Download Empirical Formula

Unit 10 – 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?
Molecular and Formula Masses
 The sum of the masses of all the atoms in a
compound
 Molecular Mass – mass in a molecule (covalent)
 Formula Mass – mass in a formula unit (ionic)
 Unit is amu for either of them
 Example: Formula Mass of CaCO3
 1 atom of Ca = 40.08 amu
=
 1 atom of C = 12.01 amu
=
 3 atoms of O = 3 x 16.00 amu =
40.08 amu
12.01 amu
48.00 amu
100.08 amu
Example:
Find the formula mass of (NH4)2SO3
2N
8H
1S
3O
= 2(14.01 amu)
= 8(1.01 amu)
= 32.07 amu
= 3(16.00 amu)
= 28.02 amu
= 8.08 amu
= 32.07 amu
= 48.00 amu
1 formula unit
= 116.17 amu
Try these problems:
1.HNO3
= 63.02 amu
2.C6H10O5
= 162.16 amu
3.Al3(PO4)2
= 270.88 amu
Mole
• A counting number (like a dozen)
• 6.02 X 1023 (in scientific notation)
• This number is named in honor
of Amedeo Avogadro (1776 –
1856)
• 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 (g/mole or g/mol)
 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
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 of Ca(OH)2
Example
Calculate the % composition of Al(NO3)3
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 vs. Molecular Formula
 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
All you need to do is reduce!!
1. C6H6
2. Fe3(CO)9
3. BaCl2
4. 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 Mass: 26.04 g/mol
 Empirical Formula: CH
 Molecular Formula: C2H2
 Molecular Mass: 380.88 g/mol
 Empirical Formula: SeO3
 Molecular Formula Se3O9
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 the 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.