Chapter 5: Atomic
Structure and the
Periodic Table
Early Models of Atoms
• Democritus (460-400B.C.) first suggested the existence of
these particles, which he called “atoms” for the Greek word
for “uncuttable”. They lacked experimental support due to the
lack of scientific testing at the time.
• John Dalton (1766-1844) performed experiments to study the
ratios in which elements combine in chemical reactions.
Formulate hypotheses and theories to explain his
observations, which became Dalton’s Atomic Theory.
– All elements are composed of tiny indivisible particles called
atoms.
– Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
– Atoms of different elements can physically mix together or
combine in simple, 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.
Size of an Atom
• Imagine grinding a copper coin (penny) into fine
dust. Each speck in the small pile of shiny red
dust would still have the properties of copper. If
by some means you could still make the dust
particles smaller you would eventually come upon
a particle known as an atom.
• An atom is the smallest particle of an element
that retains the properties of that element.
• A pure copper penny contains about 2.4 X 1022
atoms, compared to the Earth’s population of 6 X
106 people.
• If you lined 100,000,000 copper atoms up side by
side they would produce a line 1 cm long.
Atomic Structure
• Atoms are now known to be divisible as they can be broken
down to even smaller particles by atom smashers.
• J.J. Thomson (1856-1940) discovered electrons using
cathode ray tubes.
• Robert Millikan (1868-1953) carried out experiments to
determine the charge of an electron (-). He also determined
the ratio of the charge to the mass of an electron.
• In 1886, E. Goldstein observed a cathode ray tube and
found rays traveling in the opposite direction to that of the
cathode rays. He called these rays canal rays and
concluded that they must be positive particles, which are
now called protons.
• In 1932, James Chadwick confirmed the existence of yet
another subatomic particle: the neutron. Neutrons are
subatomic particles with no charge but with a mass nearly
equal to that of a proton. See simulation
• After discovering these subatomic particles,
scientists wondered how they were put together.
• JJ Thompson thought since the electrons
contributed such a small fraction of the atoms
mass, they were probably an equal fraction of it
size so it was like “Plum Pudding”.
• In 1911, Ernest Rutherford and his coworkers
performed the Gold Foil Experiment to further
study the phenomenon.
• Concluded that most of the mass of each atom
and all of its positive charge reside in a very
small, extremely dense region which is called
the nucleus. The rest of the atom is mostly
empty space.
Modern View of Atomic Structure
• Since the time of Rutherford, physicists have learned much
about the nucleus. Although many other parts have been
discovered, chemists tend to only work with three main
particles since they determine chemical behavior: Electron,
Neutron and Proton
• Electron has a charge of -1.602 X 10-19 C and a proton has a
charge of 1.602 X 10-19 C so this quantity of Coulombs is
known as one electronic charge and atomic and subatomic
particles usually have a charge that is multiples of this.
Neutrons have no charge and are electrically neutral.
• Atoms have extremely small masses so instead of using
the real numbers, atomic mass units (amus) are used.
Protons and neutrons are very similar in mass but it would
take 1836 electrons to equal 1 proton so most of an atoms
mass is in the nucleus.
• Atoms are also extremely small with diameters between 1
X 10-10 and 5 X 10-10 so they are usually expressed with
angstroms, which is 10-10.
Atomic Number
• The number of protons in the nucleus
of an atom of that element
• For an atom with no charge, this is also
the number of electrons since the
positive charge of the protons cancels
the negative charge of the electrons.
• Practice problems #7-8 pg 115
Mass Number
• Most of the mass of an atom is found in the
nucleus so the total number of protons and
neutrons equals the mass number.
• If you know the atomic umber and mass number
you can determine the composition of that atom.
• The composition can be represented by the
shorthand notation using the element symbol,
atomic number and mass number.
• For gold, Au is the symbol for the element and
the atomic number is subscript and mass
number is superscript on the left side.
197
Au
79
• Practice problems 9-11 pg 116
Isotopes
• Atoms that have the same number of
protons but different number of
neutrons.
• Affects the shorthand notation of the
element.
• Practice problems 12-13 on pg 117
Atomic Mass
• Today we can determine the masses of
individual atoms with a relative high degree of
accuracy but since they are so small atomic
mass units are used with hydrogen being 1
amu.
• The average atomic mass for an element due
to the different isotopes, the mass of those
isotopes and the natural percent abundance.
• Add up the different atomic mass of each atom
and then divide by the number of atoms.
• Or, multiply mass by % and then determine
average mass.
• Practice problems 14-15 pg 120 and 16-17 pg
121.
Mass Spectrometer
• The most direct and accurate means for
determining atomic and molecular
weights. See pg 48
Periodic Table
• The arrangement of elements in order of
increasing atomic number, with elements
having similar properties placed in vertical
column.
• Atomic number, symbol, name, atomic
weight are found in each square for each
element. Some tables have additional
information as well. Example
• Can be arranged according to metals, nonmetals and metalloids, solid liquid and
gases, and by family. Example