1. Chemical Stoichiometry
Stoichiometry - The study of quantities of materials consumed and produced in chemical reactions.
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Figure 3.1: (left) A scientist injecting a sample into a mass spectrometer. (right) Schematic diagram of a mass spectrometer.
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Atomic Masses
Elements occur in nature as mixtures of isotopes
Carbon = 98.89% 12C
1.11% 13C
<0.01% 14C
Carbon atomic mass = amu
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Figure 3.2: (a) Neon gas glowing in a discharge tube. (b) "peaks" and (c) a bar graph.
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5. Figure 3.3: Mass spectrum of natural copper.
When a sample of natural copper is vaporized and injected into a mass spectrometer, the results shown in Fig. 3.3 are obtained.
Use these data to compute the average mass of natural copper.
(The mass values for 63Cu and 65Cu are amu and amu, respectively.)

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The Mole
The number equal to the number of carbon atoms in exactly 12 grams of pure 12C.
1 mole of anything =  1023 units of
that thing
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Figure 3.4: Proceeding clockwise from the top samples containing one mole each of copper, aluminum, iron, sulfur, iodine, and (in the center) mercury.
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9. Avogadro’s number equals 6.022  1023 units
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React 1
Now Try This…
Calculate the number of copper atoms in a g sample of copper.
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Molar Mass
A substance’s molar mass (molecular weight) is the mass in grams of one mole of the compound.
CO2 = grams per mole
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QUESTION
The compound Cr2O3 (chromium (III) oxide) is one of the key components responsible for the red color of ruby gems. If you had 34.8 grams of Cr2O3, how many grams of Chromium (atomic number = 24) metal would be present?
grams of Cr
grams of Cr
grams of Cr
grams of Cr
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ANSWER
Choice 3 correctly predicts the grams of Cr found in 34.8 grams of Cr2O3.
34.8 grams Cr2O3  1 mol/152 g = moles of Cr2O3
0.229 Moles Cr2O3  2 mol Cr/1 mol Cr2O3 = 0.458moles of Cr
0.458 moles Cr  52.0 g/1 mol Cr = 23.8 grams
OR, 2  (52.0)/152  100 = 68.4 % Cr
0.684  34.8 = 23.8 g Cr
Section 3.3: Molar Mass
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Percent Composition
Mass percent of an element:
For iron in iron (III) oxide, (Fe2O3)
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React 5
Consider separate gram samples of each of the following: 
H2O N2O C3H6O2 CO2
Rank them from highest to lowest percent oxygen by mass.
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Formulas
molecular formula = (empirical formula)n
[n = integer]
molecular formula = C6H6 = (CH)6
empirical formula = CH
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17. Empirical Formula Determination
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.
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QUESTION
The dye indigo is a compound with tremendous economic importance (blue jeans wouldn’t be blue without it.) Indigo’s percent composition is: 73.27% C; 3.84% H; 10.68%N and 12.21% O. What is the empirical formula of indigo?
1. C6H4NO
2. C8H3NO
3. C8H5NO
4. I know this should be whole numbers for each atom, but I do not
know how to accomplish that.
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ANSWER
Choice 3 is the smallest whole number ratio of the atoms that make up a molecule of indigo. The percentage must be converted to a mass, then the mass is converted to moles of the atoms and finally, the smallest is divided into the others to obtain the proper ratio.
Section 3.5: Determining the Formula of a Compound
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Figure 3.6: Examples of substances whose empirical and molecular formulas differ. Notice that molecular formula = (empirical formula)n, where n is an integer.
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21. Figure 3.7: The two forms of dichloroethane.
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Chemical Equations
Chemical change involves a reorganization of the atoms in one or more substances.
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Chemical Equation
A representation of a chemical reaction:
C2H5OH + 3O2  2CO H2O
reactants products
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Chemical Equation
C2H5OH + 3O2  2CO H2O
The equation is balanced.
1 mole of ethanol reacts with 3 moles of oxygen
to produce
2 moles of carbon dioxide and 3 moles of water
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25. Calculating Masses of Reactants and Products
1. 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.
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Limiting Reactant
The limiting reactant is the reactant that is consumed first, limiting the amounts of products formed.
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27. Solving a Stoichiometry Problem
1. 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.
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