Elsa High School_DP Chemistry
10/5/2022
Unit 2 Atomic Structure
Dr. Kevin YAU
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Overview for unit 2
 2.1 The nuclear atom
 2.2 Electron configuration
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2.1 The nuclear atom
The mass of an atom is concentrated in its minute, positively
charged nucleus
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Essential ideas
 Atoms contain a positively charged dense nucleus
composed of protons and neutrons (nucleons).
 Negatively charged electrons occupy the space outside
the nucleus.
 The mass spectrometer is used to determine the relative
atomic mass of an element from its isotopic composition.
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Nature of Science (NOS)
 Evidence and improvements in instrumentation - alpha
particles were used in the development of the nuclear
model of the atom that was first proposed by Rutherford.
 Paradigm shifts - the subatomic particle theory of matter
represents a paradigm shift in science that occurred in
the late 1800s.
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Atomic Structure
 Atoms are composed of
three sub-atomic particles:
protons, neutrons and
electrons.
 Each atom consists of two
regions: the nucleus and
the electrons shells
 The nucleus contains
protons and neutrons as
known as nucleons
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Relative masses and relative charge
of protons, neutrons and electrons
 Protons have a positive charge and neutrons have NO
charge (i.e. neutral) while electrons have negative
charge
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Mass number, atomic number and
isotopes
 The atomic number (Z) is the number of protons in the
nucleus
 The mass number (A) is the total number of protons and
neutrons
 Number of neutrons = A - Z
Chemical elements (names and symbols) and their atomic
numbers are listed on P.4-5 of the IB Chemistry data booklet
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Mass number, atomic number and
isotopes (cont’d)
 Isotopes are atoms of the same element that have same
number of protons and different number of neutrons.
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Ion
 An element forms an ion by
addition/removal of electrons
from the atom
 Thus there is NO change in the
proton and neutron numbers
but the number of electron
changes
 A positive ion is formed by the
removal of electrons and a
negative ion is formed by the
additional of electrons.
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Worked example: sub-atomic
particles
p
n
e
P
15
31-15=16
15
P3-
15
16
15+3=18
p
n
e
Mg
12
12
12
Mg2+
12
12
10
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Isotopes are atoms with the same number of protons but different
number of neutrons.
Properties of isotopes
 Isotopes of the same chemical element have identical
chemical properties but slightly different physical
properties.
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Relative abundance (%)
Relative isotopic mass of Hydrogen-1 = 1;
Relative isotopic mass of Hydrogen-2 = 2;
Relative isotopic mass of Hydrogen-3 = 3;
How heavy is hydrogen?
Average = 1+2+3/3 = 2
98%
1.5%
0.5%
Weighted Average of the relative isotopic masses
Relative atomic mass = 1x90% + 2 x1.5% +3 x0.5%
= 1.008
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Radioactivity
 Elements containing unstable nuclide will break up
spontaneously with the emission of ionizing radiation.
 These unstable nuclides are described as radioactive
and are called radionuclides.
 There are three types of radiation: alpha, beta and
gamma radiation
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Radioactivity (cont’d)
 When the nucleus of a radionuclide releases a alpha or
beta particles, a new element is formed.
 E.g. Carbon-14 and Iodine-131 are beta-emitters.
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Radioactivity (cont’d)
 Radioactive decay is an
exponential process.
 The rates of radioactive decay
are compared using the halflife, which is the time taken for
half of the radioactive nuclei to
undergo decay.
 Each radionuclide has its own
unique half-life which is
unaffected by temperature or
pressure.
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Utilization: use of radioactive isotopes
Carbon-14 dating
Medical use of radioactive isotopes
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Mass spectrometer
 A mass spectrometer allows chemists to determine
accurately the relative atomic masses of atoms. It can
also be used to determine the relative molecular masses
of molecular compounds and establish their structure.
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Operation of mass spectrometer
s are
energy
ming
eld
ment of
m to enter
The processes rely upon electrostatics and electromagnetism
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Using mass spectrometry to
determine relative atomic mass
Two peaks were o
 Two isotopes we
chlorine sample
 Position of the p
is showing the r
mass of that iso
 Height of the pe
relative abunda
isotopes (The hi
the large will be
abundances)
The mass spectrum shows the
isotopes of chlorine with the relative
abundances of each of them
75%
25%
35 37
(Mass-to-charge ratio)
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Using mass spectrometry to determine
relative atomic mass (cont’d)
 The chemical elements are listed
in a special arrangement called
the periodic table (P.6; IB
Chemistry data booklet)
 Each chemical element is
placed in a box with its chemical
symbol, its atomic number
(written above) and its relative
atomic mass (written below)
90.48%
0.27%
9.25%
 The majority of chemical
elements in nature exist as a
mixture of isotopes in fixed
proportions.
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Using mass spectrometry to determine
relative atomic mass (cont’d)
 The relative atomic mass (symbol Ar ) of an element is the
weighted average of its relative isotopic masses of each
isotopes with the consideration of the relative
abundances.
Considering there are
100 atoms of chlorine
35Cl:
75%
37Cl: 25%
75 atoms are 35Cl
25 atoms are 37Cl
Total weight = 35x75 +
37x25 =3550
Relative atomic mass (Ar)
=3550/100
=35.5
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Worked example: Relative atomic
mass
Rubidium exists as a mixture of two isotopes, 85Rb and 87Rb.
The percentage abundances are 72.1% and 27.9%,
respectively. Calculate the relative atomic mass of
rubidium.
Solution
Ar of Rubidium
= (85 x 72.1 + 87 x 27.9)/100
= 85.6
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Worked example 2: Relative atomic
mass
The relative atomic mass of gallium is 69.7. Gallium is
composed of two isotopes: gallium-69 and gallium-71.
Calculate the percentage abundance of gallium-69.
Solution
Assuming the percentage abundance of Ga-69 be a%,
then the percentage abundance of Ga-71 is (100-a)%
[69a + 71(100-a)]/100 = 69.7
69a + 71(100-a) = 6970
69a + 7100 – 71a =6970
2a = 130
a=65
Therefore, the percentage abundance of Ga-69 is 65%
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2.2 Electron configuration
The electron configuration of an atom can be deduced from its
atomic number
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Light as wave and particles
 Light is an important form of energy and is often
described using a wave model.
 Light is transmitted in the form of electromagnetic waves
which consist of oscillating electric and magnetic wave
travelling together in a sinusoidal pattern.
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Common descriptions of light waves
speed of light = 3.0 x 108 ms-1
 Wavelength (λ) – the distance between two neighboring
wave crests/troughs; units: meters (m)
 Frequency (f) – the number of oscillations in one second;
units: Hertz (Hz)
 Speed (c) – the distance travelled by the wave in one
second; units: meters per second (ms-1)
where c is the speed of light in vacuum (3 x 108 ms-1)
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3 x 108 ms-1
𝑐 = 𝑓𝜆
10/5/2022
visible light
Different colours in visible light correspond to electromagnetic
waves of different wavelengths and frequencies
The electromagnetic spectrum can be found on P.3 of the IB
Chemistry data booklet
frequency ∝ penetration power
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model
0-34)
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10/5/2022
Particle model of light
 Light can also be described by particle model as a stream
of photons or tiny ‘packets’ of light energy
 The two models of light are linked by Planck’s equation:
 where E is the energy of the photon (in Joules); h is the
Planck’s constant (6.63 x 10-34); v is the frequency of light (in
Hertz)
Planck’s equation and Planck’s constant are given on page 1
and 2 of the IB Chemistry data booklet respectively
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Spectra
 The spectrum of colours formed from white
light is composed of visible light of a certain
range of wavelengths and is called a
continuous spectrum
 If gaseous atoms are excited they emit light
of certain wavelengths. Excitation occurs
when electrons in the atom are raised to a
higher energy level, and light is emitted as
they return to the unexcited state.
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Emission spectra
 If the light from atoms with excited electrons is passed
through a prism, an emission spectrum is formed.
 Emission spectra consist of a number of separate sets, or
series, of narrow colored lines on a black background.
Hence emission spectra are often called line spectra .
 Each chemical element has its own unique line spectrum
that can be used to identify the chemical element.
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Emission spectra (cont’d)
Emission spectra of hydrogen atom (Balmer series)
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a-violet
ion
man
es)
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Emission spectrum (cont’d)
 The emission spectrum for hydrogen atoms in the visible
region consists of a series of lines observed in the visible
region called the Balmer series. Similar sets of lines are
observed in the ultraviolet (Lyman series) and infrared
regions (Paschen series)of the electromagnetic spectrum.
 An emission or line spectrum differs from a continuous
spectrum in two important ways:
discrete
 An emission spectrum is made up of separate lines (coloured if
they are in the visible region), that is, it is discontinuous.
 The lines converge, becoming progressively closer as the
frequency or energy of the emission lines increases.
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Emission spectrum of white light (continuous), sodium vapour (line
spectra) and cadmium vapour (line spectra)
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Utilization: Fireworks
At the high temperature
of the flame, electrons
are promoted to higher
energy levels.
The colours observed
arise from these excited
metal atoms or ions.
When the electrons
return to the ground
state, the atoms emit
light of characteristic
frequencies.
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Worked example: emission spectrum
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(i) An emission spectrum is made up of separate/discrete
lines. The lines converge, becoming progressively closer as
the frequency or energy of the emission lines increases.
(ii) Electrons are excited to a higher energy level upon the
supply of energy. A photon with a particular frequency is
emitted as the electron transition occur from the higher
energy level to the lower energy level which produce a
line in the emission spectrum
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Energy levels and spectra
 All elements release a characteristic
mixture of light when a sample is placed
into a high voltage discharge tube
 Bohr was the first to produce a theory that
could account for the complete emission
spectrum of hydrogen
 He suggested that electrons in atoms
moved in energy levels known as orbits ,
where they had certain fixed amounts of
potential energy
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um number (n)
10/5/2022
3rd
180Hz
150Hz
100 Hz
2nd
1st excited
n=2
Ground
n=1
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Energy levels and spectra (cont’d)
 The electron must absorb energy (electrical or thermal
energy) to do work against the attraction of the positively
charged nucleus. The atom or electron is now said to be
in an excited state.
 The emission or line spectrum is formed when electrons
that have been excited drop back from orbits of high
energy to an orbit of lower energy releasing light of a
particular frequency according to the Planck’s equation:
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Energy levels and spectra (cont’d)
 Bohr labeled his energy levels or orbits with the letter n
and a number.
 An electron in the lowest energy level (nearest the
nucleus) was labeled as n = 1; an electron in this orbit is
in its ground state
 The next orbit or energy level is labeled as n = 2, and so
on.
 The energy levels or orbits correspond to electron shells
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Energy levels and spectra (cont’d)
 If an electron receives enough energy to remove it completely
from the attraction of the nucleus, the atom is ionized. The
energy required to ionize the electron is known as the ionization
energy.
 It is equivalent to the transition from n=1 to n=∞, while n is the
notation of the energy levels called the principal quantum
number
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Energy levels and spectra (cont’d)
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Energy levels and spectra (cont’d)
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The electron arrangement for atoms
and ions
 The electrons in atoms are arranged in energy shells.
 The electron enters the shell nearest the nucleus.
 The maximum number of electrons each shell (main
energy level) can hold is given by the expression 2n2.
 Chemists often use a shorthand notation to describe the
arrangement of electrons in shells. It indicates the
number of electrons in each shell without drawing the
shells. It is known as the electron arrangement
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Maximum number of electrons each
shell can accommodate
1st shell, n =1
2n2 = 2 (1)2 = 2
2nd shell, n =2
2n2 = 2 (2)2 = 8
2nd shell, n =3
2n2 = 2 (3)2 = 18
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Oc
8
From
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The electron arrangement for atoms
and ions (cont’d)
 The third shell can hold a maximum of 18 electrons.
However, when there are eight electrons in the third shell
there is a degree of stability and the next two electrons
enter the fourth shell
 The electrons in the different
shells experience different
attractive forces due to the
presence of other electrons is
called the electron shielding.
The outer electrons experience
the most shielding.
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px orbital
ell: 6
als
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pz orbital
Orbitals and energy levels
 Electrons are arranged in shells (energy levels) around
the nucleus of an atom.
 Each shell consists of a number of sub-shells (sub-levels),
labeled s, p, d or f. The existence of sub-shells is
confirmed experimentally
 The number of sub-shells is equal to the shell number
ell
s
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Orbitals and energy levels (cont’d)
 Each sub-shell contains a number of orbitals in which the
electrons are placed.
 Each orbital can hold a maximum of two electrons.
 Each orbital can be represented by a ‘square box’. An
electron is represented by an arrow 
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Shapes of orbitals
 Electrons do not occupy fixed positions within an atom,
nor do they follow orbits in the shells. Electrons occupy
volumes or regions of space called orbitals.
 The four types of orbitals, s, p, d and f, all have different
shapes.
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Shapes of orbitals (cont’d)
 All s orbitals can be represented as spheres. They differ
only in size and energy. The 3s orbital is larger than the 2s
orbital, which is larger than the 1s orbital. The larger
orbitals are described as being more diffuse since the
electron density is less.
 The p orbitals have two lobes forming a ‘dumb-bell’
shape and have different orientations in space. They are
arranged at right angles to each other and are labeled
px , py and pz to reflect their orientation. The three p
orbitals all have the same energy – the orbitals are said to
be degenerate .
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Filling atomic orbitals
The electrons are arranged in atomic orbitals according to
certain principles:
 Pauli exclusion principle: Each orbital can hold up to a
maximum of two electrons
 Aufbau principle: Electrons enter and occupy an empty
atomic orbital with the lowest energy
 Hund’s rule: Within a sub-shell, electrons experience
repulsion and hence enter two different orbitals of the
same energy.
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Filling atomic orbitals (cont’d)
 Single electrons in the same sub-shell must have the
same (parallel) spin
 Two electrons in the same orbital must have opposite
spins
Spin pair
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What’s wrong with these electron
configuration?
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Electron configuration of atoms
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Electronic configuration of ions
 To deduce the electron configuration of positive ions
(cations), electrons are removed in reverse order (that is,
the last electron is removed first).
 To deduce the electron configuration of negative ions
(anions), electrons are added with the same principles.
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Division of the periodic table into
blocks
 The long form of the periodic table is divided into four blocks: the s-, p-,
d- and f-blocks
 This division reflects the filling of the outermost orbitals with electrons
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The octet rule
 The electron arrangements of noble gases are relatively
stable and their atoms do not lose or gain electrons to
form ions.
 Atoms of noble gases, with the exception of helium, have
eight electrons in their outer shells. This arrangement is
known as an octet .
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Summary
 Atoms contain a positively charged dense nucleus composed of
protons and neutrons (nucleons).
 Negatively charged electrons occupy the space outside the
nucleus.
 The mass spectrometer is used to determine the relative atomic
mass of an element from its isotopic composition.
 The emission or line spectrum is formed when electrons that have
been excited drop back from orbits of high energy to an orbit of
lower energy releasing light of a particular frequency according to
the Planck’s equation
 The electron configuration of an atom can be deduced from its
atomic number
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