Transcript Atomic Theory of Matter

ATOMIC THEORY OF
MATTER
Chapter 2.1
2.1 The Atomic Theory of Matter
• In efforts to explain observations,
philosophers from the earliest times
have speculated about the nature of
the fundamental "stuff" from which the
world is made.
2.1 The Atomic Theory of Matter
• Democritus (460–370 BC) and other early
Greek philosophers thought that the
material world must be made up of tiny
indivisible particles that they called atomos,
meaning "indivisible."
• Plato and Aristotle thought that all matter
was made up of earth, wind, fire, and water.
2.1 The Atomic Theory of Matter
• As chemists learned to measure the
amounts of materials that reacted with one
another to make new substances, the
ground was laid for a chemical atomic
theory.
• That theory came into being during the
period 1803–1807 in the work of an English
schoolteacher, John Dalton. Reasoning
from a large number of observations,
Dalton made the following postulates:
2.1 The Atomic Theory of Matter
• Each element is
composed of extremely
small particles called
atoms.
• All atoms of a given
element are identical;
the atoms of different
elements are different
and have different
properties (including
different masses).
2.1 The Atomic Theory of Matter
• Atoms of an element are not changed
into different types of atoms by
chemical reactions; atoms are neither
created nor destroyed in chemical
reactions.
• Compounds are formed when atoms
of more than one element combine; a
given compound always has the same
relative number and kind of atoms.
2.1 The Atomic Theory of Matter
• According to Dalton's atomic
theory, atoms are the basic building blocks
of matter.
• They are the smallest particles of an
element that retain the chemical identity of
the element.
• As noted in the postulates of Dalton's
theory, an element is composed of only one
kind of atom, whereas a compound
contains atoms of two or more elements.
2.1 The Atomic Theory of Matter
• Dalton's theory explains several
simple laws of chemical combination
that were known in his time. One of
these was the law of definite
proportions: In a given compound the
relative numbers and kinds of atoms
are constant.
2.1 The Atomic Theory of Matter
• Another fundamental chemical law was
the law of conservation of mass (also
known as the law of conservation of
matter): The total mass of materials present
after a chemical reaction is the same as the
total mass before the reaction.
• Dalton proposed that atoms always retain
their identities and that during chemical
reactions the atoms rearrange to give new
chemical combinations.
THE DISCOVERY OF
THE ATOMIC
STRUCTURE
Chapter 2.2
2.2 The Discovery of the Atomic Structure
• Dalton reached his conclusion about
atoms on the basis of chemical
observations in the macroscopic world
of the laboratory.
• Today, however, we can use powerful
new instruments to measure the
properties of individual atoms and
even provide images of them.
2.2 The Discovery of the Atomic Structure
• An image of the surface
of the semiconductor
GaAs (gallium arsenide)
as obtained by a
technique called
tunneling electron
microscopy. The color
was added to the image
by computer to
distinguish the gallium
atoms (blue spheres)
from the arsenic atoms
(red spheres).
2.2 The Discovery of the Atomic Structure
• As scientists began to develop methods for more detailed
probing of the nature of matter, the atom, which was
supposed to be indivisible, began to show signs of a more
complex structure: We now know that the atom is
composed of still smaller subatomic particles.
• Before we summarize the current model of atomic
structure, we will briefly consider a few of the landmark
discoveries that led to that model.
2.2 The Discovery of the Atomic Structure
• We'll see that the atom is composed in part of electrically
charged particles, some with a positive (+) charge and
some with a negative (-) charge.
• As we discuss the development of our current model of
the atom, keep in mind a simple statement of the behavior
of charged particles with one another: Particles with the
same charge repel one another, whereas particles with
unlike charges are attracted to one another.
Cathode Rays and Electrons
• In the mid-1800s, scientists began to study electrical
discharge through partially evacuated tubes (tubes that
had been pumped almost empty of air), such as those
shown in Figure 2.3. A high voltage produces radiation
within the tube. This radiation became known as cathode
rays because it originated from the negative electrode, or
cathode. Old television picture tubes are cathode-ray
tubes; a television picture is the result of fluorescence
from the television screen.
Cathode Rays and Electrons
Figure 2.3 (a) In a cathode-ray tube, electrons move from the negative
electrode (cathode) to the positive electrode (anode). (b) A photo of a cathoderay tube containing a fluorescent screen to show the path of the cathode rays.
(c) The path of the cathode rays is deflected by the presence of a magnet.
Cathode Rays and Electrons
• Scientists held differing views about the nature of the
cathode rays.
• It was not initially clear whether the rays were a new form
of radiation or rather consisted of an invisible stream of
particles.
• Experiments showed that cathode rays were deflected by
electric or magnetic fields, suggesting that the rays
carried an electrical charge.
Cathode Rays and Electrons
• The British scientist J. J. Thomson observed many
properties of the rays, including the fact that the nature of
the rays is the same regardless of the identity of the
cathode material and that a metal plate exposed to
cathode rays acquires a negative electrical charge.
• In a paper published in 1897 he summarized his
observations and concluded that the cathode rays are
streams of negatively charged particles with mass.
Thomson's paper is generally accepted as the "discovery"
of what became known as the electron.
Cathode Rays and Electrons
• In 1909 Robert Millikan (1868–1953) of the University of
Chicago succeeded in measuring the charge of an
electron by performing what is known as the "Millikan oildrop experiment”.
• Using slightly more accurate values, the presently
accepted value for the mass of the electron is 9.11 x 10-28
g.
Radioactivity
• In 1896 the French scientist Henri
Becquerel (1852–1908) was
studying a uranium mineral
called pitchblende, when he
discovered that it spontaneously
emits high-energy radiation. This
spontaneous emission of radiation
is called radioactivity.
• At Becquerel's suggestion Marie
Curie and her husband, Pierre,
began experiments to isolate the
radioactive components of the
mineral.
The Nuclear Atom
• With the growing evidence that the atom is composed of
even smaller particles, attention was given to how the
particles fit together.
• In the early 1900s Thomson reasoned that because
electrons comprise only a very small fraction of the mass
of an atom, they probably were responsible for an equally
small fraction of the atom's size.
• He proposed that the atom consisted of a uniform positive
sphere of matter in which the electrons were embedded.
• This so-called "plum-pudding" model, named after a
traditional English dessert, was very short-lived.
The Nuclear Atom
• In 1910 Rutherford and his
coworkers performed an experiment
that disproved Thomson's model.
Rutherford was studying the angles
at which particles were scattered as
they passed through a thin gold foil a
few thousand atomic layers in
thickness (Figure 2.10). He and his
coworkers discovered that almost all
of the particles passed directly
through the foil without deflection. A
small percentage were found to be
slightly deflected.
The Nuclear Atom
• By 1911 Rutherford
was able to explain
these observations;
he postulated that
most of the mass of
the atom and all of its
positive charge reside
in a very small,
extremely dense
region, which he
called the nucleus.
• Most of the total
volume of the atom is
empty space in which
electrons move
around the nucleus.
The Nuclear Atom
• In the α -scattering experiment
most α particles pass directly
through the foil because they do
not encounter the minute nucleus;
they merely pass through the
empty space of the atom.
• Occasionally, however,
an particle comes into the close
vicinity of a gold nucleus.
• The repulsion between the highly
charged gold nucleus and
the particle is strong enough to
deflect the less massive α
particle.