Download Atomic Structure - Coronado High School

Atomic Structure
Atoms
Table of Contents
• Slide 3-8 Atomic Theory
• Slide 9 Distinguishing Between Atoms
• Slide 10 Isotopes
Atomic Theory: Democritus
• Democritus of Abdera, was a teacher who lived
in Greece before Christ.
• He suggested the existence of these particles
called atoms.
• He said they were indivisible and indestructible.
• The real nature of atoms and the connection
between observable changes and events at the
atomic level were not established for more than
2000 years.
Atomic Theory: John Dalton (1766 –
1844)
 2000 years later, an
English school teacher
performed experiments
to test and correct his
atomic theory.
 Dalton studied ratios in
which elements combine
in a chemical reactions.
This led to the
development of his
atomic theory.
Atomic Theory: Dalton’s Postulates
 All elements are composed of tiny indivisible
particles called atoms.
 Atoms of the same element are identical. Atoms of
any one element are different from those of any
other elements.
 Atoms of different elements can physically mix
together or can chemically combine with another in
simple or whole number ratios to form compounds.
 Chemical reactions occur when atoms are separated,
joined or rearranged. Atoms of one element,
however, are never changed into atoms of another
element as a result of a chemical reaction.
Atomic Theory: Rutherford’s Nuclear
Atom
• Rutherford theorized that
atoms have their positive
charge concentrated in a
very small nucleus,[5]
and thereby pioneered
the Rutherford model or
planetary model of the
atom, through his
discovery and
interpretation of
Rutherford scattering in
his gold foil experiment.
(wikipedia)
Atomic Theory: Bohr’s Nuclear Atom
•
Niels Bohr profited by following the
experimental work going on in the
Cavendish Laboratory under Sir J.J.
Thomson's guidance, at the same time as
he pursued own theoretical studies. In the
spring of 1912 he was at work in Professor
Rutherford's laboratory in Manchester,
where just in those years such an intensive
scientific life and activity prevailed as a
consequence of that investigator's
fundamental inquiries into the radioactive
phenomena.
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Contributions to Physics and Chemistry
The Bohr model of the atom, the theory
that electrons travel in discrete orbits
around the atom's nucleus.
The shell model of the atom, where the
chemical properties of an element are
determined by the electrons in the
outermost orbit.
•
Atomic Theory: Thomson’s Electron
Properties
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Discovery of the Electron
In 1897, Thomson was the first to suggest
that the fundamental unit of the atom was
over 1000 times smaller than an atom,
suggesting the sub-atomic particles now
known as electrons.
Thomson discovered this through his
explorations on the properties of cathode
rays. Thomson made his suggestion on the
30th of April 1897 following his discovery
that Lenard rays could travel much further
through air than expected for an atomicsized particle.[4
Other work
In 1905 Thomson discovered the natural
radioactivity of potassium.[7]
In 1906 Thomson demonstrated that
hydrogen had only a single electron per
atom. Previous theories allowed various
numbers of electrons.[8][9]
(wikipedia)
Distinguishing Between Atoms
• The atomic number of an element=the
number of protons and the number of
electrons.
• Mass number is the total number of protons
and neutrons in an atom.
• To get the number of neutron do this:
Mass number – Atomic number = # of
Neutrons.
Isotopes
 Isotopes are atoms that
have the same number
of protons but different
numbers of neutrons.
 Because isotopes of an
element have different
numbers of neutrons,
they also have different
mass numbers.
 Isotopes are chemically
alike because they have
identical numbers of p+
and e-.
Distinguishing Between Atoms
• Since the 1920’s, it has been possible to
determine these tiny masses (atom’s mass)
by using a mass spectrometer.
• The atomic mass of an element is a
weighted average mass of the atoms in a
naturally occurring sample of the element.
Ions
• Ions- atom or group of atoms that have a
positive or negative charge.
• Charge is due to the loss or gain of electrons
– Cation – ion with a positive charge
• Ex: Al3+
-Anion – ion with a negative charge
Ex: O2-
Atomic Mass
• Atomic mass- weighted average mass of the
atoms in a naturally occurring sample of the
element.
• Atomic mass unit (amu)- standard unit used for
indicating atomic mass.
– One -twelth the mass of a carbon-12 atom.
Calculating Average Atomic Mass
• -Involves the relative abundance of naturally occurring
isotopes of an element.
• -Relative abundance is calculated by dividing the number of
atoms of each isotope by the total # of atoms in the sample.
– Relative abundance can also be given as a %. If you are given a
percentage you must divide by 100.
• -The Relative abundance of each isotope is then multiplied
by the mass of that isotope. This is the relative mass.
• -The relative masses are added together and that equals
the average atomic mass.
• Ex: Neon has two naturally occurring isotopes. The mass
and # of atoms of each isotope are shown below. What is
the average atomic mass of neon?
Isotope
# of atoms in
sample
Mass
Neon-20
92
20.179
Neon-22
18
21.991