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Electronic Structure
Key terms in this chapter are:
Electronic Structure
Electronic Configuration
Energy Level (Shell; Principal Quantum
Level)
Quantised
Ionisation Enthalpy
Ground State
Excited State
-1
kJ mol
In this section you are going to learn:
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What is meant by ‘electronic structure'.
About electrons in atoms existing in energy levels or shells.
How the energy levels fill with electrons.
How to write electronic configurations of atoms and ions.
The meaning of the term ‘quantised’.
The electronic structure (electronic configuration) of an atom refers to the
arrangement of electrons around its nucleus.
Electrons in atoms are considered to be arranged around the nucleus in energy levels
(or shells) called Principal Quantum Levels, n. The diagram below shows how the
electronic structure of an atom is often represented. The closer the energy levels are to
the nucleus of the atom the lower they are in energy. The lowest energy levels are the
first to become occupied with electrons. Potassium is represented below.
The electronic configuration of an atom or ion is represented by showing the number
of electrons in each energy level as they become more distant from the nucleus.
Note below the electronic configurations of some atoms and ions.
Na
Na+
Cl
Cl-
2.8.1
2.8
2.8.7
2.8.8
Here is something to remember: the maximum number of electrons that can be
contained within a Principal Quantum Level, n, is given by 2n2 (2 x n2). The table
below sets this out for the first four Principal Quantum Levels, n.
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Principal Quantum Level, n
Maximum number of electrons
1
2
2
8
3
18
4
32
Immediately, this may conflict with your understanding of the concept at GCSE; in
the energy level n = 3, there is a maximum of 18 electrons rather than 8. At GCSE
only the first 20 elements are considered, and at the point when there are 8 electrons
in the n = 3 level, the next two electrons (corresponding to K and Ca) go into the n = 4
level. The electronic structure of potassium is 2.8.8.1 and that of calcium is 2.8.8.2.
This is what happens. Two electrons go into the n = 1 level, then eight in the n = 2. Eight
electrons now go into the n = 3 level, followed by two being placed in the n = 4 level. The n = 3
level now takes another ten electrons to make a maximum total of 18. This takes us to the
element Zinc with the electronic structure 2.8.18.2. The n = 4 level now continues to fill with
electrons.
It is helpful to note that for those elements (Sc to Zn) corresponding to the 10 electrons added to
the n = 3 level, when it comes to losing electrons in ion formation (e.g. Zn2+, Mn2+), the two
electrons in the n = 4 level are removed first. Also notice, for example, that Zinc with the
electronic structure 2.8.18.2 is not in Group II of the Periodic Table.
The electronic structure on manganese, Mn is 2.8.13.2.
It is important to appreciate that an electron in an atom can occupy certain energy
levels only (and can't exist in between any two energy levels). Another way of putting
this is to say that an electron in an atom can possess certain energy values only, or that
its energy is quantised.
Ionisation Enthalpy
In this section you are going to learn:
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About cations and anions.
What is meant by the ‘ground state’ of an atom.
What is meant by an ‘excited state’ of an atom.
About ‘ionisation enthalpy’ and how to define it accurately.
How to write chemical equations to satisfy the definition of ionisation
enthalpy.
Ionisation enthalpy refers to the formation of positively charged ions (cations) from
gaseous atoms of elements in their ground state. (There are also negatively charged
ions called anions.)
An ion is a single atom, or group of two or more covalently bonded atoms, which has
an overall positive or negative electrical charge. Some examples are:
Li+, Fe3+, NH4+, Br-, O2-, NO3-
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An atom is in its ground state when all the electrons are in the energy levels of
lowest energy. If, in an atom, an electron is promoted from a lower to a higher energy
level, the atom is said to be in an excited state.
Here is a definition of first ionisation enthalpy:
The energy required to remove completely the outer-most electron from a
gaseous atom in its ground state.
Note that this refers to the first ionisation enthalpy. Chemical data books often give
enthalpy values, such as ionisation enthalpy, expressed in units of kJ mol-1.
Therefore,
Alternatively, the first ionisation enthalpy is:
The energy required to remove one mole of outer-most electrons, one each,
from one mole of gaseous atoms of an element in their ground state.
Note that this definition is refined with regard to the conditions of temperature and pressure in
later units.
Having understood the definition for ionisation enthalpy you need to be able to write
equations as appropriate according to the definition. Here are two examples.
Na(g) Æ
Na+(g) +
e-
∆HIE = +496 kJ mol-1
The second ionisation enthalpy for sodium is:
Na+(g) Æ
Na2+(g) +
e-
∆HIE = +4563 kJ mol-1