Chemistry Reference
In-Depth Information
3.2
Dalton's Atomic Theory
In 1803, John Dalton (1766-1844) (Figure 3.3) proposed his atomic theory,
including the following
postulates,
to explain the laws of chemical combina-
tion discussed in Section 3.1:
1. Matter is made up of very tiny, indivisible particles called
atoms.
2. The atoms of each element all have the same mass, but the mass of the
atoms of one element is different from the mass of the atoms of every other
element.
3. Atoms combine to form
molecules.
When they do so, they combine in
small, whole-number ratios.
4. Atoms of some pairs of elements can combine with each other in different
small, whole-number ratios to form different compounds.
5. If atoms of two elements can combine to form more than one compound,
the most stable compound has the atoms in a 1 : 1 ratio. (This postulate was
quickly shown to be incorrect.)
The first three postulates have had to be amended, and the fifth was quickly
abandoned altogether. But the postulates explained the laws of chemical com-
bination known at the time, and they caused great activity among chemists,
which led to more generalizations and further advances in chemistry.
The postulates of
Dalton's atomic theory
explained the laws of chemical
combination very readily.
John Dalton
Figure 3.3
Constant
mass of
oxygen
N
O
1. The law of conservation of mass is explained as follows: Because atoms
merely exchange “partners” during a chemical reaction and are not created
or destroyed, their mass is also neither created nor destroyed. Thus, mass
is conserved during a chemical reaction.
2. The law of definite proportions is explained as follows: Because
atoms
react
in definite integral ratios (postulate 3), and atoms of each element have a
definite
mass
(postulate 2), the
mass
ratio of one element to the other(s)
must also be definite.
3. The law of multiple proportions is explained as follows: Because atoms
combine in different ratios of small whole numbers (postulate 4), for
a given number of atoms of one element, the number of atoms of the
other element is in a small, whole-number ratio. A given number of
atoms of the first element implies a given mass of that element, and a
small, whole-number ratio for the atoms of the second element (each
of the same mass) implies a small, whole-number ratio of masses of
the second element (Figure 3.4). For example, consider water (H
2
O)
and hydrogen peroxide (H
2
O
2
), two compounds of hydrogen and oxy-
gen. For a given number of hydrogen atoms (2), the numbers of oxygen
atoms in the two compounds are 1 and 2. Stated another way, for a given
mass of hydrogen (2.0160 g), the ratio of masses of oxygen in the two
compounds is 15.9994 g to 31.9988 g, a ratio of 1 to 2—a small, whole-
number ratio.
We will discuss the ways in which the first three of Dalton's postulates
have had to be amended after we learn more about the atom.
N
O
N
Two-to-one
ratio of mass
of nitrogen
Figure 3.4
Dalton's
Explanation of the Law of
Multiple Proportions
Because the atoms of each element
have a given mass, the fact that the
atomic ratio is two atoms of nitrogen
in one compound to one atom of
nitrogen in the other (for one atom of
oxygen in each) means that there is a
2 : 1 ratio of masses of nitrogen in
the two compounds (for a given mass
of oxygen).
