1. Chapter 6 Chemical Composition 6.1 Counting by Weighing Bean Lab
We can count individual units by weighing if we know the average mass of the units.
Instead of counting out 1000 jelly beans, it is easier to (1.) find the mass of 1 bean (= 5 g), (2.) multiply x (=5000 g), and (3.) measure out 5000 g of beans.
Chapter 6 Chemical Composition
6.1 Counting by Weighing

2. 6.2 Atomic Masses: Counting Atoms by Weighing Pennium Lab
When we know the average mass of the atoms of an element, we can calculate the number of atoms in any given sample of that element by weighing the sample.
1 amu = x10-24 g
Using Table 6.1, p. 157, calculate the mass, in amu of a sample of aluminum that contains 75 atoms.
1 Al atom = amu
mass of 75 atoms = 75 atoms x amu = amu
1 atom
Calculate the number of sodium atoms present in a sample that has a mass of amu, if 1 Na atom = amu.
amu x 1 Na atom = Na atoms
22.99 amu
6.2 Atomic Masses: Counting Atoms by Weighing

3. Conversions: (p. 164) 5.00 x 1020 atoms Cr Determine # of moles:
6.3 The Mole
1 dozen = 12
1 pair = 2
1 score = 20
1 mole = x 1023 objects
(a counting number, an exact number; infinite # of sig figs) aka Avogadro’s Number (p. 160)
A sample of an element with a mass equal to that element’s average atomic mass expressed in grams contains 1 mole of atoms.
Example: see Table 6.2, p. 160
Conversions: (p. 164) 5.00 x 1020 atoms Cr
Determine # of moles:
Determine mass, in grams:
6.3 The Mole

4. (use Avogadro’s #) (use 22.4 L)
Mole Map
Mass, in grams
(use molar mass)
Mole
(use Avogadro’s #) (use 22.4 L)
Number of Atoms Volume, in Liters
(of a gas at STP, standard temperature, 0oC, and standard pressure, 1 atm) (Ch. 13 Gases)
Mole Map

5. For ex., NaCl molar mass = _________ CaCl2 molar mass = _________
Ionic compounds- calculate the mass of 1 formula unit of the compound (compound formula as written).
For ex., NaCl molar mass = _________
CaCl2 molar mass = _________
For covalent compounds- calculate the mass of the molecule (compound formula as written). For ex., (p. 166, ex. 6.5)
SO2 molar mass = g
C2H3Cl molar mass = _________
CuSO4 5H2O molar mass = _________
Conversions: mass from moles, p. 168
Moles from mass, p. 168
6.4 Molar Mass

6. number of molecules (or formula units, if ionic) from mass, p. 169
Mole Map

7. 6.5 Percent Composition of Compounds
Percent (by mass) for a given element= mass of element x 100%
mass of 1 mol of compound
Example: MgCO3 molar mass = ____________
Example: penicillin, p. 173
6.5 % Composition of Compounds

8. Ex.: C6H12O6 is molecular formula; CH2O is its empirical formula.
6.6 Formulas of Compounds
Empirical formula- lowest whole-number ratio of all elements in a compound formula. For ionic compounds, ALL formulas are empirical formulas.
Molecular formula- for covalent compounds, the actual number of atoms of each type of element present in the compound.
Ex.: C6H12O6 is molecular formula; CH2O is its empirical formula.
6.6 Formulas of Compounds

9. 6.7 Calculation of Empirical Formulas
How to find:
Find actual # of moles
Find relative # of moles (must be in whole numbers); divide actual # of moles of each element by the smallest. Sometimes you must then multiply by some integer to get everything in whole numbers (integers) representing a ratio. These integers are your subscripts.
Example: p. 180
Example: p. 181
Example: p. 183 (using % composition)(assume 100 grams)
6.7 Calculation of Empirical Formulas

10. 6.8 Calculation of Molecular Formulas
(all steps are the same except for 1 additional step.) Check the molar mass of the empirical formula; if it is the same as the molecular mass given in the problem, then you are done. If not, you must find by what factor is it different. Then multiply the subscripts in the empirical formula by that factor.
Example: p. 185
6.8 Calculation of Molecular Formulas