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Transcript
Chapter 16:
The Properties of Atoms
and the Periodic Table
• Section 1 – The Structure of Atoms
• Section 2 – Masses of Atoms
Section 3 – The Periodic Table
Section 1 – The Structure of Atoms
The Development of the Atomic Model
The word atom was first used by the Greek scientist/philosopher
Democritus 2400 years ago
• He believed that the atom was the smallest particle of matter
possible
• Only four atoms: fire, water, earth, and air
 The Greeks thought that an atom of fire would look like a very
small flame, water like a very small bit of water, and so on
 All matter consisted of these atoms or mixtures of the atoms
• The problem with the Greek idea of the atom is that they could
not test to see if it was correct
In the early 1800’s, English scientist John Dalton provided the basic
theory of the scientific atomic model with four basic statements
• All matter is composed of indivisible atoms. These atoms retain
their identity during chemical reactions
• An element is a type of matter composed of only one type of
atom, each kind of atom have the same properties
• A compound is a type of matter composed of atoms of two or
more elements chemically combined in fixed proportions, Ex.:
water → 4H + O2  2H2O
• A chemical reaction consists of the rearrangement of the atoms
present in the reacting substances to give new chemical
combinations present in the substances formed by the reaction
Ex.: formation of water:
• One problem: no one could prove the existence of atoms
Section 1 – The Structure of Atoms
The Development of the Atomic Model
Dalton was wrong about one thing: atoms can be divided into
smaller particles
• In 1897 J.J., Thomson demonstrated the existence of negatively
charged particles smaller than a hydrogen atom through a series
of experiments using the cathode ray tube (CRT)
•
•
•
•
Scientists knew that charging a CRT would result in the
formation of a beam. Was beam a series of waves or a stream of
particles?
Thomson used a CRT and a magnet to show that the beam was
deflected by the magnetic field so it must consist of a stream of
particles
Next, he placed two oppositely charged plates in the CRT and
discovered that the particles were deflected towards the (+)
charged plate.
After further tests and calculations, Thomson concluded the
particles were much smaller than a hydrogen atom and carried a
(-) charge—the electron.
Section 1 – The Structure of Atoms
The Development of the Atomic Model
The Thomson atomic model consists
of a positively charged ball with
negatively charged particles evenly
embedded in the ball
In 1911 Ernest Rutherford demonstrated the presence of a
positively charge atomic nucleus
• Using alpha particles and gold
foil, Rutherford conducted
the following
experiment:
(An alpha particle is
the nucleus of a
helium atom which
has a positive
charge—at the
time they just
knew that it was a
positively charged
particle)
1. The stream of alpha particles would leave the lead block
and some would pass through a small hole in another lead
block
Section 1 – The Structure of Atoms
The Development of the Atomic Model
2. The particles would then strike a piece of gold foil
3. Most of the particles would pass through the foil,
although their path was deflected. But some particles
would rebound back toward the source.
4. Rutherford deduced the atoms of gold must have a
massive positively charged nucleus (called the proton)
The Rutherford atomic model
proposed the majority of the
mass of an atom was in the
positively charged nucleus, that
there was empty space between
the nucleus and the electrons,
and the electrons circled the
nucleus
Continued experimentation by Rutherford revealed:
• The nucleus contains 99.95% of the mass of an atom, but
occupies a very small space in the atom
• If the nucleus was the size of a golf ball, the atom would be
about 3 miles in diameter
• It would take about 2000 electrons to equal the mass of one
proton
Section 1 – The Structure of Atoms
The Development of the Atomic Model
The location of electrons was also an issue
• In 1920, Niels Bohr theorized that electrons circled the
nucleus much like the planets orbit the sun
 Bohr called these orbits
energy levels. That is, each
electron has a very specific
path, and you could
determine both the location
and motion of the electron
• In 1926, Werner Heisinberg, based on quantum mechanics,
demonstrated it was impossible to know both the motion
and location of an electron at the same time
 Heisenberg proposed that the electrons form a cloud
around the nucleus of an atom.
In the electron cloud were
regions called orbitals where the
electrons were likely to be found
Section 1 – The Structure of Atoms
The Development of the Atomic Model
There was a problem with Rutherford’s model of the atom. The
proton alone could not account for the mass of the nucleus,
there was a missing particle
• In the early 1900’s, scientists knew that hydrogen consisted
of one proton and 1 electron, and that helium contained 2
protons and 2 electrons
• The ratio of the mass of helium to the mass of hydrogen
should have been 2 to 1, but the actual ratio was 4 to 1
• In 1932 James Chadwick performed a series of experiments
similar to Rutherford’s. He found that an unidentified, high
energy radiation was given off
• Chadwick was able to prove this radiation was composed of
neutrally charged particles with about the same mass as the
proton
• Thus, the neutron was discovered
The Modern Atomic Model
Consists of the Atomic Nucleus
which contains protons and
neutron, and the Electron
Cloud which contains electrons
that are arranged in energy
levels and whose motion and
position cannot be described at
the same time.
Section 2 – Masses of Atoms
The identification of an element
• Elements are identified by the number of protons in the
nucleus of an atom of the element
• As the number of protons in the nucleus changes, so does
the element
• The number of protons in the nucleus of an atom is called its
atomic number
• The mass number of an element is the sum of the number of
protons and neutrons in the nucleus of an atom of the
element
• The number of neutrons in the nucleus of an atom can be
determined mathematically:
# of neutrons = mass number - atomic number
• Sometimes the number of neutrons does not equal the
number of protons. When this is the case, we have isotopes
of the same element
 Remember, the number of protons determine the
element, so you can more (or less) neutrons than protons
and have the same element
 Some isotopes are radioactive, and this property can be
used in different ways
 Example: Carbon-14 (C14) – Scientists know that Carbon14 decays at a regular rate. Using that fact, they can
determine the age of something by the comparing the
ratio of Carbon-12 to carbon-14
Section 2 – Masses of Atoms
• The mass of an atom can be measured using relative units
• The unit used to express the relative mass of atoms is the
atomic mass unit (amu)
• 1 amu is defined as one-twelveth (1/12) the mass of a
Carbon-12 atom
• Protons and neutrons have about the same mass: nearly 1
amu
 mP = 1.673 x 10-24-kg, mN = 1.675 x 10-24-kg
Section 3 – The Periodic Table
The Periodic Table is an representation of all the known
elements and is arranged by increasing atomic number
• When the elements are arranged in this order, the elements
in a the same group (a vertical column on the Table) have
similar properties
• Each element has its own “box” on the Periodic Table that
describes the properties of the element.
 The properties include: the atomic number, the state of
matter of the element at room temperature, the symbol
used to identify the element, the name of the element,
and the element’s average atomic mass.
• There are four basic types of elements: metals, nonmetals,
metalloids, and synthetic (man-made)
 Metals typically are solid at room temperature, and are
good conductors of heat and electricity
 Nonmetals typically are gases at room temperature and
are poor conductors of heat and electricity
 Metalloids have some of the properties of both metals
and nonmetals
 The synthetic elements are typically constructed in
particle accelerators and most do not stay in existence for
long periods – they breakdown into smaller atoms and/or
subatomic particles