1. Unit 5
Chemical Composition

2. Counting by Mass Individually mass a few
Calculate the average mass of one
Can count large numbers of by mass

3. Atomic Mass Unit (amu)
1 amu = 1.66 x g

4. Electrical charge (coulomb)
Subatomic particles
Particle
amu
mass (g)
Atomic charge
Electrical charge (coulomb)
Proton
(p+)
1.0073
1.673 x 10-24 +1
x 10-19
Electron
(e-)
9.109 x 10-28 -1
x 10-19
Neutron (n)
1.0087
1.675 x 10-24

5. C Periodic Table 6 Atomic mass 12.0111 Element symbol Atomic number
Elements with the atomic mass in ( ) means that very little of it exists at any given time. Scientists have not been able to attain an good average atomic mass.

6. When given masses and asked to find the
% Composition by Mass
When given masses and asked to find the
% compositions
% Composition of element by mass
grams of element grams of compound ) = (
100

7. In a 15. 0 gram sample of H2O, there is 1. 68 g H, and 13. 32 g O
In a 15.0 gram sample of H2O, there is 1.68 g H, and g O. Find the % composition of each element.

8. 12 grams of carbon monoxide are analyzed and found to contain 5
12 grams of carbon monoxide are analyzed and found to contain 5.14 grams of carbon and 6.86 grams of oxygen.

9. (particles: atoms, molecules, ions, electrons, etc.)
Because atoms and molecules are too small to work with, chemists have defined a large unit called a mole. The mole is defined as the number of atoms in 12 grams of Carbon x 1023 particles.
Avogadro’s number = 6.02 x 1023
(particles: atoms, molecules, ions, electrons, etc.)

10. Molar mass of an atom, is the atomic mass in grams which is 6
Molar mass of an atom, is the atomic mass in grams which is 6.02 x 1023 atoms.
What is the molar mass of iron, of sulfur.
Molar mass of a compound is the sum of the atomic mass of all the atoms in the compound in grams which is 6.02 x 1023 molecules.
Calculate the molar mass of hydrochloric acid (HCl)
Calculate the molar mass of silver nitrate (AgNO3)

11. Molar mass of water = 18.0 g/mol
Therefore, one mole of water is 18 grams of water.

12. moles = __grams_ Number of moles =
Calculate the number of moles in grams of Vanadium. Calculate the number of moles in grams of calcium carbonate (CaCO3).
moles = __grams_
1 molar mass
Number of moles =

13. Calculate the number of grams in: 0. 65 moles of barium iodide, BaI2 4
Calculate the number of grams in: moles of barium iodide, BaI moles of copper nitrate, Cu(NO3)2

15. % composition of element by molar mass
Find the molar mass of the compound
Divide the molar mass supplied by one element, by the molar mass of the compound. Multiply by 100 to obtain percentage
Repeat step 2 for each element
% composition of
element by molar mass (
total molar mass of element
total molar mass of compound ) =
100

16. Calculate the % composition of each element in sodium bicarbonate (NaHCO3)

17. Mole to mole ratio in reactions
Copper sulfate & iron yields copper & iron sulfate
CuSO4 (aq) + Fe (s)  Cu (s) + FeSO4 (aq)
Penny Lab: HCl(aq) + ___  __Cl2(aq) + H2 (g)

18. Particles = atoms, or molecules, or ions, or electrons, etc.
Number of particles =
(moles)(6.02 x 1023)
Particles = atoms, or molecules, or ions, or electrons, etc.

20. ( ) = % composition water
A hydrate is a molecule with water attached to the crystalline structure. Example: BaCl2.2H2O An anhydrous molecule has NO water attached.
% composition
water
Calculate the % water in BaCl2.2H2O (
Molar mass of all water
Molar mass of hydrate ) =
100

21. Determine the % composition of H2O in the hydrated compound CuSO4.5H2O.

22. Hydrate Lab
Determine the % composition of H2O for one of the samples.
(Teacher will assign the sample)
Bonus point for determining the number of H2O molecules attached to each molecule.
? = moles of water
moles of anhydrous

23. Example: C6H12O6 (glucose sugar)
Molecular formula is the actual formula that gives the exact composition.
Example: C6H12O6 (glucose sugar)
Empirical formula is the smallest whole number ratio of the atoms present.
Example: CH2O (empirical formula for glucose)
If multiplied by 6, it is the formula for glucose.

24. Steps for establishing formula
1. List mass of elements (if %, use mass in 100 grams)
2. Convert mass of element to moles
3. Write a tentative formula using mole values as subscripts
4. Attempt to convert subscripts to smallest whole numbers by dividing mole values by the smallest mole value
5. If decimal numbers, multiply by a small whole number to obtain empirical formula
6. To find molecular formula, compare empirical formula mass to molar mass of compound. Multiply subscripts by multiplication factor.
Write on your formula sheet!

25. A compound contains 4. 151 g if aluminum and 3. 692 g of oxygen
A compound contains g if aluminum and g of oxygen. Determine the empirical formula

26. A compound having an approximate molar mass of 168 g/mol has the following percentage composition by mass; carbon = 42.87%, hydrogen = 3.598%, oxygen = %, and nitrogen = %. Determine the empirical formula. Determine the molecular formula.

27. Review
a % Na
45.29% O
11.33 % C
b % N
24.27 % S
6.10 % H
48.44 % O
a % Ag
8.16 % N
28.32 % O
51.16 % H2O
(Empirical formula problem)
a g/mol
b g/mol
c g/mol
a mole SO2
b mole NaBr
c mole Ba(NO3)2
a g HNO3
b g CO2
c g Ca(C2H3O2)2
a x 1023 molecules
b x 1024 molecules
c x 1022 molecules
d x 1021 molecules

28. Mole to mole ratio in reactions
Copper sulfate & zinc yields copper & zinc sulfate
CuSO4 (aq) + Zn (s)  Cu (s) + ZnSO4 (aq)