1. Ch. 6 Notes -- Chemical Composition
What is a mole?

2. Ch 6 – Chemical Quantities
The Mole!!!
A counting unit
Similar to a dozen, except instead of 12, it’s
602 billion trillion…
(602,000,000,000,000,000,000,000)
___________ (in scientific notation)
This number is named in honor of Amedeo _________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present…6.02 x 1023
6.02 x 10 23
Avogadro

3. Just How Big is a Mole?
Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles.
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.

4. The Mole 1 dozen cookies = ___ cookies12
1 dozen cookies = ___ cookies
1 mole of cookies = ___________ cookies
1 dozen cars = ___ cars
1 mole of cars = __________ cars
1 dozen Al atoms = ___ Al atoms
1 mole of Al atoms = __________ atoms
Note that the NUMBER is always the same, but the ______ is very different!
Mole is abbreviated ______ .
6.02 X 1023 12
6.02 X 1023 12
6.02 X 1023
MASS
mol

5. The Mole and Mass
Mass in grams of 1 mole equal to __________ of the atomic masses.
Practice problem:
Calculate the mass of 1 mole of CaCl2.
Ca = 1 x ________ g/mol =
Cl = 2 x ________ g/mol =
= __________ g/mol CaCl2
1 mole of CaCl2 =
the sum
40.1
40.1 g/mol
35.5
71.0 g/mol
40.1 g/mol
71.0 g/mol
111.1
111.1 g/mol

6. Mole Conversion Factors that you will need to know!
1 mol = __________ atoms/molecules/etc.
1 mol = _____________ grams
1 mol = ________ Liters of gas at STP
(STP is Standard Temp. and Pressure… we will talk about what this means later!)
6.02 x 1023
? ( molar mass)
22.4

7. Ch. 6 Notes -- Chemical Composition
Practice Problems:
(1) How many atoms of hydrogen are there in each compound?
a) Ca(OH)2 ___ b) C3H8O___ c) (NH4)2HPO4 ___ d) HC2H3O2 ___
(2) Calculate the formula mass of each compound. (Add up all the
atomic masses for each atom from the Periodic Table.)
a) CaCO3 b) (NH4)2SO4
c) C3H6O d) Br2
e) H3PO4 f) N2O5 2 8 9 4
2 N’s = 2 x 14.0 = 28.0
8 H’s = 8 x 1.0 = 8.0
S = 32.1
4 O’s = 4 x 16.0 = 64.0
Ca = 40.1
C = 12.0
3 O’s =3 x 16.0 = 48.0
Add them up!
132.1 g/mol
Add them up!
100.1 g/mol
159.8 g/mol
C = 3 x 12.0 = 36.0
H = 6 x 1.0 = 6.0
O =16.0
2 Br’s = 2 x 79.9 =
Add them up!
58.0 g/mol
2 N’s = 2 x 14.0 = 28.0
5 O’s = 5 x 16.0 = 80.0
3 H’s = 3 x 1.0 = 3.0
P = 31.0
4 O’s = 4 x 16.0 =64.0
Add them up!
98.0 g/mol
Add them up!
108.0 g/mol

8. 3) Convert 835 grams of SO3 to moles.
4) How many molecules of CH4 are there in 18 moles?
5) How many grams of helium are there in 5.6 x 1023 atoms of helium?
6) How many molecules are there in 3.7 grams of H2O?
1 mole SO3
835 g SO3
10.4 moles of SO3 x =
80.1 g SO3
6.02 x 1023 molecules CH4
18 moles CH4 x =
1.08 x 1025 molecules CH4
1 mole CH4
4.0 grams He
5.6 x 1023 atoms He x
3.72 grams He =
6.02 x 1023 atoms He
6.02 x 1023 molecules H2O
3.7 grams H2O
1.24 x 1023 molecules H2O x =
18.0 grams H2O

9. Calculating Percent Composition by Mass
Step 1: Find the formula mass of the compound by adding the individual masses of the elements together.
Step 2: Divide each of the individual masses of the elements by the formula mass of the compound.
Step 3: Convert the decimal to a % by multiplying by 100.
Practice Problems:
(1) Find the % composition of the elements in each compound.
a) Na3PO b) SnCl4
= = 42.1%
3 Na’s = 3 x 23.0 = 69.0
P = 31.0
4 O’s = 4 x 16.0 =
÷ 164
Sn = 118.7
4 Cl’s = 4 x 35.5 =
÷ 260.7
= 45.5%
= = 18.9%
÷ 164 +
÷ 260.7
= 54.5%
260.7
= = 39.0% +
÷ 164
164

10. Elements in the Universe: % Composition by Mass

11. Earth’s Crust: % Composition by Mass

12. Entire Earth (Including Atmosphere): % Composition by Mass

13. Human Body: % Composition by Mass

14. Determining the Empirical Formula for a Compound
The empirical formula for a compound is the simplest __________ number __________ of the atoms in the compound.
Examples: H2O is the empirical formula for water.
_______ is the empirical formula for glucose, C6H12O6.
Practice Problems: What is the empirical formula for the following compounds? a) C6H6= ________
b) C8H14O2 = ________
c) C10H14O2 = _________
d) Ca5Br10 = ________
e) N3O9 = ________
whole
ratio
C1H2O1 CH
C4H7O
C5H7O
CaBr2
NO3

15. Determining the Molecular Formula for a Compound
The molecular formula for a compound is either the same as the empirical formula ratio or it is a “_________ _________ of this ratio. It represents the true # of atoms in the molecule.
Examples: 1) H2O is the empirical & molecular formula for water ) CH2O is the empirical formula for sugar, ethanoic acid, and methanol. The molecular formula for glucose is C6H12O6, (___times the empirical ratio!)
Practice Problems: (1) If the empirical formula for a compound is CH2, which of the following is a possible molecular formula for the compound? a) C8H16 b) C8H8 c) C4H2 d) C3H9
(2) If the empirical formula for a compound is C2H3, which of the following is a possible molecular formula for the compound? a) C2H6 b) C10H c) C6H12 d) C8H14
whole # multiple 6

16. Determining the Molecular Formula for a Compound
Find the molecular formula for C2H7 if the molecular mass of the compound is 93.0 g/mol.
Find the molecular formula for P2O5 if the molecular mass of the compound is g/mol.
C2H7 = 31.0 g/mol
93.0 g/mol
= 3
31.0 g/mol
C2H7 x 3
= C6H21
g/mol
P2O5 = g/mol
= 2
g/mol
P2O5 x 2
= P4O10