Review for Chapter 3: Mass Relationships in Chemical Reactions
1. The mass of an atom depends on the number of protons, neutrons, and electrons it contains. Atomic masses are measured in amu (atomic mass units):
1 amu is defined as 1/12 the mass of one carbon-12 atom
2. The atomic mass given in the periodic table for an element is a weighted average of the atomic masses for the naturally occurring stable isotopes of that element. That is why the atomic mass for carbon in the periodic table is 12.01 rather than 12.00 since carbon contains 98.9% of the C-12 isotope (12 amu) and 1.1% of the C-13 isotope (13 amu):
12.01 amu = 0.989 x 12.00 amu + 0.11 x 13.00 amu
3. The molecular mass of a molecule is calculated by summing the atomic masses of each of the atoms in the molecule.
For example, for H
2
O:
2 atoms hydrogen x 1.01 amu/atom hydrogen = 2.02 amu
1 atom oxygen x 16.00 amu/atom oxygen = 16.00 amu
18.02 amu per atom of H
2
0
4. A mole is 6.022 x 10 23 (Avogadro’s number) of atoms, molecules, or other particles.
5. The molar mass of a compound in grams is numerically equal to its mass in atomic mass units (amu). For example, for oxygen, the atomic mass is 16.00 amu and the molar mass is 16.00 g.
6. A mass spectrometer is a laboratory instrument that can be used to determine both atomic mass and molecular mass.
7. The molar mass of a molecule or ion is calculated by summing the masses for every mole of each element in the compound.
For example, for H
2
O:
2 moles hydrogen x 1.01 g/mole hydrogen = 2.02 g
1 mole oxygen x 16.00 g/mole oxygen = 16.00 g
18.02 g per mole of H
2
O

8. The percent composition by mass of a compound is the percent by mass of each element present.
For H
2
O:
% H: 2 moles H x 1.01 g/mole H x 100% = 11.2% H
1 mole H
2
O x 18.02 g/mole H
2
O
% O: 1 mole O x 16.00 g/mole O x 100% = 88.8% O
1 mole H
2
O x 18.02 g/mole H
2
O
9. If we know the percent composition by mass of a compound , we can deduce the empirical formula of the compound. If the approximate molar mass is known , we can also figure out the molecular formula of the compound.
10. Chemical reactions are represented by chemical equations . The reacting chemicals, called the “reactants”, are written on the left. The substances formed, called the “products”, are written to the right of the arrow. For example:
Chemical Equation: 2H
2
+ O
2

2H
2
O
Reactants: H
2
, O
2
Products: H
2
O
11. Chemical equations must be balanced so that the number of atoms of each element on the reactant side of the equation equals the number of atoms of each element on the product side of the equation.
12. Stoichiometry is the quantitative study of the reactants and products in a chemical reaction. We need to know how to convert from grams to moles of a reactant or product and from moles to grams for other reactants or products.
13. A limiting reagent is the reactant that is present in the smallest stoichiometric amount, which limits the amount of product that can be formed.
14. The amount of product that is obtained from a reaction is called the
“actual yield” and is often less than the maximum possible amount, which is called the
“theoretical yield”.
15. The percent yield of a reaction is calculated by:
% yield = actual yield x 100%
theoretical yield