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Atomic Theory:
History and Structure
Atomos is Greek for ‘indivisible’
Ancient Greeks
• Democritus (~460-370
BC) and Leucippus, his
teacher, believed that
everything is composed of
"atoms“ that are physically
indivisible,
indestructible, and are
continuously in motion
with empty space
between atoms; that there
are an infinite number
and kinds of atoms, which
differ in shape and size.
Aristotle (384-322 BC)
The heavens were
made of a fifth
element, ‘aether’.
• Aristotle believed a piece of
matter could be divided an
infinite number of times.
• He also believed that matter
was made up of 4 elements
which had 4 associated
qualities (e.g., fire was hot
and dry)
• The two forces of gravity
(the tendency for earth and
water to sink) and levity (the
tendency for air and fire to
rise) acted upon elements.
Particulate Nature of Matter
In the early 18th Century, evidence
accumulated for the view that matter was
made of particles, which led to new
models:
• Robert Boyle postulated that gases are
composed of discrete particles separated
by a void
• In 1704 Sir Isaac Newton proposed a
mechanical model of the universe, with
small solid masses in motion.
Chemical Laws of the 18th Century
Near the end of the 18th Century, three laws emerged
that governed chemical reactions.
• Antoine Lavoisier formulated the law of
conservation of mass in 1789 which
states that the total mass in a chemical
reaction remains constant (that is, the
reactants have the same mass as the
products).
• The law of definite proportions was
proven by the French chemist Joseph
Louis Proust in 1799. This law states
that if a compound is broken down into
its constituent elements, then the masses
of the constituents will always have the
same proportions, no matter what the
source was or the quantity of the original
substance.
John Dalton (1766 -1844)
• Dalton, an English
chemist, meteorologist,
and physicist, combined
these three
laws (laws of
conservation of
mass, definite
proportions, and multiple
proportions) into a single
explanation.
Dalton turned Democritus’ idea of the atom into a
scientific theory that could be tested by experiment by
relating atoms to the measurable property of mass.
Dalton’s Atomic Theory
1. All matter is composed of extremely
small particles called atoms.
2. Atoms of a given element are
identical in size, mass and other
properties; atoms of different
elements differ in size, mass, and
other properties.
3. Atoms cannot be subdivided, created,
or destroyed.
Dalton’s Atomic Theory (cont.)
4. Atoms of different elements
combine in simple whole-number
ratios to form chemical compounds.
5. In chemical reactions, atoms are
combined, separated, or
rearranged.
Dalton’s Contributions
• In 1808 Dalton published a textbook A New
System of Chemical Philosophy in which he
explained his method for estimating relative
atomic weights. His method was to compare
the mass ratios in which they combined, with
the hydrogen atom taken as unity.
• Unfortunately, Dalton did not understand that
some elements exist as diatomic molecules —
such as oxygen which occurs as O2. He also
mistakenly believed that the simplest
compound between any two elements is
always one atom of each (so he thought water
was HO, not H2O).
J.J. Thomson
• Atoms were
thought to be the
smallest possible
division of matter
until 1897 when
J.J. Thomson
discovered the
electron (which he
called a
‘corpuscle’)
through his work
on cathode rays.
Cathode-Ray Tubes
• A cathode-ray tube (CRT) is a glass tube
filled with some gas at low pressure
through which electric current is passed.
The surface of the tube directly opposite the
cathode glowed, it was theorized, due to a
stream of particles called a cathode ray which
traveled from the cathode to the anode.
CRT Experiments
• Experiments with cathode-ray tubes
showed that
1. Cathode rays are deflected by a magnetic
field just as a wire carrying an electric current
is (which they knew had a negative charge)
2. Cathode rays were deflected AWAY from a
negatively charged object.
• This led to the hypothesis that cathode
rays were made of negatively charged
particles.
Thomson’s Discovery
• Thomson’s experiments measured the ratio
of the charge of the CRT particles
compared to their mass. He learned that
this ratio was the same, regardless of which
metal he used to make the cathode or which
gas filled the tube.
• Thomson concluded that all cathode rays
are made of the same negatively charged
particles.
Electron Charge and Mass
• Furthermore, these negatively charged
particles must be present in atoms of ALL
elements and so the atom must be divisible.
• Experiments by Robert Millikan in 1909
measured the charge of this newly
discovered particle. Since Thomson had
already found the charge-to-mass ratio, now
the electron’s mass could be determined.
• The mass of an electron is 9.109 x 10-31 kg
which is 1/1837 the mass of a hydrogen
atom.
Thomson’s Atomic Model
• Scientists needed to explain (1) how
atoms are electrically neutral and
(2) since the mass of an electron is
so small compared to an atom,
where was the rest of the mass of
atom?
• Thomson proposed that the
electrons were set in a uniform sea
of positive charge (although they
could move about). This was
known as the ‘plum pudding
model.’ (small fruit pieces among the cake)
Discovery of the Nucleus
• Ernest Rutherford was a
former student of
Thomson who in 1909
performed an
experiment with
colleagues that showed
that most of the mass
and the positive
charge of an atom is
contained in a very
small volume.
Gold Foil Experiment
• A beam of alpha particles (helium
nuclei consisting of 2 protons and 2
neutrons) was shot at a thin piece
of gold foil. If Thomson’s plum
pudding model was correct, the
alpha particles were expected to
pass through the foil without being
deflected, as their momentum
would overcome both the small
mass of the electrons and the
repulsion of the uniformly
Expected result of experiment
distributed positive charge.
based on ‘plum pudding’ model
Unexpected Result
• To their surprise, about 1
in 800 alpha particles was
deflected at very large
angles!
• Rutherford’s famous
quote about this is, it was
“as if you had fired a 15inch [artillery] shell at a
piece of tissue paper and
it came back and hit you.”
Rutherford’s Planetary Model
of the Nucleus
If the nucleus were
the size of a marble,
the size of the atom
would be about the
size of a football field.
• Rutherford then proposed a
planetary atomic model in
which a cloud of electrons
surrounded a small,
compact nucleus of
positive charge. Only such
a concentration of charge
could produce an electric
field strong enough to cause
the deflection of the alpha
particle beam.
Atoms are Divisible!
• Nowadays we define an atom as the smallest
particle of an element that retains the
chemical properties of that element.
• We view atoms as consisting of two regions: the
nucleus is very small and located at the
center of an atom, consisting of one or more
protons and usually neutrons, while electrons
occupy a very large region surrounding the
nucleus.
• Protons, neutrons and electrons are referred to
as SUBATOMIC PARTICLES.
To be continued…
• Meanwhile, back at the ranch . . . .
So, what’s in the nucleus?
• We know now that, except for the simplest type
of hydrogen atom, all atomic nuclei are made
of two kinds of particles, PROTONS (p+) and
NEUTRONS (n°).
• However, neutrons were not discovered until
1932, 21 years after Rutherford’s discovery of
the nucleus!
• For the purposes of chemistry, other subatomic
particles besides protons, neutrons, and
electrons (e-) don’t play a part in chemical
properties, so we’re going to ignore them.
Composition of the Nucleus
Just to complete the set…
ELECTRONS e- or −𝟏𝟎𝒆
• Each electron has a negative charge
• Each electron has a mass equal to
9.109 x 10-31 kg
Atoms are Electrically Neutral
• This means that every atom has equal
numbers of protons and electrons
(neutrons have no charge).
• The nuclei of atoms of different
elements differ in the number of protons
they have and therefore the amount of
positive charge they possess.
• The number of protons in an atom
determines an element’s identity and is
known as its ATOMIC NUMBER.
How big are atoms?
• It is convenient to think of electrons moving
rapidly outside the nucleus in energy regions
that can be imagined as ‘clouds’ or shells of
negative charge.
• The radius of an atom is the distance from
the center of the nucleus to the outer portion
of this electron cloud.
• Atomic radii range from about 40 to 270 pm
(that’s picometers or 10-12 m) while nuclei of
atoms are much smaller, about 0.001 pm.
Isotopes
• Isotopes are atoms of an element
that have the same number of
protons but different numbers of
neutrons. Because these atoms have
the same number of electrons as all
other atoms of that element, they
behave similarly in chemical reactions.
• Almost all elements have more than
one isotope. Xenon has 9 and tin has
10. Fluorine, however, has only one.
Isotropes vs. Allotropes
• Isotopes are NOT different forms of an
element. For instance, carbon occurs as
graphite, diamond, carbon nanotubes, and
buckminsterfullerene which all have very
different properties. These are
ALLOTROPES of carbon, not isotopes
Hydrogen’s Isotopes
MASS NUMBERS OF HYDROGEN ISOTOPES
Atomic
number
(# of protons)
# of
neutrons
Mass number
(protons + neutrons)
% of
hydrogen
atoms
Protium
1
0
1+0=1
99.9885
Deuterium
1
1
1+1=2
0.0115
Tritium
1
2
1+2=3
Very rare
Tritium is radioactive. It exists in very small amounts
in nature, but it can be prepared artificially.
Designating Isotopes
• The isotopes of hydrogen have distinct
names, although isotopes are usually
identified by their mass number.
• There are two methods for specifying
isotopes that DON’T have names:
1. The mass number is written with a
hyphen after the name of the element
(tritium is hydrogen-3). This is
HYPHEN NOTATION.
Carbon’s Isotopes
• Carbon has 3 naturally occurring isotopes:
Carbon-12 (6 protons, 6 neutrons)
Carbon-13 (6 protons, 7 neutrons)
Carbon-14 (6 protons, 8 neutrons)
• Carbon-14 is unstable and undergoes
radioactive decay. This rate of decay is
known and is used to find the ages of
fossils, rocks and minerals.
• Other elements have unstable isotopes
which are also used to ‘date’ fossils, rocks
and minerals like potassium and uranium.
Designating Isotopes (cont.)
ISOTOPES OF HYDROGEN AND HELIUM
Isotope
Nuclear
Symbol
# protons
# electrons
#neutrons
Hydrogen-1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Helium-3
2
2
1
Helium-4
2
2
2
Relative Atomic Masses
• Because the masses of atoms expressed
in grams are very small, it is more
convenient to use relative atomic masses.
• Scientists have chosen the carbon-12
atom to be the standard atom for this
purpose, and it has been assigned a mass
of EXACTLY 12 atomic mass units or
amu. One amu is 1/12 of the mass of a
carbon-12 atom.
Average Atomic Masses
• Most elements occur naturally as a mixture
of isotopes with the percentage of each
isotope being nearly always the same,
regardless of where on Earth the element is
found.
• AVERAGE ATOMIC MASS is the
weighted average of the atomic masses
of the naturally-occurring isotopes of an
element.
Calculating a Weighted Average

Example problem
Problems with Rutherford’s
Planetary Model
• Remember Rutherford’s planetary model?
(a cloud of electrons, like planets, orbiting a
small, compact nucleus of positive charge).
Unfortunately, when electrical charge is
involved, it gets more complicated.
• First, planets do not have electric charge,
and orbiting charged particles would lose
energy and spiral into the nucleus in less
than a second.
• Secondly, the planetary model had no
way of explaining the highly peaked
emission and absorption spectra that
were observed for atoms.
Wavelength (nm)
Quantum Theory to the Rescue
• At the beginning of the 20th Century, Max Planck
and Albert Einstein postulated that light energy is
emitted and absorbed in discrete (distinct)
amounts called quanta which we know as
photons.
• Quantum mechanics is important for interactions
on the scale (size) of atoms.
The Bohr Model
• In 1913 Niels Bohr, with
Rutherford, incorporated the
idea of quanta into his
model of the atom, in which
an electron could only orbit
the nucleus in particular
circular orbits which
depended on the electron’s
energy and angular
momentum. This meant an
atom had fixed energy
levels for its electrons.
Bohr Model
Photons have
energies according to
this formula, E=hf,
where h is Planck’s
constant
(6.626 × 10-34 m2-kg/s)
and f is the frequency
of the light in hertz.
Bohr Model
• In this model an electron could not spiral into
the nucleus because it could not lose energy
in a continuous manner. Electrons could
only make instantaneous "quantum
leaps" between fixed energy levels, and
when this happened, light was emitted or
absorbed at a frequency that depended
upon this change in energy.
• This model successfully explained the
observed emission and absorption spectra of
light.
More Modern Theories
• Bohr’s model could only partially explain the
hydrogen atom, but it is still useful in explaining
quantized energy levels.
• However, the modern model of the atom
describes the positions of electrons in an
atom in terms of probabilities. An electron can
potentially be found at any distance from the
nucleus, but, depending on its energy level,
exists more frequently in certain regions
around the nucleus than others.
• This pattern is referred to as its atomic orbital.
Orbitals come in a variety of shapes—sphere,
dumbbell, torus, etc.—with the nucleus in the
middle.
Atoms have the same
number of protons and
electrons and so are
electrically neutral.
The electron clouds are
made up of different
energy levels. Each
energy level can only
hold a certain number of
electrons.
The electrons in the
outermost level can
interact with other atoms’
electrons and form
chemical compounds.
How many electrons in each level?
• Electrons fill the energy levels from
lowest to highest.
–The lowest level (Period 1) can hold
only 2 e–The 2nd level (Period 2) can hold 8 e–The 3rd level (Period 3) can hold 18 e–The 4th level (Period 4) can hold 32 e-
Electron orbital ring notation
Boron-10
Oxygen-16
Which element is this?
Orbitals and Quantum Numbers
• In order to completely describe orbits,
scientists use quantum numbers which
specify the properties of atomic orbitals
and the properties of electrons in those
orbitals.
• The principal quantum number n indicates
the main energy level occupied by the
electron. Values of n are positive integers
only. The lowest energy level is n = 1 and
is closest to the nucleus. Each higher value
of n is successively further out.
Orbital Shapes
• Except at the first main energy level,
orbitals of different shapes—known as
sublevels—are defined by the second
quantum number, the angular momentum
quantum number l.
ORBITAL LETTER DESIGNATIONS AND SHAPES
Value of l
Letter
designation
Shape of orbital
0
s
Spherical
1
p
Dumbbell
2
d
Dumbbell, torus,
etc.
3
f
Various
s, p, d and f orbital shapes
How many electrons can the f orbitals hold?
QUANTUM NUMBER RELATIONSHIPS IN
ATOMIC STRUCTURE
Principal
quantum
number
Sublevels in
main energy
level
Number of
electrons per
sublevel
Number of
electrons per
main energy
level
1
s
2
2
2
s
p
2
6
8
3
s
p
d
2
6
10
18
4
s
p
d
f
2
6
10
14
32
Electron Orbitals Gallery
http://en.wikipedia.org/wiki/Atomic_orbital_model
Wait a minute…
• Did the pattern in the last slide remind you
of anything else?
• Which orbital shape corresponds to the
blue ones?
• Which orbital shape corresponds to the
yellow ones?
• Which orbital shape corresponds to the
red ones?
• Which orbital shape corresponds to the
green ones?
Sublevel Blocks of the Periodic Table
Electron Configurations
• The quantum model improves on the Bohr
model because it describes the
arrangements of electrons in atoms other
than hydrogen. The arrangement of
electrons in an atom is known as the
atom’s electron configuration.
• Because atoms of different elements have
different numbers of electrons, a unique
electron configuration exists for the atoms
of each element.
• Like all systems in nature, electrons in
atoms tend to assume arrangements that
have the lowest possible energies.
• The lowest energy arrangement of electrons
for each element is called the element’s
ground-state electron configuration.
• To build up electron configurations for the
ground state of a particular atom, the energy
levels of the orbitals are found, then
electrons are added to the orbitals one by
one according to three basic rules.
Rule #1: Aufbau Principle
• This rule shows the order in which electrons
occupy orbitals.
An electron occupies the lowestenergy orbital that can receive it.
• The orbital with the lowest energy is the 1s
orbital. The 2s orbital is the next highest in
energy, then the 2p orbitals.
• Beginning with the third main energy level
(n = 3), the energies of the sublevels in
different main energy levels begin to overlap.
Sublevels in Order of Increasing Energy
Notice that the 4s energy
sublevel is lower than the 3d
energy sublevel, which is lower
than the 4p sublevel.
The same pattern holds for the
5s, 4d and 5p sublevels.
Things are even more
complicated for higher energy
sublevels. Notice that the 6s
sublevel is lower than the 4f
sublevel, which is lower than the
5d sublevel, which is lower than
the 6p sublevel.
Rule #2: The Pauli Exclusion Principle
No two electrons in the same atom can
have the same set of 4 quantum numbers.
• In addition to the two quantum numbers
already mentioned, there are two more: the
magnetic quantum number (which indicates
the orientation of an orbital around the
nucleus) and the spin quantum number
(which explains why a single orbital can
hold a maximum of two electrons with
opposite spin states).
Pauli Exclusion Principle
1s orbital for
helium
This is referred to as
an electron “pair”.
Rule #3: Hund’s Rule
• As many unpaired electrons as possible
are placed in separate orbitals in the same
sublevel.
Orbitals of equal energy are each
occupied by one electron before any
orbital is occupied by a second electron,
and all electrons in singly occupied
orbitals must have the same spin state.
Hund’s Rule
CORRECT
2s
2p
INCORRECT
2s
2p
Hund’s Rule
CORRECT
2s
2p
INCORRECT
2s
2p
Orbital Notation
• In orbital notation, an unoccupied orbital is
represented by a line, ____, with the orbital’s
name written underneath the line.
• An orbital containing one electron is written with
one arrow, shown as _ ↑ _.
• An orbital containing two electrons is written as
_↑↓ _, showing the electrons paired and with
opposite spin states.
• The lines are labeled with the element’s symbol,
the principal quantum number and sublevel letter.
H ↑_
He _↑↓ _
1s
1s
Electron Configuration Notation
No lines, no arrows—instead the number of
electrons in a sublevel is shown by adding
a superscript to the sublevel designation.
Hydrogen’s configuration is 1s1. The
superscript shows that one electron is
present in hydrogen’s 1s orbital.
Helium’s configuration is 1s2. The
superscript shows that there are two
electrons in helium’s 1s orbital.
Sample Problem
• The electron configuration of boron is
1s22s22p1. How many electrons are
present in an atom of boron? What is the
atomic number of boron? Write the orbital
notation for boron.
First draw the lines representing orbitals:
Next, add arrows showing the electron locations,
starting with the lowest energy levels:
Your turn
• The electron configuration of nitrogen is
1s22s22p3. How many electrons are
present in an atom of nitrogen? What is
the atomic number of nitrogen? Write the
orbital notation for nitrogen.
7 electrons, 7, _↑↓ _ _↑↓ _ _↑_ _↑_ _↑_
1s
2s
2p
ASPEN HILTON WS
Order of Fill of Orbitals
• Would you like a simplerto-write-and-remember
guide to filling orbitals
than the energy diagram
you colored above?
• This should help! Start
at the end of the top
arrow, continue down the
arrow to its tip then drop
down a level to the end
of the next arrow, etc.