Chapter 10
10.1

You could measure the
amount of sand in a sand
sculpture by counting each
grain of sand, but it would
be much easier to weigh
the sand. You’ll discover
how chemists measure the
amount of a substance
using a unit called a mole,
which relates the number
of particles to the mass.
10.1
10.1
 A mole (mol) of any substance contains a lot of
particles; Avogadro’s number of particles (6.02  1023).
10.1
 So a mole is a counted quantity, just like a dozen, only
a little bigger!
10.1
◦ The mass of one mole of a substance is called the
molar mass.
 The mass of one mole of an element is the atomic
mass measured out in grams.
This quantity has 6.02 x 1023 atoms of the element.
10.1
 The molar mass of carbon, sulfur, mercury, and iron
are shown below.
10.1
 The mass of one mole of a compound is the formula
mass measured out in grams.
This quantity has 6.02 x 1023 formula units or molecules of
the compound.
10.1
1 mole of SO3 has a mass of 80.1 g.
10.1
8. What is the mass of 1.00 mole of sodium
hydrogen carbonate? What is the mass of
2.00 moles?
10.2
◦ Use the molar mass to convert between the mass of
a substance and the moles of a substance.
 Ref. table T has the equation to use:
17. Calculate the mass, in grams, of 2.50
moles of iron (II) hydroxide.
19. Calculate the number of moles in 75.0 g
of dinitrogen trioxide
10.3

It helps to know the percents
of the components in a shirt
because they affect how warm
it is, whether it will need to be
ironed, and how it should be
cleaned. You will learn how
the percents of the elements
in a compound are important
in chemistry.
10.3
 The percent by mass of an element in a compound is
the number of grams of the element divided by the
mass in grams of the compound, multiplied by 100%.
10.3
◦ Reference table T has the equation to find the
percent mass of an element in a compound:
10.3
◦ If mass data isn’t given use the formula and molar
mass to find percent composition.
34. Calculate the percent composition of these
compounds.
a.Ethane (C2H6)
b.Sodium hydrogen sulfate (NaHSO4)
10.3
◦ The percent composition can be used as a
conversion factor to calculate the number of grams
of any element in a specific mass of a compound.
10.3

Propane (C3H8) is 82% carbon and 18%
hydrogen. Find the mass of carbon and
hydrogen in an 90.0 g sample of C3H8.

Water is 11% hydrogen and 89% oxygen by
mass. What is the mass of hydrogen and
oxygen in a 56.0 g sample of water?
10.3

The empirical formula gives the lowest
whole-number ratio of the atoms/ions of the
elements in a compound.
 Ionic formulas are always empirical; molecular
formulas may or may not not be.
10.3
 Ethyne (C2H2) is a gas used in
welder’s torches. Styrene (C8H8) is
used in making polystyrene.
 These two compounds of carbon
have the same empirical formula
but different molecular formulas.

What is the empirical formula of the
following:
◦
◦
◦
◦
◦
◦
C6H12O6
H2O2
NaCl
C2H4O2
CH4
C6H6
10.3

The molecular formula of a compound is
either the same as its experimentally
determined empirical formula, or it is a
simple whole-number multiple of its
empirical formula.
10.3
Methanal, ethanoic acid,
and glucose all have the
same empirical
formula—CH2O.
10.3

To find the molecular formula from the
empirical formula:
◦ Determine how many times more massive the
molecular formula is than the empirical formula
◦ Multiply the subscripts of the empirical formula by
this small whole number