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2p orbitals. These three 2p orbitals
are designated as 2px, 2py and 2pz. The third shell contains one 3s orbital,
three 3p orbitals and five 3d orbitals. Thus, the first shell can hold only two
electrons, the second shell eight electrons and the third shell up to 18
electrons, and so on. As the number of electrons goes up, the shell numbers
also increase. Therefore, electron shells are identified by the principal
quantum number, n¼ 1, 2, 3 and so on.
The electronic configuration of an atom describes the number of electrons
that an atom possesses, and the orbitals in which these electrons are
placed. The arrangements of electrons in orbitals, subshells and shells are
called electronic configurations. Electronic configurations can be repre-
sented by using noble gas symbols to show some of the inner electrons, or
by using Lewis structures in which the valence electrons are represented
by dots.
Valence is the number of electrons an atom must lose or gain to attain the
nearest noble gas or inert gas electronic configuration. Electrons in the outer
shells that are not filled are called valence electrons.
The ground-state electronic configuration is the lowest energy, and the
excited-state electronic configuration is the highest energy orbital. If
energy is applied to an atom in the ground state, one or more electrons
can jump into a higher energy orbital. Thus, it takes a greater energy to
remove an electron from the first shell of an atom than from any other
shells. For example, the sodium atom has electronic configuration of two,
eight and one. Therefore, to attain the stable configuration, the Na atom
must lose one electron from its outermost shell and become the nearest
noble gas configuration, i.e. the configuration of neon, which has the
electronic configuration of two and eight. Thus, sodium has a valence of 1.
Since all other elements of Group IA in the periodic table have one
electron in their outer shells, it can be said that Group IA elements have a
valence of 1.
At the far end on the right hand side of the periodic table, let us take
another example, chlorine, which has the electronic configuration of two,
eight and seven, and the nearest noble gas is argon, which has the electronic
configuration of two, eight and eight. To attain the argon electronic
configuration chlorine must gain one electron. Therefore, chlorine has a
valence of 1. Since all other elements of Group 7A in the periodic table have
seven electrons in their outermost shells and they can gain one electron, we
can say that the Group 7A elements have a valence of 1.
Shell Total number of shell electrons Relative energies of shell electrons
4 32 Higher
3 16
\ufffd\ufffd!2 8
1 2 Lower
Each atom has an infinite number of possible electronic configurations. We
are here only concerned with the ground-state electronic configuration,
which has the lowest energy. The ground-state electronic configuration of an
atom can be determined by the following three principles.
\ufffd The Aufbau principle states that the orbitals fill in order of increasing
energy, from lowest to highest. Because a 1s orbital is closer to the
nucleus it is lower in energy than a 2s orbital, which is lower in energy
than a 3s orbital.
\ufffd The Pauli exclusion principle states that no more than two electrons can
occupy each orbital, and if two electrons are present, their spins must be
paired. For example, the two electrons of a helium atom must occupy the
1s orbital in opposite spins.
\ufffd Hund\u2019s rule explains that when degenerate orbitals (orbitals that have
same energy) are present but not enough electrons are available to fill all
the shell completely, then a single electron will occupy an empty orbital
first before it will pair up with another electron. This is understandable, as
it takes energy to pair up electrons. Therefore, the six electrons in the
carbon atom are filled as follows: the first four electrons will go to the 1s
and 2s orbitals, a fifth electron goes to the 2px, the sixth electron to the
2py orbital and the 2pz orbital will remain empty.
The ground-state electronic configurations for elements 1\u201318 are listed
below (electrons are listed by symbol, atomic number and ground-state
electronic configuration).
Shell Number of orbitals contained each shell
4 4s, 4px, 4py, 4pz, five 4d, seven 4f
3 3s, 3px, 3py, 3pz, five 3d
2 2s, 2px, 2py, 2pz
1 1s
First period Second period Third period
H 1 1s1 Li 3 [He] 2s1 Na 11 [Ne] 3s1
He 2 1s2 Be 4 [He] 2s2 Mg 12 [Ne] 3s2
B 5 [He] 2s2 2p1 Al 13 [Ne] 3s2 3p1
C 6 [He] 2s2 2p2 Si 14 [Ne] 3s2 3p2
7 [He] 2s2 2p3 P 15 [Ne] 3s2 3p3
8 [He] 2s2 2p4 S 16 [Ne] 3s2 3p4
9 [He] 2s2 2p5 Cl 17 [Ne] 3s2 3p5
10 [He] 2s2 2p6 Ar 18 [Ne] 3s2 3p6
Let us see how we can write the ground-state electronic configurations for
oxygen, chlorine, nitrogen, sulphur and carbon showing the occupancy of
each p orbital. Oxygen has the atomic number 8, and the ground-state
electronic configuration for oxygen can be written as 1s2 2s2 2px
2 2py
1 2pz
Similarly, we can write the others as follows:
Chlorine (atomic number 17): 1s2 2s2 2px
2 2py
2 2pz
2 3s2 3px
2 3py
2 3pz
Nitrogen (atomic number 7): 1s2 2s2 2px
1 2py
1 2pz
Sulphur (atomic number 16): 1s2 2s2 2px
2 2py
2 2pz
2 3s2 3px
2 3py
1 3pz
Carbon (atomic number 6): 1s2 2s2 2px
1 2py
1 2pz
2.3 Chemical bonding theories: formation
of chemical bonds
Atoms form bonds in order to obtain a stable electronic configuration, i.e.
the electronic configuration of the nearest noble gas. All noble gases are
inert, because their atoms have a stable electronic configuration in which
they have eight electrons in the outer shell except helium (two electrons).
Therefore, they cannot donate or gain electrons.
One of the driving forces behind the bonding in an atom is to obtain a
stable valence electron configuration. A filled shell is also known as a noble
gas configuration. Electrons in filled shells are called core electrons. The
core electrons do not participate in chemical bonding. Electrons in shells
that are not completely filled are called valence electrons, also known as
outer-shell electrons, and the energy level in which they are found is also
known as the valence shell. Carbon, for example, with the ground-state
electronic configuration 1s2 2s2 2p2, has four outer-shell electrons. We
generally use the Lewis structure to represent the outermost electrons of an
2.3.1 Lewis structures
Lewis structures provide information about what atoms are bonded to each
other, and the total electron pairs involved. According to the Lewis theory,
an atom will give up, accept or share electrons in order to achieve a filled
outer shell that contains eight electrons. The Lewis structure of a covalent
molecule shows all the electrons in the valence shell of each atom; the bonds
between atoms are shown as shared pairs of electrons. Atoms are most
stable if they have a filled valence shell of electrons. Atoms transfer or share
electrons in such a way that they can attain a filled shell of electrons. This
stable configuration of electrons is called an octet. Except for hydrogen and
helium, a filled valence shell contains eight electrons.
Lewis structures help us to track the valence electrons and predict the
types of bond. The number of valence electrons present in each of the
elements is to be considered first. The number of valence electrons
determines the number of electrons needed to complete the octet of eight