Moles

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• Moles

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Counting by Weighing

• Objects do not need to have identical masses to
  be counted by weighing.
• All we need to know is the average mass of the
  objects.
• To count the atoms in a sample of a given element
  by weighing we must know the mass of the sample
  and the average mass for that element.

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Counting by Weighing

• Averaging the Mass of Similar Objects

• Example What is the mass of 1000 jelly beans?
• Not all jelly beans have the same mass.
• Suppose we weigh 10 jelly beans and find

1. Now we can find he average mass of a bean.

1. Finally we can multiply to find the mass of 1000
   beans!

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• How do you make sure you have the same number of
  Mints and jelly beans in a bag?

One scoop of Jelly Beans is 500 grams
Average Mass Jelly Beans 5 grams Mints 15 grams
So You Need 3 X 500 1500 g of Mints
Ave Mass Mints Ave Mass Jelly Beans
15g 5 g
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Counting by Weighing

• Averaging the Mass of Different Objects

• Two samples containing different types of
  components (A and B), both contain the same
  number of components if the ratio of the sample
  masses is the same as the ratio of the masses of
  the individual components.

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1. What is the Mole?

• A counting number (like a dozen)
• 1 mol 6.02 ? 1023 items

602,000,000,000,000,000,000,000
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Avagadros number

• 6.02 X1023
• Number of particles of anything that make a mole.

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A. What is the Mole?
HOW LARGE IS IT???

• 1 mole of hockey pucks would equal the mass of
  the moon!

• 1 mole of basketballs would fill a bag the size
  of the earth!

• 1 mole of pennies would cover the Earth 1/4 mile
  deep!

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Particles in Chemistry

• Ionic Compounds
• Molecular compounds
• Elements

Formula Units Molecules Atoms
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Sample

• How many molecules in 3.1 moles of CO2 gas?

How many moles in 1.32 X 1023 formula units?
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2. Molar Mass

• Mass of 1 mole of an element or compound.
• Atomic mass tells the...
• atomic mass units per atom (amu)
• grams per mole (g/mol)
• Round to 2 decimal places

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B. Molar Mass Examples

• carbon
• aluminum
• zinc

12.01 g/mol 26.98 g/mol 65.39 g/mol
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B. Molar Mass Examples

• water
• sodium chloride

• H2O
• 2(1.01) 16.00 18.02 g/mol
• NaCl
• 22.99 35.45 58.44 g/mol

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B. Molar Mass Examples

• sodium bicarbonate
• sucrose

• NaHCO3
• 22.99 1.01 12.01 3(16.00) 84.01
  g/mol
• C12H22O11
• 12(12.01) 22(1.01) 11(16.00) 342.34
  g/mol

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C. Molar Conversions
molar mass
6.02 ? 1023
(g/mol)
(particles/mol)
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C. Molar Conversion Examples

• How many moles of carbon are in 26 g of carbon?

26 g C
1 mol C 12.01 g C
2.2 mol C
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C. Molar Conversion Examples

• How many molecules are in 2.50 moles of
  C12H22O11?

6.02 ? 1023 molecules 1 mol
2.50 mol
1.51 ? 1024 molecules C12H22O11
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C. Molar Conversion Examples

• Find the mass of 2.1 ? 1024 molecules of
  NaHCO3.

2.1 ? 1024 molecules
1 mol 6.02 ? 1023 molecules
84.01 g 1 mol
290 g NaHCO3
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3. Volume

• Volume of gas changes with pressure or
  temperature
• Hot gases expand
• Increased pressure makes the volume go down

STP Standard Temp and pressure 0 degrees C and
1 atm of pressure
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Gases

• 1 mole of any gas is 22.4 liters at STP

2.5 moles of CO2 has a volume of
2.5 moles X 22.4 56.0 Liters
How many moles is 10 Liters of O2 ?
1 mole
0.45 moles
10 liters X
22.4 Liters
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If you have moles Multiple to get desired units
If you do not have moles Divide to get moles
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Sample Problems

• How many molecules of H2O in 27 grams?

5.6 liters of CO2 (at STP) ______ grams
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Percent CompositionEmpirical FormulaMolecular
Formula

• Finding Formulas

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A. Percentage Composition

• the percentage by mass of each element in a
  compound

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A. Percentage Composition

• Find the composition of Cu2S.

? 100
Cu
79.852 Cu
? 100
S
20.15 S
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A. Percentage Composition

• Find the percentage composition of a sample that
  is 28 g Fe and 8.0 g O.

? 100
78 Fe
Fe
? 100
22 O
O
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A. Percentage Composition

• How many grams of copper are in a 38.0-gram
  sample of Cu2S?

Cu2S is 79.852 Cu
(38.0 g Cu2S)(0.79852) 30.3 g Cu
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A. Percentage Composition

• Find the mass percentage of water in calcium
  chloride dihydrate, CaCl22H2O?

24.51 H2O
? 100
H2O
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B. Empirical Formula

• Smallest whole number ratio of atoms in a
  compound

C2H6
reduce subscripts
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B. Empirical Formula

• 1. Find mass (or ) of each element.
• 2. Find moles of each element.
• 3. Divide moles by the smallest to find
  subscripts.
• 4. When necessary, multiply subscripts by 2, 3,
  or 4 to get whole s.

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B. Empirical Formula

• Find the empirical formula for a sample of 25.9
  N and 74.1 O.

25.9 g
1 mol 14.01 g
1.85 mol N
1 N
74.1 g
1 mol 16.00 g
4.63 mol O
2.5 O
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B. Empirical Formula

• N1O2.5

Need to make the subscripts whole numbers
? multiply by 2
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Practice

• What is the empirical formula for a compound if
  an 8.1 g sample contains 4.9 g of magnesium and
  3.2 g of oxygen?
• Moles of element Mass / molar mass
• Given molar mass of magnesium 24.3
  g           molar mass of elemental oxygen
  16.0 g (note we use elemental, not diatomic 
  oxygen)

Mg
O
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Practice 2

• Calculate the empirical formula of a compound
  that is made from 1.67 g of cerium and 4.54 g of
  iodine.

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C. Molecular Formula

• True Formula - the actual number of atoms in a
  compound

CH3
empirical formula
?
molecular formula
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C. Molecular Formula

• 1. Find the empirical formula.
• 2. Find the empirical formula mass.
• 3. Divide the molecular mass by the empirical
  mass.
• 4. Multiply each subscript by the answer from
  step 3.

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C. Molecular Formula

• The empirical formula for ethylene is CH2. Find
  the molecular formula if the molecular mass is
  28.1 g/mol?

empirical mass 14.03 g/mol
2.00
(CH2)2 ? C2H4
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Practice

• What is the molecular formula of a compound with
  an empirical formula of CH2 and a molecular
  formula of 56.0 u?
• Find Mass of empirical
• Divide

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• Part 2 Aqueous Solutions
• Chapter 15

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A. Definitions

• Solution - homogeneous mixture containing two or
  more substances

Solute - substance being dissolved
Solvent - present in greater amount
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A. Definitions
Solute - KMnO4
Solvent - H2O
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What are solutions?

• Not possible to distinguish solute from solvent
• Solutions can exist as gas, liquid, or solid,
  depending on the state of the SOLVENT
• Ex Air (gas) oxygen in nitrogen
• Braces (solid) titanium in nickel
• Most solutions are liquids

Homogeneous mixture
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Key Terms

• Soluble describes a substance that can be
  dissolved in a given solvent
• Ex sugar soluble in water
• Insoluble describes a substance that does not
  dissolve in a given solvent
• Ex sand insoluble in water

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Key Terms

• Immiscible describes TWO LIQUIDS that can be
  mixed together but separate shortly after you
  cease mixing them
• Ex oil and vinegar in salad dressing
• Miscible describes two liquids that are soluble
  in each other
• Aqueous solution solvent is water

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Solvation in Aqueous Solutions

• Why are some substances soluble and others are
  not?
• To form a solution, solute particles must
  separate from one another and the solute and
  solvent particles must mix.

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Solvation in Aqueous Solutions

• Attractive forces exist between pure solute
  particles, between pure solvent particles, and
  between the solute and solvent particles.
• Attractive forces between solute and solvent
  particles are greater than attractive forces
  holding solute particles together.

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Solvation

• Solvation the process of dissolving

solute particles are surrounded by solvent
particles
First...
solute particles are separated and pulled into
solution
Then...
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Solvation

• Dissociation
• separation of an ionic solid into aqueous ions

NaCl(s) ? Na(aq) Cl(aq)
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Aqueous solutions of IONIC compounds
Na ion
Cl- ion
Oxygen
Water molecule
Hydrogen
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Solvation

• Molecular Solvation
• molecules stay intact

C6H12O6(s) ? C6H12O6(aq)
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Aqueous solutions of MOLECULAR compounds

• Polar molecules molecules that have positive
  regions and negative regions
• In molecular compounds, the negative region of
  the solute molecule is attracted to the positive
  region of water, and vice versa
• Oil and water are immiscible because oil is
  nonpolar

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B. Solvation
Like Dissolves Like
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Solubility
concentration
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Factors that affect solubility

• Nature of the solute and solvent
• Temperature
• Affects solubility of ALL substances
• Pressure
• Affects the solubility of GASEOUS SOLUTES and
  GASEOUS SOLUTIONS