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Simple ions are atoms that have gained or lost electrons to satisfy
the octet rule. However, not all compounds follow the octet rule.
Elements in organic compounds are joined by covalent bonds, a sharing
of electrons, and each element contributes one electron to the bond. The
number of electrons necessary to complete the octet determines the number
of electrons that must be contributed and shared by a different element in a
bond. This analysis finally determines the number of bonds that each
element may enter into with other elements. In a single bond two atoms
share one pair of electrons and form a s bond. In a double bond they share
two pairs of electrons and form a s bond and a p bond. In a triple bond two
atoms share three pairs of electrons and form a s bond and two p bonds.
Sodium (Na) loses a single electron from its 3s orbital to attain a more
stable neon gas configuration (1s2 2s2 2p6) with no electron in the outer
shell. An atom having a filled valence shell is said to have a closed shell
configuration. The total number of electrons in the valence shell of each
atom can be determined from its group number in the periodic table. The
shared electrons are called the bonding electrons and may be represented by
a line or lines between two atoms. The valence electrons that are not being
shared are the nonbonding electrons or lone pair electrons, and they are
shown in the Lewis structure by dots around the symbol of the atom. A
species that has an unpaired electron are called radicals. Usually they are
very reactive, and are believed to play significant roles in aging, cancer and
many other ailments.
In neutral organic compounds, C forms four bonds, N forms three bonds
(and a lone pair), O forms two bonds (and two lone pairs) and H forms one
bond. The number of bonds an atom normally forms is called the valence.
Lewis structure shows the connectivity between atoms in a molecule by a
number of dots equal to the number of electrons in the outer shell of an atom of
that molecule. A pair of electrons is represented by two dots, or a dash. When
drawing Lewis structures, it is essential to keep track of the number of electrons
available to form bonds and the location of the electrons. The number of
valence electrons of an atom can be obtained from the periodic table because it
is equal to the group number of the atom. For example, hydrogen (H) in Group
1A has one valence electron, carbon (C) in Group 4A has four valence
electrons, and fluorine (F) in Group 7A has seven valence electrons.
To write the Lewis formula of CH3F, first of all, we have to find the total
number of valence electrons of all the atoms involved in this structure, i.e.
C, H and F, having four, one and seven valence electrons, respectively.
4 þ 3ð1Þþ7¼14
C 3H H
The carbon atom bonds with three hydrogen atoms and one fluorine atom,
and it requires four pairs of electrons. The remaining six valence electrons
are with the fluorine atom in the three nonbonding pairs.
In the periodic table, the period 2 elements C, N, O, and F have valence
electrons that belong to the second shell (2s and three 2p). The shell can be
completely filled with eight electrons. In period 3, elements Si, P, S and Cl
have the valence electrons that belong to the third shell (3s, three 3p and
five 3d ). The shell is only partially filled with eight electrons in 3s and
three 3p, and the five 3d orbitals can accommodate an additional ten
electrons. For these differences in valence shell orbitals available to
elements of the second and third periods, we see significant differences
in the covalent bonding of oxygen and sulphur, and of nitrogen and
phosphorus. Although oxygen and nitrogen can accommodate no more
than eight electrons in their valence shells, many phosphorus-containing
compounds have 10 electrons in the valence shell of phosphorus, and many
sulphur-containing compounds have 10 and even 12 electrons in the
valence shell of sulphur.
So, to derive Lewis structures for most molecules the following sequence
should be followed.
(a) Draw a tentative structure. The element with the least number of atoms is
usually the central element.
(b)Calculate the number of valence electrons for all atoms in the compound.
(c) Put a pair of electrons between each symbol.
(d)Place pairs of electrons around atoms beginning with the outer atom until
each has eight electrons, except for hydrogen. If an atom other than
hydrogen has fewer than eight electrons then move unshared pairs to
form multiple bonds.
If the structure is an ion, electrons are added or subtracted to give the proper
charge. Lewis structures are useful as they show what atoms are bonded
together, and whether any atoms possess lone pairs of electrons or have a
formal charge. A formal charge is the difference between the number of
valence electrons an atom actually has when it is not bonded to any other
atoms, and the number of nonbonding electrons and half of its bonding
electrons. Thus, a positive or negative charge assigned to an atom is called a
formal charge. The decision as to where to put the charge is made by
calculating the formal charge for each atom in an ion or a molecule. For
example, the hydronium ion (H3O
þ) is positively charged and the oxygen
atom has a formal charge of þ1.
Assigned 5 valence electrons: 
formal charge of +1
So; formal charge ¼ ðgroup numberÞ \ufffd ðnonbonding electronsÞ \ufffd 1=2
ðshared electronsÞ
¼ 6 \ufffd 2 \ufffd 1=2ð6Þ
¼ 1:
An uncharged oxygen atom must have six electrons in its valence shell. In
the hydronium ion, oxygen bonds with three hydrogen atoms. So, only five
electrons effectively belong to oxygen, which is one less than the valence
electrons. Thus, oxygen bears a formal charge of þ1. Elements of the
second period, including carbon, nitrogen, oxygen and fluorine, cannot
accommodate more than eight electrons as they have only four orbitals (2s,
2px, 2py and 2pz) in their valence shells.
2.3.2 Various types of chemical bonding
A chemical bond is the attractive force that holds two atoms together.
Valence electrons take part in bonding. An atom that gains electrons
becomes an anion, a negatively charged ion, and an atom that loses
electrons becomes a cation, a positively charged ion. Metals tend to lose
electrons and nonmetals tend to gain electrons. While cations are smaller
than atoms, anions are larger. Atoms decrease in size as they go across a
period, and increase in size as they go down a group and increase the
number of shells to hold electrons.
The energy required for removing an electron from an atom or ion in the
gas phase is called ionization energy. Atoms can have a series of ionization
energies, since more than one electron can always be removed, except for
hydrogen. In general, the first ionization energies increase across a period
and decrease down the group. Adding more electrons is easier than
removing electrons. It requires a vast amount of energy to remove
Ionic bonds
Ionic bonds result from the transfer of one or more electrons between
atoms. The more electronegative atom gains one or more valence
electrons and hence becomes an anion. The less electronegative atom
loses one or more valence electrons and becomes a cation. A single-
headed arrow indicates a single electron transfer from the less electro-
negative element to the more electronegative atom. Ionic compounds are
held together by the attraction of opposite charges. Thus, ionic bonds
consist of the electrostatic attraction between positively and negatively