Topic 2 12 Manual 2023
Topic 2 12 Manual 2023
Topic 2 12 Manual 2023
Atomic Structure
Second edition, 2015
Table of Contents
Topic 2 & 12
Atomic Structure ii
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The atom
Atoms are the smallest particles that can take part in a chemical reaction and are thus
seen as the fundamental units of matter.
Early atomic theories (e.g. Dalton) viewed the atom as a hard indestructible unit of
matter but the discovery of the 3 fundamental sub-atomic particles completely changed
our view of atomic structure and new theories had to be devised to account for the fact
that the atom as composed of several smaller fundamental units.
Since all atoms are neutral they must contain equal numbers of protons and electrons.
There is however no simple relationship between the number of protons and neutrons in
an atom.
The protons and neutrons occupy the centre of the atom called the NUCLEUS and are
often referred to as nucleons. The electrons occupy a relatively large area of space
around the nucleus, such that the radius of the nucleus itself is only about 1/10,000 th of
the radius of the whole atom.
The nucleus of an atom consists of a number of protons and neutrons. Although there is
no fixed relationship between the numbers of these particles present in a nucleus there is
evidence to suggest that there needs to be slightly more neutrons than protons for the
nucleus to be stable.
Two important terms help quantify the number of neutrons and protons in an atom:
Atomic Number (Z) - is the number of protons in the nucleus of the atom
Mass Number (A) - is the total number of nucleons (protons + neutrons) in the
nucleus of the atom
Thus an atom of an element can be represented in the following way
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Atomic Structure 3
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Z
A X e.g.
23
11 Na
The atoms of an element must all have the same number of protons but the number of
neutrons may vary. Atoms of the same element which have different numbers of
neutrons are called ISOTOPES. For example: 1H, 2H and 3H.
Radioactive isotopes
Many isotopes are radioactive – their nuclei of these atoms are unstable and break down
spontaneously and emit radiation. The time taken for one half of the original number of
atoms to undergo radioactive decay is known as the half life. The half life of an element
can be very long (millions of years) or extremely short (tiny fractions of a second). The
half life of U-238 is 4.5 billions year and that of B-9 is 8 x 10-19s.
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Atomic Structure 4
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Assignment 2.1
56 28
21 20
11 23
Fe 57
2. Boron has an atomic number of 5. It comprises two isotopes, one with five neutrons,
the other with six.
a. Explain what is meant by the term ‘isotope’.
b. Calculate the mass number of the two isotopes and represent them in the form
yB.
x
3. Naturally occurring copper is a mixture of two isotopes. One of these has 29 protons
and 34 neutrons, the other one two more neutrons. If the relative atomic mass of
copper is 63.55, calculate the natural abundances of the two isotopes.
Number of
Species
Protons Neutrons Electrons
3
H-
26
Mg2+
13 14 10
16 18 18
4+
22 26
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Atomic Structure 5
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The existence of isotopes and the determination of the average relative atomic mass for a
sample of an element were obtained using a mass spectrometer.
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Atomic Structure 6
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The individual peaks represent particles of specific mass to charge ratios (m/z). In most
cases the charge will be +1 and hence m/z = the mass of the atom.
The height of each peak represents the abundance of each particle in the sample. Hence
the total height of all peaks = 100%
(a) The two isotopes of rubidium can be described using the following isotopic
symbols.
(b) The relative abundance of each isotope can be determined using the height (or
measurement) of each peak.
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Atomic Structure 7
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Atomic Structure 8
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Assignment 2.2
2. There are three naturally occurring isotopes of magnesium. One of them, 24Mg,
has an abundance of 79.21% and another one, 26Mg, has an abundance of 10.68%.
Determine the mass number of the third isotope.
3. (a) Look up, in a data book, the percentage abundance of the stable isotopes of
chromium.
(b) Sketch the mass spectrum that would be obtained from naturally occurring
chromium. Let 10.0cm represent 100 % on the vertical scale.
(c) Calculate the relative atomic mass, to 3 sig. fig.
(d) Label each peak on the mass spectrum using isotopic symbols.
4. Gallium has two naturally occurring isotopes, 69Ga and 71Ga. Calculate the
percentage abundance of each isotope, given the Ar quoted, to 2 d.p. in the IB data
booklet.
(a) Explain why a 20Ne+ ion is deflected more than a 22Ne+ ion by an external
magnetic field.
(b) How would a 22Ne2+ ion behave under the influence of the same magnetic
field?
6. Suggest three possible advantages of keeping the pressure very low inside a mass
spectrometer.
7. Explain how the vaporized sample inside a mass spectrometer is ionized and state
why it is necessary to form positive ions.
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Atomic Structure 9
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Electron arrangement
Electrons in an atom are arranged around the nucleus, but not in a random manner.
Each electron has a specific amount of energy and can only go in specific energy shells,
or orbits around the nucleus. There are a few basic rules for determining the electron
arrangement of electrons in an atom.
H He
1
Li Be B C N O F Ne
2,1
Na Mg Al Si P S Cl Ar
K Ca
What differences and/or similarities of electron arrangement do you notice going across
a Period (row) in the Periodic Table?
What differences and/or similarities of electron arrangement do you notice going down
a Group (column) in the Periodic Table?
How many electron shells would you expect an atom of bromine to have? How many
electrons will a bromine atom have in its outer shell?
Information about the arrangement of electrons has been obtained by ‘disturbing’ them
in two ways.
(1) By forcing them into higher energy levels and observing what happens when they
return to normal.
(2) By bombarding them with streams of fast moving particles and so detaching them
from the atom.
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Atomic Structure 10
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In the above, light is being considered as a wave, with a characteristic wavelength and
frequency. In some situations, the behaviour of light is easier to explain by thinking of it
not as waves but as particles. Here you think about electromagnetic radiation being
emitted or absorbed in packets or quanta of radiation (called photons).
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Atomic Structure 11
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Frequency
Since each frequency corresponds to a particular energy, the nature of this spectrum
implies that the atom can only lose energy in fixed amounts. This in turn leads to the
conclusion that the electron in a hydrogen atom can only exist at particular energy levels
within the atom. Each line corresponds to transitions between specific energy levels – hf
= E2 – E1 (where E2 represents the energy of the higher energy level and E1 that of the
lower energy level).
In a sample of gaseous hydrogen where there are many trillions of atoms all of the
possible electron transitions from higher to lower energy levels will take place many
times and an emission spectrum will be produced.
The diagram on the next page shows the whole range of possible transitions leading to
the different spectra. Transitions to n = 2 produce a line spectrum in the visible part of
the electromagnetic spectrum, known as the Balmer series. Transitions to n = 3, 4 or 5
will produce line spectra in the infrared part of the electromagnetic spectrum.
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Atomic Structure 12
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Within each series, the spectral lines get closer together (converge) with increasing
frequency. This suggests that the electronic energy levels get closer the more distant they
become from the nucleus of the atom.
Types of spectrum
Niels Bohr, in 1913, proposed a model of an atom that helped to explain the
observations. As spectra of large multi-electron atoms are very complex and difficult to
interpret, Bohr used the spectra of hydrogen to develop his model. He stated that
Electrons can only occupy fixed energy levels; each energy level being assigned a
principal quantum number. The energy level closest to the nucleus being n = 1 and
of lowest energy.
The energy of an electron is quantised; the electron may only have certain energies.
Electrons can only absorb or emit energy of specific frequencies.
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Atomic Structure 13
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When the electron occupies the energy level of lowest energy the atom is said to be in
its ground state. An atom can have only one ground state.
If the electron occupies one of the higher energy levels then the atom is in an excited
state. An atom has many excited states.
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Atomic Structure 14
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Assignment 2.3
1. Identify the region of the electromagnetic spectrum where the following will be
found.
2. Explain why the lines in the visible atomic emission spectrum of hydrogen
converge towards the higher energy end of the electromagnetic spectrum.
3. Draw and label (using the letters a, b and c) on the energy level diagram below
(a) the electron transition which will give the third line in the visible series of the
hydrogen emission spectrum.
(b) the electron transition which will give the third line in the ultraviolet series of
the hydrogen emission spectrum.
(c) the transition that relates to the ionisation energy of a hydrogen atom.
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Atomic Structure 15
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The Bohr model failed to explain the line spectra of multi-electron atoms because it was
based on the assumption that the electron was a solid particle of matter.
In the 1920’s, Louis de Broglie was the first to suggest that electrons could also possess
wave-like properties e.g. diffraction, the bending and spreading out of a wave on passing
through a small aperture. The realisation of the wave-particle duality of the electron
enabled the method of describing the electron structure to be based on the mathematics
of waves.
Heisenberg in 1927 put forward his famous ‘Uncertainty Principle’ in which he states
that it is impossible to measure accurately both the position and momentum of an
electron at the same time. As momentum is related to time, what this implies is that is it
is impossible to know the exact location of an electron at an exact moment in time. His
principle enabled scientists to think in terms of probabilities of finding electrons in
certain volumes of space.
Schrodinger used de Broglie’s ideas to develop his Wave equation in which he defined
regions in space or orbitals, around the nucleus, where there is a high probability
(~90%) of finding an electron of a given energy. The mathematical solution to his
equation allows each electron in an atom to be uniquely described by four different
quantum numbers.
Quantum numbers
The first or principal quantum number describes the main energy level or shell (n = 1,2
…..). The higher the number of n, the further from the nucleus it is.
The second quantum number describes the type of orbital(s) in that level. It tells you
about the shape of an orbital.
s orbitals: are all spherical in shape, but p orbitals: are dumb bell shaped.
each successive s orbital is larger than the
previous one.
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Atomic Structure 16
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The more complex shapes of the d and f orbitals are beyond the scope of the syllabus.
The third quantum number determines the number of orbitals of each type:
Thus each set of p orbitals consists of three ‘dumb bell’ shaped orbitals, 90° to each other
about the nucleus.
To summarise:
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Atomic Structure 17
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The fourth quantum number describes the spin of the electron. Electrons in one orbital
can either be spinning in one direction or in the opposite direction. The Pauli Exclusion
Principal states that no two electrons in the same atom can have exactly the same four
quantum numbers. Hence each orbital can contain a maximum of only two electrons.
1. The Aufbau Principle - which states that the energy levels of lowest energy must
be filled first. Each set of orbitals being filled before electrons may enter in to the
next set. The order being:
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Atomic Structure 18
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The Periodic Table can also be used to determine the order and will be discussed
later.
2. The Pauli Exclusion Principle - which states that there can be no more than two
electrons in any orbital, having opposite spins.
3. Hund’s Rule - which states that orbitals of the same energy (i.e., orbitals in the
same sub-level) must be occupied singly before pairing occurs.
not
Exceptions: There are two notable exceptions where the electron configurations do not
fully follow the above i.e. Cr and Cu.
Find out what the electron configurations of Cr and Cu are and suggest explanations
why they are exceptions.
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Atomic Structure 19
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or 1s2 2s2 2p6 3s2 3p6 4s2 note that the 4s fills before the 3d
or [Ar] 4s2
n=4
n=3
n=2
n= 1
2. 26 Fe
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Atomic Structure 20
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Assignment 2.4
1. Write out the electron configuration of the following elements, using the Bohr
format.
2. Write out the electron configuration of the following elements using the wave
mechanical model. Do not for this question use the abbreviated format.
(a) He (b) C
(c) Mg (d) As
(e) K (f) V
(g) Co (h) Zn
(a) Sr (b) Pt
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Atomic Structure 21
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First ionisation energy is the energy required to remove each outermost electron from
one mole of atoms of an element in its gaseous state.
The second IE is the energy required to remove the next outermost electron and so
on.....
Draw a graph of these values. Put ionisation energy on the vertical axis.
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Atomic Structure 22
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Write an equation representing the change occurring for each successive ionisation
energies.
Explain, using your graph, which Periodic Table Group you would place boron.
Explain why the values for the first three ionisation energies increase: 807, 2430, 3670.
Explain the very large jump between the third and fourth ionisation energies.
5.5
4.5
4
log10 IE
3.5
2.5
2
0 2 4 6 8 10 12
# electrons rem oved
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Atomic Structure 23
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Unknown Elements
This table gives some, but not all, of the successive ionisation energies (in kJ/mol) of
several fictitious elements.
Place each element in its proper group. Explain how you arrived at each answer.
The graph below shows the IE of the second energy level electrons of sodium. What
does this graph suggest about the arrangement of these electrons?
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Atomic Structure 24