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Chapter 13 Notes

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Chapter 13 NotesNotes,notes

    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 NaHCO. 3

    2. % NaHCO in baking powder. 3

    3. The Empirical Formula of a compound.

SECTION PAGES TOPIC Assignments # #

     Page # Questions #

    3.1 56-60 Atomic masses 76 -77 2-26 even, packet 3.2 60-62 The Mole

    Mass relations in Chemical 3.3 a 62-63 78 28-32, packet Formulas: percent composition

    Empirical and Molecular 3.3 b 64-66 78 36, 38, packet Formula

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 74, 78, 80 81

    Packet: as assigned.

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

Formula Mass

How many atoms of each element are in the formula?

1. CuSO ______________ 5. (NH)PO ______________ 4434

    2. NaHCO ______________ 6. Ba(OH) ______________32

3. HCHO ______________ 7. CH(NO) ______________ 2323533

    4. CHCHCOOH ______________ 32

What is the molecular mass of each of the following compounds?

8. NaS ______________ 2

    9. Ba(NO) ______________ 32

    10. (NH)P ______________ 43

    1. C1H(NO)______________ 3533

    12. (NH)PO ______________ 434

    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?

     10716. The element silver, Ag, has two naturally occurring isotopes: Ag with a mass of 106.905 amu, and 109107Ag. Silver consists of 51.82% Ag and has an average atomic mass of 107.868 amu. Calculate 109the mass of Ag.

    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 amu 9. 261.3 amu 10. 85.0 amu 11. 227.0 amu 12. 149.0 amu

    13. a. 1.840 b. 5.45 14. 179 15. 28.1 u 16. 108.9 u 17. 35.450 u

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

Moles

How many grams of each is needed?

1. 2.40 moles of NaOH

    2. 0.600 moles of Al(SO) 243

How many moles of each?

3. 40.0 g of KO 4. 100. g of Ni(CO) 2233

5. How many molecules are there in 0.400 mol of NO? How many atoms? 25246. How many moles are contained in 1.20 x 10 molecules of CO? 2

    7. How many ammonium ions are in 0.036 moles of ammonium phosphate?

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

    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?

     o11. A flask containing hydrogen gas at 0 C 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 H present. How many 2

    atoms does this represent?

     23 12. How many ions are in 3.01 x 10formula units of sodium hydroxide?

     -2213. Determine the formula mass of a substance in which one molecule has a mass of 1.06 x 10 grams

1. 96.0 g 2. 205g

    3. 0.423 moles 4. 0.336mol 23245. 2.41 x 10 molecules NO1.687 x 10 atoms6. 1.99 mole CO 25, 222+7. 6.50 x 10 NH 8. a. 834g b. 1.55 g c. 0.343 mol 4-2410. 1.66 x 10 mol 16. 5.4 mole/mile 27272311. 2256 moles; 1.358 x 10molec; 2.716 x 10 atoms 12. 6.02 x 10 ions

    13. 20. 64 g

    ch3part1.doc 3 2014-1-28

    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

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?

     233. What volume in liters at STP is occupied by 1.80 x 10 molecules of oxygen gas?

     194. What volume, in liters at STP, is occupied by 2.32 x 10 molecules of helium gas?

5. What volume would 8.0 g of methane (CH) occupy at STP? 4

6. Calculate the mass of 112 L of O at STP. 2

7. Calculate the density, in grams per liter at STP, of ethylene gas, CH. 24

8. Calculate the molar mass of the gas that has a mass of 3.74 g and a volume of 2.464 liters.

     -39. What is the volume, in liters, occupied by 1.00 x 10 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.

     -41. 105 L 2. 2.00 moles 3. 6.70 liters 4. 8.63 x 10 liters

    5. 11.2L 6. 160. g 7. 1.25 g/L 8. 34.0 g/L 239. 0.0224 L 10. 3.01 x 10 molec

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

    11. 16.8 L 12. 66.0 g 13. 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. AgCrO b. Al(SO) 227243

     .2. Calculate the mass of copper in 45 g of CuCO Cu(OH) 32

    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, CH 38

     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 BaCl 2

     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

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

    4. 25.4% Ca; 30.4% C; 3.8% H; 40.5% O

    5. a. 64 g b. 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(NO), forms two different hydrated salts. One contains 24.7% water; the other 32

    30.4% water. What are the formulas for these two hydrated salts?

8. A sample of a compound of carbon and hydrogen, H, is decomposed to produce 0.0314 g of solid 2

    carbon and 0.0728 L (at STP) of gaseous hydrogen (remember the diatomics). What is the empirical

    formula of the compound?

     39. An unknown compound decomposes to produce nitrogen gas, N, and oxygen gas, O. If 100.0 cm of 22

    each gas is formed at STP, what is the empirical formula of the compound?

1. a. CaCl 2

    2. KMnO3. AgO 4. Al(SO)4 2243

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

    .. .5. CH 6. CuSO 5HO 7. Ca(NO) 3HO; Ca(NO) 4HO 3842322322

    8. CH 9. NO 25

Molecular Formula

1. The simplest formula of vitamin C is found by analysis to be CHO. From another experiment, the 343

    molar mass is found to be about 180g/mol. What is the molecular formula of vitamin C?

2. The simplest formula of hexane is CH. Its molecular mass is about 86 u. What is the molecular 37

    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 NO. The density of the gas at STP is 4.11 g/L. What is the 2

    molecular mass of this gas? What is its molecular formula?

6. The empirical formula of a gas is CH. The density of the gas at STP is 2.14 g/L. What is the 4

    molecular formula of this gas?

     -227. If a molecule of a compound has a mass of 2.425 x 10 g. What is its gram molecular mass? If its

    empirical formula is SF, what is its molecular formula? 6

     -228. If a molecule of a compound has a mass of 1.46 x 10 g and its simplest formula is CHO, what is 24

    its

     a. gram molecular mass 9molar mass) b. its molecular formula?

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

    HO 2. CH 3. CH, CH 4. FeS FeS 1. C6866143761423, 23

    5. NONO 6. CH 7. SF, 146.0 g 8. CHO, 87.9 g 2, 243126482

    Mathematics of Formulas (Review 1)

1. Calculate the percent composition of CsClO. 2

     .2. How much phosphorus is contained in 5.00 g of the compound CaCO3Ca(PO)? 3342

    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?

     224. A sample of pure compound contains 2.04 g of sodium, 2.65 x 10 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), was found to contain 19.0% water. What is the 3

    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 CH. Derive its molecular 2

    formula.

     -2310. A molecule of a compound has a mass of 4.32 x 10 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, H, and O. When the compound was 2222completed, the following had been consumed: 2.92 g C, 7.32 x10 molecules of O, 5.45 L of H at 22

    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?

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

1. Cs, 66.3%; Cl , 17.7%; O, 16.0% 2. 0.900 g 3. ZnO 4. NaCO 235. KCrO6. Fe(SCN).3HO 7. CH 8. CHO227 324242 9. CH10CH 11. NaHPO . HO 36 . 22242

    12. CH 13. a. 40.0% C, 53.3% O, 6.67 %H b. COH c. CHO 81826126

    Stoichiometry of Formulas: Review #2

1. How many grams of fluorine, F, can be obtained form 25.7 g of silicon tetrafluoride, SiF? 4

2. Calculate the mass percent of hydrogen in morphine, CH NO. 17193

    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 MgI . xHO is analyzed to determine the value of x. A 1.557 g sample of the 22

    compound is heated to remove all the water. 1.0254 g of MgI remains after heating. What is the 2

    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?

     236. A given sample of pure compound contains 9.81 g of zinc, 1.8 x 10 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 X, a compound TiX is formed. Given the following data 2n

    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% HO. 2

    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 O. Calculate the simplest formula of 2

    this compound.

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    Chapter 3 - Mass Relationship in Chemistry: Stoichiometry of Formulas

     .HO 4. MgI 8HO 5. CN 1. 18.8 g 2. 6.67 % 3. C9842222

    6. ZnCrO 7. CH, 28.0 g/mol 8. TiCl 9. PO 2424346.10. CuSO 3HO 11. XO424

    ch3part1.doc 10 2014-1-28

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