1. Unit 3 Section 1 Chemical Composition
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2. Goals of Unit Moles and Avogadro's number Molar Mass
Convert between moles and mass
Mass percent of elements in compounds
Empirical formulas
Calculating molecular formulas

3. One of most important chemical activities: Synthesis of new substances
Nylon, aspartame, Kevlar (bulletproof vests), PVC, Teflon
All originated in chemist's laboratory
Once they make it – they must determine what it is
What is it's composition?
What is it's chemical formula?

4. What Is a Mole?
Recall that matter is composed of atoms, molecules, and ions.
These particles are much, much smaller than grains of sand, and an extremely large number of them are in a small sample of a substance.
Obviously, counting particles one by one is not practical.
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5. We usually are dealing with three types of representative particles
Molecules
H2O and H2 are molecules, covalently bonded
Atoms
Al and Na are atoms, not bonded
Formula units
CaCl2 and NaOH are formula units, ionicly bonded

6. Converting Grams Moles
Steps:
Find the molar mass of the compound you will be converting to moles.
Use the molar mass to convert from grams to moles and vice versa.

7. Converting Grams Moles
Ex. 1: Convert 7.5 grams of NaCl to moles.
Step 1 - Find molar mass of NaCl
Step 2 - Convert to moles.
MM = = g NaCl
1 mol
7.5 g NaCl
1 mol NaCl
58.44 g NaCl
0.13 mol NaCl x =

8. Converting Grams Moles
Ex. 2: Convert 4 moles of BaF2 to grams.
Ex. 3: Convert 75.5 grams of Fe(NO2)2 to moles.

10. Since atoms are so small, we need a quantity with a larger amount
We use the mole (mol)
1 mole = 6.02 X 1023 representative particles (atoms, molecules, formula units)
It is known as Avogadro’s number
Amedeo Avogadro

11. 1. How many particles can be found in 10 grams of neon?
10 g Ne
1 mol
20.18 g Ne
6.02 x 1023 atoms Ne
1 mol
X X =
= 2.98 x 1023 Ne atoms

12. 2) If you had a bottle that contained 5
2) If you had a bottle that contained 5.69x1024 particles of water, how many grams of water does the bottle hold?
5.69 x 1024 molecules H2O
1 mol
6.02 x 1023 molecules H2O
18.02 g H2O
1 mol
X X =
= g H2O

13. 3. How many particles are in 50 g of Ca(OH)2?
50 g Ca(OH)2 1 mol 6.02 x 1023 f.u. Ca(OH) g Ca(OH)2 1 mol
X X =
= 4.06 x 1023 f.u. Ca(OH)2

14. Key Concepts
The mole allows chemists to count the number of representative particles in a substance.
The atomic mass of an element expressed in grams is the mass of a mole of the element.
To calculate the molar mass of a compound, find the number of grams of each element in one mole of the compound. Then add the masses of the elements in the compound.
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15. 9.5 Percent Composition
Percent composition of a compound is the relative amount of each element in the compound.

16. Which of the following compounds contains the highest percent of Iron?
Iron III Acetate, Iron II Hydroxide or Iron II nitride.

17. Hydrates
Sometimes, percent composition can include compounds called hydrates
Hydrates are compounds that bind water molecules to their structure
BaCl2 • 6H2O Name -> barium chloride
hexahydrate
MgSO4 • 3H2O Name ->
6 moles of water
3 moles of water

18. Hydrates Find the percentage of water in barium chloride hexahydrate
H2O
108.12g/316.35g x = %H2O

19. Find the percent composition of water in magnesium sulfate heptahydrate.

20. Find the percent composition of water in magnesium sulfate heptahydrate.
MgSO4  7H2O MM = (16) + 7(18.02)
= g/mol
%H2O = 7(18.02 g/mol)
246.51g/mol
= % H2O
x 100

21. Do Now = 29.26 % H2O %H2O = 4(18.02 g/mol) x 100 246.34 g/mol
Find the percent composition of water in potassium sulfate tetrahydrate.
K2SO4  4H2O MM =2(39.10) (16) + 4(18.02)
= g/mol
%H2O = 4(18.02 g/mol) x 100
g/mol
= % H2O

22. Empirical formula
Empirical formula is the simplest whole number ratio between elements in a compound

23. Example 1: You have 26.56% K, 35.41% Cr, and 38.03% O .
You are given % of an element. Change the % to grams (the number does not change).
Example 1: You have 26.56% K, 35.41% Cr, and % O .
Therefore, you have: g K, g Cr, and
Now change Grams to Moles

24. Divide each molar quantity by the smaller number of moles to get 1 mol for the element with the smaller number of moles.
Multiply each part of the ratio by the smallest whole number that will convert both subscripts to whole numbers.

25. Guidelines for Empirical Formula
1. Change % to grams 2. Change grams to moles 3. Divide all moles by the lowest number of moles 4. If all the numbers are not whole numbers, you must multiply the lowest number possible to get a whole number.

26. Example 2: 63.52% Fe and 36.48% S Find the Empirical Formula.

27. Answer: K2CO3 Calculate the empirical formula with the following:
56.4% potassium, 8.7% carbon, 34.9% oxygen
Answer: K2CO3

28. 9.8 Molecular Formula
Molecular Formula – A multiple of an empirical formula.
Can also be the same as empirical in some cases

29. Example: Empirical formula N2O – 2:1 ratio
Molecular Formula of N2O could be:
N4O2 – A multiple of (N2O) = N4O2
N6O3 – A multiple of (N2O) = N6O3
How to determine the multiple you need:
Molecular Formula Mass = Multiple
Empirical Formula Mass
Then multiply the multiple to the Empirical Formula

30. Example 1: Calculate the molecular formula of a compound whose molar mass is 60.0 g/mol and empirical formula is CH4N.
First calculate the empirical formula mass.

31. Use the formula Molecular Formula Mass = Multiple
Empirical Formula Mass
Multiply the formula subscripts by this multiple to determine the molecular formula.

32. Example 2
What is the molecular formula of a compound with the empirical formula CClN and a molar mass of g/mol.

33. Find the molecular formula of ethylene glycol, which is used as antifreeze. The molar mass is g/mole, and the empirical formula is CH3O