There are molecules and ions for which drawing a single Lewis structure is not possible. For example, we can write two structures of O3.
In (A) the oxygen-oxygen bond on the left is a double bond and the oxygen-oxygen bond on the right is a single bond. In B the situation is just the opposite. The experiment shows, however, that the two bonds are identical.
Therefore neither structure A nor B can be correct. One of the bonding pairs in ozone is spread over the region of all three atoms rather than localized on a particular oxygen-oxygen bond. This delocalized bonding is a type of chemical bonding in which bonding pair of electrons are spread over a number of atoms rather than localized between two.
Structures (A) and (B) are called resonating or canonical structures and (C) is the resonance hybrid. This phenomenon is called resonance, a situation in which more than one canonical structure can be written for a species. The chemical activity of an atom is determined by the number of electrons in its valence shell. With the help of the concept of chemical bonding, one can define the structure of a compound and is used in many industries for manufacturing products in which the true structure cannot be written at all.
Some other examples:
Carbon-oxygen bond lengths in carboxylate ion are equal due to resonance.
Greater the size of the atom, greater is the bond length and less is the bond dissociation enthalpy i.e. less is the bond strength during chemical bonding.
Multiplicity of Bonds
Greater is the multiplicity of the bond, greater is the bond dissociation enthalpy.
Number of Lone Pair of Electrons Present
More the number of lone pair of electrons present on the bonded atoms, greater is the repulsion between the atoms and thus less is the bond dissociation enthalpy of the chemical bond.
Bond Angle
A bond is formed by the overlap of atomic orbitals. The direction of overlap gives the direction of the bond. The angle between the lines representing the direction of the bond i.e. the orbitals containing the bonding electrons is called the bond angle.
Bond Order
In Lewis representation, the number of bonds present between two atoms is called the bond order. Greater the bond order, greater is the stability of the bond during chemical bonding i.e. greater is the bond enthalpy. Greater the bond order, shorter is the bond length.
During chemical bonding, when the atoms come closer to each other, the attraction takes place between them and the potential energy of the system keeps on decreasing till a particular distance at which the potential energy is minimum. If the atoms come more closer, repulsion starts and again the potential energy of the system begins to increase.
At equilibrium distance, the atoms keep on vibrating about their mean position. The equilibrium distance between the centres of the nuclei of the two bonded atoms is called its Bond length.
It is expressed in terms of an angstrom (A0) or picometer (pm). It is determined experimentally by x-ray diffraction or electron diffraction method or spectroscopic method. The bond length in chemical bonding is the sum of their ionic radii, in an ionic compound. In a covalent compound, it is the sum of their covalent radii. For a covalent molecule AB, the bond length is given by d= ra + rb
Factors Affecting the Bond length
Size of the atoms: The bond length increases with increase in the size of the atom. HI > HBr > HCl > HF
The multiplicity of Bond: The bond length decreases with an increase in bond order.
Type of hybridization: A‘s’ orbital is smaller in size, greater the ‘s’ character, shorter is the bond length.
Bond Enthalpy
When atoms come close together the energy is released due to the chemical bonding between them. The amount of energy required to break one mole of bonds of a type so as to separate the molecule into individual gaseous atoms is called bond dissociation enthalpy or Bond enthalpy. Bond enthalpy is usually expressed in KJ mol-1.
Greater is the bond dissociation enthalpy, greater is the bond strength. For diatomic molecules like H2, Cl2, O2, N2, HCl, HBr, HI the bond enthalpies are equal to their dissociation enthalpy.
In the case of polyatomic molecules, bond enthalpies are usually the average values, because the dissociation energy varies with each type of bond.
In H20, first O-H bond enthalpy = 502 KJ/mol; Second bond enthalpy = 427 KJ/mol Average bond enthalpy = (502 + 427) / 2 = 464.5 KJ/mol
The bond formed as a result of strong electrostatic forces of attraction between a positively and negatively charged species is called an electrovalent or ionic bond. The positively and negatively charged ions are aggregated in an ordered arrangement called the crystal lattice which is stabilized by the energy called the Lattice enthalpy.
Conditions for the formation of an Ionic Bond
The low ionization energy of the atom forming the cation.
High electron gain enthalpy of the atom forming the anion.
High negative lattice enthalpy of the crystal formed.
Generally, the ionic bond is formed between a metal cation and non-metal anion.
Writing Lewis Structures
The following steps are adopted for writing the Lewis dot structures or Lewis structures:
Step 1: Calculate the number of electrons required for drawing the structure by adding the valence electrons of the combining atoms. For Example, in methane, CH4 molecule, there are 8 valence electrons (in which 4 belongs to carbon while other 4 to H atoms).
Step 2: Each negative charge i.e. for anions, we add an electron to the valence electrons and for each positive charge i.e. for cations we subtract one electron from the valence electrons.
Step 3: Using the chemical symbols of the combining atoms and constructing a skeletal structure of the compound, divide the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.
Step 4: The central position in the molecule is occupied by the least electronegative atom. Hydrogen and fluorine generally occupy the terminal positions.
Step 5: After distributing the shared pairs of electrons for single bonds, the remaining electron pairs are used for multiple bonds or they constitute lone pairs.
The basic requirement is that each bonded atom gets an octet of electrons.
Example 1: Lewis formula for carbon monoxide, CO
Step 1: Counting the total number of valence electrons of carbon and oxygen atoms: C (2s22p2) + O (2s22p4) 4 + 6 = 10 that is, 4(C) + 6(O) = 10
Step 2: The skeletal structure of carbon monoxide is written as CO
Step 3: Drawing a single bond between C and O and completing octet on O, the remaining two electrons are lone pair on C.
Step 4: This does not complete the octet of carbon, and hence we have a triple bond.
Example 2: Lewis Structure of nitrite, NO2–
Step 1: Counting the total number of valence electrons of one nitrogen atom, two oxygen atoms and the additional one negative charge (equal to one electron). Total Number of valence electrons is: N (2s22p3) + 2O (2s22p4) + 1 (negative charge) => 5+ 2(6) +1=18e–
Step 2: The skeletal structure of nitrite ion is written as O-N-O
Step 3: Drawing a single bond between nitrogen and each oxygen atom: O – N – O
Step 4: Complete the octets of atoms.
This structure does not complete octet on N if the remaining two electrons constitute of a lone pair on it. Therefore, we have a double bond between one N and one of the two O atoms. The Lewis structure is
Compared to ionic and covalent bonding, Hydrogen bonding is a weaker form of chemical bonding. It is a type of polar covalent bonding between oxygen and hydrogen wherein the hydrogen develops a partial positive charge. This implies that the electrons are pulled closer to the more electronegative oxygen atom.
This creates a tendency for the hydrogen to be attracted towards the negative charges of any neighbouring atom. This type of chemical bonding is called a hydrogen bond and is responsible for many of the properties exhibited by water.
A covalent bond indicates the sharing of electrons between atoms. Compounds that contain carbon (also called organic compounds) commonly exhibit this type of chemical bonding. The pair of electrons which are shared by the two atoms now extend around the nuclei of atoms, leading to the creation of a molecule.
Polar Covalent Bonding
Covalent bonds can be either be Polar or Non-Polar in nature. In Polar Covalent chemical bonding, electrons are shared unequally since the more electronegative atom pulls the electron pair closer to itself and away from the less electronegative atom. Water is an example of such a polar molecule.
A difference in charge arises in different areas of the atom due to the uneven spacing of the electrons between the atoms. One end of the molecule tends to be partially positively charged and the other end tends to be partially negatively charged.
Ionic bonding is a type of chemical bonding which involves a transfer of electrons from one atom or molecule to another. Here, an atom loses an electron which is in turn gained by another atom. When such an electron transfer takes place, one of the atoms develops a negative charge and is now called the anion.
The other atom develops a positive charge and is called the cation. The ionic bond gains strength from the difference in charge between the two atoms, i.e. the greater the charge disparity between the cation and the anion, the stronger the ionic bond.
When substances participate in chemical bonding and yield compounds, the stability of the resulting compound can be gauged by the type of chemical bonds it contains.
The type of chemical bonds formed vary in strength and properties. There are 4 primary types of chemical bonds which are formed by atoms or molecules to yield compounds. These types of chemical bonds include:
Ionic Bonds
Covalent Bonds
Hydrogen Bonds
Polar Bonds
These types of bonds in chemical bonding are formed from the loss, gain, or sharing of electrons between two atoms/molecules.
Albrecht Kössel and Gilbert Lewis were the first to explain the formation of chemical bonds successfully in the year 1916. They explained chemical bonding on the basis of the inertness of noble gases.
Lewis Theory of Chemical Bonding
An atom can be viewed as a positively charged ‘Kernel’ (the nucleus plus the inner electrons) and the outer shell.
The outer shell can accommodate a maximum of eight electrons only.
The eight electrons present in the outer shell occupy the corners of a cube which surround the ‘Kernel’.
The atoms having octet configuration, i.e. 8 electrons in the outermost shell, thus symbolize a stable configuration.
Atoms can achieve this stable configuration by forming chemical bonds with other atoms. This chemical bond can be formed either by gaining or losing an electron(s) (NaCl, MgCl2) or in some cases due to the sharing of an electron (F2).
Only the electrons present in the outer shell, also known as the valence electrons take part in the formation of chemical bonds. Gilbert Lewis used specific notations better known as Lewis symbols to represent these valence electrons.
Generally, the valency of an element is either equal to the number of dots in the corresponding Lewis symbol or 8 minus the number of dots (or valence electrons).
Lewis symbols for lithium (1 electron), oxygen (6 electrons), neon (8 electrons) are given below:
Kossel’s theory of Chemical Bonding
Noble gases separate the highly electronegative halogens and the highly electropositive alkali metals.
Halogens can form negatively charged ions by gaining an electron. Whereas alkali metals can form positively charged ions by losing an electron.
These negatively charged ions and positively charged ions have a noble gas configuration that is 8 electrons in the outermost shell. The general electronic configuration of noble gases (except helium) is given by ns2np6.
As unlike charges attract each other these unlike charged particles are held together by a strong force of electrostatic attraction existing between them. For example, MgCl2, the magnesium ion, and chlorine ions are held together by force of electrostatic attraction. This kind of chemical bonding existing between two unlike charged particles is known as an electrovalent bond.
Explanation of Kossel Lewis Approach
In 1916 Kossel and Lewis succeeded in giving a successful explanation based upon the concept of an electronic configuration of noble gases about why atoms combine to form molecules. Atoms of noble gases have little or no tendency to combine with each other or with atoms of other elements. This means that these atoms must be having stable electronic configurations.
Due to the stable configuration, the noble gas atoms neither have any tendency to gain or lose electrons and, therefore, their combining capacity or valency is zero. They are so inert that they even do not form diatomic molecules and exist as monoatomic gaseous atoms.
Covalent bonds have certain characteristics that depend on the identities of the atoms participating in the bond. Two characteristics are bond length and bond polarity.
Bond length
The covalent bond in the hydrogen molecule (H2) has a certain length (about 7.4 × 10−11 m). Other covalent bonds also have known bond lengths, which are dependent on both the identities of the atoms in the bond and whether the bonds are single, double, or triple bonds. Table 1.1. lists the approximate bond lengths for some single covalent bonds. The exact bond length may vary depending on the identity of the molecule but will be close to the value given in the table.
Table 1.1. Approximate bond lengths of some single bonds
Bond
Length (× 10−12 m)
H–H
74
H–C
110
H–N
100
H–O
97
H–I
161
C–C
154
C–N
147
C–O
143
N–N
145
O–O
145
Table b. compares the lengths of single covalent bonds with those of double and triple bonds between the same atoms. Without exception, as the number of covalent bonds between two atoms increases, the bond length decreases. With more electrons between the two nuclei, the nuclei can get closer together before the internuclear repulsion is strong enough to balance the attraction.
Table b. Comparison of bond lengths for single and multiple bonds
Bond
Length (× 10−12 m)
C–C
154
C=C
134
C≡C
120
C–N
147
C=N
128
C≡N
116
C–O
143
C=O
120
C≡O
113
N–N
145
N=N
123
N≡N
110
O–O
145
O=O
121
Electronegativity and bond polarity
Polar versus nonpolar covalent bonds. (a) The electrons in the covalent bond are equally shared by both hydrogen atoms. This is a nonpolar covalent bond. (b) The fluorine atom attracts the electrons in the bond more than the hydrogen atom does, leading to an imbalance in the electron distribution. This is a polar covalent bond.
Although we defined covalent bonding as electron sharing, the electrons in a covalent bond are not always shared equally by the two bonded atoms. Unless the bond connects two atoms of the same element, there will always be one atom that attracts the electrons in the bond more strongly than the other atom does,“Polar versus Nonpolar Covalent Bonds”. When such an imbalance occurs, there is a resulting buildup of some negative charge (called a partial negative charge and designated δ−) on one side of the bond and some positive charge (designated δ+) on the other side of the bond. A covalent bond that has an unequal sharing of electrons, as in part (b) of Figure , is called a polar covalent bond. A covalent bond that has an equal sharing of electrons (part (a) is called a nonpolar covalent bond.
Any covalent bond between atoms of different elements is a polar bond, but the degree of polarity varies widely. Some bonds between different elements are only minimally polar, while others are strongly polar. Ionic bonds can be considered the ultimate in polarity, with electrons being transferred rather than shared. To judge the relative polarity of a covalent bond, chemists use electronegativity, which is a relative measure of how strongly an atom attracts electrons when it forms a covalent bond. There are various numerical scales for rating electronegativity. Figure b. “Electronegativities of Various Elements” shows one of the most popular—the Pauling scale. The polarity of a covalent bond can be judged by determining the difference in the electronegativities of the two atoms making the bond. The greater the difference in electronegativities, the greater the imbalance of electron sharing in the bond. Although there are no hard and fast rules, the general rule is if the difference in electronegativities is less than about 0.4, the bond is considered nonpolar; if the difference is greater than 0.4, the bond is considered polar. If the difference in electronegativities is large enough (generally greater than about 1.8), the resulting compound is considered ionic rather than covalent. An electronegativity difference of zero, of course, indicates a nonpolar covalent bond.
When a molecule’s bonds are polar, the substance is usually found to be polar. The polarity of water has an enormous impact on its physical and chemical properties. (For example, the boiling point of water [100°C] is high for such a small molecule and is due to the fact that polar molecules attract each other strongly.)
Figure b. Electronegativities of Various Elements. A popular scale for electronegativities has the value for fluorine atoms set at 4.0, the highest value.
