Chapter 10

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# Chapter 10 - PowerPoint PPT Presentation

Chapter 10. Chemical Quantities. Molar Mass. MOLAR MASS—# of grams of an element in one mole of that element. = atomic mass of an element with the units of g/mol Get it from the periodic table. Molar Mass. Elements. Find the molar mass on the periodic table round to 0.1 g/mol.

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### Chapter 10

Chemical Quantities

Molar Mass
• MOLAR MASS—# of grams of an element in one mole of that element.

= atomic mass of an element with the units of g/mol

• Get it from the periodic table

Molar Mass

Elements

• Find the molar mass on the periodic table round to 0.1 g/mol

Compounds

• ADD the molar mass of every atom present in the compound
If you know the MOLAR MASS …
• You can convert between grams and moles!

Volume

of gas at STP

1 Mole

Molar mass (g)

Mass

Representative Particles

Mole ↔ Mass Examples
• How many moles in 51.0 g of Na3PO4?
• Find Molar Mass: (23.0 x 3) +31.0+(16.0 x 4)

= 164.0 g/mol

• 51.0 g Na3PO4 x 1 mole Na3PO4= 0.311 mol Na3PO4

164 g Na3PO4

• What is the mass of 0.70 mol of NH4Cl?
• Find Molar Mass: 14.0+(1.0 x 4) +35.5

= 53.5g/mol

• 0.70 mol NH4Cl x 53.5 g NH4Cl = 37.45 g NH4Cl

1 mol NH4Cl

1 mole= Avogadro’s # = 6.02x1023 representative particles

• How big is 6.02x1023? 602 000 000 000 000 000 000 000
• A representative particle: atom, molecule, or formula unit
• 1 mole of nitrogen gas (N2) contains 6.02 x 1023molecules.
• Because N2 is a molecule (Covalent)
• 1 mole of calcium fluoride (CaF2) contains 6.02 x 1023formula units
• Because CaF2 is a formula unit (ionic)
• How many atoms of He are in 1 mole?

Volume

of gas at STP

6.02 x 10 23

1 Mole

Molar mass

Mass

Representative Particles

Unit: Grams (g)

Type of Substance:Unit:

Elements Atoms

ionic Formula Units (f.un)

Covalent Molecules (mocs)

By COUNT: Avogadro's NumberHow big is 6.02x1023?602 000 000 000 000 000 000 000
• If you had Avogadro's number of un-popped 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.
• A mole of coke cans would cover the surface of the earth to a depth of over 200 miles.
Measuring Matter
• Some units of measurement indicate specific numbers.
• A pair means 2
• A dozen means 12
• Knowing how count, mass, & volume relate allows you to convert between them.
• If 1 dozen apples = 12 apples, and 1 dozen apples has a mass of 2.0 kg, What is the mass of 90 apples?
• 90/12 = 7.5 dozen 7.5 dozen x 2.0 kg
• = 15 kg
What is a Mole?
• When dealing with tiny particles (atoms, ions, compounds), the sample size is usually very large.
• Counting is not practical.
• Just as a dozen eggs represents 12 eggs, a mole (mol) of a substance represents 6.02 x 1023 representative particles of that substance.
• The number 6.02 x 1023 is known as Avogadro’s number
Moles
• Determining the number of atoms in a mole of a compound:
• How many moles are in a representative particle (formula unit or molecule) of the substance?
• This can be determined from the formula.
• Example: each molecule of CO2 contains 3 atoms: 1 carbon atom and 2 oxygen atoms.
• A mole of CO2 contains 1 mole of carbon atoms and 2 moles of oxygen atoms.
The Mass of a Mole
• When dealing with atoms, it is often easier to work with mass.
• Gram atomic mass (gam)—the atomic mass of an element expressed in grams.
• The atomic mass of carbon is 12 amu
• The gam of carbon is 12 g.
• The gam is the mass of 1 mole of atoms of any element.
The Mass of a Mole of a Compound
• To determine the mass of a mole of a compound you need to know the formula and the gam of each atom in the compound.
• Add the masses of each atom to get the mass of the compound
• Sulfur trioxide (SO3)
• 1 mole of S (32.1g) & 3 moles of O (16.0 g each)
• 32.1g + 16.0g + 16.0g + 16.0g = 80.1 g
• Gram molecular mass (gmm)—the mass of 1 mole of a molecular compound
• Gram formula mass (gfm)—the mass of 1 mole of an ionic compound
Sec. 2 Mole-Mass-Volume
• We learned about gam, gmm, & gfm last time. We can use 1 broad term to tell the mass of a substance.
• Molar Mass—the mass (in grams) of 1 mole of a substance.
• Why do we have the 3 terms then? Sometimes the term molar mass in unclear. What is the molar mass of oxygen? Do you mean oxygen gas (O2)? Then the molar mass is 32.0g (2 x 16.0g). Or do you mean oxygen atoms (O)? Then the molar mass is 16.0 g.
The Volume of a Mole of Gas
• The volumes of 1 mole of different solid and liquid substances are not the same.
• The volumes of 1 mole of different gases are the same under the same conditions.
• To keep things under the same conditions, gases are measured at standard temperature and pressure (STP)
• Standard temp is 0°C
• Standard pressure is 101.3kPa (or 1 atm)
• At STP, 1 mole of any gas has a volume of 22.4 L
Converting Between moles, particles, mass, and volume

Volume

of gas at STP

Note: to convert between particles, mass, and volume, you have to go through moles.

1 mol

22.4 L

22.4 L

1 mol

Mole

6.02 x 1023 particles

1 mol

Molar mass

1 mol

1 mol .

molar mass

1 mol .

6.02 x 1023 particles

Mass

Representative Particles

Sec. 3: Percent Composition & Chemical Formulas
• Percent Composition—The relative amount of each element in a compound
• The % of all elements in the compound must equal 100%
• % mass of element = grams of element x 100

grams of compound

• Or
• % mass of element = molar mass of element x 100

molar mass of compound

Example
• An 8.20 g piece of magnesium combines completely with 5.40 g of oxygen to form a compound. What is the % composition?
• First—add 8.20 g & 5.40 g to get the mass of the compound. 8.20 + 5.40 = 13.60 g
• % Mg = (mass of Mg/mass of compound) x 100
• % Mg = 8.20g/13.60g x 100 = 60.3%
• % O = (mass of O / mass of compound) x 100
• % O = 5.40/13.60 x 100 = 39.7%
• Does this make sense?
• 60.3 + 39.7 = 100
Empirical Formulas
• Empirical formula—gives the lowest whole-number ratio of the elements in a compound
• An empirical formula may or may not be the same as a molecular formula.
• If the formulas are different, the molecular formula is a simple multiple of the empirical formula.
• Examples:
• The empirical formula for H2O2 is HO
• For CO2, the empirical & molecular formula are the same.
• C6H6 and C2H2 have the same empirical formula: CH
Molecular Formulas
• You can determine the molecular formula if you know empirical formula and molar mass.
• Divide the molar mass by the empirical formula mass.
• Multiply this number by all subscripts in the empirical formula to get the molecular formula.
Example:
• Find molecular formula with a molar mass of 60.0g and empirical formula of CH4N
• 1st find the empirical formula mass
• 1 C, 4 H, 1 N
• 12 + (4 x 1) + 14 = 30 g
• Then, divide molar mass by empirical mass
• 60.0g / 30 g = 2
• Multiply each element subscript by the this.
• C2H8N2