mass CO mol CO mol C mass C Mass of H can be derived from amount of H O mass H O mol H O mol H mass H Mass of oxygen is obtained by difference mass O = mass sample (mass C + mass H) 55

56 Ex. Indirect or Combustion Analysis The combustion of a 5.17 g sample of a compound of C, H, and O in pure oxygen gave g CO and 4.51 g of H O. Calculate the empirical formula of the compound. C H H O CO MM (g/mol) Calculate mass of C from mass of CO. mass CO mole CO mole C mass C g CO =.01 g C 1 mol CO g CO 1 mol C 1 mol CO g C 1 mol C 56

57 Ex. Indirect or Combustion Analysis The combustion of a 5.17 g sample of a compound of C, H, and O gave g CO and 4.51 g of H O. Calculate the empirical formula of the compound.. Calculate mass of H from mass of H O. mass H O mol H O mol H mass H 4.51 g H = g H 1 mol H O O g HO mol H 1 mol HO 3. Calculate mass of O from difference g sample.01 g C g H g H 1 mol H =.691 g O 57

58 Ex. Indirect or Combustion Analysis mol C C H O At. mass g Calculate mol of each element mol mol H O g C MM C g H MMH g O MM O.01 g g/mol g g/mol.691 g g/mol = mol C = mol H = mol O 58

59 Ex. Indirect or Combustion Analysis Preliminary empirical formula C H O Calculate mol ratio of each element C H0.5009O = C 1.00 H.97 O Because all values are close to integers, round to Empirical Formula = CH 3 O 59

60 Your Turn! The combustion of a g sample of a compound of C, H, and S in pure oxygen gave g CO and g of H O. Calculate the empirical formula of the compound. A. C 4 H 1 S B. CH 3 S C. C H 6 S D. C H 6 S 3 E. CH 3 S 60

61 19.35 g CO Your Turn! Solution 1. Calculate mass of C from mass of CO. mass CO mole CO mole C mass C = 5.81 g C 1 mol CO 44.01g CO 1mol C 1mol CO g C 1 mol C. Calculate mass of H from mass of H O. mass H O mol H O mol H mass H g H = g H 1mol H O O g HO mol H 1mol HO g H 1mol H 61

62 Your Turn! Solution Continued 3. Calculate mass of S from difference g sample 5.81 g C g H = g S mol C g C MM C 5.81g g/mol = mol C mol H g H MM H g g/mol = mol H mol S g S MM S g g/mol = mol S 6

63 Your Turn! Solution Continued Preliminary empirical formula C H O Calculate mol ratio of each element C H O = C.03 H 6.00 O 1.00 Because all values are close to integers, round to Empirical Formula = C H 6 S 63

64 Determining Molecular Formulas Empirical formula Accepted formula unit for ionic compounds Molecular formula Preferred for molecular compounds In some cases molecular and empirical formulas are the same When they are different, the subscripts of molecular formula are integer multiples of those in empirical formula If empirical formula is A x B y Molecular formula will be A n x B n y 64

65 Determining Molecular Formula Need molecular mass and empirical formula Calculate ratio of molecular mass to mass predicted by empirical formula and round to nearest integer molecular mass n = empirical formula mass Example: Glucose Molecular mass is g/mol Empirical formula = CH O Empirical formula mass = g/mol n = g g = 6 Molecular formula = C 6 H 1 O 6 65

66 Learning Check The empirical formula of a compound containing phosphorous and oxygen was found to be P O 5. If the molar mass is determined to be 83.9 g/mol, what is the molecular formula? n = Step 1: Calculate empirical mass empirical mass P O g/mol P O 5 molecular mass empirical mass mass P 5 mass O g/mol g/mol g/mol Step : Calculate ratio of molecular to empirical mass n = 83.9 g / mol g/mol = Molecular formula = P 4O 10 66

67 Your Turn! The empirical formula of hydrazine is NH, and its molecular mass is 3.0. What is its molecular formula? Atomic Mass: N = ; H = 1.008; O = A. NH B. N H 4 Molar mass of NH = C. N 3 H 6 D. N 4 H 8 n = (3.0/16.0) = E. N 1.5 H 3 ( ) g + ( 1.008) g = g 67

68 Balanced Chemical Equations Useful tool for problem solving Prediction of reactants and products All atoms present in reactants must also be present among products Coefficients are multipliers that are used to balance equations Two step process 1.Write unbalanced equation Given products and reactants Organize with plus signs and arrow. Adjust coefficients to get equal numbers of each kind of atom on both sides of arrow 68

69 Guidelines for Balancing Equations 1. Start balancing with the most complicated formula first Elements, particularly H and O, should be left until the end. Balance atoms that appear in only two formulas: one as a reactant and the other as a product Leave elements that appear in three or more formulas until later 3. Balance as a group those polyatomic ions that appear unchanged on both sides of the arrow 69

70 Balancing Equations Use the inspection method Step 1. Write unbalanced equation Zn(s) + HCl(aq) ZnCl (aq) + H (g) unbalanced Step. Adjust coefficients to balance numbers of each kind of atom on both sides of arrow. Since ZnCl has two Cl on the product side, HCl on reactant side is needed to balance the equation. Zn(s) + HCl(aq) ZnCl (aq) + H (g) One Zn each side Two H each side Two Cl each side Balanced! 70

71 Learning Check: Balancing Equations AgNO 3 (aq) + Na 3 PO 4 (aq) Ag 3 PO 4 (s) + NaNO 3 (aq) Count atoms Reactants Products 1 Ag 3 Ag 3 Na 1 Na Adjust coefficients of both AgNO 3 + NaNO 3 to three. 3AgNO 3 (aq) + Na 3 PO 4 (aq) Ag 3 PO 4 (s) + 3NaNO 3 (aq) Now check polyatomic ions 3 NO 3 3 NO 3 1 PO PO 3 4 Equation is balanced 71

72 Balance by Inspection C 3 H 8 (g) + O (g) CO (g) + H O(l ) Assume 1 in front of C 3 H 8 3 C 1 C 3 8 H H 4 1C 3 H 8 (g) + O (g) 3CO (g) + 4H O(l ) O 5 =10 O = (3 ) + 4 = 10 8 H H = 4 = 8 1C 3 H 8 (g) + 5O (g) 3CO (g) + 4H O(l ) 7

73 Your Turn! Balance each of the following equations. What are the coefficients in front of each compound? 1 Ba(OH) (aq) + 1 Na SO 4 (aq) 1 BaSO 4 (s) + NaOH(aq) KClO 3 (s) KCl(s) + 3 O (g) H 3 PO 4 (aq) + 3 Ba(OH) (aq) Ba 1 3 (PO 4 ) (s) + H 6 O(l ) 73

74 Using Balanced Equations: Reaction Stoichiometry Balanced equation Critical link between substances involved in chemical reactions Gives relationship between amounts of reactants used and amounts of products formed Numeric coefficient tells us The mole ratios for reactions How many individual particles are needed in reaction on microscopic level How many moles are necessary on macroscopic level 74

75 Stoichiometric Ratios Consider the reaction Could also be read as: N + 3H NH 3 When 1 mole of nitrogen reacts with 3 moles of hydrogen, moles of ammonia are formed. Molar relationships 1 mole N mole NH 3 3 mole H mole NH 3 1 mole N 3 mole H 75

76 Stoichiometric Ratios Consider the reaction Could be read as: N + 3H NH 3 When 1 molecule of nitrogen reacts with 3 molecules of hydrogen, molecules of ammonia are formed. Molecular relationships 1 molecule N molecule NH 3 3 molecule H molecule NH 3 1 molecule N 3 molecule H 76

77 Stoichiometry Mass balance of all formulas involved in chemical reactions Stoichiometric Calculations Conversions from one set of units to another using dimensional analysis Need to know: 1. Equalities to make conversion factors. Steps to go from starting units to desired units 77

78 Using Stoichiometric Ratios Example: For the reaction N + 3 H NH 3, how many moles of N are used when.3 moles of NH 3 are produced? Assembling the tools moles NH 3 = 1 mole N.3 mole NH 3 =? moles N 1 mol N.3 mol NH 3 = 1. mol N mol NH 3 78

79 Your Turn! If mole of CO is produced by the combustion of propane, C 3 H 8, how many moles of oxygen are consumed? The balanced equation is C 3 H 8 (g) + 5O (g) 3CO (g) + 4H O(g) A mole B..88 mole C mole D mole E mole mol CO Assembling the tools mole CO =? moles O 3 moles CO = 5 mole O 5 mol O 3 mol CO = mol O 79

80 Your Turn! How many moles of hydrochloric acid are required to completely react with 1.43 moles of iron(ii) chloride according to the reaction below? 14HCl + Na Cr O 7 + 6FeCl A. 0.0 mole B mole C mole D mole E mole CrCl 3 + 7H O + 6FeCl 3 + NaCl Assembling the tools 1.43 mole FeCl =? moles HCl 6 moles FeCl = 14 mole HCl 1.43 mol FeCl = 3.34 mol HCl 14 mol HCl 6 mol FeCl 80

81 Mass-to-Mass Conversions Most common stoichiometric conversions that chemists use involve converting mass of one substance to mass of another. Use molar mass A to convert grams A to moles A Use chemical equations to relate moles A to moles B Use molar mass B to convert to moles B to grams B 81

82 Stoichiometry Calculations Using a Balanced Chemical Equation Example: What mass of O will react with 96.1 g of propane (C 3 H 8 ) gas, to form gaseous carbon dioxide and water? Strategy 1. Write the balanced equation C 3 H 8 (g) + 5O (g) 3CO (g) + 4H O(g). Assemble the tools 96.1 g C 3 H 8 moles C 3 H 8 moles O g O 1 mol C 3 H 8 = 44.1 g C 3 H 8 1 mol O = 3.00 g O 1 mol C 3 H 8 = 5 mol O 8

83 Stoichiometry Calculations Using a Balanced Chemical Equation Example: What mass of O will react with 96.1 g of propane in a complete combustion? C 3 H 8 (g) + 5O (g) 3CO (g) + 4H O(g) 3. Assemble conversions so units cancel correctly 1 mol C3H8 5 mol O 96.1 g C3H g C H 1 mol C H g O 1 mol O = 349 g of O are needed 83

84 Your Turn! How many grams of Al O 3 are produced when 41.5 g Al react? Al(s) + Fe O 3 (s) Al O 3 (s) + Fe(l ) A g B. 157 g C. 314 g D..0 g E g 41.5 g Al = 78.4 g Al O 3 1 mol Al 6.98 g Al 1 mol AlO mol Al g Al O 1 mol AlO3 3 84

85 Your Turn! How many grams of sodium dichromate are required to produce 4.7 g iron(iii) chloride from the following reaction? 14HCl + Na Cr O 7 + 6FeCl CrCl 3 + 7H O + 6FeCl 3 + NaCl A g Na Cr O 7 B g Na Cr O 7 C g Na Cr O 7 D g Na Cr O 7 E g Na Cr O g FeCl 3 1mol Na CrO 6 mol FeCl3 1mol FeCl g FeCl g Na CrO 1mol Na Cr O = 6.64 g Na Cr O

86 Molecular Level of Reactions Consider industrial synthesis of ethanol C H 4 + H O C H 5 OH 3 molecules ethylene + 3 molecules water react to form 3 molecules ethanol 86

87 Molecular Level of Reactions What happens if these proportions are not met? 3 molecules ethylene + 5 molecules of water All ethylene will be consumed and some water will be left over 87

88 Reactant that is completely used up in the reaction Present in lower number of moles It determines the amount of product produced For this reaction the limiting reactant is ethylene Excess reactant Limiting Reactant Reactant that has some amount left over at end Present in higher number of moles For this reaction it is water 88

89 Limiting Reactant Calculations 1. Write the balanced equation. Identify the limiting reagent Calculate amount of reactant B needed to react with reactant B mass reactant A have mol reactant A mol reactant B Mass reactant B need Compare amount of B you need with amount of B you actually have. If need more B than you have, then B is limiting If need less B than you have, then A is limiting 89

90 Limiting Reactant Calculations 3. Calculate mass of desired product, using amount of limiting reactant and mole ratios. mass limiting reactant mol limiting reactant mol product mass product 90

91 Ex. Limiting Reactant Calculation How many grams of NO can form when 30.0 g NH 3 and 40.0 g O react according to: Solution: Step 1 4NH 3 + 5O 4NO + 6H O mass NH 3 mole NH 3 mole O mass O Assembling the tools 1 mol NH 3 = g 1 mol O = 3.00 g Only have 40.0 g O, O limiting reactant 4 mol NH 3 5 mol O 1 mol NH3 5 mol O 30.0 g NH g NH 4 mol NH = 70.5 g O needed mol g O O 91

92 Ex. Limiting Reactant Calculation How many grams of NO can form when 30.0 g NH 3 and 40.0 g O react according to: Solution: Step 4 NH O 4 NO + 6 H O mass O mole O mole NO mass NO Assembling the tools 1 mol O = 3.00 g 1 mol NO = g 5 mol O 4 mol NO 1 mol O 40.0 g O 3.00 g O = 30.0 g NO formed 4 mol NO 5 mol O Can only form 30.0 g NO g NO 1 mol NO 9

93 Your Turn! If 18.1 g NH 3 is reacted with 90.4 g CuO, what is the maximum amount of Cu metal that can be formed? NH 3 (g) + 3CuO(s) N (g) + 3Cu(s) + 3H O(g) (MM) (17.03) (79.55) (8.01) (64.55)(18.0) (g/mol) A. 17 g B. 103 g C. 7. g D. 108 g E g 18.1g NH 1mol NH g NH 3 mol CuO mol NH g CuO needed Only have 90.4 g so CuO limiting g CuO 1mol CuO 1mol CuO 3 mol Cu g Cu 90.4 g CuO g CuO 3 mol CuO 1mol Cu = 7. g Cu can be formed 93

94 Your Turn! If 8.51 g C H 5 SH reacts with.4 g O, what is the maximum amount of sulfur dioxide that can form? C H 5 SH(l) + 9O (g) 4CO (g) + 6H O(g) + SO (g) (6.13) (3.00) (44.01) (18.0) (64.06) A g B g C g D g E g 1mol CH5SH 9 mol O 3.00g O 8.51g CH5SH 6.13g C H SH mol C H SH 1mol O 19.7 g O needed Have.4 g so C H 5 SH is limiting 5 5 1mol CH5SH molso g SO 8.51g CH5SH 6.13g C H SH mol C H SH 1molSO 5 5 = 8.77 g SO can form 94

95 Reaction Yield In many experiments, the amount of product is less than expected Losses occur for several reasons Mechanical issues sticks to glassware Evaporation of volatile (low boiling) products. Some solid remains in solution Competing reactions and formation of by-products. Main reaction: P(s) + 3 Cl (g) PCl 3 (l ) Competing reaction: PCl 3 (l ) + Cl (g) PCl 5 (s) By-product 95

96 Theoretical vs. Actual Yield Theoretical Yield Amount of product that must be obtained if no losses occur Amount of product formed if all of limiting reagent is consumed Actual Yield Amount of product that is actually isolated at end of reaction Amount obtained experimentally How much is obtained in mass units or in moles 96

97 Percentage Yield Useful to calculate percentage yield Percentage yield Relates the actual yield to the theoretical yield It is calculated as: ercentage yield actual theoretical yield yield 100% Example: If a cookie recipe predicts a yield of 36 cookies and yet only 4 are obtained, what is the percentage yield? percentage yield % 67% 97

98 Ex. Percentage Yield Calculation When 18.1 g NH 3 and 90.4 g CuO are reacted, the theoretical yield is 7. g Cu. The actual yield is 58.3 g Cu. What is the percent yield? NH 3 (g) + 3CuO(s) N (g) + 3Cu(s) + 3H O(g) percentage yield = 58.3 g Cu 7. g Cu 100% = 80.7% 98

99 Learning Check: Percentage Yield A chemist set up a synthesis of solid phosphorus trichloride by mixing 1.0 g of solid phosphorus with 35.0 g chlorine gas and obtained 4.4 g of solid phosphorus trichloride. Calculate the percentage yield of this compound. Analysis: Write balanced equation P(s) + Cl (g) PCl 3 (s) Determine Limiting Reagent Determine Theoretical Yield Calculate Percentage Yield 99

100 Learning Check: Percentage Yield Assembling the Tools: 1 mol P = g P 1 mol Cl = g Cl 3 mol Cl mol P Solution 1. Determine Limiting Reactant 1 mol P 3 mol Cl g Cl 1.0 g P = 41. g Cl g P mol P 1 mol Cl But you only have 35.0 g Cl, so Cl is limiting reactant 100

101 Learning Check: Percentage Yield Solution. Determine Theoretical Yield 1 mol Cl mol PCl g Cl g Cl 3 mol Cl g PCl 1 mol PCl Determine Percentage Yield Actual yield = 4.4 g percentage yield = 4. g PCl 45. g PCl 3 = 45. g PCl 3 100% 3 = 93.8 % 101

102 Your Turn! When 6.40 g of CH 3 OH was mixed with 10. g of O and ignited, 6.1 g of CO was obtained. What was the percentage yield of CO? CH 3 OH + 3O CO + 4H O MM(g/mol) (3.04) (3.00) (44.01) (18.0) A. 6.1% 1 mol CH3OH 3 mol CO 6.40 g CH3OH B. 8.79% 3.04 g CH3OH mol CH3OH C. 100% = 9.59 g O needed; CH 3 OH limiting 1 mol CH3OH mol CO D. 14%.40 g CH3OH 3.04 g CH3OH mol CH3OH E. 69.6% 6 = 8.79 g CO in theory 6.1 g CO actual 100 % = 69.6% 8.79 g CO theory 3.00 g O 1 mol g CO 1 mol O CO 10

103 Stoichiometry Summary 103