1. Seating Chart

2. Seating Chart

3. Seating Chart

4. Imagine your boss wants you to count the number of Styrofoam packing peanuts in a large bag. Suppose you found the mass of 100 peanuts to be 5.5 g. How could you use this to find the number of peanuts in the bag?
Warm Up 1/5-6

5. Unit 5
Counting Particles

6. Relative mass compares the mass of the same number of particles of different elements.
The periodic table gives the relative mass of each element in atomic mass units (amu’s).
Relative Mass

7. Moles A mole is a counting number. Also called Avogadro’s number (NA)
1 mole = 6.02 x 1023
602,000,000,000,000,0 00,000,000
Moles

8. Molar Mass One mole of an element = the atomic mass in grams.
C = g/mol
One mole of a compound = the sum of the atomic masses of each atom present.
Sodium Oxide: Na2O
Na = 23.0 amu
O = 16.0 amu
Na2O = 2(23.0g) + 1(16.0g) = 62.0g/mol
Molar Mass

9. Molar Mass Examples Ex: How many moles are in 200g of copper?
200𝑔 𝐶𝑢 ( 𝑚𝑜𝑙 𝐶𝑢 63.5𝑔 )=3.15 𝑚𝑜𝑙 𝐶𝑢
Molar Mass Examples

10. Ex: What is the mass of 5.50 moles of sodium chloride (NaCl)?
NaCl = 23.0g g = 58.5g/mol
5.50 𝑚𝑜𝑙 𝑁𝑎𝐶𝑙 ( 58.5 𝑔 𝑁𝑎𝐶𝑙 𝑚𝑜𝑙 )=322𝑔 𝑁𝑎𝐶𝑙
Molar Mass Examples

11. Ex: How many moles are in 2.75g of gold? How many atoms is this?
2.75𝑔 𝐴𝑢 𝑚𝑜𝑙 𝐴𝑢 𝑔 = 𝑚𝑜𝑙 𝐴𝑢
𝑚𝑜𝑙 𝐴𝑢 ( 6.02 𝑥 𝑎𝑡𝑜𝑚𝑠 𝐴𝑢 𝑚𝑜𝑙 𝐴𝑢 )=8.40 x 1021 𝑎𝑡𝑜𝑚𝑠 𝐴𝑢
Molar Mass Examples

12. Empirical formula (EF) represents the simplest whole number ratio of elements in a compound.
Molecular formula (MF) represents the actual number of atoms of each element in a compound.
Empirical Formula
Molecular Formula
CO2
CH2O
C6H12O6 HO
H2O2
Empirical Formulas

13. Determining Empirical Formulas
Ex: An unknown solid is determined to be 85% Ag and 15% F. Find the empirical formula.
1. Assume a 100g sample
2. Find moles of each
3. Write ratio and divide by least amount
4. Get whole number ratio
85𝑔 𝐴𝑔 𝑚𝑜𝑙 𝐴𝑔 𝑔 =0.788 𝑚𝑜𝑙 𝐴𝑔
15𝑔 𝐹 𝑚𝑜𝑙 𝐹 19.0𝑔 =0.789 𝑚𝑜𝑙 𝐹
Ag F = AgF
Determining Empirical Formulas

14. Determining Empirical Formulas
Ex: An unknown solid is found to be 37.5g C and 12.5g H. Find the empirical formula.
1. Find moles of each
37.5𝑔 𝐶 𝑚𝑜𝑙 𝐶 12.0 𝑔 =3.13 𝑚𝑜𝑙 𝐶
12.5𝑔 𝐻 𝑚𝑜𝑙 𝐻 1.0𝑔 =12.5 𝑚𝑜𝑙 𝐻
2. Write ratio and divide by least amount
C 3.13 H 12.5
3. Get whole number ratio
CH4
Determining Empirical Formulas

15. Determining Empirical Formulas
Ex: A sample of aluminum sulfate is found to be 15.8g Al, 28.1g S and 56.1g O. Find the empirical formula.
1. Find moles of each
15.8𝑔 𝐴𝑙 𝑚𝑜𝑙 𝐴𝑙 27.0 𝑔 =0.585 𝑚𝑜𝑙 𝐴𝑙
28.1𝑔 𝑆 𝑚𝑜𝑙 𝑆 32.1𝑔 =0.875 𝑚𝑜𝑙 𝑆
56.1𝑔 𝑂 𝑚𝑜𝑙 𝑂 16.0𝑔 =3.51 𝑚𝑜𝑙 𝑂
2. Write ratio and divide by least amount
Al S O 3.51
3. Get whole number ratio
Al1S1.5O6 x 2 = Al2S3O12 = Al2(SO4)3
Determining Empirical Formulas

16. Percentage Composition
Percent composition is the percentage by mass of each element in a compound.
% 𝑐𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛= #𝑎𝑡𝑜𝑚𝑠 𝑥 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 𝑥 100%
Percentage Composition

17. Finding Percent Composition
Sucrose C12H22O11
𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠= 𝑎𝑚𝑢 𝑎𝑚𝑢 𝑎𝑚𝑢 =342.0 𝑎𝑚𝑢
%𝐶= 12 𝑥 12.0 𝑎𝑚𝑢 𝑎𝑚𝑢 𝑥 100%=42.1% 𝐶
%𝐻= 22 𝑥 1.0 𝑎𝑚𝑢 𝑎𝑚𝑢 𝑥 100%=6.4% 𝐻
%𝑂= 11 𝑥 16.0 𝑎𝑚𝑢 342.𝑜 𝑎𝑚𝑢 𝑥 100%=51.5% 𝑂
Finding Percent Composition

18. Finding Molecular Formulas
Formula Mass (FM)
Molecular Mass (MM)
Water H20
18.0u
Acetic acid C2H4O2
30.0u
60.0u
Glucose C6H12O6
180.0u
Molecular formula (MF) is a multiple of the empirical formula (EF).
𝐸𝐹 𝑥=𝑀𝐹
𝑥= 𝑀𝑀 𝐹𝑀
Finding Molecular Formulas

19. Finding Molecular Formulas
Ex: Find the molecular formula for a compound that contains 5.88g hydrogen and 94.1g oxygen. The compound has a molecular mass of amu.
1. Find EF 𝑔 𝐻 𝑚𝑜𝑙 𝐻 1.0𝑔 =5.88 𝑚𝑜𝑙 𝐻
94.1𝑔 𝑂 𝑚𝑜𝑙 𝑂 16.0𝑔 =5.88 𝑚𝑜𝑙 𝑂
H 5.88 O = HO
2. Find formula mass
FMHO = 1(1.0) + 1(16.0) = 17.0 amu
3. Molecular mass multiple of FM
𝑥= 𝑀𝑀 𝐹𝑀 = 34.0 𝑎𝑚𝑢 17.0 𝑎𝑚𝑢 =2
4. Find the molecular formula
(𝐸𝐹) 𝑥 =𝑀𝐹
(𝐻𝑂) 2 = 𝐻 2 𝑂 2
Finding Molecular Formulas