Atomic Structure
Textbook Pages 4-5
Lesson Objectives
• Appreciate that there are various models to
illustrate atomic structure
• Know that early models of atomic structure
predicted that atoms and ions with noble gas
electron arrangements should be stable
• Describe the properties of protons, neutrons
and electrons in terms of relative charge and
relative mass
• Understand the importance of elementary
particles in the structure of the atom
• Be able to recall the meaning of mass number
(A) and atomic (proton) number (Z)
The History of the Atom
As early as 400 B.C., some
Greek philosophers proposed
that matter is made of
indivisible building blocks
known as atomos. (Atomos in
Greek means indivisible.) To
these early Greeks, matter
could not be continuously
broken down and divided
indefinitely. Rather, there was a
basic unit or building block that
was indivisible and
foundational to its structure.
• Boyle's studies (1661) of gaseous substances
promoted the idea that there were different types
of atoms known as elements. He suggested that
elements were types of substances that could not
be made simpler.
• Dalton (1803) conducted a variety of experiments
to show that different elements can combine in
fixed ratios of masses to form compounds. He
suggested that elements were composed of
indivisible atoms. All the atoms had the same
mass.
• Henri Becquerel discovered radioactivity (1896).
This showed that particles could come from inside
the atom.
1897 English scientist J.J.
Thomson's cathode ray
experiments led to the
discovery of the negatively
charged electron and the
first ideas of the structure of
these indivisible atoms.
Thomson proposed the Plum
Pudding Model, suggesting
that an atom's structure
resembles the favourite
English dessert - plum
pudding. The raisins
dispersed amidst the plum
pudding are analogous to
negatively charged
electrons immersed in a sea
of positive charge.
1911, Ernest Rutherford's famous
gold foil experiments led to the
nuclear model of atomic
structure. Rutherford's model
suggested that the atom
consisted of a densely packed
core of positive charge known
as the nucleus surrounded by
negatively charged electrons.
While the nucleus was unique
to the Rutherford atom, even
more surprising was the
proposal that an atom
consisted mostly of empty
space. Most of the mass was
packed into the nucleus that
was abnormally small
compared to the actual size of
the atom.
Neils Bohr improved upon
Rutherford's nuclear model
(1913) by explaining that the
electrons were present in
orbits outside the nucleus.
The electrons were confined
to specific orbits of fixed
radius, each characterized
by their own discrete levels
of energy. While electrons
could be forced from one
orbit to another orbit, it
could never occupy the
space between orbits.
• 1916 Gilbert Lewis proposed that the inertness of
noble gases was down to having a full outer shell
of 8 electrons. Ions were formed by atoms gaining
or losing electrons to attain full outer shells and
thus become stable and that atoms could also
bond by sharing electrons to form full outer shells.
• 1926 Erwin Schrodinger devised an equation
incorporating the idea that electrons had some
of the properties of waves as well as those of
particles
• 1932 James Chadwick discovered the neutron
Bohr's view of
quantized energy
levels was the
precursor to modern
quantum mechanical
views of the atoms.
Quantum mechanics
suggests that an
atom is composed of
a variety of
subatomic particles.
The modern idea of what an
atom looks like (or what the
exam board want you to
describe it as)
The properties of subatomic
particles
Extra notes
• A nucleon is a subatomic particle that is found in the
nucleus (i.e. a proton or neutron)
• The force of attraction between the electrons and the
protons is known as an electrostatic force
• Although you might expect the protons to repel each
other as they are all positively charged, they are held
together by the strong nuclear force
• We use relative masses and charges to describe
subatomic particles as otherwise the numbers are
ridiculously small