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Atomic Structure
Atoms: The Building Blocks
of Matter
Early Atomic Theory
400 BC to the early 1800’s
The Ancient Greeks
• Democritus – the • Aristotle –
particle theory
succeeded
of matter,
Democritus and
nature’s basic
did not believe
particle is the
in atoms. He
atom from the
thought all
Greek word
matter was
atomos which
continuous and
means
made of a
indivisible
substance
called hyle.
Antoine Lavoisier – the
Father of Chemistry
• Lavoisier used the
balance and measured
everything which lead to
the Law of conservation of
Mass
• He named oxygen which
means “acid former” and
helped determine that air
is a mixture not an
element
John Dalton’s Atomic
Theory 1808
• All matter is composed
of small particles called
atoms and empty space
• All atoms of an element
are identical
• Atoms of different
elements are different
• Atoms combine in
simple whole number
ratios to form
compounds
Modern Atomic Theory
The late 1800’s to Present
J. J. Thompson in 1898
discovered the electron
• The atom is a
mass of positive
charge which
contains regions
of negative
charge
(electrons) like
the plums in
plum pudding
• Plum Pudding
Model
Rutherford, Geiger, and
Marsden – discovered
the nucleus
• The center of an atom is dense
and postively charged
• The positive “nucleus” of the
atom is very small compared to
the total volume of the atom
• Rutherford suggested that the
electrons surround the nucleus
like planets around the sun
4.2
The Atomic Nucleus
•Rutherford’s Gold-Foil Experiment
4.2
The Atomic Nucleus
•Alpha particles scatter from the gold foil.
4.2
The Atomic Nucleus
•
The Rutherford Atomic Model
•Rutherford concluded that the atom is
mostly empty space. All the positive
charge and almost all of the mass are
concentrated in a small region called the
nucleus.
•The nucleus is the tiny central core of an
atom and is composed of protons and
neutrons.
Rutherford-Bohr Model
Electron
Proton
Neutron
6.1
The Periodic Law
•In the modern periodic
table, elements are
arranged in order of
increasing atomic number.
Atomic Number (Z)
• Is the number of protons in
the nucleus of an atom
• Is the whole number in the
periodic table
• Is used the calculate the
number of neutrons in an
atom by its subtraction from
the Mass Number
6.2
Squares in the Periodic Table
4.3
Mass Number
The total number of protons and
neutrons in a single atom of an
element
•Au is the
chemical
symbol for
gold.
79 Protons
79 Electrons
118 Neutrons
for Sample Problem 4.1
4.3
Atomic Mass
•It is useful to compare the relative
masses of atoms to a standard
reference isotope. Carbon-12 is the
standard reference isotope. Cabon12 has a mass of exactly 12 atomic
mass units.
•An atomic mass unit (amu) is
defined as one twelfth of the mass
of a carbon-12 atom.
4.3
Distinguishing Among
Atoms
•Just as apples
come in different
varieties, a
chemical element
can come in
different
“varieties” called
isotopes.
The Nucleus of the Atom
• All elements are
identified by
their number of
protons
• Atoms with the
same number of
protons but
different
numbers of
neutrons are
called isotopes
ISOTOPES
Average
Atomic Mass
Number
• The average of the
total number of
protons and
neutrons in the
nuclei of all the
isotopes of an
element
• The decimal
number in the
periodic table
Relative Abundance
• Tells how common a particular
isotope of an element is in
nature
• It is used to calculate the
average atomic mass of an
element
4.3
Average Atomic Mass
•Some Elements and Their Isotopes
4.3
Average Atomic Mass
•To calculate the average
atomic mass of an element,
multiply the mass of each
isotope by its natural
abundance, expressed as a
decimal, and then add the
products.
4.3
Average Atomic Mass
•For example, carbon has two stable
isotopes:
− Carbon-12, which has a natural
abundance of 98.89%, and
− Carbon-13, which has a natural
abundance of 1.11%.
for Conceptual Problem 4.3
Conceptual Problem