Week 1
Chapter 2: A Particle View of
Matter
• Because atoms are too small to see, even with
the most powerful microscopes, chemists
used models to represent them.
• These models have changed with time as
more is discovered about atoms.
• In 1897, Thomson took the latest
experimental evidence for the existence of
positive matter and negative particles and
proposed an atomic model.
Thomson’s Model
• He imagined that the atom consisted of a
sphere of positive matter in which negative
particles were randomly embedded, like
plums in a pudding…thus becoming the “Plum
Pudding” model.
Ernest Rutherford
• Rutherford completed experiments that
produced results that could not be explained
by the “Plum Pudding” model.
• His work bought us even closer to the atomic
model that we use today.
• He proposed the most of the mass of an atom
and all of the positive charge was located in
the centre of the atom, the nucleus.
Rutherford continued…
• He also said the most of the volume of an
atom is empty space occupied by the
electrons.
• And, the force of attraction between the
positive nucleus and negative electrons is
electrostatic.
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James Chadwick’s Discovery
• In 1932, James Chadwick discovered an
uncharged particle that had a mass a little
greater than the proton.
• He called this particle a neutron and reasoned
that they were held in the nucleus with the
protons.
• With this discovery, scientists had discovered
three subatomic particles, making it possible
to identify individual atoms.
Outcomes of Chadwick’s
Discovery…
• Chadwick’s discovery explained a problem
that Rutherford had faced earlier.
• When he was working with Thorium,
Rutherford identified a radioactive material
chemically similar to Thorium. He called this
Thorium-X.
• Other chemists also found similar substances
for the elements in the periodic table between
Pb and U.
Explaining Isotopes
• British scientist Frederick Soddy decided to
place similar substances in the same position
in the periodic table.
• He suggested that varieties of particular
elements existed and called them isotopes.
• Isotopes = Greek for same place.
• With the discovery of neutrons, Chadwick
could now explain isotopes.
• They had the same number of protons and
electrons but a varying number of neutrons.
• Rutherford’s model proposed that the
electrons move in circular orbits around the
nucleus with an electrostatic force of
attraction operating between positive and
negative particles.
Conflicting Ideas…
• While Rutherford’s model worked for many aspects
of an atom, there were some limitations relating to
the arrangement of electrons around the nucleus:
– The laws of physics that had been used to explain the
behaviour of large objects to indicate that electrons
moving in a circular orbit should continually emit
electromagnetic radiation (light).
– As the radiation is emitted, electrons should lose energy
and spiral into the nucleus.
– The model did not explain why, when an element is
heated, only light of certain specific energies is emitted,
and not the light of every energy.
Refining Rutherford’s Model
• In 1913, Niels Bohr, suggested that the laws of
physics that scientists used to explain the
behaviour of large objects did not apply to the
motion of small objects, such as electrons.
• He proposed the electrons:
– Circle the nucleus without losing energy
– Could move in certain fixed orbits of particular
energies.
Energy and Heating
• The orbit that an electron moves depends on
the energy of the electron. A low energy orbits
closer to the nucleus while a high energy
orbits further out.
• Heating an element can result in an electron
absorbing energy and ‘jumping’ to a higher
energy state.
• After this the electron emits this excess
energy, as light and returns to its original
state.
Extending Bohr’s Model
• Scientists quickly extended on Bohr’s ideas.
Evidence for the existence of energy levels in
other atoms were obtained from studies of
successive ionisation energies of elements.
• For example, the ionisation energy for the
removal of the first electron is fairly low. The
next eight require more energy but are all
removed at similar levels.
Extending Bohr’s Model cont…
• This conclusion of successive ionisation
energies led scientists to conclude that
electrons were grouped in different energy
levels, called shells.
• Electrons in the same shell:
– Are about the same distance from the nucleus
– Have about the same energy.
• This means that a shell can be described
based on the amount of energy needed to
move an electron out of this region.
Extending Bohr’s Model cont…
• Each energy level or shell, can hold different
numbers of electrons.
Shell Number
Maximum Number of
Electrons
1
2
2
8
3
18
4
32
n
2n2
Electronic Configuration
• The way in which electrons are arranged is
known as the electronic configuration of an
atom.
• In a stable atom the electrons are as close to
the nucleus as possible, meaning that an
electron fills up from the inside out.
• After the second shell, the electron fills in
stages. The outermost shell can never contain
more than 8 electrons, regardless of its ability
to hold more than that.
Electronic Configuration cont…
• For example, despite the fact the the third
shell can hold 18 electron, it will not hold
more than 8 unless the fourth shell has at
least two electrons in it.
• The electrons in the outermost shell of an
atom are known as the valence electrons.
• These are the ones that require the least
amount of energy to be removed.
• It is these electrons that are involved in
chemical reactions.
• In 1926, Erwin Schrodinger proposed that
electrons behaved as waves around the
nucleus.
• His description is known as wave mechanics or
quantum mechanics.
• Scientists now believe that small particles
such as protons, electrons and neutrons
behave according to the laws of quantum
mechanics.
Quantum Mechanics
• The electron is now seen as a vague object
that behaves like a cloud of negative charge.
• Unlike the Bohr model, there are no definite
orbits for electrons.
• They are instead thought to move in regions
of space, called orbitals.
Shells and Subshells
• Quantum mechanics shows the following:
– Within an atom there are major energy level
which are known as shells. These are numbered 1,
2, 3 etc.
– Within these shells are energy levels of similar
energy called subshells. These are lettered as
s,p,d,f
– The energy of the subshells increases in this order
– Subshells are made up of orbitals.
Energy Levels within an Atom…
Shell Number
Number of
Subshells
Subshell in
Shell
Number of
Orbitals in
Subshell
Maximum
Number of
Electrons
1
1
1s
1
2
2
2
2s
2p
1
3
2
6
3
3
3s
3p
3d
1
3
5
2
6
10
4
4
4s
4p
4d
4f
1
3
5
7
2
6
10
14
5
5
5s
5p
5d
5f
5g
1
3
5
7
9
2
6
10
14
18
Energy Levels