Chemical bonding and Molecular structure

Chemical bond is the physical and chemical process responsible for the attractive interactions between atoms molecules, ions and that which confers stability to diatomic and polyatomic chemical entities. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics. In general, chemical bonding is associated with the sharing or transfer of electrons between the participating atoms.

The study on the “nature of forces that hold or bind atoms or ions or molecules or complex molecules together to form a molecule” is required to gain knowledge of the following

i)    To know about how atoms of same element form different compounds combining with different elements.

ii)   To know why particular shapes are adopted by molecules.

iii)  To understand the specific properties of molecules or ions and the relation between the specific type of bonding in the molecules.

Kossel-Lewis approach to Chemical Bonding

In 1916, W.Kossel and G.N.Lewis, separately developed theories of chemical bonding inorder to understand why atoms combined to form molecules. Lewis introduced simple symbols to denote the electron present in the outer orbit of atom , these electrons are known as valence electrons. These symbols are known as electron dot symbols and the structure of compound is known as Lewis dot structure.

The electron dot structure of on then is given as  The number of dots around the symbol equal to number of electrons.

INTRODUCTION

Introduction

Kossel-Lewis Approach to Chemical Bonding

Kossel-Lewis Approach to Chemical Bonding

In order to explain the formation of chemical bond in terms of electrons, a number of attempts were made, but it was only in 1916 when Kössel and Lewis succeeded independently in giving a satisfactory explanation. They were the first to provide some logical explanation of valence which was based on the inertness of noble gases.

Lewis pictured the atom in terms of a positively charged ‘Kernel’ (the nucleus plus the inner electrons) and the outer shell that could accommodate a maximum of eight electrons. He, further assumed that these eight electrons occupy the corners of a cube which surround the ‘Kernel’. Thus the single outer shell electron of sodium would occupy one corner of the cube, while in the case of a noble gas all the eight corners would be occupied. This octet of electrons, represents a particularly stable electronic arrangement.

Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. In the case of sodium and chlorine, this can happen by the transfer of an electron from sodium to chlorine there by giving the Na⁺ and Cl– ions. In the case of other molecules like Cl₂, H₂, F₂, etc., the bond is formed by the sharing of a pair of electrons between the atoms. In the process each atom attains a stable outer octet of electrons.

Lewis Symbols: In the formation of a molecule, only the outer shell electrons take part in chemical combination and they are known as valence electrons. The inner shell electrons are well protected and are generally not involved in the combination process.

G.N. Lewis, an American chemist introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols. For example, the Lewis symbols for the elements of second period are as under:

Significance of Lewis Symbols : The number of dots around the symbol represents 

the number of valence electrons. This number of valence electrons helps to calculate the common or group valence of the element. The group valence of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.

Kössel, in relation to chemical bonding, drew attention to the following facts:

1. In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases;

2. The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of  an  electron by the respective atoms;

3. The negative and positive ions thus formed attain stable noble gas electronic configurations. The noble gases (with the exception of helium which has a duplet of electrons) have a particularly stable outer shell configuration of eight (octet) electrons, ns²np⁶.

4. The negative and positive ions are stabilized by electrostatic attraction.

5.For example, the formation of NaCl from sodium and chlorine, according to the above scheme, can be explained as :

Na                   \(\to\)         Na⁺ +   e^–

[Ne] 3s₁                       [Ne]

Cl + e^–          \(\to\)         Cl^–

[Ne] 3s² 3p⁵              [Ne] 3s² 3p⁶ or [Ar]

Na+ + Cl^–     \(\to\)            NaCl or Na⁺Cl^–

Similarly the formation of CaF₂ may be shown as:

Ca                  \(\to\)      Ca²⁺ + 2e^– [Ar]

4s²                          [Ar]

F + e^–              \(\to\)       F^–

[He] 2s² 2p⁵           [He] 2s² 2p⁶ or [Ne]

Ca²⁺ + 2F^–     \(\to\)       CaF     or Ca²⁺(F^– )

The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the electrovalent bond. The electrovalence is thus equal to the number of unit charge(s) on the ion. Thus, calcium is assigned a positive electrovalence of two, while chlorine a negative electrovalence of one.

Kössel’s postulations provide the basis for the modern concepts regarding ion-formation by electron transfer and the formation of ionic crystalline compounds. His views have proved to be of great value in the understanding and systematisation of the ionic compounds. At the same time he did recognise the fact that a large number of compounds did not fit into these concepts.

Octet Rule and Covalent Bond

Octet Rule and Covalent Bond

Octet

Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding. According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells.This is known as octet rule.

Covalent Bond

Langmuir (1919) refined the Lewis postulations by abandoning the idea of the stationary cubical arrangement of the octet, and by introducing the term covalent bond.The Lewis-Langmuir theory can be understood by considering the formation of the chlorine molecule,Cl₂. The Cl atom with electronic configuration, [Ne]3s² 3p⁵, is one electron short of the argon configuration.The formation of the Cl2 chlorine atoms, each chlorine atom contributing one electron to the shared pair. In the process both chlorine atoms attain the outer shell octet of the nearest noble gas (i.e., argon).

Covalent bond between two Cl atoms

The dots represent electrons. Such structures are referred to as Lewis dot structures.

The Lewis dot structures can be written for other molecules also, in which the combining atoms may be identical or different. The important conditions being that:

1. Each bond is formed as a result of sharing of an electron pair between the atoms.

2. Each combining atom contributes at least one electron to the shared pair.

3. The combining atoms attain the outer- shell noble gas configurations as a result of the sharing of electrons.

4. Thus in water and carbon tetrachloride molecules, formation of covalent bonds can be represented as:

Thus, when two atoms share one electron pair they are said to be joined by a single covalent bond. In many compounds we have multiple bonds between atoms.The formation of multiple bonds envisages sharing of more than one electron pair between two atoms.If two atoms share two pairs of electrons, the covalent bond between them is called a double bond.For example, in the carbon dioxide molecule, we have two double bonds between the carbon and oxygen atoms.Similarly in ethene molecule the two carbon atoms are joined by a double bond.

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When combining atoms share three electron pairs as in the case of two nitrogen atoms in the N₂ molecule and the two carbon atoms in the ethyne molecule,a triple bond is formed.

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Lewis Representation of Simple Molecules (The Lewis Structures)

Lewis Representation of Simple Molecules (The Lewis Stuctures)

The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule. While such a picture may not explain the bonding and behaviour of a molecule completely, it does help in understanding the formation and properties of a molecule to a large extent. Writing of Lewis dot structures of molecules is,therefore,very useful.The Lewis dot structures can be written by adopting the following steps:

1. The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms. For example, in the CH₄ molecule there are eight valence electrons available for bonding (4 from carbon and 4 from the four hydrogen atoms).

2. For anions, each negative charge would mean addition of one electron.For cations, each positive charge would result in subtraction of one electron from the total number of valence electrons. For example, for the CO^2– ion, the two negative charges indicate that there are two additional electrons than those provided by the neutral atoms. For NH⁺ ion, one positive charge indicates the loss of one electron from the group of neutral atoms.

3. Knowing the chemical symbols of the combining atoms and having knowledge of the skeletal structure of the compound (known or guessed intelligently), it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.

4. In general the least electronegative atom occupies the central position in the molecule/ion.For example in the NF₃ and CO₃^2–, nitrogen and carbon are the central atoms whereas fluorine and oxygen occupy the terminal positions.

5. After accounting for the shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs. The basic requirement being that each bonded atom gets an octet of electrons.

Lewis representations of a few molecules/ ions are given in Table 4.1.

                   Table 4.1 The Lewis Representation of Some Molecules

PROBLEM 1

Write the Lewis dot structure of CO molecule.

SOLUTION

Step 1. Count the total number of valence electrons of carbon and oxygen atoms.The outer (valence) shell configurations of carbon and oxygen atoms are: 2s² 2 p² and 2 s² 2p ⁴, respectively. The valence electrons available are 4 + 6 =10.

Step 2. The skeletal structure of CO is written as: C O

Step 3. Draw a single bond (one shared electron pair) between C and O and complete the octet on O, the remaining two electrons are the lone pair on C.

This does not complete the octet on carbon and hence we have to resort to multiple bonding (in this case a triple bond) between C and O atoms. This satisfies the octet rule condition for both atoms.

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PROBLEM 2

Write the Lewis structure of the nitrite ion,\(% MathType!MTEF!2!1!+- % feaagKart1ev2aaatCvAUfeBSjuyZL2yd9gzLbvyNv2CaerbuLwBLn % hiov2DGi1BTfMBaeXatLxBI9gBaerbd9wDYLwzYbItLDharqqtubsr % 4rNCHbGeaGqiVu0Je9sqqrpepC0xbbL8F4rqqrFfpeea0xe9Lq-Jc9 % vqaqpepm0xbba9pwe9Q8fs0-yqaqpepae9pg0FirpepeKkFr0xfr-x % fr-xb9adbaqaaeGaciGaaiaabeqaamaabaabaaGcbaGaamOtaiaad+ % eadaqhaaWcbaGaaGOmaaqaaiabgkHiTaaaaaa!3973! NO_2^ - \).

SOLUTION

Step 1. Count the total number of valence electrons of the nitrogen atom, the oxygen atoms and the additional one negative charge (equal to one electron).

N(2s² 2p³), O (2s² 2p⁴)

5 + (2 × 6) +1 = 18 electrons.

Step 2. The skeletal structure of \(% MathType!MTEF!2!1!+- % feaagKart1ev2aaatCvAUfeBSjuyZL2yd9gzLbvyNv2CaerbuLwBLn % hiov2DGi1BTfMBaeXatLxBI9gBaerbd9wDYLwzYbItLDharqqtubsr % 4rNCHbGeaGqiVu0Je9sqqrpepC0xbbL8F4rqqrFfpeea0xe9Lq-Jc9 % vqaqpepm0xbba9pwe9Q8fs0-yqaqpepae9pg0FirpepeKkFr0xfr-x % fr-xb9adbaqaaeGaciGaaiaabeqaamaabaabaaGcbaGaamOtaiaad+ % eadaqhaaWcbaGaaGOmaaqaaiabgkHiTaaaaaa!3973! NO_2^ - \) is 

written as : O   N   O

Step 3. Draw a single bond (one shared electron pair) between the nitrogen and each of the oxygen atoms completing the octets on oxygen atoms. This, however, does not complete the octet on nitrogen if the remaining two electron constitute lone pair on it.

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Hence we have to resort to multiple bonding between nitrogen and one of the oxygen atoms (in this case a double bond). This leads to dot structures.

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Formal Charge

Formal Charge

Lewis dot structures, in general, do not represent the actual shapes of the molecules. In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom. It is, however, feasible to assign a formal charge on each atom. The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It is expressed as :

Formal charge (F.C.) on an atom in a Lewis structure =[total number of valence electrons in the free atom]-[total number of non bonding (lone pair) electrons]-(1/2)[total number of bonding(shared) electrons]

The counting is based on the assumption that the atom in the molecule owns one electron of each shared pair and both the electrons of a lone pair.

Let us consider the ozone molecule (O₃). The Lewis structure of O3 may be drawn as :

The atoms have been numbered as 1, 2 and 3.The formal charge on:

1. The central O atom marked 1

 \(% MathType!MTEF!2!1!+- % feaagKart1ev2aaatCvAUfeBSjuyZL2yd9gzLbvyNv2CaerbuLwBLn % hiov2DGi1BTfMBaeXatLxBI9gBaerbd9wDYLwzYbItLDharqqtubsr % 4rNCHbGeaGqiVu0Je9sqqrpepC0xbbL8F4rqqrFfpeea0xe9Lq-Jc9 % vqaqpepm0xbba9pwe9Q8fs0-yqaqpepae9pg0FirpepeKkFr0xfr-x % fr-xb9adbaqaaeGaciGaaiaabeqaamaabaabaaGcbaGaeyypa0JaaG % OnaiabgkHiTiaaikdacqGHsisldaWcaaqaaiaaigdaaeaacaaIYaaa % amaabmaabaGaaGOnaaGaayjkaiaawMcaaiabg2da9iabgUcaRiaaig % daaaa!40C5! = 6 - 2 - \frac{1}{2}\left( 6 \right) = + 1\)

2. The end O atom marked 2

\(% MathType!MTEF!2!1!+- % feaagKart1ev2aaatCvAUfeBSjuyZL2yd9gzLbvyNv2CaerbuLwBLn % hiov2DGi1BTfMBaeXatLxBI9gBaerbd9wDYLwzYbItLDharqqtubsr % 4rNCHbGeaGqiVu0Je9sqqrpepC0xbbL8F4rqqrFfpeea0xe9Lq-Jc9 % vqaqpepm0xbba9pwe9Q8fs0-yqaqpepae9pg0FirpepeKkFr0xfr-x % fr-xb9adbaqaaeGaciGaaiaabeqaamaabaabaaGcbaGaeyypa0JaaG % OnaiabgkHiTiaaisdacqGHsisldaWcaaqaaiaaigdaaeaacaaIYaaa % amaabmaabaGaaGinaaGaayjkaiaawMcaaiabg2da9iaaicdaaaa!3FE2! = 6 - 4 - \frac{1}{2}\left( 4 \right) = 0\)

3. The end O atom marked 3

\(% MathType!MTEF!2!1!+- % feaagKart1ev2aaatCvAUfeBSjuyZL2yd9gzLbvyNv2CaerbuLwBLn % hiov2DGi1BTfMBaeXatLxBI9gBaerbd9wDYLwzYbItLDharqqtubsr % 4rNCHbGeaGqiVu0Je9sqqrpepC0xbbL8F4rqqrFfpeea0xe9Lq-Jc9 % vqaqpepm0xbba9pwe9Q8fs0-yqaqpepae9pg0FirpepeKkFr0xfr-x % fr-xb9adbaqaaeGaciGaaiaabeqaamaabaabaaGcbaGaeyypa0JaaG % OnaiabgkHiTiaaiAdacqGHsisldaWcaaqaaiaaigdaaeaacaaIYaaa % amaabmaabaGaaGOmaaGaayjkaiaawMcaaiabg2da9iabgkHiTiaaig % daaaa!40D0! = 6 - 6 - \frac{1}{2}\left( 2 \right) = - 1\)

Hence, we represent O₃ along with the formal charges as follows:

We must understand that formal charges do not indicate real charge separation within the molecule. Indicating the charges on the atoms in the Lewis structure only helps in keeping track of the valence electrons in the molecule. Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species. Generally the lowest energy structure is the one with the smallest formal charges on the atoms. The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighbouring atoms.

Limitations of the Octet Rule

Limitations of the Octet Rule

The octet rule, though useful, is not universal. It is quite useful for understanding the structures of most of the organic compounds and it applies mainly to the second period elements of the periodic table. There are three types of exceptions to the octet rule.

The incomplete octet of the central atom

In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. Examples are LiCl, BeH₂ and BCl₃.