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Atomic Structure page 1 of 18
ATOMIC STRUCTURE AND THE PERIODIC TABLE
Theoretical change with respect to Dalton’s atomic theory
1. In 1803, atomic theory was revived by John Dalton
a) matter is made up of tiny particles called atoms which cannot be created,
destroyed or split
b) all atoms of one element are identical:- same mass and same chemical
properties
c) a chemical reaction consists of rearranging atoms from one combination to
another.
d) When elements combine to form compounds, small whole numbers of atoms form molecules.
However this was proved to be not entirely correct. Atoms have been split as
well as created i.e. nuclear reactions. Also there are isotopes, meaning that not
all atoms of an element are identical.
Therefore theory was forced to CHANGE in regards to these observations
contradicting to the theory put forward by Dalton.
The distribution of charge and mass in an atom
Particle Location Mass Charge
Electron Orbitals 1/1837 unit -1 unit
Proton Nucleus 1 unit +1 unit
Neutron Nucleus 1 unit 0
A unit is one atomic mass unit = 1.67 x 10-27
kg
Atomic Structure page 2 of 18
Atomic Structure page 3 of 18
Terminology
Term Definition
Atomic/proton number Number of protons in a nucleus of an atom
Nucleon/mass number Sum of the number of protons and neutrons in the
nucleus of an atom
Nuclide Any atomic species of which the proton number and
nucleon number are specified e.g. 126C and 9
4B are
nuclides
Isotopes Nuclides of the same element or atoms of the same
element with different mass numbers
NB isotopes have the same chemical properties but
different physical properties
Relative atomic mass Mass of an atom based on a scale such that the C-12
isotope has a mass of 12.00 units
relative atomic mass
= mass of 1 atom of an element x 12
mass of 1 atom of carbon-12
Phenomenon of radioactivity
Radiation is the spontaneous decay of unstable atoms with the emission of either
alpha, beta or gamma radiation.
Alpha decay is a type of radioactive decay in which an atomic nucleus emits an
alpha particle (two protons and two neutrons bound together into a particle
identical to a helium nucleus) and transforms (or 'decays') into an atom with a
mass number 4 less and atomic number 2 less.
For example:
although this is typically written as:
Beta decay is a type of radioactive decay in which a fast moving electrons is emitted. The new atom has no change in mass number but an atomic number
increases by 1.
Atomic Structure page 4 of 18
Gamma rays or gamma-ray (denoted as γ) are forms of electromagnetic
radiation (EMR) or light emissions of a specific frequency produced from sub-
atomic particle interaction, such as electron-positron annihilation and
radioactive decay. There is no change in atomic or mass number of the
atom.
Band of stability (n/p ratio)
Uses of radioisotopes
1. radiocarbon dating 2. smoke detectors
3. pacemakers 4. medical uses i.e. tracers or chemotherapy
5. irradiation in pest control
Calculations of relative atomic mass from isotopic data
Ar of an element = sum of (abundances x mass number of all of the isotopes of
an element)
e.g. what is the relative atomic mass of zirconium (Zr) Zr-90 51.5% ,
Zr-91 11.2%, Zr-92 17.1%, Zr-94 17.4 % and Zr-96 2.8%
Ar Zr = (51.5x 90) + (11.2x91) + (17.1x92) + (17.4x94) + (2.8x96) = 9131.8
The average mass of these 100 atoms would be 9131.8 / 100 = 91.3
(to 3 significant figures).
91.3 is the relative atomic mass of zirconium.
Most elements have isotopes. For stable isotopes, an interesting plot arises when the number of neutrons is plotted versus the number of protons.
Because the plot shows only the stable isotopes, this graph is often called the Nuclear Belt of Stability. The plot indicates that lighter nuclides (isotopes) are most stable when the neutron/proton ratio is 1/1. This is the case with any nucleus that has up to 20 protons.
As the atomic number increases beyond 20, a different trend becomes apparent. In this range, it appears that a stable nucleus is able to accommodate more neutrons. Stable isotopes have a higher neutron to proton ratio, rising to 1.5/1 for elements having atomic
numbers between 20 and 83.
Atomic Structure page 5 of 18
Terminology
Quantum number Definition
First principal quantum number (n) This corresponds to the shell number
e.g. the 1st shell has n=1, 2nd shell has
n=2
1. Orbital – volume of space in which there is a 95% chance of finding an
electron
2. Subshell – a group of orbitals with the same energy i.e. they are
“degenerate”
e.g. 3p subshell which has 3 orbitals of the same energy
3. Shell – a group of orbitals and/or subshells with the same principal quantum
number. n =1 shell sometimes shells can be the K shell where n=1, the L
shell where n=2, the M shell where n=3, etc.
Principal quantum number Types of orbitals/subshells
present in the shell
n=1 1s orbital
n=2 2s orbital and 2p subshell (which
contain THREE 2p orbitals)
n=3 3s orbital, 3p subshell (THREE p
orbitals) and 3d subshell (FIVE d
orbitals)
The relative energies of s, p and d orbitals up to principal quantum
number 4
Note: The 4s orbital is LOWER than the 3d orbital. Therefore electrons
will enter the 4s orbital first before the 3d
Atomic Structure page 6 of 18
Shapes of atomic orbitals (s orbital and p orbital respectively)
Order of filling electrons in orbitals & sub-shells (electronic
configuration)
1s 2s 2p 3s 3p 4s 3d
Remember for each p sub-
shell, there are 3 p orbitals in
x, y and z axis called
px, py and pz. They are
perpendicular to each other.
They are of the same energy
level and are called
“degenerate”.
Due to Pauli’s Exclusion
Principle (no two electrons can have the same 4 quantum states),
this implies that each orbital can
hold only TWO electrons, one
spinning “up” and the other spinning “down”. The Aufbau
Principle states that electrons enter
and fill orbitals of lower energy levels before going to higher
energy levels. Hund’s Rule states
that electrons entering sub-shells
containing 2 or more orbitals must enter and occupy the orbitals
SINGLY before pairing.
Atomic Structure page 7 of 18
Element 1s 2s 2px 2py 2px 3s 3px 3py 3pz 4s 3d
1H 1
2He 2
3Li 2 1
4Be 2 2
5B 2 2 1
6C 2 2 1 1
7N 2 2 1 1 1
8O 2 2 2 1 1
9F 2 2 2 2 1
10Ne 2 2 2 2 2
11Na 2 2 2 2 2 1
19K 2 2 2 2 2 2 2 2 2 1
20Ca 2 2 2 2 2 2 2 2 2 2
Element 1s 2s 2px 2py 2pz 3s 3px 3py 3pz 4s 3d
21Sc 2 2 2 2 2 2 2 2 2 2 1
22Ti 2 2 2 2 2 2 2 2 2 2 2
23V 2 2 2 2 2 2 2 2 2 2 3
24Cr* 2 2 2 2 2 2 2 2 2 1 5
25Mn 2 2 2 2 2 2 2 2 2 2 5
26Fe 2 2 2 2 2 2 2 2 2 2 5
27Co 2 2 2 2 2 2 2 2 2 2 7
28Ni 2 2 2 2 2 2 2 2 2 2 8
29Cu* 2 2 2 2 2 2 2 2 2 1 10
30Zn 2 2 2 2 2 2 2 2 2 2 10
* Indicates that the electronic configuration is not what is expected.
For Cr what would have been expected would be 3d4 4s
2, however, half
filled and totally filled shells/orbitals are very stable and thus more preferred than any other configuration. Therefore the electrons half fill the
3d subshell with 5 electrons and half fill the 4s orbital with 1 electron.
For Cu, what would be expected was 3d
9 4s
2, again the combination of a
totally filled subshell and a half filled orbital is more stable than just a filled
orbital and a partly filled subshell. Therefore the electrons adopt the more
stable configuration.
Atomic Structure page 8 of 18
Evidence of discrete energy levels using emission spectra
Based on information given above, it is shown that energy occupy different
orbitals or even-subshells and in essence occupy discrete energy levels.
When elements undergo emission spectroscopy and produce an emission
spectrum, a series of lines are shown like the emission spectrum of hydrogen
shown below.
But how are these lines explained and how do they show evidence of
discrete energy levels?
When the electrons in the element absorb energy, they move to higher energy
levels and are no longer in the “ground” state (lowest energy state), they are now in an “excited” state (state of higher energy). As the electrons release the
energy absorbed and leave the excited state to return to the ground state, the
excess energy is emitted, the wavelength of which refers to the discrete lines of
the emission spectrum. If discrete energy levels were NOT present, no lines
could EVER be formed in an emission spectrum.
Emission spectrum of hydrogen When a gaseous hydrogen atom in its ground state is excited by an input of
energy, its electron is 'promoted' from the lowest energy level to one of higher
energy (similar from moving from a lower rung in a ladder to a higher
rung). The atom does not remain excited but re-emits the excess energy as
electromagnetic radiation. This is as a result of an electron 'falling' from a higher
energy level to one of lower energy. This electron transition results in the release
of a photon from the atom of an amount of energy (E = hv) equal to the difference in energy of the electronic energy levels involved in the transition.
Note E = energy, h = Planck’s constant and v = frequency of wavelength of
radiation emitted
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. A prism can now be used to separate the emitted
electromagnetic radiation into its component frequencies (wavelengths or
energies). These are then represented as spectral lines along an increasing
frequency scale to form an atomic emission spectrum.
Atomic Structure page 9 of 18
Atomic Structure page 10 of 18
In each spectra a group of lines are see together which is classified as a series.
There are 3 series which are of significance.
The Lyman series occurs when electrons drop from higher energy levels to the
ground state (n = 1), in this series, the most amount of energy is released and
thus the smallest wavelength and highest frequency. This is why Lyman series
corresponds to the ultra-violet region (high energy)
The Balmer series occurs when electrons drop from higher energy levels to the
n = 2 level, here energies released are not as high as in the Lyman series. This
corresponds to the visible region of electromagnetic (EM) spectrum.
The Paschen series occurs when electrons drop from higher energy levels to the
n = 3 level. This corresponds to the infra-red region of electromagnetic (EM)
spectrum.
No two elements have the same atomic emission spectrum; the atomic
emission spectrum of an element is like a fingerprint.
Ionisation energy
It can be quoted more accurately as either 1
st, 2
nd, 3
rd, 4
th etc ionisation
energy. For our purposes, we will deal with the 1st ionisation energy.
The 1st ionisation energy is the energy required to remove a mole of
electrons from a mole of gaseous atoms to form a mole of gaseous
univalent ions. A (g) A+ (g) + e
-
Trend of 1
st ionisation energies
Ionisation energies generally increase going across a period Remember two factors must be considered: (1) proton number increases
sequentially going across a period i.e. greater nuclear attraction for the
outermost electron(s) and (2) number of electrons are also increasing.
Although the addition of electrons into the shell causes repulsion and thus
would increase the atomic radius, the predominant factor is the increased
“effective nuclear charge” (which is the residual attraction of the nucleus
and the outermost electron(s) after shielding of the inner electrons) i.e. more
energy would be needed to remove the outermost electron(s). Thus
ionisation energy increases from left to right of a period
Atomic Structure page 11 of 18
Ionisation energies decreases going down a group Although nuclear charge increases, the dominant factor is the increasing number
of shells between the nucleus and the outermost electron(s). This results in
increased shielding of the nuclear charge, therefore less attraction of the nucleus
and the outermost electron(s) i.e. less energy needed to remove an outermost
electron. Thus ionisation energy decreases down a group.
Atypical behaviour seen in period 3 for Mg & Al AND P & S (period 3) In period 3, Mg has E.C. of [Ne] 3s2, while Al has E.C. of [Ne] 3s2 3px
1, in Al
the outermost electron (3p) is at a higher energy level than the outermost
electron in Mg (3s), therefore less energy is needed to remove it. Or using a
different explanation, the valence electron in Al experiences more shielding
i.e. less nuclear attraction than one of the valence electrons in Mg i.e. less
energy needed to remove it from Al than for Mg.
In period 3 for P and S, the explanation needed is somewhat different. For P,
the E.C. is [Ne] 3s2 3p3, while for S the E.C. is [Ne] 3s2 3p4. In the 3p subshell
of P, the half-filled subshell represents a very stable configuration since it
represents a system of minimum repulsion as each electron occupies one orbital
singly. A lot of energy would be needed to disrupt this configuration. While in
S, 3p subshell experiences electron-electron repulsion in one of its orbital which
raises the energy of the system, therefore it is LESS stable and LESS energy
would be needed to remove one of the valence electrons.
Below is a diagram for the 1st ionisation energy of period 3 elements. Note
the circles show the areas of atypical behaviour
Atomic Structure page 12 of 18
Evidence of sub-shells using ionization data
The graph below shows the successive ionization energies for an atom of
sodium:
The electronic structure for sodium is 1s2 2s
2 2p
6 3s
1. The energy required to
remove the first electron is relatively low. This corresponds to the loss of
one 3s electron. To remove the second electron needs a much greater energy
because this electron is closer to the nucleus in a 2p orbital. There is a steady
increase in energy required as electrons are removed from 2p and then 2s
orbitals.
The removal of the tenth and eleventh electrons requires much greater
amounts of energy, because these electrons are closer to the nucleus in the 1s
orbital.
Large jumps in energy shown by the circles, indicate moving from one
principal quantum number to another. The smaller, more gradual increases
indicate going moving within subshells as the energies of the electrons will
slowly decrease resulting in more and more energy needed to remove them.
Atomic Structure page 13 of 18
How to derive group number of an element from successive ionization
energies A large jump (usually an increase of 3 or more times the amount) between
two successive ionisation energies is typical of suddenly breaking in to an
inner level. You can use this to work out which group of the Periodic Table
an element is in from its successive ionisation energies.
Example 1 Magnesium (1s22s
22p
63s
2) is shown with the following
successive ionisation energies:
Here the big jump occurs after the second ionisation energy. It means that
there are 2 electrons which are relatively easy to remove (the 3s2 electrons),
while the third one is much more difficult (because it comes from an inner
level - closer to the nucleus and with less screening). Mg is therefore in
group II
Example 2 Silicon (1s22s
22p
63s
23px
13py
1) is shown with the following
successive ionisation energies:
Here the big jump comes after the fourth electron has been removed. The
first 4 electrons are coming from the 3rd
shell orbitals; the fifth from. Silicon
is therefore in group IV
To try on your own
Decide which group an atom is in if it has the following successive
ionisation energies:
END OF ATOMIC STRUCTURE
Atomic Structure page 14 of 18
Worksheet
1. Write the electronic configurations of the following atoms or ions
a) 20Ca…………………………………..
b) 7N
3-……………………………..
c) 26Fe2+
………………………..
d)29Cu…………………………..
2.
Atomic Structure page 15 of 18
3.
Atomic Structure page 16 of 18
4.Below are incomplete nuclear equations, fill in any missing
information below i.e. any atomic and/or mass numbers missing as
well as any symbols of atoms that were not included.
5.
………………………………………………………………………
…………………..…………………………………………………
…………………………………………..…………………………
………………………………………………………………….
………………………………………………………………………
…………………….………………………………………………
…………………………………………….………………………
…………………………………………………………………….
………………………………………………………………………
……………………..………………………………………………
……………………………………………..………………………
Atomic Structure page 17 of 18
6. Using the same atoms from question 1, use the orbital diagrams
to illustrate the electronic configurations.
a) Mg
b) C
c) Ne
d) Li
e) K
f) Sc
Atomic Structure page 18 of 18
g) Fe
h) Co
i) Cu
j) Mn