Download 3.1 Early History of Atomic Theories

3.1
Figure 1
In Dalton’s atomic model, an atom is
a solid sphere, similar to a billiard
ball. This simple model is still used
today to represent the arrangement
of atoms in molecules.
DID YOU
KNOW
?
William Crookes (1832–1919)
William Crookes was the eldest of
sixteen children and inherited his
father’s fortune, made in real estate.
This enabled him to lead a leisurely
life, and also to conduct scientific
research in many areas of chemistry
and physics. Crookes is best known
for his cathode ray tube, which was
made possible by his improvements
to the vacuum pump and Volta’s
invention of the electric cell. His
vacuum techniques later made mass
production of the light bulb practical.
Early History of Atomic Theories
The history of atomic theories is full of success and failure stories for hundreds of
chemists. In textbooks such as this one, only the success of a few is documented. However,
the success of these chemists was often facilitated by both the success and failure of
many others.
Recall that by the use of deductive logic the Greeks (for example, Democritus) in
about 300 B.C. hypothesized that matter cut into smaller and smaller pieces would eventually reach what they called the atom — literally meaning indivisible. This idea was
reintroduced over two thousand years later by an English chemist/schoolteacher named
John Dalton in 1805. He re-created the modern theory of atoms to explain three important scientific laws — the laws of definite composition, multiple proportions, and conservation of mass. The success of Dalton’s theory of the atom was that it could explain
all three of these laws and much more. Dalton’s theory was that the smallest piece of
matter was an atom that was indivisible, and that an atom was different from one element to another. All atoms of a particular element were thought to be exactly the same.
Dalton’s model of the atom was that of a featureless sphere — by analogy, a billiard ball
(Figure 1). Dalton’s atomic theory lasted for about a century, although it came under
increasing criticism during the latter part of the 1800s.
SUMMARY
Creating the Dalton Atomic Theory
(1805)
Table 1
Key experimental work
Theoretical explanation
Atomic theory
Law of definite composition:
elements combine in a
characteristic mass ratio
Each atom has a particular
combining capacity.
Law of multiple proportions:
there may be more than
one mass ratio
Some atoms have more
than one combining
capacity.
Matter is composed of
indestructible, indivisible
atoms, which are identical
for one element, but
different from other
elements.
Law of conservation of
mass: total mass remains
Atoms are neither created
nor destroyed constant
in a chemical reaction.
The Thomson Atomic Model
The experimental studies of Svante Arrhenius and Michael Faraday with electricity and
chemical solutions and of William Crookes with electricity and vacuum tubes suggested
that electric charges were components of matter. J. J. Thomson’s quantitative experiments with cathode rays resulted in the discovery of the electron, whose charge was later
measured by Robert Millikan. The Thomson model of the atom (1897) was a hypothesis that the atom was composed of electrons (negative particles) embedded in a positively charged sphere (Figure 2(a)). Thomson’s research group at Cambridge University
in England used mathematics to predict the uniform three-dimensional distribution of
162 Chapter 3
NEL
Section 3.1
the electrons throughout the atom. The Thomson model of the atom is often communicated by using the analogy of a raisin bun, with the raisins depicting the electrons and
the bun being the positive material of his atom (Figure 2(b)).
(a)
SUMMARY
INVESTIGATION 3.1.1
The Nature of Cathode Rays
(p. 209)
The discovery of cathode rays led to
a revision of the Dalton atomic
model. What are their properties?
Figure 2
(a) In Thomson’s atomic model, the
atom is a positive sphere with
embedded electrons.
(b) This model can be compared to
a raisin bun, in which the
raisins represent the negative
electrons and the bun represents the region of positive
charge.
(b)
Creating the Thomson Atomic Theory
(1897)
Table 2
Key experimental work
Theoretical explanation
Atomic theory
Arrhenius: the electrical
nature of chemical
solutions
Atoms may gain or lose
electrons to form ions in
solution.
Faraday: quantitative
work with electricity and
solutions
Particular atoms and ions
gain or lose a specific
number of electrons.
Crookes: qualitative
studies of cathode rays
Electricity is composed
of negatively charged
particles.
Matter is composed of
atoms that contain
electrons (negative
particles) embedded in a
positive material. The kind
of element is characterized
by the number of electrons
in the atom.
Thomson: quantitative
studies of cathode rays
Electrons are a
component of all matter.
Millikan: charged oil
drop experiment
Electrons have a specific
fixed electric charge.
DID YOU
KNOW
?
Rutherford Quotes
• “You know it is about as incredible as if you fired a 350-mm
shell at a piece of tissue paper
and it came back and hit you.”
• “Now I know what the atom
looks like.” 1911
• The electrons occupy most of the
space in the atom, “like a few
flies in a cathedral.”
• The notion that nuclear energy
could be controlled is “moonshine.” 1933
The Rutherford Atomic Theory
One of Thomson’s students, Ernest Rutherford (Figure 3), eventually showed that some
parts of the Thomson atomic theory were not correct. Rutherford developed an expertise
with nuclear radiation during the nine years he spent at McGill University in Montreal.
He worked with and classified nuclear radiation as alpha (), beta (), and gamma ()
— helium nuclei, electrons, and high-energy electromagnetic radiation from the nucleus,
respectively. Working with his team of graduate students he devised an experiment to test
the Thomson model of the atom. They used radium as a source of alpha radiation,
which was directed at a thin film of gold. The prediction, based on the Thomson model,
was that the alpha particles should be deflected little, if at all. When some of the alpha
particles were deflected at large angles and even backwards from the foil, the prediction
NEL
Figure 3
Rutherford’s work with radioactive
materials at McGill helped prepare
him for his challenge to Thomson’s
atomic theory.
Atomic Theories
163
ACTIVITY 3.1.1
Rutherford’s Gold Foil
Experiment (p. 210)
Rutherford’s famous experiment
involved shooting “atomic bullets” at
an extremely thin sheet of gold. You
can simulate his experiment.
Prediction
was shown to be false, and the Thomson model judged unacceptable (Figure 4).
Rutherford’s nuclear model of the atom was then created to explain the evidence gathered in this scattering experiment. Rutherford’s analysis showed that all of the positive
charge in the atom had to be in a very small volume compared to the size of the atom.
Only then could he explain the results of the experiment (Figure 5). He also had to
hypothesize the existence of a nuclear (attractive) force, to explain how so much positive charge could occupy such a small volume. The nuclear force of attraction had to be
much stronger than the electrostatic force repelling the positive charges in the nucleus.
Even though these theoretical ideas seemed far-fetched, they explained the experimental
evidence. Rutherford’s explanation of the evidence gradually gained widespread acceptance in the scientific community.
alpha particles
metal foil
SUMMARY
Evidence
alpha particles
Creating the Rutherford Atomic Theory
(1911)
Table 3
metal foil
Figure 4
Rutherford’s experimental observations were dramatically different
from what he had expected based
on the Thomson model.
Key experimental work
Theoretical explanation
Atomic theory
Rutherford: A few positive
alpha particles are
deflected at large angles
when fired at a gold foil.
The positive charge in the
atom must be concentrated
in a very small volume
of the atom.
Most materials are very
stable and do not fly
apart (break down).
A very strong nuclear
force holds the positive
charges within the nucleus.
Rutherford: Most alpha
particles pass straight
through gold foil.
Most of the atom is
empty space.
An atom is composed of a
very tiny nucleus, which
contains positive charges
and most of the mass of
the atom. Very small
negative electrons occupy
most of the volume of the
atom.
Protons, Isotopes, and Neutrons
nucleus
atom
Figure 5
To explain his results, Rutherford suggested that an atom consisted mostly
of empy space, explaining why most
of the alpha particles passed nearly
straight through the gold foil.
proton ( 01p or p+) a positively
charged subatomic particle found in
the nucleus of atoms
164 Chapter 3
The Thomson model of the atom (1897) included electrons as particles, but did not
describe the positive charge as particles; recall the raisins (electrons) in a bun (positive
charge) analogy. The Rutherford model of the atom (1911) included electrons orbiting a
positively charged nucleus. There may have been a hypothesis about the nucleus being
composed of positively charged particles, but it was not until 1914 that evidence was gathered to support such a hypothesis. Rutherford, Thomson, and associates studied positive
rays in a cathode ray tube and found that the smallest positive charge possible was from
ionized hydrogen gas. Rutherford reasoned that this was the fundamental particle of positive charge and he named it the proton, meaning first. (Again Rutherford showed his
genius by being able to direct the empirical work and then interpret the evidence theoretically.) By bending the hydrogen-gas positive rays in a magnetic field they were able to
determine the charge and mass of the hypothetical proton. The proton was shown to have
a charge equal to but opposite to that of the electron and a mass 1836 times that of an
electron. All of this work was done in gas discharge tubes that evolved into the version of
the mass spectrometer (Figure 6) developed by Francis Aston during the period 1919–1925.
Evidence from radioactivity and mass spectrometer investigations falsified Dalton’s
theory that all atoms of a particular element were identical. The evidence indicated that
NEL
Section 3.1
magnet
slit
gas discharge tube
sample inlet
slit
detector
magnet
beam of positive ions
there were, for example, atoms of sodium with different masses. These atoms of different mass were named isotopes, although their existence could not yet be explained.
Later, James Chadwick, working with Rutherford, was bombarding elements with
alpha particles to calculate the masses of nuclei. When the masses of the nuclei were
compared to the sum of the masses of the protons for the elements, they did not agree.
An initial hypothesis was that about half of the mass of the nucleus was made up of
proton–electron (neutral) pairs. However, in 1932 Chadwick completed some careful
experimental work involving radiation effects caused by alpha particle bombardment.
He reasoned that the only logical and consistent theory that could explain these results
involved the existence of a neutral particle in the nucleus. According to Chadwick, the
nucleus would contain positively charged protons and neutral particles, called neutrons.
The different radioactive and mass properties of isotopes could now be explained by
the different nuclear stability and different masses of the atom caused by different numbers of neutrons in the nuclei of atoms of a particular element.
SUMMARY
Figure 6
A mass spectrometer is used to
determine the masses of ionized
particles by measuring the deflection
of these particles as they pass
through the field of a strong magnet.
isotope ( AZ X) a variety of atoms of
an element; atoms of this variety
have the same number of protons as
all atoms of the element, but a different number of neutrons
neutron ( 01n or n) a neutral
(uncharged) subatomic particle
present in the nucleus of atoms
Rutherford Model
• An atom is made up of an equal number of negatively charged electrons and postively
charged protons.
• Most of the mass of the atom and all of its positive charge is contained in a tiny core region
called the nucleus.
• The nucleus contains protons and neutrons that have approximately the same mass.
• The number of protons is called the atomic number (Z).
• The total number of protons and neutrons is called the mass number (A).
NEL
Atomic Theories
165
SUMMARY
Creating the Concepts of Protons,
Isotopes, and Neutrons
Table 4
Key experimental work
Theoretical explanation
Atomic theory
Rutherford (1914): The
lowest charge on an
ionized gas particle is
from the hydrogen ion
The smallest particle of
positive charge is the
proton.
Soddy (1913): Radioactive
decay suggests different
atoms of the same
element
Isotopes of an element
have a fixed number of
protons but varying
stability and mass.
Aston (1919): Mass
spectrometer work
indicates different
masses for some atoms
of the same element
The nucleus contains
neutral particles called
neutrons.
Atoms are composed of
protons, neutrons, and
electrons. Atoms of the
same element have the
same number of protons
and electrons, but may
have a varying number of
neutrons (isotopes of the
element).
Radiation is produced
by bombarding elements
with alpha particles.
Section 3.1 Questions
Understanding Concepts
1. Summarize, using labelled diagrams, the evolution of
atomic theory from the Dalton to the Rutherford model.
2. Present the experimental evidence that led to the
Rutherford model.
3. How did Rutherford infer that the nucleus was
(a) very small (compared to the size of the atom)?
(b) positively charged?
4. (a) State the experimental evidence that was used in the
discovery of the proton.
(b) Write a description of a proton.
5. (a) State the experimental evidence that was used in the
discovery of the neutron.
(b) Describe the nature of the neutron.
Applying Inquiry Skills
6. What is meant by a “black box” and why is this an appro-
priate analogy for the study of atomic structure?
7. Theories are often created by scientists to explain scientific
laws and experimental results. To some people it seems
strange to say that theories come after laws. Compare the
scientific and common uses of the term “theory.”
8. What is the ultimate authority in scientific work (what kind
of knowledge is most trusted)?
166 Chapter 3
Making Connections
9. State some recent examples of stories in the news media
that mention or refer to atoms.
10. Describe some contributions Canadian scientists and/or
scientists working in Canadian laboratories made to the
advancement of knowledge about the nature of matter.
GO
www.science.nelson.com
Extension
11. Rutherford’s idea that atoms are mostly empty space is
retained in all subsequent atomic theories. How can solids
then be “solid”? In other words, how can your chair support
you? Why doesn’t your pencil go right through the atoms
that make up your desk?
12. When you look around you, the matter you observe can be
said to be made from electrons, protons, and neutrons.
Modern scientific theories tell us something a little different
about the composition of matter. For example, today protons are not considered to be fundamental particles; i.e.,
they are now believed to be composed of still smaller particles. According to current nuclear theory, what is the composition of a proton? Which Canadian scientist received a
share of the Nobel Prize for his empirical work in verifying
this hypothesis of sub-subatomic particles?
NEL