Chapter 6 Chemical Composition
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1 Chemistry 25 Chapter 6 Chemical Composition Why Is Knowledge of Composition Important? Everything in nature is either chemically or physically combined with other substances. To know the amount of a material in a sample, you need to know what fraction of the sample it is. Some Applications: The amount of sodium in sodium chloride for diet. The amount of iron in iron ore for steel production. The amount of hydrogen in water for hydrogen fuel. The amount of chlorine in freon to estimate ozone depletion. 1
2 The Mole Concept We can calculate the numbers of mole By numbers of atoms/molecules By mass of the atoms/molecules Avogadro s number (N A ). The number of atoms in a 12-gram sample of carbon 12 The value of Avogadro s number is 6.02 x Mole Calculation moles of "A" 2
3 Example A Silver Ring Contains 1.1 x Silver Atoms. How Many Moles of Silver Are in the Ring? Practice Calculate the Number of Atoms in 2.45 Mol of Copper. 3
4 Relationship Between Moles and Mass The mass of one mole of atoms is called the molar mass. The molar mass of an element, in grams, is numerically equal to the element s atomic mass, in amu. The lighter the atom, the less a mole weighs. the more atoms there are in 1 g. Mole and Mass Relationships Substance Pieces in 1 mole Weight of 1 mole Hydrogen x atoms g Carbon x atoms g Oxygen x atoms g Sulfur x atoms g Calcium x atoms g Chlorine x atoms g Copper x atoms g 4
5 Example Calculate the Moles of Sulfur in 57.8 gm of Sulfur. Practice Calculate the Moles of Carbon in g of Pencil Lead. 5
6 Practice How Many Copper Atoms Are in a Penny Weighing 3.10 g? Practice How Many Copper Atoms Are in a Penny Weighing 3.10 g? 6
7 Molar Mass of Compounds The relative weights of molecules can be calculated from atomic weights. Formula mass = 1 molecule of H 2 O Since 1 mole of H 2 O contains 2 moles of H and 1 mole of O. Molar mass = 1 mole H 2 O Example Calculate the Mass of 1.75 Mol of H 2 O. 7
8 Practice How Many Moles Are in 50.0 g of PbO 2? (Pb = 207.2, O = 16.00) Example What Is the Mass of 4.78 x NO 2 Molecules? 8
9 Practice How Many Formula Units Are in 50.0 g of PbO 2? (PbO 2 = 239.2) Chemical Formulas as Conversion Factors 1 spider 8 legs. 1 chair 4 legs. 1 H 2 O molecule 2 H atoms 1 O atom. 9
10 Counting Parts If we know how many parts are in the whole unit, by counting the number of whole units, we can effectively count the parts. All the desks in the room have 4 legs, if there are 30 desks in the room, Since every H 2 O molecule has 2 H atoms, in 100 H 2 O molecules, in 1 mole of H 2 O molecules, Mole Relationships in Chemical Formulas Since we count atoms and molecules in mole units, we can find the number of moles of a constituent element if we know the number of moles of the compound. Moles of compound 1 mol NaCl 1 mol H 2 O Moles of constituents 1 mol CaCO 3 1 mol C 6 H 12 O 6 10
11 Example Calculate the Moles of Oxygen in 1.7 Moles of CaCO 3. Find the Mass of Carbon in 55.4 g C 10 H 14 O. Practice Find the Mass of Sodium in 6.2 g of NaCl (MM: Na = 22.99, Cl = 35.45) 11
12 Percent Composition Percentage of each element in a compound. By mass. Can be determined from: The total percentage Find the Mass Percent of Cl in C 2 Cl 4 F 2. 12
13 Practice Determine the Mass Percent Composition of the Following: CaCl 2 (Ca = 40.08, Cl = 35.45) Mass Percent as a Conversion Factor The mass percent tells you the mass of a constituent element in 100 g of the compound. The fact that NaCl is 39% Na by mass means that 100 g of NaCl contains 39 g Na. This can be used as a conversion factor. 100 g NaCl 39 g Na 13
14 Empirical Formulas The simplest, whole-number ratio of atoms in a molecule is called the empirical formula. Can be determined from percent composition or combining masses. The molecular formula is a multiple of the empirical formula. Empirical Formulas, Continued Hydrogen Peroxide Molecular formula = Empirical formula = Glucose Molecular formula = Empirical formula = Benzene Molecular formula = Empirical formula = 14
15 Practice Determine the Empirical Formula of Benzopyrene, C 20 H 12. Find the greatest common factor (GCF) of the subscripts. Divide each subscript by the GCF to get the empirical formula. Finding an Empirical Formula 1. Convert the percentages to grams. a. Skip if already grams. 2. Convert grams to moles. a. Use molar mass of each element. 3. Write a pseudoformula using moles as subscripts. 4. Divide all by smallest number of moles. 5. Multiply all mole ratios by number to make all whole numbers, if necessary. a. If ratio?.5, multiply all by 2; if ratio?.33 or?.67, multiply all by 3, etc. b. Skip if already whole numbers after Step 4. 15
16 Example 1. Find grams. 2. Convert grams to moles. 3. Use moles as subscripts. 4. Divide all by smallest number of moles. 5. Whole numbers mole ratios Finding an Empirical Formula from Experimental Data A laboratory analysis of aspirin determined the following mass percent composition. Find the empirical formula. C = 60.00% H = 4.48% O = 35.53% Example 1. Find grams. 2. Convert grams to moles. 3. Use moles as subscripts. 4. Divide all by smallest number of moles. 5. Whole numbers mole ratios Finding an Empirical Formula from Experimental Data A 3.24-g sample of titanium reacts with oxygen to form 5.40 g of the metal oxide. What is the formula of the oxide? 16
17 Practice 1. Find grams. 2. Convert grams to moles. 3. Use moles as subscripts. 4. Divide all by smallest number of moles. 5. Whole numbers mole ratios Determine the Empirical Formula of Stannous Fluoride, which Contains 75.7% Sn (118.70) and the Rest Fluorine (19.00). Practice 1. Find grams. 2. Convert grams to moles. 3. Use moles as subscripts. 4. Divide all by smallest number of moles. 5. Whole numbers mole ratios Determine the Empirical Formula of Hematite, which Contains 72.4% Fe (55.85) and the Rest Oxygen (16.00). 17
18 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. Example Determine the Molecular Formula of Cadinene if it has a Molar Mass of 204 g and an Empirical Formula of C 5 H 8. 18
19 Practice Benzopyrene has a Molar Mass of 252 g and an Empirical Formula of C 5 H 3. What is its Molecular Formula? (C = 12.01, H=1.01) Practice Determine the Molecular Formula of Nicotine, which has a Molar Mass of 162 g and is 74.0% C, 8.7% H, and the Rest N. (C=12.01, H=1.01, N=14.01) 19
20 Vocabularies Avogadro s number The mole concept Percent composition Empirical formula Molecular formula 20
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