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 1000 (=5000 g), and (3.)
measure out 5000 g of beans.
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 = 1.66 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 = 26.98 amu
mass of 75 atoms = 75 atoms x 26.98 amu = 2024 amu
1 atom
Calculate the number of sodium atoms present in a sample that has a
mass of 1172.49 amu, if 1 Na atom = 22.99 amu.
1172.49 amu x 1 Na atom = 51.00 Na atoms
22.99 amu
6.3 The Mole
1 dozen = 12
1 pair = 2
1 score = 20
1 mole = 6.02 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:
Mole Map
Mass, in grams
(use molar mass)
Mole
(use Avogadro’s #)
Number of Atoms
(use 22.4 L)
Volume, in Liters
(of a gas at STP, standard
temperature, 0oC, and
standard pressure, 1 atm)
(Ch. 13 Gases)
6.4 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 = 64.07 g
C2H3Cl
molar mass = _________
CuSO4 5H2O molar mass = _________
Conversions: mass from moles, p. 168
Moles from mass, p. 168
number of molecules (or formula units, if ionic) from mass, p. 169
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
6.6 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.
6.7 Calculation of Empirical Formulas
How to find:
1. Find actual # of moles
2. 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)
6.8 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