Ch. 5 Notes
The mole (mol) is one of the seven base units in the SI system. It measures the amount of substance.
The form in which a substance exists is its “representative particle”. Representative particles can be
atoms, ions, molecules, formula units, or anything else. Just as one dozen is 12 representative particles,
a mole is 6.02 x 1023 representative particles. 602,000,000,000,000,000,000,000!!!!!
Examples:
1 mole Fe = 6.02 x 1023 atoms of Fe
1 mole H2O molecules = 6.02 x 1023 molecules of water
1 mole Na+ ions = 6.02 x 1023 Na+ ions
1 mole eggs = 6.02 x 1023 eggs
6.02 X 1023 Watermelon Seeds:
Would be found inside a melon slightly larger than the moon.
6.02 X 1023 Donut Holes:
Would cover the earth and be 5 miles (8 km) deep.
6.02 X 1023 Pennies:
Would make at least 7 stacks that would reach the moon.
6.02 X 1023 Grains of Sand:
Would be more than all of the sand on Miami Beach.
6.02 X 1023 Blood Cells:
Would be more than the total number of
blood cells found in every human on earth.
Diatomic Elements
Certain elements are only stable in pairs or with other elements in a compound. These elements are
called the diatomic elements. There are 7 diatomic elements: Hydrogen, bromine, oxygen, nitrogen,
chlorine, iodine, and fluorine
(Memory trick: HOBrFINCl twins or 7th Heaven)
Avogadro’s Number
6.02 x 1023 is called Avogadro’s number. It is named after Amadeo Avogadro who did work in the 1800’s
that allowed 6.02 x 1023 to be calculated.
The mole is the “chemist’s dozen”. It is a convenient way to count extremely large numbers of atoms,
molecules or ions.
New Conversion Factor!
1 mole = 6.02 x 1023 representative particles
We work these problems using dimensional analysis.
Gram atomic mass (gam)
-atomic mass of an element in grams
-mass of one mole of atoms of a monatomic element
-use the periodic table and take masses to 0.1 g
Example:
C =12.0 g = mass of 6.02 x 1023atoms
12.0 g/mol is the gram atomic mass of carbon
Gram molecular mass (gmm)
-mass of one mole of a molecule
-sum of the atomic masses of each atom in the molecule
1 mol H2O:
2 mol H
1 mol O
=2 x 1.0 g H/mol
=1 x 16.0 g O/mol
= 2.0 g
= 16.0 g
18.0 g H2O
Gram formula mass (gfm)
-mass of one mole of an ionic compound
-sum of the atomic masses of each atom in a formula unit
Gram formula mass is a generic term and can be used for either ionic or molecular compounds.
Molar mass or molecular weight are terms also used to mean the same thing.
Mole – Mass Conversions
New Conversion factor!
1 mol = gfm
Molar Volume of a Gas
The volume of a gas is usually measured at 0oC and 1 atmosphere of pressure.
This is called standard temperature and pressure (STP).
-At STP, one mole of any gas has a volume of 22.4 L.
-22.4 L is called the molar volume of a gas.
-22.4 L of a gas at STP contains 6.02 x 1023 particles of the gas.
-22.4 L of a gas has a mass equal to the gfm of the gas.
New conversion factor!
1 mol of any gas at STP = 22.4 L (for gases only!)
Gas Density
If we know the density of a gas in g/L at STP, we can calculate the gfm of the gas (or visa versa)
22.4 L/mol x density(g/L) = gfm
Percent Composition
(remember: percent = part divided by total x 100)
percent by mass of each element in a compound
# of grams of the element per 100 grams of the compound
OR solve for percent composition from the chemical formula:
% mass = grams of element in 1 mol of cmpd x 100
gfm of compound
CALCULATING EMPIRICAL FORMULAS
Empirical formula - lowest whole number ratio of the elements in a compound
- may or may not be the same as the molecular formula.
Ex. H2O is both the empirical & molecular formula.
H2O2 is a molecular formula.
HO is the empirical formula for H2O2.
Steps to calculate empirical formula:
1. Find moles of each element.
2. Set up mole ratio.
3. Simplify mole ratio (divide by smallest). If your answers are not in whole numbers, you must
multiply by 2,3,4,or 5 to get whole numbers.
4. Use mole ratio as subscripts in the formula.
If given % composition, assume 100 g of compound.
Poem:
Percent to mass,
Mass to moles,
Divide by small,
Multiple ‘til whole!
Molecular formulas are the actual formulas. They may be the same as the empirical formula or a
multiple of it. To find the multiple (n), take the gram formulas mass (gfm) and divide by the empirical
formula mass (efm):
n = gfm
efm
Multiply each subscript in the empirical formula by n to get the molecular formula.
Ex. Complete combustion (reaction with O2) of a sample of propane gas (contains only C & H) produced
2.641 g of CO2 and 1.442 g of water as the only products. Find the empirical formula for propane.
2.641g CO2 | 1 mol CO2 | 1 mol C
| 44.0 g CO2 | 1 mol CO2
= 0.0600 mol C
1.442g H2O | 1 mol H2O | 2 mol H =
| 18.0g H2O | 1 mol H2O
C:H 0.0600 : 0.160 = 1 : 2.67 (*3) = 3 : 8
0.0600 0.0600
C3H8
0.160 mol H