Atomic Masses. Chapter 3. Stoichiometry. Chemical Stoichiometry. Mass and Moles of a Substance. Average Atomic Mass


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1 Atomic Masses Chapter 3 Stoichiometry 1 atomic mass unit (amu) = 1/12 of the mass of a 12 C atom so one 12 C atom has a mass of 12 amu (exact number). From mass spectrometry: 13 C/ 12 C = amu So the mass of a 13 C atom is ( )(12 amu) = amu 1 amu = x kg So the mass of a 13 C atom is: ( amu)( x kg/amu) = x kg Chemical Stoichiometry Figure 3.1: (left) A scientist injecting a sample into a mass spectrometer. (right) Schematic diagram of a mass spectrometer. Stoichiometry  The study of quantities of materials consumed and produced in chemical reactions. Mass and Moles of a Substance Chemistry requires a method for determining the numbers of molecules in a given mass of a substance. This allows the chemist to carry out recipes for compounds based on the relative numbers of atoms involved. The calculation involving the quantities of reactants and products in a chemical equation is called stoichiometry. Average Atomic Mass Elements occur in nature as mixtures of isotopes We will call the average atomic mass the atomic weight Atomic weight is based on the relative abundance of isotopes Carbon = 98.89% 12 C = x 12 amu = % 13 C = x amu = <0.01% 14 C = neglect Average carbon atomic weight = amu
2 Figure 2.11: Diagram of a simple mass spectrometer. Conceptual Problem 2.47: Green and brown spheres. X23: (5/20) x amu = Figure 2.12: The mass spectrum of neon. 20 Ne = Ne =.003 X25: (15/20) x amu = Average atomic weight = amu 22 Ne =.090 Atomic Weight of Elements Calculate the average atomic weight of boron, B, from the following data: ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE B B The Mole One mole is the number equal to the number of carbon atoms in exactly 12 grams of pure 12 C. The number = x mole of anything = units B10: x = B11: x = = amu
3 Avogadro s number equals units Figure 3.2: One mole each of various substances. Photo courtesy of American Color. S 8 C 8 H 17 OH HgI 2 CH 3 OH Molar Mass A substance s molar mass (or molecular weight) is the mass in grams of one mole of the compound. CO 2 = grams per mole g/mole 1 C = O 2 x = g/mole Determining the Moles of Atoms How many moles of Se atoms are in 20.0 g of Se? 20.0 g Se x 1 mol Se = mol Se atoms g Se molar mass of Se How many atoms of Se are in 20.0 g of Se? mol Se atoms x x Se atoms = 1.53 x Se atoms mol Se atoms Write the molecular formula and calculate the formula weight for hydrogen peroxide and nitric acid a) H 2O 2 2 x at. wgt. H = 2.02 amu 2 x at. wgt. O = amu amu b) HNO 3 1 x at. wgt. H = 1.01 amu 1 x at. wgt. N = amu 3 x at. wgt. O = amu amu
4 Determining Molar Mass What is the molar mass of Al 2 (SO 4 ) 3? 1 mol Al 2 (SO 4 ) 3 x 2 mol Al x g = g mol Al 2 (SO 4 ) 3 mol Al 1 mol Al 2 (SO 4 ) 3 x 3 mol S x g = g mol Al 2 (SO 4 ) 3 mol S Figure 3.7: A welder welding using an oxyacetylene torch. Photo courtesy of PhotoDisc. CH 1 mol Al 2 (SO 4 ) 3 x 12 mol O x g = g mol Al 2 (SO 4 ) 3 mol O Molar mass of Al 2 (SO 4 ) 3 = g = g acetylene Percent Composition Mass percent of an element: mass of element in compound mass % = 100% mass of compound For iron in iron (III) oxide, (Fe 2 O 3 ) Molar mass of Fe 2 O 3 = g Mass of Fe = 1 mol Fe 2 O 3 x 2 mol Fe x g Fe = g mol Fe 2 O 3 mass % Fe = 100% = % Figure 3.7: Photo of benzene in lab glassware. Photo courtesy of American Color. CH benzene Determining the formula of a compound from the percent composition. The percent composition of a compound leads directly to its empirical formula. An empirical formula (or simplest formula) for a compound is the formula of the substance written with the smallest integer (whole number) subscripts.
5 Figure 3.5: A schematic diagram of the combustion device used to analyze substances for carbon and hydrogen. molecular formula = (empirical formula) n [n = integer] molecular formula = C 6 H 6 = (CH) 6 (benzene) molecular formula = C 2 H 2 = (CH) 2 (acetylene) empirical formula = CH Another example: hydrogen peroxide OH = empirical formula H 2 O 2 = molecular formula Benzoic acid Determining the empirical formula from the percent composition. Benzoic acid is a white, crystalline powder used as a food preservative. The compound contains 68.8% C, 5.0% H, and 26.2% O by mass. What is its empirical formula? In other words, give the smallest wholenumber ratio of the subscripts in the formula C x H y O z Benzoic Acid is a white, crystalline powder used as a food preservative which is 68.8% C, 5.00 % H, and 26.2% O by mass. What is its empirical formula? For g of benzoic acid: 68.8 g C x 1 mol C = 5.73 mol C 12.0 g C 5.00 g H x 1mol H = 4.95 mole H 1.01 g H 26.2 g O x 1 mol O = mol O 16.0 g O divide by 1.638, then empirical formula is C 3.5H 3.0O or C 7H 6O 2 Determining the empirical formula from the percent composition. For the purposes of this calculation, we will assume we have grams of benzoic acid. Then the mass of each element equals the numerical value of the percentage. Since x, y, and z in our formula represent molemole ratios, we must first convert these masses to moles. C x H y O z Empirical Formula Determination (when you are given mass percentages of elements) 1. Base calculation on 100 grams of compound. 2. Determine moles of each element in 100 grams of compound. 3. Divide each value of moles by the smallest of the values. 4. Multiply each number by an integer to obtain all whole numbers.
6 Another type of empirical formula problem g of a compound containing only C, H and O are combusted to form CO 2 and H 2 O. If g of CO 2 and g of H 2 O are formed, what is the empirical formula of the compound? 1. Calculate the masses of C and H formed: g CO 2 x (1 mol CO 2 /44.01 g CO 2 )x(1 mol C/mol CO 2) ) x (12.01 g C/mol C) = g C g H 2 O x (1 mol H 2 O/18.01 g) x ( 2 mol H/mol H 2 O) x (1.008 g H/mole H) = g H 2. Determine the mass of O by subtracting the masses of C and H from the mass of the compound reacted: g ( g C g O) = g O Determining the molecular formula from the empirical formula. For example, suppose the empirical formula of a compound is CH 2 O and its molar mass is 60.0 g/mol. The molar mass of the empirical formula (the empirical weight) is only 30.0 g/mol. This would imply that the molecular formula is actually the empirical formula doubled, or (CH 2 O) 2 = C 2 H 4 O 2 Empirical formula calculation continued. 3. Now calculate the moles of C, H and O in the compound: moles of C = g C x (1 mol C/12.01 g) = mol C moles of H = g H x (1 mol H/1.008 g)= mol H moles of O = g O x (1 mol O/16.00 g) = mol O Figure 3.6: Examples of substances whose empirical and molecular formulas differ. Notice that molecular formula = (empirical formula) n, where n is an integer. 4. Now divide mol C & mol H by mol O to get ratio of mol C and H to O since it is present in smallest no. moles = = and substitute the coefficients into C x H y O z to give C 3.5 H 3 O or C 7 H 6 O  the same empirical formula as benzoic acid! Determining the molecular formula from the empirical formula. An empirical formula gives only the smallest wholenumber ratio of atoms in a formula. The molecular formula should be a multiple of the empirical formula (since both have the same percent composition). To determine the molecular formula, we must know the molar mass of the compound. Chemical Equations Chemical change involves a reorganization of the atoms in one or more substances.
7 Chemical Equation Figure 3.11: Steps in a stoichiometric calculation. A representation of a chemical reaction: C 2 H 5 OH(g) + 3O 2 (g) 2CO 2 (g) + 3H 2 O(g) reactants products 1 mole C 2 H 5 OH + 3 moles O 2 º 2 moles CO moles H 2 O 1 molecule C 2 H 5 OH + 3 molecules O 2 º 2 molecules CO molecules H 2 O Practice Problem 3.79 Calculating Masses of Reactants and Products Practice Problem Balance the equation. 2.Convert mass to moles. 3.Set up mole ratios. 4.Use mole ratios to calculate moles of desired substituent. 5.Convert moles to grams, if necessary.
8 Figure 3.10: A mixture of CH4 and H2O molecules. Limiting Reactant Figure 3.11: Methane and water have reacted to form products according to the equation CH4 + H2O º 3H2 + CO. The limiting reactant is the reactant that is consumed first, limiting the amounts of products formed. Figure 3.9: Three different stoichiometric mixtures of methane and water, which react onetoone. Figure 3.12: Hydrogen and nitrogen react to form ammonia according to the equation N2 + 3H2 º 2NH3. CH4 + H2O º 3H2 + CO. 8
9 Solving a Stoichiometry Problem Practice Problem Balance the equation. 2. Convert masses to moles. 3. Determine which reactant is limiting. 4. Use moles of limiting reactant and mole ratios to find moles of desired product. 5. Convert from moles to grams. Conceptual Problem 3.13 Solving stoichiometry problems flow chart. Conceptual Problem 3.18 Calculating Percent Yield Actual Yield x 100% = % Yield Theoretical Yield What is the percent yield of menthol if the theoretical yield is 30.0 g and the actual yield is 20.0 g? 20.0 g 30.0 g x 100% = 66.7% 9
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