Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas
Objectives
Part I: Sections 3.1-3.3
1.Define atomic, molecular, and formula mass.
2.Explain the importance of relative atomic mass scale.
3.Given the masses and abundance of the isotopes of an element, calculate its atomic mass. Given the atomic mass of an element and the masses of its isotopes, calculate their abundance.
4.Given the gram atomic mass and Avogadro’s number calculate the mass of an individual atom.
3.Calculate the molecular or formula mass of any substance.
Define the mole.
9.Do simple conversions to find the number of moles or molecules in a given mass or vice versa.
10.Convert mass to volume, and/or number of molecules for gases at STP.
5.Relate gram, atomic mass unit, and the Avogadro constant.
11.Define percentage composition and calculate the percentage composition for any given compound ( including hydrates).
12.Define and give examples of an empirical formula.
13.Determine the empirical formula of a substance from experimental percentage composition data, or from a given number of grams.
14.Define and give examples of molecular formula.
15.Given the molecular mass and the empirical formula, calculate the molecular formula.
16.Determine the percentage of water and the empirical formula in a hydrate.
Part II: Section 3.4
1.Define chemical reaction, and list the reactants and products in a given reaction.
2.Use the correct symbols for the physical state of each substance involved in a chemical equation.
3.Distinguish subscripts and coefficients and write a balanced equation given names and /or formulas for reactants and products.
4.Use a balanced equation to relate the numbers of moles or grams of reactants and products.
5.Given the numbers of moles or grams of each reactant, determine the limiting reactant and calculate the theoretical yield of product.
6.Relate the actual yield of product to the theoretical yield and percent yield.
Labs:
1.Decomposition of NaHCO3.
2.% NaHCO3 in baking powder.
3.The Empirical Formula of a compound.
SECTION # / PAGES # / TOPIC / AssignmentsPage # / Questions #
3.1 / 56-60 / Atomic masses / 76 -77 / 2-26 even, packet
3.2 / 60-62 / The Mole
3.3 a / 62-63 / Mass relations in Chemical Formulas: percent composition / 78 / 28-32, packet
3.3 b / 64-66 / Empirical and Molecular Formula / 78 / 36, 38, packet
WRITTEN ASSIGNMENTS:
TOPIC / PAGE # / QUESTION(S) #Balancing equations / 79 / 48-52
Mass relations in Reactions / 79 / 54-60; 64-68 all even
Unclassified and Conceptual / 81 / 74, 78, 80
Packet: as assigned.
Formula Mass
How many atoms of each element are in the formula?
1.CuSO4 ______5. (NH4)3PO4______
2.NaHCO3______6.Ba(OH)2______
3.HC2H3O2______7.C3H5(NO3)3______
4.CH3CH2COOH ______
What is the molecular mass of each of the following compounds?
8.Na2S______
9.Ba(NO3)2______
10.(NH4)3P______
11.C3H5(NO3)3______
12.(NH4)3PO4______
Relative Atomic Masses and Abundance
13.Zinc, Zn, has atomic mass of 65.35. The atomic mass of chlorine, Cl, is 35.45. A Zn atom is how many times as heavy as
a. Cl atom?b. as a C-12 atom?
14.What is the atomic mass of hafnium, Hf, if out of every 100 atoms, 5 have mass 176 u, 19 have mass 177 u, 27 have mass 178u, 14 have mass 179 u, and 35 have mass of 180 u?
15.What is the average atomic mass of silicon if 94.21% of its atoms have mass of 27.977 U, 4.70% have a mass of 28.976 u, and 1.09% have a mass of 29.974 u?
16.The element silver, Ag, has two naturally occurring isotopes: 107Ag with a mass of 106.905 amu, and 109Ag. Silver consists of 51.82% 107Ag and has an average atomic mass of 107.868 amu. Calculate the mass of 109Ag.
17.Chlorine has two isotopes, Chlorine -35 has an actual mass of 34.968 u and chlorine-37 has a mass of 36.9659 u. In any sample of chlorine atoms, 75.771% will be chlorine-35 and 24.229% will be chlorine-37. Calculate the average atomic mass of chlorine.
8. 78.0 amu9. 261.3 amu10. 85.0 amu11. 227.0 amu12. 149.0 amu
13. a. 1.840 b. 5.45 14. 17915. 28.1 u16. 108.9 u17. 35.450 u
Moles
How many grams of each is needed?
1.2.40 moles of NaOH
2.0.600 moles of Al2(SO4)3
How many moles of each?
3.40.0 g of K2O4.100. g of Ni2(CO3)3
5.How many molecules are there in 0.400 mol of N2O5? How many atoms?
6.How many moles are contained in 1.20 x 1024 molecules of CO2?
7.How many ammonium ions are in 0.036 moles of ammonium phosphate?
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Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas
8.Find the mass in grams of each quantity.
a. 5.08 moles of calcium nitrate
b. 0.0112 mol of potassium carbonate
c. 27.4 g of titanium(IV) oxide
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Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas
9.How many moles are there in one atom?
10.An automobile traveling 10.0 miles per hour produces 0.33 Lb of CO gas per mile. How many moles of molecules of CO are produced per mile?
11.A flask containing hydrogen gas at 0 oC was sealed at a pressure of 1 atm and the gas was found to weigh 4512 g. Calculate the number of moles and the number of molecules of H2 present. How many atoms does this represent?
12.How many ions are in 3.01 x 1023 formula units of sodium hydroxide?
13.Determine the formula mass of a substance in which one molecule has a mass of 1.06 x 10-22 grams
1.96.0 g2. 205g
3.0.423 moles4. 0.336mol
5.2.41 x 10 23 molecules N2O5, 1.687 x 1024 atoms6.1.99 mole CO2
7.6.50 x 1022 NH4+8.a. 834g b. 1.55 g c. 0.343 mol
10. 1.66 x 10-24 mol16. 5.4 mole/mile
11. 2256 moles; 1.358 x 1027molec; 2.716 x 1027 atoms12. 6.02 x 1023 ions
13. 20.64 g
Gases and the Mole(STP)
Solve the following problems. Show all work. Express your answers in the correct units with the appropriate number of significant figures.
1.What volume, in liters, will be occupied by 4.70 moles of helium gas at STP?
2.How many moles are present in 44.8 L of chlorine gas at STP?
3.What volume in liters at STP is occupied by 1.80 x 1023 molecules of oxygen gas?
4.What volume, in liters at STP, is occupied by 2.32 x 1019 molecules of helium gas?
5.What volume would 8.0 g of methane (CH4) occupy at STP?
6.Calculate the mass of 112 L of O2 at STP.
7.Calculate the density, in grams per liter at STP, of ethylene gas, C2H4.
8.Calculate the molar mass of the gas that has a mass of 3.74 g and a volume of 2.464 liters.
9.What is the volume, in liters, occupied by 1.00 x 10-3 mole of ammonia gas at STP?
10.How many molecules of nitrogen dioxide are present in 11.2 L of nitrogen dioxide gas at STP?
11.What volume would 24.0 g of oxygen gas occupy at STP?
12.What is the mass of 33.6 L of carbon dioxide at STP?
13.Calculate the density, in grams per liter, of dinitrogen oxide.
14.Calculate the molar mass of each of the following:
a. a gas, 5.00 liters of which has a mass of 5.85 g.
b. a gas, 500. mL of which has a mass of 0.98 g.
1. 105 L2. 2.00 moles3. 6.70 liters4. 8.63 x 10-4 liters
5. 11.2L6. 160. g7. 1.25 g/L8. 34.0 g/L
9. 0.0224 L10. 3.01 x 1023 molec
11. 16.8 L12. 66.0 g13. 1.96 g/L
14. a. 26.4 g/mol; 44 g/mol
Percent Composition
1.What is the % composition of each element in the following compounds?
a. Ag2Cr2O7b. Al2(SO4)3
2.Calculate the mass of copper in 45 g of CuCO3. Cu(OH)2
3.Calculate the percent composition of the compounds that is formed when 29.0 g of silver reacts completely with 4.30 g of sulfur.
4.Calculate the percent composition of calcium acetate.
5.Calculate the amount of hydrogen in the following amounts of these compounds:
a. 350 g of propane, C3H8
b. 378 g sodium hydrogen sulfate.
6.In a laboratory experiment, barium chloride dihydrate is heated to remove completely its water of hydration. Calculate
a. the experimental % of water;b. the percent of BaCl2
c. the percent error in this experiment
1. empty crucible and cover ...... 20.286 grams
2. crucible, cover, and contents before heating ...... 21.673 grams
3. crucible, cover and contents after heating ...... 21.461 grams
1. a. Ag, 50%; Cr, 24%; O, 26%b. Al, 15.8%; S, 28.1%; S, 56.1%
2.
3 a. 87.1% Ag, 12.9% S
4. 25.4% Ca; 30.4% C; 3.8% H; 40.5% O
5. a. 64 gb. 3.15 g
6. a. 15.3 %b. 84.7 %c. 4.08% check
Empirical Formulas
1.Find the empirical formulas for the compound formed from the following elements:
3.611 g Ca; 6.389g Cl
2.Find the empirical formula for a compound that has the following composition:
24.74%K; 34.76% Mn; 40.50% O
3.A compound of silver and oxygen decomposes when heated. Given the data table, calculate the empirical formula for this compound.
Mass of crucible, cover and compound before heating ...... 22.89 g
Mass of crucible, cover and compound after heating ...... 22.70 g
Mass of crucible and cover ...... 20.15 g
4.A compound was found to contain 2.16 g of Al, 3.85 g of S, and 7.68 g of O. Find its empirical formula.
5.An 0.884 g sample of a compound was found to contain 0.722 g of carbon. The rest of the sample was hydrogen. What is the empirical formula of this compound?
6.A 5.00 gram hydrated copper sulfate was heated until all the water was driven off. After heating, the remaining 3.19 g sample contained 1.27 g of copper, 0.64 g of sulfur, and 1.28 g of oxygen. Find the empirical formula of the hydrate. ( note: water is part of the formula).
7.Calcium nitrate, Ca(NO3)2, forms two different hydrated salts. One contains 24.7% water; the other 30.4% water. What are the formulas for these two hydrated salts?
8.A sample of a compound of carbon and hydrogen, H2, is decomposed to produce 0.0314 g of solid carbon and 0.0728 L (at STP) of gaseous hydrogen (remember the diatomics). What is the empirical formula of the compound?
9.An unknown compound decomposes to produce nitrogen gas, N2, and oxygen gas, O2. If 100.0 cm3 of each gas is formed at STP, what is the empirical formula of the compound?
1. a. CaCl2
2. KMnO43. Ag2O4. Al2(SO4)3
5. C3H86. CuSO4 . 5H2O7. Ca(NO3)2. 3H2O; Ca(NO3)2. 4H2O
8. C2H59. NO
Molecular Formula
1.The simplest formula of vitamin C is found by analysis to be C3H4O3. From another experiment, the molar mass is found to be about 180g/mol. What is the molecular formula of vitamin C?
2.The simplest formula of hexane is C3H7. Its molecular mass is about 86 u. What is the molecular formula of hexane?
3.Known hydrocarbon contains 83.6% C and 16.4% H by mass. What is its simplest formula? If its molecular mass is about 86 u, what is its molecular formula?
4.Iron reacts with sulfur to form iron sulfide. If 2.561 g of iron reacts with 2.206 g of sulfur, what is the simplest formula of the sulfide? If its molecular mass is about 208 g/mol, what is its molecular formula?
5.The empirical formula of a gas is NO2. The density of the gas at STP is 4.11 g/L. What is the molecular mass of this gas? What is its molecular formula?
6.The empirical formula of a gas is CH4. The density of the gas at STP is 2.14 g/L. What is the molecular formula of this gas?
7.If a molecule of a compound has a mass of 2.425 x 10-22 g. What is its gram molecular mass? If its empirical formula is SF6, what is its molecular formula?
8.If a molecule of a compound has a mass of 1.46 x 10-22 g and its simplest formula is C2H4O, what is its
a. gram molecular mass9molar mass)b. its molecular formula?
1.C6H8O62. C6H143. C3H7, C6H144. Fe2S3, Fe2S3
5. NO2, N2O46. C3H127. SF6, 146.0 g8. C4H8O2, 87.9 g
Mathematics of Formulas (Review 1)
1.Calculate the percent composition of CsClO2.
2.How much phosphorus is contained in 5.00 g of the compound CaCO3.3Ca3(PO4)2?
3.When 1.010 g of zinc vapor burned in air, 1.257 g of the zinc oxide is produced. What is the empirical formula of the zinc oxide?
4.A sample of pure compound contains 2.04 g of sodium, 2.65 x 1022 atoms of carbon, and 0.132 mol of oxygen atoms. Find the empirical formula.
5.A compound gave on analysis the following percent composition: K = 26.57% ; Cr = 35.36%; O = 38.07%. Derive the empirical formula.
6.A hydrate of iron(III) thiocyanate, Fe(SCN)3, was found to contain 19.0% water. What is the empirical formula for the hydrate?
7.A sample of a compound of carbon and hydrogen is decomposed to produce 0.0500L of gaseous hydrogen (remember the diatomics), and 0.0134 g of solid carbon. What is the empirical formula of the compound?
8.A compound has the following percent composition: C = 40.0%; H = 6.67%; O = 53/.3%. Its molecular mass is 60.0 u . Derive its molecular formula.
9.A gaseous compound has a density of 1.875 g/l. Its empirical formula is CH2. Derive its molecular formula.
10.A molecule of a compound has a mass of 4.32 x 10-23 g. Its empirical formula is CH. What is its molecular formula?
11.What is the molecular formula of a hydrated salt which has a formula mass of about 268 and contains 46.9% water of hydration? An analysis reveals the following composition: Na = 17.18%; P = 11.57%; H = 5.60%; and O = 65.70%.
12.Octane, a compound of hydrogen and carbon found in gasoline, has a molecular mass of 114.26 u. If the percentage of hydrogen in octane is 15.75, what is its molecular formula?
13.A compound was synthesized in a lab from the elements C, H2, and O2. When the compound was completed, the following had been consumed: 2.92 g C, 7.32 x1022 molecules of O2, 5.45 L of H2 at STP. Further lab work found the new compound’s molecular mass to be 180 u.
a. What is its percentage composition by mass?
b. What is its empirical formula?
c. What is its molecular formula?
1. Cs, 66.3%; Cl , 17.7%; O, 16.0%2. 0.900 g3. ZnO4. Na2CO3
5. K2Cr2O76. Fe(SCN)3.3H2O7. CH48. C2H4O2
9. C3H610. C2H211. Na2HPO4 . H2O
12. C8H1813. a. 40.0% C, 53.3% O, 6.67 %H b. COH2c. C6H12O6
Stoichiometry of Formulas: Review #2
1.How many grams of fluorine, F, can be obtained form 25.7 g of silicon tetrafluoride, SiF4?
2.Calculate the mass percent of hydrogen in morphine, C17H19 NO3.
3.The elemental analysis of acetylsalicylic acid (aspirin) is 60.0% C; 4.48% H, and 35.5% oxygen atoms. If the molecular mass of this substance is 180.2 amu, what is its molecular formula?
4.The compound MgI2 . xH2O is analyzed to determine the value of x. A 1.557 g sample of the compound is heated to remove all the water. 1.0254 g of MgI2 remains after heating. What is the value of x?
5.Cyanogen is 46.2% C and 53.8% N by mass. At STP 1.05 g of cyanogen occupies 0.452 L. What is the molecular formula of cyanogen?
6.A given sample of pure compound contains 9.81 g of zinc, 1.8 x 1023 atoms of chromium and 0.30 mol of oxygen gas. What is the simplest formula of this compound?
7.If the density of ethylene is 1.25 g/L at STP, and the ratio of carbon to hydrogen atoms is 1:2, what is the molar mass and formula of ethylene?
8.When the metal Ti is heated in halogen X2, a compound TiXn is formed. Given the following data calculate the simplest formula of this new compound.
Mass of crucible + cover= 28.35 g
Mass of crucible + cover + titanium= 29.35 g
Mass of crucible + cover + final product= 31.57 g
9.When 10.00 g of phosphorus was reacted with oxygen, it produces 17.77 of the oxide. This oxide of phosphorus was found to have a molecular mass of approximately 220 amu in the vapor phase. Determine its molecular formula.
10.A certain hydrate analyzes as follows: 29.97% copper, 15.0% sulfur, 2.8% hydrogen, and 52.5% oxygen. Determine the empirical formula of this hydrate from these percentages knowing that it contains 25.3% H2O.
11.A 0.240 g sample of a compound of oxygen and element X, which has atomic mass of 42.8 amu, was found by analysis to contain 0.192 g of X and 0.00897 mol of O2. Calculate the simplest formula of this compound.
1. 18.8 g2. 6.67 %3. C9H8O44. MgI2.8H2O5. C2N2
6. ZnCr2O47. C2H4, 28.0 g/mol8. TiCl39. P4O6
10. CuSO4. 3H2O11. XO4
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