Cape Chemistry
Topic: The Atomic Structure
Objectives:
1. Describe the structure of the atom
2. Define the following terms: mass number, isotopes, and, relative atomic
3. Explain the phenomenon of radioactivity
4. Cite the use of radioisotopes
5. Calculate the relative atomic mass of an element, given isotopic masses and abundances
6. Explain how data from emission spectra provide evidence for discrete energy levels
within the atom
7. Describe the atomic orbitals
8. Describe the shapes of the s and p orbitals
9. Determine the electronic configurations of atoms and ions in terms of s, p and d orbitals
10. State the factors which influence the first ionization energy of elements
11. Explain how ionization energy data provide evidence for sub-shells
12. Derive the electronic configuration of an element from data on successive ionization
energies
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The Atomic Structure
Evolution of the atomic model
Dalton’s atomic theory
John Dalton thought of atoms as a
hard sphere. He suggested that:
•
All atoms of the same
element are exactly alike
•
Atoms cannot be broken
down any further
•
Atoms of different
elements have different
masses
•
Atoms combine to form
complex structures
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Masses and charges of sub-atomic particles
Subatomic particles and electrical fields
What would happen if protons and neutrons were fired
through the electrical field?
Subatomic particles in magnetic field
The direction of the movement
of an electron in a magnetic
field can be found by using
Fleming left hand rule as shown
above. Still don’t get it!!,
follow this link to a YouTube
video:
https://www.youtube.com/watch?v=8li1Vp8vLaI&t=275s&ab_channel=It%27sAumSumTime
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Isotopes and radioactivity
Mass number = Proton + Neutrons
Relative atomic is the weighted average mass of naturally occurring atoms of an element.
Diagram of a mass spectrometer
Another way of determining the atomic mass of an element is to use a mass spectrometer.
You do not need to know how the instrument works, only that it produces positive ions of
atoms or fragments of molecules and separates them according to their masses. On placing a
sample of an element in a mass spectrometer, atoms of the element become positively
charged and separated according to their masses. Many elements are made up of atoms with
the same number of protons but different numbers of neutrons. This means that they have
different masses. The data can be used to calculate the average atomic mass of the sample.
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Ar Mg = (24 × 0.79) + (25 × 0.10) + (26 × 0.11) = 24.32
Isotopes have the same number of protons but different number of neutrons
Isotopes of hydrogen
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Radioactivity
Isotopes that have nuclei that break down (decay) spontaneously are called radioactive isotopes.
As the nuclei decays, rays or particles are given out. These are called emissions.
Name of emission
Types of particles/ rays
Stopped by
emitted
Alpha (α)
Helium nuclei (positively
Thin sheet of paper
charged particle)
Beta (β)
Electrons
6 mm thick aluminium foil
Gamma (ℽ)
Very high frequency
Thick lead sheet
electromagnetic radiation
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Equations for each decay
α decay
the isotope produced has a mass number of 4 units lower and a nuclear charge of 2 units lower
than the original atom:
223
88𝑅𝑎
→
219
86𝑅𝑛
+ 42𝐻𝑒2+
Note the α particle is not He atoms but the He2+ nuclei.
β decay
The mass number stays the same, but the protons increase by one. This is because the neutron
breaks down into a proton and an electron.
1
0𝑛
→ 11𝑝 + −10𝑒
14
For example: 6𝐶
→
14
7𝑁
+
0
−1𝑒
ℽ decay
A proton is converted to a neutron through a process called electron capture. The mass number
stays the same, but the proton number decreases by one.
37
18𝐴𝑟
+
0
−1𝑒
→
37
17𝐶𝑙
+ℽ
Uses of radioisotopes: Tracer, cancer treatment, dating objects, generating power
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Energy level and emission spectra
Electrons are arranged outside the nucleus in energy levels or quantum shells. When electrons
gain specific quanta of energy they move from lower energy level to higher energy level. They
become excited. When excited electrons lose their energy, they fall back to the lower energy
level emitting radiation of characteristic frequency. This is the origin of the emission spectrum.
The difference in energy between two energy level is related to the frequency of radiation by the
relationship
∆𝐸 = ℎ𝑣
∆E is the energy in joules
h is plank’s constant; 6.63*10-34J.s
V is frequency of the radiation in (s-1)
Page 8 of 17
Emission Spectra
When electrical or thermal energy is passed through a gaseous sample of an element the
radiation is emitted only at certain wavelengths or frequencies.
In the hydrogen emission spectrum electrons are moving from higher energy levels to lower
energy level. The lines in the emission spectra represent the energy levels that the electron will
transition when it absorbs a particular wavelength of light.
Video link:
https://www.youtube.com/watch?v=6rHerkru60E&t=324s&ab_channel=RichardThornley
Page 9 of 17
The points where the lines converge and come together is called the convergence limit. This
represents the electrons falling from the highest possible energy level. If the electron has more
energy than this, it becomes free from the pull of the nucleus of the atom. The atom is converted
to an ion.
Among the several series of lines seen in the emission spectra there are:
Lyman series (ultraviolet region) excited electron fall to n=1 energy level. They have the
highest frequency and the most energy.
Balmer series (visible region) excited electrons fall to the n=2 energy level.
Page 10 of 17
Niels Bohr discovered that electrons were actually found in energy levels with discrete quanta
of energy. His discovery lead to the atomic model we use today.
Sub-Shells and Atomic orbitals
The principal quantum shells, apart from
the first, are split into subshells
(sublevels). Each principal quantum
shell contains a different number of
subshells. The subshells are
distinguished by the letters s, p, d or f.
There are also f subshells for elements
with more than 57 electrons.
The maximum number of electrons that
are allowed in each subshell is: s = 2
electrons, p = 6 electrons, d = 10, f = 14
electrons.
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The number of electrons that can occupy an energy level can be calculated using 2n2 where n is
the energy level.
➢ The first principal quantum level, n = 1, can hold a maximum of 2 electrons in an s
subshell.
➢ The second principal quantum level, n = 2, can hold a maximum of 8 electrons: 2
electrons in the s subshell and 6 electrons in the p subshell.
➢ The third principal quantum level, n = 3, can hold a maximum of 18 electrons: 2 electrons
in the s subshell, 6 electrons in the p subshell and 10 electrons in the d subshell.
Atomic Orbitals
Each subshell contains one or more atomic orbitals. An atomic orbital is a region of space
around the nucleus of an atom that can be occupied by one or two electrons. As each orbital can
only hold a maximum of two electrons, the number of orbitals in each subshell must be:
s – one orbital
p – three orbitals
d – five orbitals
f – seven orbitals
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Shapes of the orbitals
Each orbital has a three-dimensional shape. Within this shape there is a high probability of
finding the electron or electrons in the orbital.
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Electronic configuration
A detailed way of writing the electronic configuration of an atom that includes information about
the number of electrons in each subshell is shown below.
Use the order of filling to write the electronic configuration
Order of Filling
for the following:
1. K
2. Cl3. Ca
4. Cr = 24
5. Cu = 29
6. Ti = 22
7. Ti 2+
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The arrangement of elements in the Periodic Table reflects the electronic structure of the
elements. The Periodic Table can be split into blocks of elements.
Ionization energy
Ionization energy is the energy required to remove an electron from an atom. This results in the
formation of an ion. Ionization energy increases across the period but decreases down the group.
Factor that affect ionization energy
1. Positive pull from
nucleus
2. Atomic radius
3. Shielding (Maximum
and minimum)
Page 15 of 17
Table of successive ionization energies from hydrogen to sodium
Interpreting successive ionization energies using a Graph of logarithm (log) of ionization
energy
✓ The electronic configuration
✓ Number of electrons on the
atom
✓ Number of shells
✓ Number of electrons on each
shell
✓ The electrons nearest or
furthest from the nucleus
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References
Norris, Roger, et al. Cape Chemistry Unit 1 . Cheltenham: Nelson Thornes Ltd, 2012.
Ryan, Lawrie and Roger Norris. Cambridge International AS and A Level Chemistry Courebook.
2nd edition. Cambridge University Press, 2012, 2014.
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