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**Chapter 10 Worksheet Examples**

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Worksheet: Molar mass

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**Example: Calculate the molar mass (gram molecular weight) of a mole of iodine, I2.**

Unit for molar mass I = 126.9(2) = 253.8 From the formula of the compound From the mass off of the periodic table.

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**---------------------------------------- add together = 342.17**

Example: Calculate the molar mass of a mole of aluminum sulfate (Al2(SO4)3). From the masses off of the periodic table. From the formula of the compound Al = (2) = 53.96 S = (3) = 96.21 O = (12) = add together = Unit for molar mass

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Worksheet: Mole as a unit of mass

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**Example: What is the mass of 5.00 moles of water?**

Molar Mass Calculation H = 1.01 (2) = 2.02 O = (1) = 16.00 = g/mol 90.1 g H2O 18.02 g/mol 5.00 mol X Molar mass From the problem

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**Example: What is the mass of 0.50 moles of calcium carbonate?**

Molar Mass Calculation Ca = (1) = 40.08 C = (1) = 12.01 O = (3) = 48.00 = g/mol g CaCO3 g/mol 0.50 mol X Molar mass From the problem

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**Example: How many moles of calcium chloride are in 333 grams of calcium chloride**

From the problem Molar Mass Calculation Ca = (1) = 40.08 Cl = (2) = 70.9 = g/mol 333 g CaCl2 g/mol 3.00 mol ÷ Molar mass

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Worksheet: Avogadro’s Number

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**Example: How many molecules of water are there in 3.00 moles of water?**

1.806x1024 mcl H2O 6.02x1023 mcl/mol 3.00 mol X The number of atoms/molecules in one mole is always 6.02x1023 From the problem

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**Example: How many moles of neon are there in 2.408x1024 atoms of neon?**

From the problem 2.408x1024 atm Ne 6.02x1023 atm/mol 3.99 mol ÷ The number of atoms/molecules in one mole is always 6.02x1023

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Worksheet: Molar Volume of a gas

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**Example: What is the volume, in liters, of a 2**

Example: What is the volume, in liters, of a 2.00 mole sample of methane (CH4) at STP? 44.8 L CH4 22.4 L/mol 2.00 mol X The number of liters in one mole is always 22.4 From the problem

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**Example: How many moles of ethane (C2H6) are there in 5**

Example: How many moles of ethane (C2H6) are there in 5.60 liters of ethane? From the problem 5.6 L C2H6 22.4 L/mol 0.25 mol ÷ The number of liters in one mole is always 22.4

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Worksheet: Mixed Mole Problems

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**Molar mass is g/mol and you find grams from the periodic table.**

We’ve done problems of one step. Moles to a unit or unit to moles. Now we can do multi step problems using this picture to help us see where we need to go next. This picture is a summary of all of the problems we have done up to this point along with helpful hints to units and numbers. The better you understand this the easier mole conversions will be.

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**Example: What would be the volume in liters of 40**

Example: What would be the volume in liters of grams of neon at STP? From the problem From the periodic table 40.36 g Ne 20.18 g/mol 2.00 mol 44.8 L Ne 22.4 L/mol 2.00 mol ÷ X The number of liters in one mole is always 22.4 We’re not to liters yet…

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**56 L 1.505x1024 mcl CO2 CO2 22.4 L/mol 6.02x1023 mcl/mol 2.5 mol**

Example: How many molecules would there be in 56 liters of carbon dioxide at STP? From the problem The number of liters in one mole is always 22.4 56 L CO2 22.4 L/mol 2.5 mol 1.505x1024 mcl CO2 6.02x1023 mcl/mol 2.5 mol ÷ X The number of atoms/molecules in one mole is always 6.02x1023 We’re not to molecules yet…

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**Percent Composition & Formulas**

Worksheet: Percent Composition & Formulas

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**Example: Calculate the percent composition of sodium hydrogen carbonate (NaHCO3)?**

Step 1 – find molar mass of the compound. Step 2 – divide mass of each element by molar mass. Step 3 – multiply answers by 100%. Molar Mass Calculation Na = (1) = 22.99 H = 1.01 (1) = 1.01 C = (1) = 12.01 O = 16.00(3) = 48.00 = g/mol

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**Example: Which pair of molecules has the same empirical formula?**

Molecular formula Empirical formula C2H4O2 c) NaCrO4 C6H12O6 d) Na2Cr2O7 CH2O CH2O These two can not be “reduced”. An empirical formula is similar to a reduced fraction.

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**Example: Calculate the empirical formula for a compound with 67**

Example: Calculate the empirical formula for a compound with 67.6% Hg, 10.8% S, 21.6% O? Step 1 – change percents to grams (assume we have 100 gram sample then percents are the number of grams). Step 2 – change the grams to moles. Step 3 – divide all answers by the smallest number (of the answers). Step 4 – the results become the subscripts for the formula. Step 4 HgSO4 Step 1 Step 2 Step 3 ÷ ÷

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Example: Find the molecular formula of ethylene glycol, which is used as antifreeze. The molecular mass is 62g/mol and the empirical formula is CH3O? Step 1 – Find the empirical formula mass (EFM) Step 2 – Divide the molar mass (from problem) by the EFM to get the multiplier. Step 3 – Use the multiplier to determine the subscripts by multiplying each subscript by the multiplier. Step 3 C2H6O2 Step 1 EFM Calculation C = (1) = 12.01 H = 1.01 (3) = 3.03 O = 16.00(1) = 16.00 =31.04 g/mol Step 2 MULTIPLIER

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