Page 51 - Theory and Problems of BEGINNING CHEMISTRY

40 ATOMS AND ATOMIC MASSES [CHAP. 3

A great deal of difficulty was encountered at first, because Dalton’s fifth postulate gave an incorrect ratio
of numbers of atoms in many cases. Such a large number of incorrect results were obtained that it soon became
apparent that the fifth postulate was not correct. It was not until some 50 years later than an experimental method
was devised to determine the atomic ratios in compounds, at which time the scale of relative atomic masses was
determined in almost the present form. These relative masses are called atomic masses, or sometimes atomic
weights.
Atomic masses are so small that an appropriate unit was developed to report them—an atomic mass unit
(amu). 1 amu = 1.66 × 10 −24 g. The atomic mass of the lightest element, hydrogen, was originally taken to
be 1 amu. The modern values of the atomic masses are based on the most common kind of carbon atom, called
“carbon-12” and written 12 C, as the standard. The mass of 12 C is measured in the modern mass spectrometer,
12
and Cis defined to have an atomic mass of exactly 12 amu. On this scale hydrogen has an atomic mass of 1.008
amu.
The atomic mass of an element is the relative mass of an average atom of the element compared with 12 C,
which by definition has a mass of exactly 12 amu. Thus, since a sulfur atom has a mass 8/3 times that of a carbon
atom, the atomic mass of sulfur is
8
12 amu × = 32 amu
3
A complete list of the modern values of the atomic masses of the elements is given in the Table of Elements,
page 349, and in the periodic table, page 350.

3.4. ATOMIC STRUCTURE
From 50 to 100 years after Dalton proposed his theory, various discoveries were made that show that the
atom is not indivisible, but really is composed of parts. Natural radioactivity and the interaction of electricity with
matter are two different types of evidence for this subatomic structure. The most important subatomic particles
are listed in Table 3-1, along with their most important properties. The protons and neutrons are found in a very
tiny nucleus (plural, nuclei). The electrons are found outside the nucleus. The information in Table 3-1 must be
memorized, but only the whole number part of the masses (1 amu, 1 amu, 0 amu) must be remembered.

Table 3-1 Subatomic Particles
Charge (e) Mass (amu) Location

Proton +1 1.00728 In nucleus
Neutron 0 1.00894 In nucleus
Electron −1 0.0005414 Outside nucleus

There are two types of electric charges that occur in nature—positive and negative. Charges of these two
types are opposite one another, and cancel the effect of the other. Bodies with opposite charge types attract one
another; those with the same charge type repel one another. If a body has equal numbers of charges of the two
types, it has no net charge and is said to be neutral. The charge on the electron is a fundamental unit of electric
charge (equal to 1.6 × 10 −19 Coulomb) and is given the symbol e.

EXAMPLE 3.4. Using the data of Table 3-1, find the charge on a nucleus that contains (a) 8 protons and 8 neutrons and
(b) 8 protons and 10 neutrons.
Ans. (a)8(+1) + 8(0) =+8(b)8(+1) + 10(0) =+8
Both nuclei have the same charge. Although the nuclei have different numbers of neutrons, the neutrons have no
charges, so they do not affect the charge on the nucleus.

EXAMPLE 3.5. To the nearest integer, calculate the mass (in amu) of a nucleus that contains (a) 17 protons and 18 neutrons
and (b) 17 protons and 20 neutrons.