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Elements – different types of atom
Elements are the simplest substances. There are about
100 different elements.
Each element is made up of very tiny particles called atoms,
and each element is made up of just one particular type of
atom, which is different to the atoms in any other element.
Gold is an element
made up of only
gold atoms.
Carbon is an element
made up of only
carbon atoms.
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Atoms – the building blocks
John Dalton had the first
ideas about the existence of
atoms over 200 years ago.
However, it is only relatively
recently that special microscopes
(called electron microscopes)
have been invented that can
actually ‘see’ atoms.
This image is highly magnified.
What could it be showing?
The yellow blobs are individual
gold atoms, as seen through
an electron microscope.
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Dalton’s Atomic Theory:
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How was atomic structure discovered?
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What particles are atoms made of?
For some time, people thought that atoms were the
smallest particles and could not be broken into anything
smaller.
Scientists now know that atoms are actually made from
even smaller particles. There are three types:
proton
neutron
electron
How are these particles arranged inside the atom?
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How did our understanding change?
J.J Thomson discovered electrons, and proposed the
existence of a (+) particle.
It wasn’t until Rutherford’s famous
gold foil experiment that the + proton
was discovered, and atoms were
thought to me mostly empty space. He
named the centre of atoms the
nucleus.
Bohr improved on this model proposing
that electrons move around the nucleus
in specific layers called shells.
It was James Chadwick who discovered particles
with no charge, which he named neutrons.
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What is the structure of an atom?
Protons, neutrons and electrons are not evenly distributed in
an atom.
The protons and neutrons
exist in a dense core at the
centre of the atom. This is
called the nucleus.
The electrons are
spread out around the
edge of the atom. They
orbit the nucleus in
layers called shells.
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Labelling the atom
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Mass and electrical charge
There are two properties of protons, neutrons and electrons
that are especially important:
 mass
 electrical charge.
Particle
Mass
Charge
proton
1
+1
neutron
1
0
electron
almost 0
-1
The atoms of an element contain equal numbers of protons
and electrons and so have no overall charge.
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How many protons?
The atoms of any particular element always contain the
same number of protons. For example:
 hydrogen atoms always contain 1 proton
 carbon atoms always contain 6 protons
 magnesium atoms always contain 12 protons.
The number of protons in an atom
is known as the atomic number
or proton number.
It is the smaller of the two numbers
shown in most periodic tables.
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What is the atomic number?
What are the atomic numbers of these elements?
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sodium
11
iron
26
tin
50
fluorine
9
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What is mass number?
Electrons have a mass of almost zero, which means that the
mass of each atom results almost entirely from the number
of protons and neutrons in the nucleus.
The sum of the protons and neutrons
in an atom’s nucleus is the mass number.
It is the larger of the two numbers shown
in most periodic tables.
Atoms
Protons
Neutrons
Mass
number
hydrogen
1
0
1
lithium
3
4
7
aluminium
13
14
27
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What’s the mass number?
mass number = number of protons + number of neutrons
What is the mass number of these atoms?
Atoms
Protons
Neutrons
Mass
number
helium
2
2
4
copper
29
35
64
cobalt
27
32
59
iodine
53
74
127
germanium
32
41
73
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How many neutrons?
number of neutrons = mass number - number of protons
= mass number - atomic number
How many neutrons are there in these atoms?
Atoms
Mass
number
Atomic
number
Neutrons
helium
4
2
2
fluorine
19
9
10
strontium
88
38
50
zirconium
91
40
51
uranium
238
92
146
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Building a nucleus
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How many electrons?
Atoms have no overall electrical charge and are neutral.
This means atoms must have an equal number of positive
protons and negative electrons.
The number
of electrons
is therefore
the same as
the atomic
number.
Atoms
Protons
Neutrons
Electrons
helium
2
2
2
copper
29
35
29
iodine
53
74
53
Atomic number is the number of protons rather than the
number of electrons, because atoms can lose or gain
electrons but do not normally lose or gain protons.
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From atoms to ions
How can reactive metal atoms become stable positive ions?
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How do atoms form ions?
An ion is an atom or group of atoms that has
an electrical charge, either positive or negative.
Atoms have an equal number of protons and electrons and
so do not have an overall charge.
Atoms with incomplete outer electron
shells are unstable. By either gaining or
losing electrons, atoms can obtain full
outer electron shells and become stable.
When this happens, atoms have an unequal number of
protons and electrons and so have an overall charge.
This is how atoms become ions.
How does an atom become a positive or negative ion?
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Positive and negative ions?
An atom that loses electrons has more protons
than electrons and so has a positive overall charge.
This is called a positive ion.
An atom that gains electrons has more electrons
than protons and so has a negative overall charge.
This is called a negative ion.
The electron configuration of an atom shows how many
electrons it must lose or gain to have a filled outer shell.
 Atoms with a nearly empty outer shell, will lose electrons
to obtain a full outer shell.
 Atoms with a nearly full outer shell, will gain electrons
to obtain a full outer shell.
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How do atoms form positive ions?
An atom that loses one or more electrons
forms a positive ion.
Metal atoms, such as sodium, magnesium
and iron, form positive ions.
Positive ions have a small ‘+’ symbol and a number by this
to indicate how many electrons have been lost.
This number is usually the same as the number of electrons
in the atom’s outer shell. For example:
lithium atom 2.1
lithium ion [ 2 ]
= Li+
magnesium atom 2.8.2
magnesium ion [ 2.8 ] = Mg2+
aluminium atom 2.8.3
aluminium ion [ 2.8 ] = Al3+
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How is a sodium ion formed?
Sodium atom:
11 protons
= +11
11 electrons = -11
Sodium ion:
11 protons
= +11
10 electrons = -10
Total charge =
Total charge =
0
+1
+
Na
loses
1 electron
2.8.1
(partially full outer shell)
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Na
[2.8]
(full outer shell)
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How is a magnesium ion formed?
Magnesium atom:
12 protons
= +12
12 electrons = -12
Magnesium ion:
12 protons
= +12
10 electrons = -10
Total charge =
Total charge =
0
+2
2+
Mg
loses
2 electrons
2.8.2
(partially full outer shell)
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Mg
[2.8]2+
(full outer shell)
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How do atoms form negative ions?
An atom that gains one or more electrons
forms a negative ion.
Non-metal atoms, such as chlorine,
oxygen and nitrogen, form negative ions.
Negative ions have a small ‘-’ symbol and a number by this
to indicate how many electrons have been gained to fill their
outer shell. For example:
chlorine atom 2.8.7
chloride ion [ 2.8.8 ] = Cl-
oxygen atom 2.6
oxide ion
[2]
= O2-
nitrogen atom 2.5
nitride ion
[2]
= N3-
The name of the ion is slightly different to the atom’s name.
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How is a fluoride ion formed?
Fluorine atom:
9 protons
= +9
9 electrons
= -9
Fluoride ion:
9 protons
= +9
10 electrons = -10
Total charge =
Total charge = -1
0
-
F
gains 1
electron
2.7
(partially full outer shell)
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F
[2.8](full outer shell)
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How is a sulfide ion formed?
Sulfur atom:
16 protons
= +16
16 electrons = -16
Sulfide ion:
16 protons
= +16
18 electrons = -18
Total charge =
Total charge = -2
0
2-
S
gains 2
electrons
2.8.6
(partially full outer shell)
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S
[2.8.8]2(full outer shell)
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Subatomic particles
Atoms are composed of three subatomic particles: protons,
neutrons and electrons. The two important properties of
these particles are mass and charge:
Particle
Relative
mass
Relative
charge
proton
1
+1
neutron
1
0
electron
1/1840
-1
The mass of electrons is negligible when compared to the
mass of protons and neutrons, so their mass is not included
when calculating the mass of the atom.
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Atomic number and mass number
The number of protons in an atom is known as the
atomic number or proton number and is represented
by the symbol Z.
The mass number of an atom is the number of protons
plus the number of neutrons, and is represented by the
symbol A.
When an atom is
represented by its
symbol, the mass
number, and
sometimes the atomic
number, are shown.
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mass
number (A)
atomic
number (Z)
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What are isotopes?
Isotopes are atoms of the same element that contain
different numbers of neutrons.
mass number
is different
atomic number
is the same
carbon-12
carbon-13
The reactivity of different isotopes of an element is
identical because they have the same number of electrons.
The different masses of the atoms means that physical
properties of isotopes are slightly different.
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Isotopes of chlorine
About 75% of naturally-occurring chlorine is chlorine-35 (35Cl)
and 25% is chlorine-37 (37Cl).
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17 protons
17 protons
18 neutrons
20 neutrons
17 electrons
17 electrons
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Isotopes of carbon
There is also more than one isotope of carbon:
Isotope
Protons
Neutrons
12C
6
6
13C
6
7
14C
6
8
All isotopes of carbon have 6 protons and so have 6
electrons.
Because chemical reactivity depends on the number of
electrons the reactivity of the isotopes of carbon is identical.
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What’s the number?
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Isotopes & Mass Spectrometry
The number of naturally occurring isotopes of an element
and their relative masses and abundances may be
determined by a mass spectrometer, an instrument based
on the same principles employed by JJ Thomson to
discover the electron.
Mass Spectrometry uses the fact that charged particles
moving through a magnetic field are deflected from their
original path based on their charge-to-mass ratio (e/m
or e/z)

The greater the ratio, the
more deflection occurs.
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‘Weighing’ atoms
Mass spectrometry is an accurate instrumental technique
used to determine the relative isotopic mass (mass of each
individual isotope relative to carbon-12) and the relative
abundance for each isotope. From this, the relative atomic
mass of the element can be calculated.
Some uses of mass
spectrometry include:
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
carbon-14 dating

detecting illegal drugs

forensic science

space exploration.
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Mass spectrometry
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Parts of the mass spectrometer
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Process of mass spectrometry
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Mass spectra of monatomic elements
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Mass spectra of diatomic elements
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Process of mass spectrometry
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What is relative atomic mass?
The relative atomic mass (Ar) of an element is the mass of
one of its atoms relative to 1/12 the mass of one atom of
carbon-12.
relative atomic mass
average mass of an atom × 12
=
(Ar)
mass of one atom of carbon-12
Most elements have more
than one isotope. The Ar of
the element is the average
mass of the isotopes taking
into account the abundance of
each isotope. This is why the
Ar of an element is frequently
not a whole number.
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Using mass spectra to calculate Ar
The mass spectrum of an element indicates the mass and
abundance of each isotope present. For example, the mass
spectrum of boron indicates two isotopes are present:
abundance (%)
100
11B
80
(80%)
60
10B
40
(20%)
20
0
0
2
4
6
m/z
8
10
12
How can this be used to calculate the Ar of boron?
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Calculating Ar
Most elements have more than one isotope. The relative
atomic mass of the element is the average mass of the
isotopes taking into account the abundance of each isotope.
Example: what is the Ar of boron?
In a sample of boron, 20% of the atoms are 10Br and 80%
are 11Br.
If there are 100 atoms, then 20 atoms would be 10Br and
80 atoms would be 11Br.
The relative atomic mass is calculated as follows:
Ar of Br = (20 × 10) + (80 × 11)
100
Ar of Br = 10.8
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Calculating Ar of magnesium
In a sample of magnesium, 79.0% of the magnesium
atoms are 24Mg, 10.0% are 25Mg and 11.0% are 26Mg.
Example: What is the Ar of magnesium?
1. Calculate mass × abundance
of each isotope
24 × 79.0
25 × 10.0
26 × 11.0
2. Add these values,
and divide by 100
(1896 + 250 + 286) / 100
Ar of Mg = 24.3
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Ar calculations
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How are atoms arranged?
Where are the electrons found in the atom?
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How are electrons arranged?
Electrons are not evenly spread but exist in layers called
shells. (The shells can also be called energy levels).
The arrangement of electrons in these shells is often
called the electron configuration.
1st shell
2nd shell
3rd shell
Note that this diagram is not drawn to scale – the atom is
mostly empty space. If the electron shells were the size
shown, the nucleus would be too small to see.
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How many electrons per shell?
Each shell has a maximum number of electrons that it can
hold. Electrons will fill the shells nearest the nucleus first.
1st shell holds
a maximum of
2 electrons
2nd shell holds
a maximum of
8 electrons
3rd shell holds
a maximum of
8 electrons
This electron arrangement is written as 2,8,8.
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Models of atoms
The model of the atom states that
a nucleus is surrounded by shells
of electrons. Each shell holds a
different maximum number of
electrons:

1st shell = 2 electrons

2nd shell = 8 electrons

3rd shell = 8 electrons.
At the Chemistry 11 level, this model is slightly different.
Instead of electrons being arranged in shells that are a
different distance from the nucleus, they are arranged in
energy levels.
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Waves behaving like particles? Well hit
me with a Planck!
A serious challenge to
Rutherford’s atomic model
arose almost immediately.
By the end of the 1800s the
physics available stated that
accelerating charges should
radiate energy.
This meant that atoms should collapse in a fraction of a
second as their electrons lost energy and spiralled around
the nucleus.
Rutherford’s conclusions were correct, the problems was
with the physics of the day…
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The Bohr model of the atom
In 1900, Max Planck (right)
developed his ‘Quantum
theory’, which states that
energy could be shown to
behave like particles in fixed
amounts he called quanta.
In 1913, Niels Bohr (left)
applied Plank’s theory to
electrons, and improved upon Rutherford’s model. He
proposed that electrons could only exist in fixed energy levels.
The main energy levels are called principal energy levels
and are given a number called the principal quantum
number (n) with the lowest in energy being 1.
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The Bohr Model
Each electron has a
fixed energy = an
energy level.
Electrons can jump from
one energy level to
another.
Electrons can not be or
exist between energy
levels.
A quantum of energy is the amount of energy needed to
move an electron from one energy level to another
energy level.
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Bhor’s experiments…
Bhor knew that when high
voltage was applied across
the electrodes of a sealed
glass tube containing a gas
such as hydrogen, the gas
was heated and emitted
light…he looked at this light
through a spectroscope.
He saw that the spectroscope separated the light into its
component wavelengths, and for hydrogen he saw a series
of coloured lines against a black background.
Bhor applied quantum principles to explain the bright-line
spectrum he saw for hydrogen.
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Quantum Theory rescues the Nuclear Model
To move from one level to
another, the electron must
gain or lose the right
amount of energy.
The higher the energy level,
the farther it is from the
nucleus.
• Gain energy to move to higher
energy levels (away from
nucleus)
• Lose energy to move to lower
energy levels (closer to nucleus)
The degree to which they move from level to level
determines the frequency of light they give off.
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Bhor proposed…
The amount of energy required to go from one
energy level to another is the not same for the
electrons.
Higher energy levels are closer together.
This means it takes less energy to change
levels in the higher energy levels.
Once in a higher energy orbit, any electron
could then return to a lower energy orbit by
emitting a specific amount of energy
corresponding to the energy difference. If the
frequency of that emitted energy corresponding
to any part of the visible spectrum, then a bright
line of that specific colour would be seen.
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An Explanation of Atomic Spectra
• The three groups of lines in the hydrogen
spectrum correspond to the transition of
electrons from higher energy levels to lower
energy levels.
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The shortcomings…
Bohr’s model of the hydrogen atom was successful in
explaining the mystery of bright line spectra.
His calculations and predictions worked for hydrogen and he
even calculated the radius of the orbit for hydrogen’s electron
in its ground state.
BUT his model failed to explain the energies absorbed
and emitted by atoms with more than one electron.
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If an excited electron emits energy and drops to n = 2 from a higher energy orbit,
the wavelength of the emitted energy corresponds to a particular colour of visible
light.
If an electron drops from n = 3 to n = 2, the energy difference between the two orbits
(and therefore the energy emitted) corresponds to that of red light. Hence the red line
appears in the emission spectrum. The blue-green line results from an electron
transition from n = 4 to n = 2, the blue line from an electron transition from n = 5 to n
= 2, and the violet line from an electron transition from n = 6 to n = 2.
This series of four bright lines in the visible spectrum is called the Balmer series,
named after the Swiss schoolteacher who first derived a mathematical relationship
between the lines in hydrogen’s visible emission spectrum.
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Did you know that an element can be identified by
its emission spectra?
When atoms absorb energy, electrons move
into higher energy levels. These electrons then
lose energy by emitting light when they return to
lower energy levels.
Mercury
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Nitrogen
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“Fingerprints”
of certain
atoms
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What is ionization energy?
Ionization is a process in which atoms lose or gain electrons
and become ions.
The first ionization (I1) energy of an element is the energy
required to remove one electron from a gaseous atom.
M(g) → M+(g) + e-(g)
The second ionization (I2) energy involves the removal of a
second electron:
M+(g) → M2+(g) + e-(g)
Looking at trends in ionization energies can reveal useful
evidence for the arrangement of electrons in atoms and ions.
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Ionization energy definitions
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ionization energy
Evidence for energy levels
Plotting the successive ionization energies of magnesium
clearly shows the existence of different energy levels, and
the number of electrons at each level.
6
Successive ionization
energies increase as more
5
electrons are removed.
Large jumps in the ionization
energy reveal where electrons
are being removed from the
next principal energy level,
such as between the 2nd and
3rd, and 10th and 11th ionization
energies for magnesium.
4
3
2
electron removed
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More evidence for energy levels
first ionization energy
(kJ mol-1)
The first ionization energies of group 2 elements also show
evidence for the existence of different principal energy levels.
Even though the nuclear
charge increases down the
group, the first ionization
energy decreases.
900
800
700
600
500
400
Be
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Mg
Ca
Sr
element
This means electrons are
being removed from
successively higher
energy levels, which lie
further from the nucleus
and are less attracted to
Ba
the nucleus.
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Trends in first ionization energies
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Energy levels
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Beyond Bhor: Particles Behaving Like Waves
The beginnings of quantum mechanics was in
1923 with Louis de Broglie: “If waves behaved
like particles, could it also be possible for
particles to have wave properties?”
• Applied waveparticle theory to
electrons
n=4
n=5
• electrons exhibit
wave properties
n=6
Forbidden
n = 3.3
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Heisenberg: Quantum Mechanics
• Treated the electron as a pure particle with
quantum behavior.
• Heisenberg Uncertainty Principle: impossible to
know exactly both where any particle was located
and where it was going as the same time (velocity)
•
Werner Heisenberg
~1926
Ie: It is impossible to describe precisely both the location and
the speed of particles that exhibit wavelike behavior.
g
Microscope
Electron
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The Quantum Mechanical Model
• Rutherford’s and Bohr’s model focused on describing the
path of the electron around the nucleus like a particle
(like a small baseball).
• Austrian physicist Erwin Schrödinger rejected the idea
of an electron as a particle, he treated the electron as
a wave.
The modern description of the electrons in atoms, the
quantum mechanical model, comes from the
mathematical solutions to the Schrödinger equation.
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Electrons as Waves
EVIDENCE: DIFFRACTION PATTERNS
VISIBLE LIGHT
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ELECTRONS
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The electron as a WAVE
Schrödinger’s wave equation
Used to determine the probability of finding
the H electron at any given distance from
the nucleus
Electron best described as a cloud
• Effectively covers all points at the same time
(fan blades)
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Quantum Mechanics
• Schrödinger Wave Equation (1926)
finite # of solutions  quantized energy
levels which defines probability of finding
an electron
Ψ1s 

Erwin Schrödinger
~1926
1 Z 3/2 σ
π a0
e
The regions in 3D space around the nucleus in which
there is the highest probability of finding electrons are
called atomic orbitals.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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Atomic Orbitals
• (fuzzy cloud) = An atomic orbital is often
thought of as a region of space in which there
is a high probability of finding an electron.
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first ionization energy
(kJ mol-1)
Evidence for sub-levels
1600
1400
1200
The first ionization
energies for the elements
in period 3 show a
general increase.
1000
However, aluminium’s
800
value is below that of
magnesium. This suggests
600
that the third principal
400
energy level is not one
Na Mg Al Si P S Cl Ar
single energy level.
element
All principal energy levels contain one or more sub-levels,
with different but exact energy values.
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The sub-levels
There are four sub-levels,
labelled in order of
increasing energy: s, p, d
and f. Each holds a
different number of
electrons.
Each principal
energy level
contains a
different
number of
sub-levels.
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sub-level max no. electrons
s
2
p
6
d
10
f
14
principal energy
level, n
1
2
3
4
sub-levels
max no.
electrons
1s
2
2s, 2p
8
3s, 3p, 3d
18
4s, 4p, 4d, 4f
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Levels and sub-levels
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Atomic Orbitals
• The numbers and kinds of atomic orbitals
depend on the energy sublevel.
Energy
Level, n
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# of
sublevels
Letter of
sublevels
# of
# of
Total
orbitals per electrons in electrons in
sublevel
each orbital
energy
level
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Atomic Orbitals
• The numbers and kinds of atomic orbitals
depend on the energy sublevel.
Energy
Level, n
# of
sublevels
Letter of
sublevels
1
1
s
1
2
2
2
2
s
p
1
3
2
6
8
3
s
p
d
1
3
5
2
6
10
18
4
s
p
d
f
1
3
5
7
2
6
10
14
32
3
4
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# of
# of
Total
orbitals per electrons in electrons in
sublevel
each orbital
energy
level
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Blocks of the periodic table
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Order of sub-levels
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The Aufbau principle
As part of his work on electron configuration, Niels Bohr
developed the Aufbau principle, which states how
electrons occupy sub-levels.
The Aufbau principle states that the
lowest energy sub-levels are occupied first.
This means the 1s sub-level
is filled first, followed by 2s,
2p, 3s and 3p.
However, the 4s sub-level is
lower in energy than the 3d,
so this will fill first.
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Writing electron configuration
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Electron configuration: true or false?
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Electron configuration of transition metals
Although the 3d sub-level is in
a lower principal energy level
than the 4s sub-level, it is
actually higher in energy.
This means that the 4s sub-level
is filled before the 3d sub-level.
Example: what is the electron structure of vanadium?
1. Count number of electrons in atom
23
2. Fill sub-levels, remembering
4s is filled before 3d
1s22s22p63s23p64s23d3
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Electronic configuration: atoms
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Electron configuration of ions
When writing the electron
configuration of ions, it is
important to add or subtract the
appropriate number of electrons.
For non-transition metals, the sublevels are then filled as for atoms.
For negative ions
add electrons.
For positive ions
remove electrons.
Example: what is the electron structure of O2-?
1. Count number of electrons in atom
8
2. Add or remove electrons due to charge
8 + 2 = 10
3. Fill sub-levels as for uncharged atom
1s22s22p6
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Electronic configuration of transition metal ions
When transition metals form ions, it is the 4s electrons that
are removed before the 3d electrons.
Example: what is the electron structure of Ni2+?
1. Count number of electrons
in atom
28
2. Fill sub-levels, remembering
4s is filled before 3d
1s22s22p63s23p64s23d8
3. Count number of electrons
to be removed
2
4. Remove electrons starting
with 4s
1s22s22p63s23p63d8
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Electron configuration: ions
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Electron orbitals
As Heisenberg suggested, it is impossible to exactly locate
the position of an electron within an energy sub-level.
But Schrödinger’s wave equation showed by measuring the
electron density around the nucleus, it is possible to define
regions where electrons are most likely to be found at any one
time. These regions are called orbitals.
sub-level no. orbits max no.
Each energy sub-level
electrons
has one or more orbitals,
s
1
2
each of which can
p
3
6
contain a maximum of
two electrons.
d
5
10
f
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14
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Shapes of electron orbitals
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Atomic Orbitals
• Different atomic orbitals are denoted by letters.
• The s orbitals are spherical, and p orbitals are
dumbbell-shaped.
• Four of the five d orbitals have the same shape but
different orientations in space.
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The Pauli exclusion principle and spin
The Pauli exclusion principle states that each orbital
may contain no more than two electrons.
It also introduces a property of electrons called spin, which
has two states: ‘up’ and ‘down’. The spins of electrons in the
same orbital must be opposite, i.e. one ‘up’ and one ‘down’.
A spin diagram shows
how the orbitals are
filled. Orbitals are
represented by squares,
and electrons by arrows
pointing up or down.
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spin diagram
for
magnesium,
1s22s22p63s2
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Rules for filling electrons
When two electrons
occupy a p sub-level,
they could either
completely fill the same
p orbital or half fill two
different p orbitals.
Hund’s rule states that single electrons
occupy all empty orbitals within a sub-level
before they start to form pairs in orbitals.
If two electrons enter the same orbital there is repulsion
between them due to their negative charges. The most stable
configuration is with single electrons in different orbitals.
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Evidence for Hund’s rule
first ionization energy
(KJ mol-1)
The first ionization energies for the elements in period 3 show
a general increase.
1600
1400
1200
1000
800
600
However, sulfur’s value
is below that of
phosphorus. As the
highest energy
electrons of both are in
the 3p sub-level this is
evidence for Hund’s
rule.
400
Na Mg Al Si P S Cl Ar
element
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Evidence for Hund’s rule: P vs. S
Phosphorus has three electrons in its 3p sub-level and
sulfur has four.
The lower first ionisation energy for sulfur is because it has
a pair of electrons in one of the 3p orbitals. Mutual repulsion
between these two electrons makes it easier to remove one
of them.
phosphorus
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Electron configuration of Cr and Cu
The electron configurations of chromium and copper are
exceptions to the normal rules of orbital filling:
chromium
copper
1s22s22p63s23p64s13d5
1s22s22p63s23p64s13d10
In each case the 4s orbital contains one electron. This is
because the 4s and 3d sub-levels lie very close together in
energy, and the 3d being either half full or completely full is a
lower energy arrangement.
With larger atoms like this it can be useful to shorten the
electron arrangement. Copper can be shortened to
[Ar]4s13d10.
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Creating spin diagrams
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Summary of Quantum Numbers
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Glossary
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What’s the keyword?
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Energy levels and electrons
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Multiple-choice quiz
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What is periodicity?
The term periodicity describes a repeating pattern in
properties of elements across periods of the periodic table.
The Russian chemist Dmitry
Mendeleev is credited with being the
creator of the first version of the
periodic table. He observed that when
the elements are arranged in order of
atomic mass, there are recurring
patterns in certain properties.
The modern periodic table can be used
to analyse trends in properties such as
atomic radius across periods and
down groups.
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What is atomic radius?
The atomic radius of an element is difficult to precisely define
because of the uncertainty over the size of the electron cloud.
Several definitions are used.
One definition is half the shortest internuclear distance found
in the structure of the element.
For non-metallic elements, the
covalent radius is often used
as the atomic radius. This is half
the internuclear distance
between two identical atoms in
a single covalent bond.
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covalent
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More on atomic radius
For non-bonded adjacent atoms (e.g. in a covalent crystal of
a non-metallic element), the van der Waals radius is used as
a value for atomic radius. This is half the shortest internuclear
distance between two similar non-bonded atoms.
van der
Waals radius
For metallic elements, the metallic radius is often used as
the atomic radius. This is half the shortest internuclear
distance between two adjacent atoms in a metallic bond.
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Trends in atomic radius in period 3
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Trends in atomic radius in period 3
Atomic
Element
radius (nm)
Na
0.190
Mg
0.145
Al
0.118
Si
0.111
P
0.098
S
0.088
Cl
0.079
Ar
0.071
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The atomic radius of the elements
across period 3 decreases.
This might seem counterintuitive, because as the
numbers of sub-atomic particles
increase, the radius might be
expected to also increase.
However, more than 99% of the
atom is empty space – the
nucleus and electrons themselves
occupy a tiny volume of the atom.
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Increase in proton number
The number of protons in the nucleus of the atoms
increases across period 3.
proton
number
Element
11Na 12Mg 13Al
14Si
15P
16S
17Cl
18Ar
This increase in the number of protons increases the
nuclear charge of the atoms. The nucleus has stronger
attraction for the electrons, pulling them in closer and so
the atomic radius decreases across the period.
increased
nuclear charge
pulls electrons
closer
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What is shielding?
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Explaining atomic radius in period 3
Atomic
Proton
Element number radius
(nm)
Na
11
0.190
Mg
12
0.145
Al
13
0.118
Si
14
0.111
P
15
0.098
S
16
0.088
Cl
17
0.079
Ar
18
0.071
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Proton number increases
across period 3, but
shielding remains
approximately constant.
This causes an increase
in effective nuclear
charge, leading to a
greater attraction
between the nucleus and
the outermost electrons.
This pulls these electrons
closer to the nucleus and
results in a smaller radius.
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Atomic radius in period 3
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Atomic radius: true or false?
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What is first ionization energy?
Ionization is a process in which atoms lose or gain electrons
and become ions.
The first ionization energy of an element is the energy
required to remove one electron from a gaseous atom.
M(g) → M+(g) + eThe first ionization energy is therefore a measure of the
strength of the attraction between the outermost electrons
and the nucleus.
The first ionization energies of the elements in periods 2
or 3 can give information about their electronic structure.
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Plot of the first ionization energies
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General trend in first ionization energy
1600
ionization energy
(kJ mol-1)
There is a general
increase in the first
ionization energies
across period 3.
1400
1200
1000
800
Across period 3, the
600
proton number
400
increases but the
Na Mg Al Si P S Cl Ar
amount of shielding
element
does not change significantly.
The effective nuclear charge therefore increases.
The greater attraction between the nucleus and the
outermost electrons means that more energy is required to
remove an electron.
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Trend in first ionization energy: exceptions
ionization
energy (kJ mol-1)
There are two exceptions to the general trend in first
ionization energy: both aluminium and sulfur have lower
ionization energies than might be expected.
1600
1400
1200
1000
800
600
400
Na Mg Al
lower
ionization
energies
than
expected
Si
P
S
Cl
Ar
element
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First ionization energy of Al vs. Mg
The first ionization energy of aluminium is less than that
of magnesium, even though aluminium has a higher
nuclear charge.
The electron removed when aluminium is ionized is in a 3p
sub-level, which is higher in energy than the 3s electron
removed when magnesium is ionized. Removing an
electron from a higher energy orbital requires less energy.
magnesium
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First ionization energy of S vs. P
The first ionization energy of sulfur is less than that of
phosphorus, even though sulfur has a higher nuclear charge.
The highest energy electron in both phosphorus and sulfur
is in the 3p sub-level. However, in sulfur this electron is
paired, while in phosphorus each 3p orbital is singly
occupied. Mutual repulsion between paired electrons
means less energy is required to remove one of them.
phosphorus
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Ionization energy in period 3
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Ionization energy in period 3
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Plot of the melting and boiling points
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Na, Mg and Al: melting and boiling points
The melting and boiling points increase for the three metallic
elements from sodium to aluminium.
temperature (K)
3000
2500
2000
boiling point
1500
1000
500
0
Na
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melting point
Mg
element
Al
This is because the
strength of the
metallic bonds
increases. More
energy is needed to
break the stronger
metallic bonds, so
melting and boiling
points are higher.
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Na, Mg and Al: metallic bond strength
The increase in metallic bond strength from sodium to
aluminium is due to two factors:
1. Charge density. This is the ratio of an
ion’s charge to its size. Na+ ions are
large with a small charge, so have a
low charge density. Al3+ ions are
smaller with a larger charge, and so
have a higher charge density. They are
therefore more strongly attracted to the
delocalized electrons.
2. Number of free electrons. Sodium has one
free electron per metal ion, whereas
aluminium has three. This leads to more
attractions that must be broken in aluminium.
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Silicon
Silicon has a macromolecular
structure similar to that of
diamond.
Each silicon atom is
bonded to four
neighbouring silicon atoms
by strong covalent bonds.
These must be broken in
order for silicon to melt.
This requires a lot of
energy, so silicon's melting
and boiling points are high.
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Period 3 non-metals
This is because
they have a
simple molecular
structure with
weak van der
Waals forces
holding the
molecules together.
temperature (K)
The melting and boiling points of phosphorus, sulfur and
chlorine are much lower than those of silicon.
3500
3000
2500
2000
1500
1000
500
0
Na
Mg
Al
Si
P
element
S
Cl
Breaking these forces of attraction requires much less
energy than breaking covalent bonds.
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Ar
Period 3 non-metals: structure
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Melting points in period 3
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Melting points in period 3
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Glossary
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What’s the keyword?
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Multiple-choice quiz
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