UNIT 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE
SYLLABUS
• Valence electrons
• Ionic bond
• Covalent bond
• Bond parameters
• Lewis structure
• Polar character of covalent bond
• Covalent character of ionic compound
• Valence bond theory
• Resonance
• Geometry of covalent molecules
• VSEPR theory
• Concept of hybridization involving s, p and d orbitals
• Shapes of some simple molecules
• Molecular orbital theory of homomolecular diatomic molecules (qualitative ides only)
• Hydrogen bond
An atom , by itself, is rarely capable of free existence. Most of the substances exist in the form of clusters or aggregates of atoms, called molecules. The attractive force which holds together the constituting atoms with in a molecule is called a chemical bond.
Examples : Hydrogen gas exists as molecules ; each molecule consists of two atoms of hydrogen bonded together. Water molecule consists of two atoms of hydrogen bonded to one atom of oxygen. Hydrogen chloride molecule contains one atom of hydrogen bonded to one atom of chlorine.
CAUSE OF CHEMICAL COMBINATION
The following two tendencies are known to be responsible for chemical bonding :
1. Tendency to acquire noble gas configuration
Atoms of noble gases are known to be most stable. They have very little tendency to form chemical bonds. The atoms of noble gases are not bonded even among themselves to give polyatomic molecules. The result is that molecules of noble gases are monoatomic. This means that noble gas atoms have stable electronic configurations. These are represented in TABLE.
Electronic configurations of Noble gases
Noble gas Electronic configuration Confn. of valency shell
He 1s2 s2
Ne 1s22s2p6 s2p6
Ar 1s22s2p63s2p6 s2p6
Kr 1s22s2p63s2p6d104s2 p6 s2p6
Xe 1s22s2p63s2p6d104s2d105s2p6 s2p6
Rn 1s22s2p63s2p6d104s2p6d10f145s2p6d106s2p6 s2p6
From the table we find that s and p-orbitals of the outermost shell (also called valence shell) of each gas have their full quota of 2 and 6 electrons, respectively. (In helium, the solitary s-orbital of the valence shell contains full quota of 2 electrons). Thus, we see that s2p6 configuration in the valence shell constitutes a structure of maximum stability. Due to stable configurations, the atoms of these elements neither have a tendency to lose nor gain electrons and therefore they are not bonded amongst themselves or with atoms of other elements.
Atoms of all other elements do not have the stable s2p6 valence shell configurations. Therefore they have a tendency to get bonded with one another or with other atoms to attain stable noble gas configurations. The tendency of atoms of various elements to acquire stable noble gas configurations is responsible for chemical bonding. The rule of attaining maximum of eight (s2p6) electrons in the valence shell of atoms is called the octet rule.
2. Tendency to acquire minimum energy
A bonded state of atoms is more stable than an unbonded state. Since a more stable state has lower potential energy, it is obvious that a bonded state has lower potential energy than an unbonded state. Hence, when two atoms approach each other, a chemical bond is formed between them only if potential energy of the system constituting the two atoms decreases. If there is no decrease in potential energy of the system, no bonding is possible.
KOSSEL LEWIS APPROACH TO CHEMICAL BONDING
W. Kossel and G.N Lewis visualised that noble gases do not take part in chemical reactions under ordinary sets of conditions and concluded that the electronic configurations of noble gas atom are stable. On the basis of this conclusion they proposed a theory of valence known as Electronic theory of valence in 1916.
According to Kossel and Lewis, during the formation of a chemical bond, atoms combine together by gaining , losing or sharing electrons in such a way that they acquire stable noble gas configuration.
OCTET RULE
The noble gases (except He) possess ns2np6 configuration for their valence shells and exist as monoatomic gases. They do not enter into chemical combination under normal conditions. These observations imply that a valence shell containing eight electrons(ns2np6) is particularly stable. Since all other atoms have a basic urge to attain noble gas configuration, they gain , lose or share electrons until their valence shells contain eight electrons. This led Lewis to form a rule known as Octet rule. The rule can be stated as follows. During the formation of a molecule, an atom of a particular element gains, loses or shares electrons until it acquires a stable configuration of eight electrons in its valence shell.
The octet rule is found to be very useful in explaining the normal valence of elements and in the study of the chemical combinations of atoms leading to the formation of molecules. However, it is to be noted that octet rule is not valid for H and Li atoms. These atoms tend to acquire only two electrons in their valence shells similar to that in helium.
Valence electrons and Lewis structure
During the formation of a molecule, two atoms are held together by an attractive force. The attractive forces responsible for holding atoms together are the electrical forces between the electrons and the nuclei. However , all electrons of an atom do not take part in the formation of chemical bonds between the combining atoms. The inner shell electrons are well protected and they do not take part in the formation of chemical bonds. The formation of a chemical bond between two atoms involves only the electrons present in the outermost shells. The electrons present in the outermost shell (valence shell) of an atom are called valence electrons.
Gilbert Newton Lewis introduced a simple notation to denote the valence electrons in an atom. These notations are called Lewis symbols or electron dot symbols. In these symbols, only valence shell electrons are shown as dots surrounding the symbol of the given element.
For example, the Lewis symbols of the elements of first and second periods can be written as follows.
Lewis Symbols and valence of elements
Element Lewis symbol No of dots in Lewis symbol Valence
Hydrogen
1
2 1 = 1
Helium
2
2 2 = 0
Lithium
1
1
Beryllium
2 2
Boron 3 3
Carbon 4 4
Nitrogen 5 8 5 = 3
Oxygen 6 8 6 = 2
Fluorine 7 8 7 = 1
Neon 8 8 8 = 0
Significance of Lewis Symbols
The number of dots in a Lewis symbol represents the number of valence electrons possessed by the atom. Since the valence electrons are related to the valence of atom, an idea of the valence of the element can be had by the number of dots present in the Lewis symbol. The valence of the elements of the first period is equal to two minus the number of dots ; whereas for other elements , it is either equal to the number of dots in Lewis symbol or is eight minus the number of dots. These are illustrated in the Table (above).
Kossel , in relation to chemical bonding drew attention to the following facts :
In the periodic table, the highly electronegative halogens and highly electropositive alkali metals are separated by noble gases.
The formation of negative ion from the halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of an electron by respective atoms.
The negative and positive and positive ions thus formed attain stable noble gas electronic configurations. The noble gases (with exception of helium which has a duplet of electrons) have particularly stable outer shell configuration of eight (octet) electrons , ns2p6.
The negative and positive ions are stabilized by electrostatic attraction.
IONIC BOND OR ELECTROVALENT BOND
When the electronegativity difference between two atoms is large, they can attain octets by the complete transfer of one or more electrons from one atom to the other and can be linked together by an attractive force known as ionic bond or electrovalent bond. During formation of this type of bond, one atom loses electrons while the other gains electrons. The atom which loses electrons is called an electropositive atom while the one which gains electron is known as electronegative atom. The metallic atoms are usually electropositive while non metallic atoms are usually electronegative. Thus an ionic bond is usually established between a metallic and non metallic atoms.
During the formation of an ionic bond, the electropositive atom loses one or more electrons and changes into a cation while the electronegative atom gains the electrons lost by the electropostive atom and changes into an anion. The oppositely charged ions thus formed are held together by electrostatic force of attraction existing between them. This electrostatic force of attraction is the ionic bond and may be defined as follows.
Ionic bond between two atoms is the electrostatic force of attraction which holds together the ions of the combining atoms formed by the complete transfer of one or more electrons from electropositive to electronegative atom.
Following examples illustrate the formation of ionic bond.
1. Formation of sodium chloride (NaCl)
The electronic configuration of sodium and chlorine are :
Na(Z=11) : 1s22s22p63s1 or 2, 8, 1
Cl(Z= 17) : 1s22s22p63s23p5 or 2, 8, 7
Obiously, sodium has one electron in its valence shell. It is an electropositive element and has a tendency to lose electrons. It can acquire the configuration of nearest noble gas Ne(2,8) by losing its valence electron. Chlorine has seven electrons in its valence shell. It is an electronegative element and it has a tendency to gain electrons. It can acquire the stable electronic configuration of nearest noble gas Ar(2,8,8) by gaining one electron. Therefore , during the combination of sodium and chlorine atoms, the sodium transfers its valence electron to chlorine atom. Sodium atom changes into Na+ ion while the chlorine atom changes to Cl . The two ions get linked together by strong electrostatic force of attraction. The formation of ionic bond between Na and Cl can be shown as follows.
2. Formation of calcium chloride (CaCl2)
The electronic configuration of Ca is 1s22s22p63s2 3p64s2 i.e., 2, 8, 8, 2. It has two valence electrons. Since it is an electropositive element, it can easily lose its valence electrons to acquire a stable configuration similar to that of its nearest inert gas Ar (2,8, 8). It transfers its two electrons , one to each chlorine atom which is short of one valence electron as compared to Ar. Thus , CaCl2 molecule is formed as shown below.
It is to be noted that calcium atom acquires two positive charges because it loses two electrons while each chlorine atom acquires only one negative charge because each chlorine gains only one electron.
The number of electrons that an atom gains or loses during the formation of an ionic bond is called electrovalence. The electrovalence is either equal to the number of electrons lost to form the positive ion or to the number of electrons gained to form the negative ion. For example, the electrovalence of Na is + 1 because it loses one electron to form Na+ ion. Similarly, the electrovalence of Cl is 1 because it gains one electron to form Cl ion.
The common monoatomic ions of some representative elements are given below.
Group 1 2 13 14 15 16 17
Li+ Be2+
Na+ Mg2+ Al3+ N3 O2 F
K+ Ca2+ P3 S2 Cl
Rb+ Sr2+ Br
Cs+ Ba2+ I
If the electrovalence of the combining atoms are known , the empirical formula of the ionic compound can be easily derived. For example, suppose we wish to know the ionic compound formed by the combination of Mg and F. Mg forms Mg2+ while F forms F. Thus, the electrovalence of Mg is + 2 while that of F is 1.Hence the compound formed will be MgF2.
Factors influencing the formation of Ionic bond
The tendency of atoms to form ionic bonds between them depends upon the following factors :
1. Ionisation Energy : It is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom of an element. The lesser the ionisation energy, the greater the ease of formation of a cation. Alkali metals and alkaline earth metals have low ionisation energies and therefore , they form metal cations very easily.
2. Electron affinity or electron gain enthalpy : It is defined as the amount of energy released when an electron is added to an isolated gaseous atom of an element. The formation of a non-metal anion occurs with the addition of one or more electrons to the non-metal atom. The higher the energy released during this process, the easier will be the formation of the anion. Thus high electron affinity of a non-metal favours the formation of an anion. Thus , low ionisation energy of a metal atom and high electron affinity of a non-metal atom facilitate the formation of an ionic bond between them.
3. Lattice energy or lattice enthalpy : It is defined as the amount of energy released when cations and anions are brought from infinity to their respective equilibrium sites in the crystal lattice to form one mole of the ionic compound. The higher the magnitude of the lattice energy, the greater is the tendency of formation of an ionic bond.
For lattice energy to be high, the force of electrostatic attraction between the constituent ions should be high. The force of attraction between two oppositely charged ions in air (vaccum) is given by Coulomb’s law as :
where q1 and q2 are the respective charges of the ions and ‘r’ is the distance between them. The distance ‘r’ is obtained by adding the radii of positive and negative ions.
Hence the value of lattice energy depends upon the following factors: :
a. Charge on the ions : The higher the charge on the ions, greater is the force of attraction and hence larger is the amount of energy released.
b. Size of the ions : As highly charged species is rare, the other factor , i.e., inter nuclear distance between the ions becomes more important. If the size of the ions is large, internuclear distance will be more and force of attraction will be less, while in the case of small ions, internuclear distance is less and so force of attraction is greater.
Note
The formation of a positive ion involves ionization , i.e., removal of electron(s) from the neutral atom and that of the negative ion involves the addition of electron(s) to the neutral atom.
M(g) M+(g) + e : ionization enthalpy
X(g) + e X (g) : electron gain enthalpy
M+(g) + X (g) MX(s) : Lattice enthalpy
The electron gain enthalpy , egH , is enthalpy change , when a gas phase atom in its ground state gains an electron. The electron gain process may be exothermic or endothermic. The ionization on the other hand is always endothermic. Electron affinity is the negative of energy change accompanying electron gain.
Ionic bonds will be formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy.
Most ionic compounds have cations derived from metallic elements and anions from non-metallic elements. The ammonium ion , NH4+ (made up of two non-metallic elements) is an exception. It forms cation of a number of ionic compounds.
Ionic compounds in the crystalline state consists of cations and anions held together by coulombic interaction energies. These compounds crystallize in different crystal structures determined by the size of the ions, their packing arrangements and other factors. The crystal structure of sodium chloride , NaCl (rock salt) , for example is shown below.
Rock salt structure
In ionic solids, the sum of the electron gain enthalpy and ionization enthalpy may be positive but still the crystal structure gets stabilized due to energy released in the formation of the crystal lattice. For example, the ionization enthalpy for Na+(g) formation from Na(g) is 495.8 kJmol1 ; while the electron gain enthalpy for the change Cl(g) + e Cl (g) is , 348.7 kJmol1 only. The sum of the two , 147.1 kJmol1 is more than compensated for by the enthalpy of lattice formation of NaCl(s) (788 kJmol1 ). Therefore , the energy released in the process is more than the energy absorbed. Thus, a quantitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving octet of electrons around the ionic species in gaseous state.
Lattice enthalpy
The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJmol1 . This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+(g) and one mole of Cl (g) to an infinite distance.
This process involves both attractive forces between ions of opposite charges and repulsive forces between ions of like charge. The solid crystal being three-dimensional, it is not possible to calculate to calculate lattice enthalpy directly from the interaction of forces of attraction and repulsion only.
Note
i) The greater the lattice enthalpy, more stable is the ionic compound.
ii) The lattice enthalpy is greater, for ions of higher charge and smaller radii.
iii) The lattice enthalpies affect the solubilities of ionic compounds.
Properties of Ionic compounds
1. Conduction of electricity : The ionic compounds consists of ions. They do not conduct electricity in the solid state. The reason is that the anions and cations on account of electrostatic forces, remain intact occupying fixed positions in the crystal lattice. The ions therefore are unable to move to any large extent when electric field is applied. Hence no current flows. As the temperature is raised, there is an increase in kinetic energy of the ions. Ultimately, when the substance goes into molten state, the kinetic energy of the ions increases to such a large extent that they overcome the attractive forces and becomes mobile and are thus free to move under the influence of applied voltage. The ionic substances are thus able to conduct electricity in the molten state.
2. Melting and boiling points : Since in ionic solids, the ions are held together tightly, a considerable amount of energy is needed to dislodge them from the crystal lattice. Hence they melt at high temperatures to undergo change of state. The strong electrostatic force existing in ionic solids also explain their low volatility and high boiling points.
3. Solubility of ionic compounds : Ionic compounds are freely soluble in water, but only slightly soluble in organic solvents. Since water has a high dielectric constant, the electrostatic forces of attraction between ions decreases and they get separated from one another to give a solution. The ions in solution acquire greater freedom of motion. Another favourable factor is that water acts as a dipole, as a result of which the positive end of dipole interacts with negative ion, while the negative end of the dipole interacts with positive ion. Most of the organic solvents have low dielectric constants and are not dipoles. Hence ionic solids are only slightly soluble in organic solvents.
4. Ionic reactions : Ionic compounds exhibit ionic reactions, as these compounds form free ions in solution. These ionic reactions take place instantaneously.
5. Non-directional character : The ionic compounds are made of oppositely charged ions. Around each ion is present uniformly distributed electric field. The electric field is non-directional. So it imparts non-directional property to the ionic compound.
Problems
1. Draw the Lewis symbol for the following :
(i) Sodium 11Na (v) Xenon 54Xe
(ii) Calcium 20Ca (vi) Arsenic 33As
(iii) Boron 5B (vii) Germanium 32Ge
(iv) Bromine 35Br
2. Deduce the empirical formulae and draw the Lewis structures for ionic compounds formed by the following pairs of elements :
(i) Na, O (v) Al, F
(ii) K, S (vi) Ca, O
(iii) Na, P (vii) Li, S
(iv) Mg, Br
THE COVALENT BOND
Lewis in 1916 suggested that when both the atoms taking part in a chemical combination are short of electrons than the nearest noble gas configuration, they can share their electrons in order to complete their octets. Each atom contributes the same number of electrons to form common pairs which are then shared by both atoms.
The bond formed between the two atoms by mutual sharing of electrons between them so as to complete their octets or duplets in case of elements having only one shell is called covalent bond or covalent linkage and the number of electrons contributed by each atom is known as covalency.
Example to illustrate the formation of covalent bond
Let us examine the formation of chlorine molecule. In this case two chlorine atoms combine to produce chlorine molecule. Each chlorine atom is short of one electron to attain stable configuration of argon. Each of them contributes one electron to form a common shared pair. By doing so both of them acquire argon structure.
The dots represent electrons. Such structures are referred to as Lewis structures.
The Lewis dot structures can be written for other molecules also, in which the combining atoms may be alike or different. The important conditions being that :
i) Each bond is formed as a result of sharing of electron pair between the atoms.
ii) Each combining atom contributes one electron to the shared pair.
iii) The contributing atoms attain the outer shell noble gas configurations as a result of sharing of electrons.
The formation of bonds in water and carbon tetrachloride molecule can be represented as :
The two atoms joined by one electron pair 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 , covalent bond between them is called a double bond. For example in carbon dioxide molecule, we have two double bonds between carbon and oxygen atoms and in ethylene molecule the two carbon atoms are joined by a double bond.
When two combining atoms share three electron pairs as in the case of two nitrogen atoms in the N2 molecule and two carbon atoms in the ethyne molecule a triple bond is formed.
Bond pairs and Lone pairs
The shared pairs of electrons are present between atoms are called bond pairs because they are responsible for the bonding between atoms. On the other hand, the valence electrons not involved in bonding (i.e., sharing ) are shown as such and are called non-bonding electrons or lone pairs or unshared pairs.
LEWIS STRUCTURES AND COUPER STRUCTURES
The structures drawn on the left in which valence electrons are represented by dots are Lewis structures, whereas structures drawn on the right side in which bonds have been represented by lines or dashes are known as Couper structures.
LEWIS REPRESENTATION OF SIMPLE MOLECULES (THE LEWIS STRUCTURES)
The Lewis dot structures provide a picture of bonding in molecules in terms of the shared pair of electrons and octet rule. While such a picture may not explain the bonding and behaviour of molecules completely, it does help in understanding the formation and properties of the molecules to a large extent. Writing of Lewis structures of molecule is therefore very useful. The Lewis structures can be written by adopting the following basic steps.
i) 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).
ii) For anions, each negative charge would mean addition of one electron and for cations, each positive charge result in subtraction of one electron from the total number of valence electrons. For example for CO32 ion we add two electrons , because the two negative charges indicate that there are two additional electrons than are provided by the neutral atoms. For NH4+ ion we subtract one electron because the + 1 charge indicates the loss of one electron from the group of neutral atoms.
iii) Knowing the chemical symbols of the combining atoms and having the knowledge of the skeletal structure(what atoms are bonded to what other atoms) , it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to total bonds.
iv) In general the least electronegative atom occupies the central position in the molecule. Hydrogen and fluorine usually occupy the terminal positions, for example in the NF3 and CO32 , nitrogen and carbon respectively are the central atoms.
v) After counting the number of shared pairs of electrons for single bonds, remaining electron pairs are utilized either for multiple bonding or they constitute , 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 the following TABLE.
Problems
3. Depict the Lewis structures for the formation of :
(i) CO (vi) NO2
(ii) H2O (vii) CH4
(iii) H2S (viii) C2H6
(iv) Cl2 (ix) C2H4
(v) NH3 ( x) C2H2
4. Draw the Lewis structures for each of the following molecules and ions.
(i) F2 (v) F2O
(ii) PH3 (vi) Na+
(iii) SiCl4 (vii) Br
(iv) C3H8 (viii) H2S
5. Elements having the following Lewis symbols
(i) Place the elements in the appropriate group of the periodic table.
(ii) Which of the following is likely to form ions ? What is the expected charge on the ions ?
(iii) Write the formulae and Lewis structures of the covalent bonds between :
(a) A and B and
(b) A and C
6. Identify the atoms in each of the following compounds which do not obey the Octet rule ?
(i) SO2 (ii) SF4 (iii) OF2 (iv) SF2
(v) SF6 (vi) BCl3 (vii) PCl5
FORMAL CHARGE ON AN ATOM IN LEWIS STRUCTURE
Lewis dot structures , in general do not represent the actual shapes of molecules. Although in the case of polyatomic ions the net charge is possessed by the ion as a whole and not by a particular atom, for some purposes a formal charge (F.C) is assigned to each atom. The formal charge of an atom in a polyatomic ion/molecule is defined as :
The formal charge is the difference between the number of valence electrons in an isolated atom (i.e., free atom) and the number of electrons assigned to that atom, in a Lewis structure. The counting being 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.
Illustration
The concept of formal charge can be illustrated by ozone molecule, O3. Experimental data shows that in O3 the central O atom bonded to two other O atoms in the arrangement shown below:
Using the relationship given above , we can calculate the formal charges on the O atoms as follows.
• The central atom marked 1
The central atom marked 1 has 6 valence electrons, one lone pair ( or two non-bonding electrons) and three bonds (or six bonding electrons). Therefore its formal charge is = 6 2 ½ (6) = + 1
• The end O atom marked 2
This atom has 6 valence electrons, two lone pair(or 4 non-bonding electrons) and two bonds (or 4 bonding electrons). Therefore its formal charge = 6 4 ½ (4) = + 0
• The end O atom marked 3
This atom has 6 valence electrons, three lone pairs ( or six non-bonding electrons) and one bond (or two bonding electrons). Thus its formal charge = 6 6 ½ (2) = 1
Therefore the formal charges on the oxygen atoms in Lewis structures of O3 given above are written as :
The 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 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 neighboring atoms.
LIMITATIONS OF THE OCTET RULE
The octet rule , though useful is , is not universal. It is quite useful , for understanding the structures of most organic compounds and it applies mainly to the second period elements of periodic table. There are three types of exceptions to the octet rule.
i) The incomplete octet of the central atom : In some compounds , the number of electrons surrounding the central atom is less than eight. This is specially the case with elements having less than four valence electrons. Examples are LiCl, BeH2 and BCl3.
Li, Be and B have 1, 2 and 3 valence electrons only. Other such compounds are AlCl3 and BF3.
ii) Odd Electron molecules : In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxideNO2, the octet rule is not satisfied for all atoms.
iii) The expanded octet : Elements in and beyond the third period of the periodic table have apart from 3s and 3p orbitals, 3d orbitals also available for bonding. In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet. Obiously the octet rule does not apply in such cases. Some examples of such compounds are PF5, SF6, H2SO4 and a number of co-ordination compounds.
Sulphur also forms many compounds in which the octet rule is obeyed. In sulphur chloride the S atom has an octet of electrons around it.
Other drawbacks of the Octet theory
i) Octet rule is based up on the chemical inertness of noble gases. However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form number of compounds like XeF2, KrF2, XeOF2 etc.
ii) This theory does not account for the shape of molecules.
iii) It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.
BONDING PARAMETERS
Covalent bonds are characterized by the following parameters :
(i) Bond length (ii) Bond angle
(iii) Bond enthalpy (iv) Bond order
Bond Length
The bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule. Bond lengths are measured by spectroscopic , X-ray diffraction and electron-diffraction techniques. Each atom of the bonded pair contributes to the bond length.
The bond length in a covalent molecule AB.
R = rA + rB ( R is the bond length and rA and rB are the covalent radii of atoms A and B respectively.
In the case of a covalent bond , the contribution from each atom is called the covalent radius of that atom.
The covalent radius is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation. The covalent radius is half of the distance between two similar atoms joined by a covalent bond in the same molecule.
Bond length is usually expressed in Angstrom units(Å) or picometer (pm) and it can be determined by X-ray diffraction and other spectroscopic techniques. Bond lengths depends upon the sizes of the atoms and nature of bonds.
(i) Bond length increases with increase in size of atoms.
(ii) Bond length decreases with multiplicity of bonds.
The van der Waals radius represents the overall size of the atom which includes its valence shell in a bonded situation. The van der Waaals radius is half the distance between two similar atoms in separate molecules in a solid. Covalent and van der Waals radii of chlorine are depicted in Fig.
Covalent and van der Waals radii in a chlorine molecule
The inner circles corresponds to the size of the chlorine atom .
(rvw and rc are van der Waals and covalent radii respectively)
Some typical average bond lengths and bond enthalpies of some bonds are shown below.
Bond Bond energy kJ mol1 Bond length
pm
HH 435 74
HC 414 110
HN 389 100
HO 464 97
HS 368 132
HF 569 101
HCl 431 136
HBr 368 151
HI 297 170
NN 163 145
N=N 418 123
NN 949 109
NO 230 136
N=O 590 115
CC 347 154
C=C 611 134
CC 837 120
CN 305 147
C=N 615 128
CN 891 116
CO 360 143
C=O 728 123
CCl 326 177
OO 142 145
O=O 498 121
FF 159 128
ClCl 243 199
BrBr 192 288
II 151 266
Bond angle
Covalent bonds are formed by the overlapping of atomic orbitals. Due to directional character of atomic orbitals, the covalent bonds are oriented in specified directions. It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. It gives some idea regarding the distribution of orbitals around the central atom in a molecule/complex ion and helps to to determine its shape.
Bond angle is expressed in degree/minutes/seconds. For example, the HCH bond angle in methane is 109028’. Similarly , FBF bond angle in BF3 is 1200 HNH bond angle in NH3 molecule is 106045’. The bond angles in CH4, NH3, H2O and BF3 molecules are shown below.
Bond angles in CH4, NH3, H2O and BF3
Bond Enthalpy
It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in gaseous state. The unit of bond energy is kJ mol1. For example, the the HH bond enthalpy in hydrogen molecule is 435.8 kJ mol1.
H2(g) 2 H(g) : Hө = 435.8 kJ mol1
Similarly the bond enthalpy for molecules containing multiple bonds , for example O2 and N2 will be as under :
O=O (g) O(g) + O(g) ; Hө = 498 kJ mol1
NN(g) N(g) + N(g) ; Hө = 946.0 kJ mol1
Larger the bond dissociation energy, stronger will be the bond in the molecule.
For a hetero-nuclear diatomic molecules like HCl, we have ,
HCl(g) H(G0 + Cl(g) ; Hө = 431.0 kJ mol1
In the case of polyatomic molecules , the measurement of bond strength is more complicated. For example in the case of H2O molecule, the enthalpy needed to break the two OH bonds is not the same.
H2O(g) H(g) + OH (g) ; Hө1 = 502 kJ mol1
OH(g) O(g) + H(g) ; Hө2 = 427 kJ mol1
The difference in the Hө value shows that the second OH bond undergoes some change because of changed chemical environment. This is the reason for some difference in energy of the same OH bond in different molecules like C2H5OH and water. Therefore in polyatomic molecules the term mean or average bond enthalpy is used. It is obtained by dividing total bond dissociation enthalpy by number of bonds broken as explained below in case of water molecule.
Factors on which Bond enthalpy depends
(i) Size of the participating atoms : Larger the size of the atoms involved in bond formation, lesser is the extent of overlapping and consequently, smaller is the value of bond energy.
(ii) Multiplicity of bonds : The magnitude of the bond energy increases with the multiplicity of bonds, even though the atoms involved in the bond formation is the same. It is because of the fact that the multiplicity of the bonds , the number of shared electrons between atoms increases. As a result, the attractive force between nuclei and electrons also increases and consequently, the magnitude of bond energy increases.
Bond Order
In the Lewis description of covalent bond, the Bond Order is given by the number of bonds between two atoms in a molecule. The bond order for example in H2 (with single shared electron pair) , in O2 (with two shared electron pairs) and in N2 (with three shared electron pairs) is 1, 2, 3 respectively. Similarly in a CO (three shared electron pairs between C and O ) the bond order is 3. For N2, bond order is 3 .
Isoelectronic molecules and ions have identical bond orders ; for example, F2 and O22 have bond order 1. N2 , CO and NO+ have bond order of 3.
A general correlation useful for understanding the stabilities of of molecule is that : with increase in bond order , bond enthalpy increases and bond length decreases.
RESONANCE
It is often observed that a single Lewis structure is inadequate for representation of a molecule in conformity with its experimentally determined parameters. For example ozone O3, molecule can be equally well represented by the structures I and II shown below :
In both structures we have O = O double bond. The normal O O and O = O bond length are 148 pm and 121 pm respectively. Experimentally determined oxygen oxygen bond lengths in O3 molecule are same (128 pm ). Thus oxygen-oxygen bonds in O3 molecule are intermediate between a double bond and a single bond. Obiously, this cannot be represented by either of the two Lewis structures shown above.
The concept of resonance was introduced to deal with the type of difficulty experienced in the depiction of accurate structures of molecules like O3. According to the concept of resonance , whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy , positions of nuclei, bonding and non-bonding pairs of electrons are taken as cannonical structures of the resonance hybrid describes the molecule accurately. Thus, for O3 , the two structures shown above constitute the cannonical structures and their hybrid represents the structure of O3 more accurately. Resonance is represented by a double headed arrow .
Resonance in the O3 molecule, I and II represent the two cannonical forms
Some other examples of resonance structures have been provided by the carbonate ion and the carbon dioxide molecule.
The structure of CO32
The single Lewis structure based on the presence of two single bonds and one double bond between carbon and oxygen atoms is inadequate to represent the molecule accurately as it represents unequal bonds. According to the experimental findings , all carbon to oxygen bonds in CO32 are equivalent. Therefore the carbonate ion is best described as a resonance hybrid of the cannonical forms I, II and III as shown below :
Resonance in CO32 , I, II and III represent the three canonical structures
Resonance in CO2 molecule
The experimentally determined carbon to oxygen bond length in CO2 is 115 pm. The lengths of a normal carbon to oxygen double bond (C=O) and carbon to oxygen triple bond (C O) are 121 pm and 110 pm respectively. The carbon-oxygen bond lengths in CO2 (115 pm) lie between the values for C=O and C O. Obiously , a single Lewis structure cannot depict this position and it becomes necessary to write more than one Lewis structures and to consider that the structure of CO2 is best described as a resonance hybrid of the cannonical forms, I, II and III.
Resonance in CO2 molecule, I, II and III represent the three cannonical forms.
In general , it may be said that :
i) Resonance stabilizes the molecule as the energy of resonance hybrid is less than the energy of any single cannonical structure ; and
ii) Resonance averages the bond characterisics as a whole.
Thus the energy of the O3 resonance hybrid is lower than either of the two cannonical forms I and II.
The individual structures of similar energy contribute equally to the resonance hybrid. Cannonical structures with some what higher energies may also contribute to the resultant hybrid, but the higher the energy of a particular structure, the lesser is its contribution. For example, in the example of benzene molecule (C6H6) the cannonical forms are :
Kekule Dewar
The two equivalent Kekule forms and three equivalent Dewar forms (The energy of the Kekule forms is lower than the energy of the three Dewar forms)
The contribution from the three equivalent Dewar high energy cannonical forms to the resonance hybrid representing benzene is lesser as compared to the contribution from the two equivalent lower energy Keule forms.
NOTE
i) The cannonical forms have no real existence.
ii) The molecule does not exist for a certain fraction of time in one cannonical form and other fractions of time in other cannonical forms.
iii) There is no such equilibrium between the cannonical forms.
iv) The molecules as such has a single structure which is the resonance hybrid of the cannonical forms and which cannot as such be depicted by a single Lewis structure.
Conditions for writing Resonance Structures
The following are essential conditions for writing resonating structures :
(i) The contributing structures should have the same atomic positions.
(ii) The contributing structures should have the same number of unpaired electrons.
(iii) The contributing structures should have nearly the same energy.
(iv) The structures should be so written that negative charge is present on the electronegative atom and positive charge is present on the electropositive atom.
(v) In contributing structures, the like charges should not reside on adjacent atoms.
Resonance Energy
It is the difference between the actual bond energy of the molecule and that of the most stable of the resonating structures (having least energy).
Problems
07. H3PO3 can be represented by structures 1 and 2 as shown below. Can these structures can be taken as the canonical forms of the resonance hybrid representing H3PO3 ? If not cite reasons for the same.
08. The skeletal structure of CH3COOH as shown below is correct, but some of the bonds are wrongly shown. Write the correct Lewis structure for acetic acid.
09. Out of the following resonating structures for CO2 molecule, what are the important for describing the bonding in the molecule and why ?
10. Write resonance structures for SO3 , NO2 and NO3
POLARITY OF BONDS
The existence of a hundred percent ionic or covalent bond represents an ideal situation. In reality no bond or a compound is either completely covalent or ionic. Even in case of covalent bond between two hydrogen atoms, there is some ionic character.
A covalent bond between two identical or similar atoms is said to be a non-polar covalent bond but if formed between two dissimilar atoms, the bond formed is said to be a polar covalent bond. The reason is that in the former case the shared electron pair is attracted equally by both atoms and lies exactly midway between them as in hydrogen molecule, H : H. The molecule formed is said to be a non-polar molecule. The examples are H2, F2, Cl2.
In the case of a covalent bond formed between two dissimilar atoms, one of the atoms generally has a greater tendency to attract the electrons towards itself (i.e., it is more electronegative). The electron pair is , therefore pulled closer to that atom, as in hydrogen fluoride molecule
in which the electron pair shared between hydrogen and fluorine remain closer to the fluorine atom. This unsymmetric distribution of electrons leads to charge separation, i.e., development of partial positive charge near fluorine end and partial positive charge near hydrogen end. This is represented in HF and HCl molecules as follows :
The molecule formed is said to be polar molecules. The polar covalent bond, therefore, has partial ionic character. The two opposite charges at the ends are called electrical poles and the molecules are called dipoles.
Distinguishing between a polar and a non-polar molecule
A polar molecule can be distinguished from a non-polar molecule by a simple experiment.
Behaviour of polar molecules in the electric field
of the a condenser
We place the substance under examination , say HCl between two charged parallel plates. If the substance is polar, the molecules will orient in such a manner that the positive(hydrogen) ends are directed towards negatively charged plate and the negative (chlorine) ends are directed towards the positively charged plate. As a result of this, the intensity of charge on each plate will diminish. If, however, the molecule is non-polar, there will be no such effect.
Covalent Character of Ionic Bond : Polarisation of Ions
Till recently, the ionic and covalent bonds were considered to be distinct from each other. But, now we know that it is not always so. It is a common experience that many times an ionic bond has some covalent character and a covalent bond has some ionic character. The formation of a bond intermediate between an ionic bond and covalent bond occurs through a phenomenon known as polarisation of ions. When oppositely charged ions approach each other , the attraction between the positive charge of cation and negative charge of anion and also the simultaneous repulsion between their nuclei results in the distrortion, deformation or polarisation of the anion. The electronic charge cloud of the anion no longer remains spherical but gets distorted, i.e., polarised, towards the cation, as shown below :
The electronic charge cloud of the cation also get distorted, i.e., polarised by the anion through a similar process but the polarisation of the cation is far less pronounced because of its small size. The polarisation of ions results in a high electron charge concentration between the two nuclei. This results in the formation of a covalent bond with a high degree of polarisation. This type of bond is called polar covalent bond. A polar covalent bond is more stable than a pure covalent or a pure ionic bond. The higher is the degree of ionic polarisation, the greater is the stability of the polar covalent bond. The extent of polarisation depends , evidently, on the polarising power of the cation and the polarisability of the anion. The rules regarding polarisation are known as Fajan’s rules. These are given below :
FAJAN’S RULES
1. The cations with smaller size have higher polarising power, i.e., they cause polarisation of electronic charge cloud of an anion to a greater extent. Such cations have positive charge concentrated over a small surface area , i.e., they have a high charge density and thus distort the electronic charge cloud of the anion efficiently. The cations with large size have low polarising power.
2. The anions with large size have high polarizability, i.e., their electron cloud can be deformed easily. The grip of the nucleus on the orbital electrons in large anion will be weak and hence such anions will get polarised by a cation relatively easily.
3. For effective polarisation, there should be high charge on the cation or anion or both. The electrostatic forces which cause polarisation would increase with increase in charge on the ions.
4. Cations with pseudo inert gas configuration, viz., ns2p6d10 or with inert pair configuration, viz., ns2p6d10 (n+1)s2 , have high polarising power, while cations with inert gas configuration ns2p6 have low polarising power. This is due to greater effective nuclear charge in the former cases and smaller effective nuclear charge in the latter cases.
To sum up
1. Cations with small size, high charge and possessing pseudo inert gas or inert pair configuration have high polarising power, while cations with inert gas configuration viz., ns2p6 , have low polarising power.
2. Anions with large size and high charge are easily polarizable whereas anions with small size and small charge are not easily polarizable.
Effects of Polarisation
1. Bromides and iodides have higher lattice energies(and hence higher stabilities) than expected from theoretical calculations. The extra stability is due to polarisation of the anions resulting in the formation of a polar covalent bond which is more stable than a pure ionic bond.
2. The solubility of ionic compounds in polar solvents decreases with increase in degree of polarisation, i.e., with increase in the degree of covalent bonding.
3. The hardness of ionic compounds decreases with increase in the degree of polarisation, i.e., with increase in covalent character.
Percent Ionic Character of a Polar Covalent Bond
The percent ionic character of a polar covalent bond depends upon two factors.:
(i) The electronegativity difference of the bonded atoms
(ii) Dipole moment of the compound.
1. Electronegativity Difference and Percent ionic character : The degree of ionic character of a bond is measured from the electronegativity difference between the two bonded atoms. Greater the difference in electronegativities , greater is the percentage ionic character in a bond. It has been observed that the bond has 50% ionic character and 50% covalent character, if the difference in electronegativities of the participating atoms is 1.7. On the other hand , the covalent character dominates if the difference of electronegativities is less than 1.7, while ionic character dominates if difference of electronegativities is greater than 1.7. A bond having ionic character more than 50% is usually called ionic bond while that having ionic character less than 50% is called covalent bond. This is illustrated in the following TABLE
TABLE
Compound Electronegativity of one element Electronegativity of second element Electronegativity difference Nature of bond
CsF Cs(0.7) F(4) 3.3 Ionic
HCl H(2.1) Cl(3) 0.9 Covalent
BrCl Br(2.8) Cl(3) 0.2 Covalent
Several empirical equations have been proposed to calculate the percent ionic character of covalent bond from the electronegativities of the bonding atoms. Two of the equations are described below :
Pauling Equation : Pauling proposed the following equation for determining the percent ionic character of a covalent bond :
Percent Ionic character = 1 e ¼ ( A B)
where A and B are the the electronegativities of A and B. He established the following empirical relationship between A B and percent ionic character :
A B
0.6
1.0
1.4
2.0
2.4
3.0
Percentage Ionic character 9 22 39 63 76 89