Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed.
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1 Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed. by Steven S. Zumdahl & Donald J. DeCoste University of Illinois
2 Chapter 8 Chemical Composition
3 Atomic Masses Balanced equations tell us the relative numbers of molecules of reactants and products. C + O 2 CO 2 1 atom of C reacts with 1 molecule of O 2 to make 1 molecule of CO 2 Copyright Houghton Mifflin Company. All rights reserved. 8 3
4 Atomic Masses (cont.) If you want to know how many O 2 molecules you will need, or how many CO 2 molecules you can make, you will need to know how many C atoms are in the sample of carbon you are starting with. You can either count the atoms (!) or weigh the carbon and calculate the number of atoms if you knew the mass of one atom of carbon. Copyright Houghton Mifflin Company. All rights reserved. 8 4
5 Atomic Masses (cont.) Atomic masses allow us to convert weights into numbers of atoms. Unit is the amu (atomic mass unit) 1 amu = 1.66 x g Copyright Houghton Mifflin Company. All rights reserved. 8 5
6 Atomic Masses (cont.) Copyright Houghton Mifflin Company. All rights reserved. 8 6
7 Atomic Masses (cont.) If our sample of carbon weighs 3.00 x amu, we will have 2.50 x atoms of carbon. Since our equation tells us that 1 C atom reacts with 1 O 2 molecule, if we have 2.50 x C atoms, we will need 2.50 x molecules of O 2 Copyright Houghton Mifflin Company. All rights reserved. 8 7
8 Example #1 Calculate the mass (in amu) of 75 atoms of Al. Determine the mass of 1 Al atom 1 atom of Al = amu Use the relationship as a conversion factor Copyright Houghton Mifflin Company. All rights reserved. 8 8
9 Moles The mass of 1.0 mole of an element is equal to the atomic mass in grams. A mole is the number of particles equal to the number of carbon atoms in 12 g of C-12. One mole = x units. This number is called Avogadro s number. 1 mole of C-12 atoms weighs 12.0 g and has 6.02 x atoms. 1 atom of Carbon-12 weighs 12.0 amu. Copyright Houghton Mifflin Company. All rights reserved. 8 9
10 Example #2 Compute the number of moles and the number of atoms in 10.0 g of Al. Copyright Houghton Mifflin Company. All rights reserved. 8 10
11 Example #2 (cont.) Use the periodic table to determine the mass of 1 mole of Al. 1 mole Al = g Use this as a conversion factor for grams-to-moles. Copyright Houghton Mifflin Company. All rights reserved. 8 11
12 Example #2 (cont.) Use Avogadro s Number to determine the number of atoms in 1 mole. 1 mole Al = 6.02 x atoms Use this as a conversion factor for moles-to-atoms. Copyright Houghton Mifflin Company. All rights reserved. 8 12
13 Example #3 Compute the number of moles in and the mass of 2.23 x atoms of Al. Copyright Houghton Mifflin Company. All rights reserved. 8 13
14 Example #3 (cont.) Use Avogadro s Number to determine the number of atoms in 1 mole. 1 mole Al = 6.02 x atoms Use this as a conversion factor for atoms-to-moles. Copyright Houghton Mifflin Company. All rights reserved. 8 14
15 Example #3 (cont.) Use the periodic table to determine the mass of 1 mole of Al. 1 mole Al = g Use this as a conversion factor for moles-to-grams. Copyright Houghton Mifflin Company. All rights reserved. 8 15
16 Molar Mass Molar mass: the mass in grams of one mole of a compound The relative weights of molecules can be calculated from atomic masses water = H 2 O = 2(1.008 amu) amu = amu 1 mole of H 2 O will weigh g, therefore the molar mass of H 2 O is g 1 mole of H 2 O will contain g of oxygen and 2.02 g of hydrogen Copyright Houghton Mifflin Company. All rights reserved. 8 16
17 Percent Composition Percentage by mass of each element in a compound Can be determined from the formula of the compound or by experimental mass analysis of the compound The percentages may not always total to 100% due to rounding. Copyright Houghton Mifflin Company. All rights reserved. 8 17
18 Example #4 Determine the percent composition from the formula C 2 H 5 OH. Copyright Houghton Mifflin Company. All rights reserved. 8 18
19 Example #4 (cont.) Determine the mass of each element in 1 mole of the compound. 2 moles C = 2(12.01 g) = g 6 moles H = 6(1.008 g) = g 1 mol O = 1(16.00 g) = g Determine the molar mass of the compound by adding the masses of the elements. 1 mole C 2 H 5 OH = g Copyright Houghton Mifflin Company. All rights reserved. 8 19
20 Example #4 (cont.) Divide the mass of each element by the molar mass of the compound and multiply by 100% Copyright Houghton Mifflin Company. All rights reserved. 8 20
21 Empirical Formulas Empirical formula: the simplest, whole-number ratio of atoms in a molecule Can be determined from percent composition or by combining masses Molecular formula: a multiple of the empirical formula 100g MM A % A mass A (g) moles A moles A 100g MM B moles B % B mass B (g) moles B Copyright Houghton Mifflin Company. All rights reserved. 8 21
22 Example #5 Determine the empirical formula of benzopyrene, C 20 H 12 Copyright Houghton Mifflin Company. All rights reserved. 8 22
23 Example #5 (cont.) Find the greatest common factor (GCF) of the subscripts. factors of 20 = (10 x 2), (5 x 4) factors of 12 = (6 x 2), (4 x 3) GCF = 4 Divide each subscript by the GCF to get the empirical formula. C 20 H 12 = (C 5 H 3 ) 4 Empirical Formula = C 5 H 3 Copyright Houghton Mifflin Company. All rights reserved. 8 23
24 Example #6 Determine the empirical formula of acetic anhydride if its percent composition is 47% carbon, 47% oxygen, and 6.0% hydrogen. Copyright Houghton Mifflin Company. All rights reserved. 8 24
25 Example #6 (cont.) Convert the percentages to grams by assuming you have 100 g of the compound. Step can be skipped if given masses Copyright Houghton Mifflin Company. All rights reserved. 8 25
26 Example #6 (cont.) Convert the grams to moles Copyright Houghton Mifflin Company. All rights reserved. 8 26
27 Example #6 (cont.) Divide each by the smallest number of moles. For this example, 2.9 is the smallest. Copyright Houghton Mifflin Company. All rights reserved. 8 27
28 Example #6 (cont.) If any of the ratios is not a whole number, multiply all the ratios by a factor to make it a whole number. If ratio is?.5, then multiply by 2; if?.33 or?. 67 then multiply by 3; if?.25 or?.75, then multiply by 4 Multiply all the Ratios by 3 Because C is 1.3 Copyright Houghton Mifflin Company. All rights reserved. 8 28
29 Example #6 (cont.) Use the ratios as the subscripts in the empirical formula. C 4 H 6 O 3 Copyright Houghton Mifflin Company. All rights reserved. 8 29
30 Molecular Formulas The molecular formula is a multiple of the empirical formula. To determine the molecular formula you need to know the empirical formula and the molar mass of the compound. Copyright Houghton Mifflin Company. All rights reserved. 8 30
31 Example #7 Determine the molecular formula of benzopyrene if it has a molar mass of 252 g and an empirical formula of C 5 H 3 Copyright Houghton Mifflin Company. All rights reserved. 8 31
32 Example #7 (cont.) Determine the empirical formula - May need to calculate it as previous C 5 H 3 Determine the molar mass of the empirical formula 5 C = g, 3 H = g C 5 H 3 = g Copyright Houghton Mifflin Company. All rights reserved. 8 32
33 Example #7 (cont.) Divide the given molar mass of the compound by the molar mass of the empirical formula Round to the nearest whole number Copyright Houghton Mifflin Company. All rights reserved. 8 33
34 Example #7 (cont.) Multiply the empirical formula by the calculated factor to give the molecular formula (C 5 H 3 ) 4 = C 20 H 12 Copyright Houghton Mifflin Company. All rights reserved. 8 34
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