Chapter 11.1 The Mole

Mole SI base unit to measure the amount of a substance  Equal to Avogadro's number (NA) (6.02 x 1023) of things i.e. atoms, pennies, pencils, etc. 

The Mole  

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A counting unit Similar to a dozen, except instead of 12, it’s 602 billion trillion 602,000,000,000,000,000,000,000 6.02 X 1023 (in scientific notation) This number is named in honor of Amedeo Avogadro (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present

Just How Big is a Mole? 

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Enough pop cans to cover the surface of the earth to a depth of over 200 miles. If you had Avogadro's number of unpopped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.

Converting Moles to Particles # of moles x 6.02 x 1023 particles 1 mole

Practice Problems p. 311

Assignment 

Determine the number of representative particles for each of the following:     

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11.5 mol Ag 18 mol water 0.15 mol NaCl 3.75 mol Fe 12.5 mol CaCO3

How many moles of CaCl2 contains 1.26 x 1024 formula units of CaCl2 How many moles of Ag contains 4.59 x 1025 formula units of Ag

Chapter 11.2

Molar Mass

Molar Mass 

The Mass of 1 mole (in grams)

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Equal to the numerical value of the average atomic mass (get from periodic table) 1 mole of C atoms = 12.0 g 1 mole of Mg atoms =

24.3 g

Some people used to call this a (gram

formula mass) when used with ionic compounds

Find the molar mass (round to the tenths place)

A.1 mole of Br atoms = 79.9 g/mole B.1 mole of Sn atoms= 118.7 g/mole

Molar Mass of Molecules and Compounds Mass in grams of 1 mole equal numerically to the sum of the atomic masses 1 mole of CaCl2

= 111.1 g/mol

1 mole of N2O4 = 92.0 g/mol

Converting Moles to Grams 

grams = moles X molecular mass 2

moles of CuSO4 = _____g  G = 2 x 159.6 = 319.2 g

Converting Grams to Moles 

moles =

 779

grams molecular mass

g CuSO4 = _____ mol  Mol = 779/159.6 = 4.88 moles

Atoms/Molecules and Grams Since 6.02 X 1023 particles = 1 mole AND 1 mole = molar mass (grams)  You can convert atoms/molecules to moles and then moles to grams! (Two step process) 

Atoms/Molecules and Grams You can’t go directly from atoms to grams!!!! You MUST go thru MOLES.  That’s like asking 2 dozen cookies weigh how many ounces if 1 cookie weighs 4 oz? You have to convert to dozen first! 

Converting Grams to Atoms 

# grams/molar mass x NA

Atoms/Molecules and Grams How many atoms of Cu are present in 35.4 g of Cu? 35.4 g Cu

1 mol Cu 63.5 g Cu

6.02 X 1023 atoms Cu 1 mol Cu

= 3.4 X 1023 atoms Cu

How many atoms of K are present in 78.4 g of K? 78.4 g K

1 mol K 39.1 g K

6.02 X 1023 atoms K 1 mol K

= 1.20 X 1024 atoms K

Converting Atoms to Grams 

# atoms/NA x Molar Mass

 1.2

x 1025 atoms of N = _____ g

x 1025 x 14 = 279.1 6.02 x 1023

 1.2

Assignment  

Get a book Questions 11-19

Chapter 11.3

Molar Relationships

Molar Mass Sum of the molar masses of the component elements  In one mole of a compound, the ratio of moles of each element is the same as for one molecule (p.321 20-24) 

Chapter 11.4 Percent Composition and Formulas

Percent Composition 

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Percent by mass of each element in a compound Mass of Element x100=%mass Mass of Compound

Percent Composition Fe2O3 = 160 g Fe= 56 x 2 = 112 g 112/160x100 = 70% 48/160x100 = 30% O = 16 x 3 = 48 g CaCl2 H3PO4

Chemical Formulas of Compounds 

Formulas give the relative numbers of atoms or moles of each element in a formula unit always a whole number ratio (the law of definite proportions).

NO2; 2 atoms of O for every atom of N

1 mole of NO2 = 2 moles of O atoms to every 1 mole of N atoms

Chemical Formulas of Compounds 

If we know or can determine the relative number of moles of each element in a compound, we can determine a formula for the compound.

Empirical Formula The formula of a compound that expresses the smallest whole number ratio of the atoms present.  Ionic formula are always empirical formula  May or may not be the same as the molecular formula 

To obtain an Empirical Formula 1. Determine the mass in grams of each element present, if necessary. If the problem gives you a percent composition assume a 100 g sample

2. Calculate the number of moles of each element from the masses.

To obtain an Empirical

Formula

3. Divide each by the smallest number of moles to obtain the

simplest whole number ratio.

4. If whole numbers are not obtained* in step 3, multiply through by the smallest number that will give all whole numbers *

Be careful! Do not round off numbers prematurely

Assignment 

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A blue solid is 36.8% nitrogen and 63.2% oxygen what is the empirical formula? Determine the empirical formula for a compound that contains 36% aluminum and 64% Sulfur Propane is 81.8% carbon and 18.2% hydrogen, what is the empirical formula for propane? Aspirin is 60% carbon, 4.4% hydrogen, and 35.6% oxygen, what is the empirical formula of aspirin? Determine the empirical formula for a compound that contains 10.9% magnesium, 31.8% chlorine, and 57.3% oxygen

Molecular Formula 

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The formula that states the actual number of each kind of atom found in the compound Necessary since multiple compounds can have the same empirical formula Molecular Formula = Empirical Formula x factor by which everything is multiplied

Determining Molecular Formula 1. 2.

3.

4.

Calculate the empirical formula Calculate the empirical formula mass by multiplying number of moles by molar mass Divide the given molar mass by the mass of the empirical formula Multiply subscripts in the empirical formula by the answer in #3

Assignment 

Page 335 &337, #’s 51-62

Chapter 11.5

Hydrates

Hydrate Solids that contain water molecules  Each hydrate has a specific number of water molecules 

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Anhydrous – Water has been removed from a hydrate

Uses for Hydrates Used to store solar energy  Anhydrous materials are used as desiccants 

Naming Hydrates 

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# of water molecules follows the solid formula and a dot MgSO4 • 4H2O Prefixes for water are on page 338 Mass of water must be included in all calculations

Determining the Formula of a Hydrate  

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Write the formula for the solid Determine the mass of the water (Original mass – anhydrous mass) Calculate the moles of water (mass/18g) Calculate the moles of compound Determine the amount of water (mol H2O/mol compound) Rewrite formula

Determining the Formula of a Hydrate If 2.50 g of copper sulfate reduces to1.59 g after heating what is the formula for the copper sulfate Hydrate?

1. 2. 3.

4. 5. 6.

CuSO4 • _ H2O 2.50 g – 1.59 g = 0.91 g 0.91g/18g = 0.05 mols 1.59g/160g = 0.01 mols 0.05/0.01 = 5:1 CuSO4 • 5 H2O

Assignment 

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Find the formula for the following hydrates: 48.8% MgSO4 and 51.2% H2O An 11.75 g sample of cobalt (II) chloride hydrate is heated, 9.25 g of anhydrous cobalt chloride remains.

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What is a hydrate? Describe the experimental procedure for finding the formula of a hydrate. Name the compound having the formula SrCl26H2O Arrange these hydrates in order of increasing percent water content: MgSO47H2O, Ba(OH)28H2O, CoCl26H2O