Download Molecular Mass - Teacher Notes

Unit 8 – The Mole
Essential Questions:
•What is the relationship between a mole of a
substance and its mass?
•How can the mole of a substance be
calculated?
•How can the percent composition of a
compound be determined?
•How does the molecular formula of a
compound compare with the empirical formula?
Formula Mass
•The sum of the average atomic mass for all atoms in
represented in a formula
•Unit is atomic mass units (amu)
1 atom of C
=
12.01 amu
1 atom of Mg
=
24.31 amu
1 atom of Cu
=
63.55 amu
Molecular Mass – the sum of the masses of
all the atoms in a molecule of a substance
The unit is amu.
CaCO3
1 atom of Ca = 40.08 amu
1 atom of C = 12.00 amu
3 atoms of O = 3 x 16.00 amu
100.08 amu
Example:
Find the molecular mass of NH4SO2
1N
4H
1S
2O
= 14.01 amu
= 4(1.01 amu)
= 32.07 amu
= 2(16.00 amu)
1 molecule
= 120.7 amu
Try these
problems:
1. HNO3
= 63.01 amu
2. C6H10O5
= 162.16 amu
3. H2SO4
= 98.08 amu
Mole
•A counting unit
•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 in the same
volume
Mole – 6.02 x 1023
particles
1 mole C
= 6.02 x 1023 C atoms
1 mole H2O
= 6.02 x 1023 H2O molecules
1 mole NaCl
= 6.02 x 1023 NaCl formula units
6.02 x 1023 Na+ ions and
6.02 x 1023 Cl– ions
Avogadro’s Number as Conversion Factor
Particles =
Moles
6.02 x 1023 particles
X
1 mole
Or
Moles =
Particles
X
1 mole
6.02 x 1023 particles
Note that a particle could be an atom OR a
molecule!
You MUST use dimensional analysis for conversions!
Examples:
How many molecules are in 3.5 moles of H2O?
How many moles are present in 465 molecules of NO2?
How many atoms of nitrogen are in 3.15 moles of NH3?
How many atoms of chlorine are in .862 moles of MgCl2?
Molar Mass
 Molar Mass- the mass of one mole of a
substance
 Unit is grams/mole
 Equivalent to the molecular mass in amu
 Ex: molar mass of Iron = 55.85 g /mole
molecular mass of Iron = 55.85 amu
Mass and Mole Relationships
Examples:
1. Find the number of moles present in 56.7 g of HNO3.
2. Find the number of grams present in 4.5 moles of C6H10O5.
3. Find the number of moles present in 12.31 g of H2SO4.
Percent Composition
•Finding what percent of the total weight of a
compound is made up of a particular element
Formula for calculating % composition:
Total amu of the element in the compound
Total formula amu
X 100%
Example:
 Calculate the % composition of BeO
Example:
Calculate the % composition Ca(OH)2
Example:
 Calculate the % composition of Al(NO3)2
Chemical Formulas
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).
1 molecule NO2 : 2 atoms of O for every 1 atom of
N
1 mole of NO2 : 2 moles of O atoms to every 1 mole
of N atoms
Law of Multiple Proportions
When any two elements, A and B, combine to
form more than one compound, the different
masses of B that unite with a fixed mass of A
bear a small whole-number ratio to each other
Example:
In H2O, the proportion of H:O = 2:16 or 1:8
In H2O2, H:O is 2:32 or 1:16
Empirical Formula - The formula of a compound that
expresses the smallest whole number ratio of the
atoms present.
Ionic formulas are always empirical formulas
Molecular Formula - The formula that states the
actual number of each kind of atom found in one
molecule of the compound.
Determine the Empirical Formula
From the Molecular Formula
 Reduce!!
 C6H6
 Fe3(CO)9
 BaCl2
 P4O10
Determine the Molecular Formula
from the Empirical Formula
 Calculate the molar mass of the Empirical
Formula.
 Divide the molar mass of the Molecular
Formula by the molar mass of the Empirical
Formula
 Multiply the numbers of each type of atom by
that number
Determine the Molecular Formula
from the Empirical Formula
 Examples:
 Molecular Formula: 26.04 g/mol
 Empirical Formula: CH
 Molecular Formula: 380.88 g/mol
 Empirical Formula: SeO3
To Obtain Empirical Formula
1. Assume the percent is out of 100
grams. That means you can change
the % sign to grams.
2. Calculate the number of moles of
each element.
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
**Remember
this**
Percent to mass
Mass to mole
Divide by small
Multiply 'til whole
Calculating Empirical Formula Example:
1. Given that a compound is composed of 60.0% Mg
and 40.0% O, find the empirical formula.
Calculating the Empirical Formula
Example #2:
 A compound is analyzed and is found to
contain 13.5g of calcium, 10.8g of oxygen,
and 0.675g of hydrogen. Calculate the
empirical formula of this compound.
Calculating the Empirical Formula
Example #3:
 NutraSweet is a zero calorie sweetener used
in many food products. A sample is analyzed
and it’s percent composition is as follows;
57.14% carbon, 6.16% hydrogen, 9.52%
nitrogen, and the rest is oxygen. Calculate
the empirical formula of NutraSweet.
Try this!
A compound is found to contain 68.5% carbon, 8.63%
hydrogen, and 22.8% oxygen. The molecular weight
of this compound is known to be approximately
140.00 g/mol. Find the empirical and molecular
formulas.