UNIT D READING NOTES
D-1 Reading
Historical Background
The continuous theory of matter is a view that a solid body can be divided and subdivided into
smaller and smaller pieces. This idea came from Aristotle.
Greek philosophers, like Democritus and Leucippus (from more than 2000 years ago) suggested
the discontinuous theory of matter where matter is made up of particles so small and
indestructible that they cannot be divided into anything smaller. They called these indestructible
particles atomos, which means indivisible.
The Greeks weren’t able to prove their theories—evidence to support the atomic hypothesis was
not found until the 18th century.
Antoine Lavoisier provided the first experimental evidence for the law of conservation of mass:
MATTER IS NEITHER CREATED NOR DESTROYED.
Joseph Proust worked to determine the chemical composition of compounds. He observed that
the proportion by mass (% composition) of the elements in a SINGLE given compound is
always the same. This observation, called the law of definite proportions (or definite
composition), applies to all pure compounds. We will see this time and again in our future unit on
percent composition.
Finally, John Dalton worked to see if there was some sort of relationship between compounds
formed from the same elements. Dalton found that if two or more different (read multiple)
compounds are composed of the same two elements, the ratio of the masses of the second element
that combines with a certain mass of the first element will always be a whole number ration. For
example, carbon and oxygen can combine chemically to form both carbon monoxide (CO) and
carbon dioxide (CO2). If we look at the respective masses of oxygen in each of these compounds
that will combine with one gram of carbon, we find the following:
mass _ Oxygen _ that _ combines _ with _ 1 _ g _ C _ in _ CO2
mass _ Oxygen _ that _ combines _ with _ 1 _ g _ C _ in _ CO

2.66 g _ O
1.33 g _ O
 12 , which is a whole-number ratio.
This law is called the law of multiple proportions.
In the early 1800’s, Dalton combined all of this information into the first unified atomic theory
that was supported by evidence:
1. All matter is composed of extremely small particles called atoms, which cannot be
subdivided, created, or destroyed.
2. Atoms of a given element are identical in size, mass, and other properties
3. atoms of different elements will have different size, mass, and other properties.
4. Atoms of different elements combine in simple whole-number rations to form chemical
compounds
5. In chemical reactions atoms are combined, separated, or rearranged, but never created,
destroyed, or changed.
Part 1 is an expression of the law of conservation of mass. Dalton explains in Part 5 how
formation of new substances does not violate the law of conservation of mass. An illustration of
the law of multiple proportions is shown in Figure 9.17 on page 275.
Just as Dalton based his early atomic theory on the evidence available at the time, we now know
that atoms are, in fact, not indestructible. Experiments discussed in D-2 Reading detail how
scientists discovered the presence of subatomic particles, and early nuclear physics research
provided evidence that atoms can be split apart into smaller atoms, or fused together to make
larger atoms. In addition, although we still believe that all atoms of the same element have the
Page 1 of 6
same chemical properties, we now have evidence that suggests that atoms of a given element can
have different masses.
Atoms Today
Atoms are the fundamental building blocks that make up all matter. All atoms of the same
element have the same chemical properties.
D-2 Reading
Discovery of the Electron
The first evidence of electrons came with the work of Crookes. In the late 1800’s he was trying
to see how much he had to lower the pressure (by evacuation) of a gas in a tube in order to get a
spark to jump from a negatively charged cathode to a positively charged anode. When the tube
was almost completely evacuated, he observed a greenish glow in the tube that originated at the
cathode (Figure 4.4, page 104). He called the glow cathode rays, and the entire apparatus was
called a cathode ray tube, or CRT.
Crookes and other researchers, like J.J. Thomson, conducted further experiments that showed the
following:
1. An object placed in the cathode ray glow cast a shadow on the back wall of the glass tube;
2. A paddle wheel placed on rails began to move when the current was turned on—this
indicated that cathode rays might have mass that were sufficient to turn the wheel;
3. Cathode rays were deflected (bent) by a magnetic field like a wire carrying an electric current
(which was known to have negative charge)—see Figure 4.5 (a) on page 105;
4. Cathode rays were deflected away from a negatively charged object—because like charges
repel, this made it clear that cathode rays were negatively charged—see Figure 4.5 (b) on
page 105.
5. Thomson found that the ratio of the charge of the cathode ray particles to the mass of the
particles was always the same, even when he changed the gas in the tube, or the metal out
of which he made the cathode. From this, he concluded that cathode ray particles, or
electrons, were present in all atoms, and that each one carried an identical negative charge.
Although Thomson used a Crookes tube as the basis for his experiments, Thomson is given credit
for the discovery of the electron.
Robert Millikan, an American physicist, conducted experiments with charged oil droplets that,
combined with Thomson’s work, demonstrated both the charge and mass of an electron. The
1 _ amu
mass of the electron is 1837 , (1amu is called an atomic mass unit, and is equivalent to 121 the
mass of a carbon atom, or the mass of a hydrogen atom), or 9.109 1031 kg. These discoveries
led to some important new ideas that needed to be tested:
1. Because atoms are electrically neutral, they must contain a positive charge to balance the
negative electrons;
2. Because electrons have such a small mass compared to the mass of the atom, they must also
contain other particles that have much greater mass to account for this.
Thomson came up with a new model of the atom in which he described an atom as being a
positively-charged globular mass dotted with “raisins” or spots of negative charge which he
dubbed electrons. He called this the “Plum Pudding” model of atomic structure (Figure 5.2 on
page 128).
Discovery of the Atomic Nucleus
Next up was the work of Geiger, Marsden and Rutherford and the Gold Foil Experiments.
Radioactivity had already been discovered, and they had characterized alpha particles, as very
Page 2 of 6
small, dense (about 4 times the mass of a hydrogen atom), positively charged particles that
traveled at high speeds as they were ejected by radioactive material.
In this experiment, they fired a stream of alpha particles at a very thin piece (a foil) of gold (See
Figure 4.7 (a) on page 107). They set up this experiment to verify the plum pudding model,
which said that all mass and charge was distributed evenly throughout the atoms of the gold. If
the model was correct, they would have seen all of the alpha particles pass through with slight
deflections from their original path. What they saw was something very different. Astonishingly,
they found that NEARLY ALL of the alpha particles passed through the foil with no deflection at
all. Very few were slightly deflected from their original path, and an even smaller amount (1 in
8000) were widely deflected (kicked straight back at the alpha particle source). They found
exactly the same results even when they changed the foil target material.
These observations called for a modification of Thomson’s atomic model. Because most of the
alpha particles passed straight through the foil with no deflection, Rutherford, et al. were forced
to conclude that atoms are mostly empty space. They interpreted the infrequent slight deflections
of the positively charged alpha particles as evidence that they had come very close to, and been
repelled by, a very small, dense, positive charge. They viewed the very rare wide deflections of
alpha particles as the repulsion of a positively charge alpha particle as it approached headon a very small, dense, positively charged body that they thought was located at the center
of each atom. The like charges of the alpha particles and the so-called nucleus caused the
alpha particles to be repelled. You can see a physical model of Rutherford’s hypothesis of what
was happening in the gold foil at the atomic level in Figure 4.7(b) on page 107. Rutherford had
found the nucleus, but he had no evidence as to where the electrons were located. He suggested
that they surrounded the positively-charged nucleus like planets around the sun, but he could not
explain why they stayed in motion around the nucleus and did not spiral in toward the nucleus.
Eugen Goldstein did further experiments with cathode ray tubes and found that as electrons
traveled from the cathode to the anode, positively charged particles (protons) moved in the
opposite direction. Protons were found to explain the strong positive charge of the nucleus, and
uncharged neutrons were later identified in the 1930’s by James Chadwick to explain the mass of
an atom being greater than the sum of the masses of protons and electrons in an atom.
The nucleus is located at the center of the atom, and there are 3 major types of particles that make
up atoms:
SUBATOMIC PARTICLES AND THEIR CHARACTERISTICS
Characteristic
Symbol
Proton
p+,
1
1
p
Neutron
1
n, 0 n
Electron
e ,
0
1
e
Location in the atom
nucleus
nucleus
in motion outside of nucleus
Charge
+1
0
1
Relative Mass
1 amu (1.673  1027
kg)
1 amu
1 _ amu
(much less than others)
1837
Page 3 of 6
The net charge of a neutral atom is zero, because the number of protons and electrons is equal.
When the number of protons and electrons in an atom are not equal, the particle becomes
charged:
Sample Problem 1
What is the net charge on a particle that consists of 15 protons, 16 neutrons, and 18 electrons?
Calculation of Net Charge
Particle
Number of Particles
Charge per particle



Total charge
proton
neutron
electron
15
16
18-
1+
0
1-
=
15+
=
0
=
18Net charge
3We have since found out that there are other subatomic particles, but because they don’t affect the
chemical properties of the elements, we do not discuss them.
Scientific models help scientists interpret what they observe. They are mental pictures that help
us understand something we cannot see or experience directly. The atomic model that has
been developed over time has explained the characteristics and behavior of atoms. Only recently
have scientists finally been able to “see” atoms under extremely high magnification, and even
then we have not yet viewed any internal structures that we theorize here. Models change as more
is learned, so the atomic model is sure to be refined. The following is a summary of the atomic
model starting in 1803:
Date
Scientist
Description/name of model
1803
Dalton
Indivisible particle theory
1897
Thomson
Plum-pudding model
1909
Rutherford
Positively charged nucleus surrounded by mostly empty space
We will discuss more recent models in our Electronic Structure unit.
D-3 Reading:
The Major Nucleons
The particles that make up the nucleus are called nucleons. The most important ones are called
protons and neutrons.
Protons:
 Single positive charge that is the same size as the negative charge on the electron.


Mass of a proton = 1.67  1027 kg.
The number of protons in the nucleus of an atom is called its atomic number. All atoms of
the same element will have the same atomic number (and therefore the same number of
protons in the nucleus. Atoms of different elements have a different number of protons in
their nuclei.
 The atomic number indicates the number of electrons in a electrically neutral atom (number
of positive charge must balance negative charge).
 The symbol Z stands for the atomic number of an element.
Neutrons:
 The neutron was discovered by James Chadwick (a British physicist) in 1932.
Page 4 of 6
 It has almost the same mass as a proton, but has no electrical charge (neutral)
Isotopes:
 Not all atoms of the same element have the same mass. The mass of an atom is the sum of the
neutrons and the protons. The number of protons cannot vary and still have the same element,
but the number of neutrons can vary. Thus isotopes are atoms of the same elements with
different numbers of neutrons in the nucleus. Isotopes of the same element will have different
masses.
An example of isotopes of an element is provided in the text:
Hydrogen Isotope
# protons
# neutrons
Atomic mass
Symbol
Protium
1
0
1 amu
1
1H
Deuterium
1
1
2 amu
2
1
Tritium
1
2
3 amu
3
1H
H

In a natural sample of hydrogen, about one atom in every 6000 is an atom of deuterium.
Tritium is not found in natural samples of hydrogen—it is made artificially. It is an unstable,
or radioactive isotope.
 Only the isotopes of hydrogen have been given special names. All others are identified
according to their atomic mass: uranium with a mass of 235 amu is called U-235, and the
isotope with a mass of 238 amu is called U-238.
Concept of Atomic Mass
Until the discovery of isotopes, chemists assumed that all the atoms of a particular element have
the same mass. They weren’t able to determine the mass of individual atoms, but they were able
to determine relative masses with respect to the element with the smallest mass, hydrogen.
Hydrogen was assigned the atomic mass of 1 atomic mass unit (1 amu). Masses of the other
elements were then assigned based on their mass relative to hydrogen.
Mass Number
 The mass of a proton is approximately 1 amu. So is the mass of the neutron. As we
previously discussed, the mass of an electron is substantially less than that of a proton or a
neutron, so it contributes little to the mass of an atom. If you add together the masses of the
neutrons and protons, you will find the total mass of an isotope, which scientists call the
mass number. We designate the mass number of an isotope as the letter A.
 The number of protons in an atom determines the type of element. The number of protons in
a nucleus is called the atomic number. Scientists use the symbol Z to represent the atomic
number.
 As we have said before, there is often more than one isotope of each element. Because of this,
we need to have a symbolic way to represent different isotopes of the same element that tells
as much information as possible. Let’s first discuss the isotope of carbon that has a mass of
12 amu. We know that every atom of carbon has an atomic number of 6 atoms because it has
6 protons. We represent C-12 symbolically:
Mass number, A
12
6
C
Atomic number, Z
Page 5 of 6
It should be clear that if the mass number is the sum of the protons and neutrons, then if you
know the mass number and the atomic number you should be able to find the number of
neutrons of an isotope:
# of neutrons = A-Z, so there are 6 neutrons in C-12
The isotope for C-14 (a radioactive isotope of carbon) would be written in the following way:
14
6
C
Carbon-14 has 6 protons, and an atomic mass of 14 amu. The number of neutrons is therefore
14 amu – 6 amu = 8 amu, or 8 neutrons in carbon-14.
Sample Problem
What is the number of neutrons in the nucleus of a sodium-23 atom? Represent this atom in
symbol form.
From the chemical element table (page 817), we see that the atomic number, A, is 11, which
is the number of protons. The mass number is 23, so
# of neutrons = 23 – 11 = 12 neutrons
23
We write this isotope symbolically as 11
Na, where the left superscript, 23, is the mass
number, and the left subscript, 11, is the atomic number.
D-4 Reading
Modern Standard of Atomic Mass
 Recall that 19th century chemists used measured masses relative to the mass of hydrogen.
They found that chlorine had a mass 35.5 times as large as hydrogen, but that doesn’t fit
with today’s understanding of mass and atomic structure. You cannot have half of a
neutron, and it is certainly impossible according to the theory to have a fraction of proton.
 We explain it by saying that different isotopes of an element will have different masses,
and the early scientists measured an average atomic mass that is a weighted average of
the isotopes of chlorine.

We now use carbon-12, ( 12
6 C), as the basis for mass. One atomic mass unit, amu, is now
defined as
1
12
the mass of
12
6
C.
Determining Atomic Masses from Weighted Averages
Sample Problem (copper, page 80):
Copper has two naturally occurring isotopes. 69.17% of them are copper-63, which has an
atomic mass of 62.929598 amu, and 30.83% of them are copper-65, which has an atomic
mass of 64.927793 amu. We determine the average atomic mass of a naturally occurring
sample of copper by setting up the following table:
Copper-63:
Copper-65
(0.6917)
(0.3083)
 (62.929598 amu)
 (64.927793 amu)
Total average atomic mass
=
=
=
43.53 amu
20.02 amu
63.55amu
PREVIEW OF PERIODIC TABLE
We use the periodic table to organize elements according to their chemical and physical
properties. Vertical columns on the table are called groups or families, because they have similar
characteristics. Horizontal rows are called periods, because their properties vary in predictable
patterns, and this pattern is repeated in every row.
Page 6 of 6