Notes of Chemical Bonding

Quality Education At Every level

 Chapter #04 

Chemical Bonding 

 Prepared by:

Lecturer. S.Fayyaz Hussain

City Of Knowledge

              (Science Campus)

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CHEMICAL BOND

The attractive forces due to energy barrier that hold atoms together in a compound is known as chemical bond. There are various kinds chemical bonds like ionic covalent but all involve the stable configuration of atom or ion.

Ionic bond:

This type of chemical bond is proposed by kossel in 1916.

“A chemical bond which is formed as a result of complete transfer of electron from one atom to another so that both atom acquire inert gas configuration is called ionic or electro covalent bond”

The atom, which can loose electrons, is called electropositive and atom which can gain electron is called electronegative. These appositively charged ion are held together electrostatic force of attraction.

Condition For Ionic Bonding

1-         The electro negativity difference between the combing elements should be greater than 1.7.

2-         The energy evolved in the third step is greater than the energy absorb in the first step.

Ionic or electrovalent bond is generally formed between elements of group 1A, 11A (metal) and group V1A ,V11A(non metal).this is due to low ionization potential of metal and high E.A of non metals for e.g.NaCl,CaBr2,MgO,KI.

Explanation:

For the formation of ionic bond between metal and nonmetal, let us consider the energy change involve in the formation of sodium chloride. The hole process may occur in three steps.

1-Ionisation of Atom.

Gaseous sodium atoms lose an electron and formed cations by absorbing energy equal to ionization potential.

                     

 Na_(g)      +   I.P                                       Na⁺_(g)     +   e^-     

2-Absorption of Electron.

Chlorine atom gains an electron and converted into anions in this process equal to electron affinity is released.

                       

                        Cl_(g)     +   e^-                                              Cl^-_(g)   +   E.A

3-Formation of Lattice.

The gaseous cations and anions due to the electrostatic attraction combine together to give stable ionic crystal.

                         Na⁺_(g)  +  Cl^-_(g)                                                             NaCl_(s)

Because the oppositively charged ions are brought to their position in the crystal lattice from infinite distance, enough energy is released to make the over all process energetically favorable, this is known as lattice energy.

Energy of new system =495-348-788=-641 kj/mole.

The higher the value lattice energy of the resulting ionic compound the greater the ease of its formation.

Lattice energy:

“The energy released when gaseous cation and anion are brought together from infinite distance to form 1gm mole of solid crystal is called lattice. It is denoted by “U” and measured kj/mole.”

PROPERTIES OF IONIC COMPOUNDS :

1.         Hard solid: Ionic compounds consist of large no. of oppositively charge ions, which are arrange in definite pattern at the strong electrostatic forces between ions act in all direction through the crystal that is why ionic compounds are hard.

2.         High melting point due to strong ionic forces in the crystal, high heat is required to break these forces hence they possess high melting point.

3.         Solubility: Most of the ionic compound is soluble in water and some other polar solvents because of the strong electrostatics attractions between the ions and the polar molecules of solvent causing a release of energy known as “solvation energy” which overcome the high lattice of ionic compounds in the solution that’s why those are soluble in water.

4.         Electrolytic nature: These are invariable electrolytes and conduct electricity in molten as well as aqueous state; this is due to the movement of free ions under the influence of electric current.

5.         An unusual behavior: Some ionic compounds like sulphate, phosphate, and fluoride of Ca, Sr and Ba are not soluble even in polar solvent, this is due to their very high value of lattice energy which is not over come by solvation energy and ions thus do not separate in free state.

COVALENT BOND:

“It is a chemical bond which is formed by mutual sharing of electrons between two atoms.”

The concept of covalent bond was introduced by G.N Lewis in 1916 and was later explained by pauling, in term of wave function.

Covalent bond is generally represented by short line (-) between the bonded atoms or (:) electron pair. Covalent bond can be classified into single double or triple bond on the basis of number of bonded electron pairs. If there is only one electron pair between the two atoms, this is called single bond(-) similarly if 2 or 3 bonded electron pair are present by (=) or (=) these are called double or triple bond respectively.

Example:

                   Cl – Cl,        H-H          H-Cl

                   O=O,            N =N 

PROPERTIES OF COVALENT COMPOUNDS:

1-         They exist in separate covalent molecule because particles are electrically neutral and have less attractive forces.

2-         These are volatile in nature and most of them are gases or liquid

3-         They posses low boiling point and melting point.

4-         They are usually insoluble in water and soluble in organic solvent.

IONIC CHARACTER OF COVALENT BOND::

When a covalent bond exists between two non identical atom the bonded electron pair will be attracted more towards high electronegative atom thus one of its end is relatively negative and other end is relatively positive; in this way a negative and positive pole appears on the molecules this is called ionic character of covalent bond or polarization of covalent bond.

In HCl, the bonded electron pair id distorted due to the high E.N of Cl atom and probability of finding the electrons near Cl is greater, hence chloride is slightly negative and hydrogen is slightly positive.

The covalent bond in HCl is called partially ionic bond or polarized ionic bond or polarized covalent bond.

The ionic character of a covalent molecule depends upon the difference in E.N. of the bonded atom.

HF is more polarized and ionic due to big difference in EN of H and F atoms.

POLAR COVALENT BOND:

“It is a covalent bond which exists, between non identical atoms and is based on the distortion in shared pair of electron between bonded atoms”.

In a covalent bond between two unlike atoms, the shared electrons pair will be attracted more towards high electronegative atom. This permanent displacement of electron paired towards one atom in a covalent bond. It will develop a fraction of negative charged (-δ) on one atom and a fraction of positive charge (+δ) on another atom.

The molecule AB has some ionic character and is called polar molecule and such type of bond is called as polar polarized covalent bond.

The extent of ionic character of polar covalent bond depends on the difference of bonded atoms for e.g. HF is more polar than HCl because the electron cloud of HF molecule is shifted towards F atom to a greater extent.

A polar covalent bond is more stable and stronger due to the polarization of ionic character on the molecule, which makes the bond distance, and short pulls the atoms closer. Thus high bond energy is needed to break the bond.

The presence of partial ionic character affects the physical properties of molecules thus boiling point, melting point, solubility and density increases.

NON-POLAR BONDS:

“A covalent bond which exists between two identical atoms and the electron density of the shared atoms is equally distributed to around the two nuclei is called non polar or non polarized or true covalent bond”

When a covalent bond exists between two similar atoms A and A, the electron will equally be shared between the two atoms hence bond will have no ionic character.

e.g. H₂  ,  Cl₂   ,   O₂    &    N₂

A non-polar bond is relatively week and unstable due to the long bond distance and less value of bond energy. Non-polar compound are generally insoluble in water and have low B.P.

The extent of ionic character of a polar covalent bond depends on the difference of EN of bonded atoms.

Reason:             Why polar bond is more stronger than non polar?

A polar covalent bond is much stronger and stable due to the ionic character or polarization of molecule. The polarity decreases the bond distance and pulls the atoms closer. Thus high bond energy is required for bond breaking. Non-polar bond on the other hand is comparatively week because of non-polarization of molecules and greater bond distance.

HCl is a polar molecule and its bond dissociation energy is 431 Kj/mole therefore it is more stronger than Cl₂.

Reason:          Bond energy of  “HF” is very high.

Partial ionic character of HF molecule shortens its bond length and due to the high EN value of fluorine, the electronic cloud is greatly shifted towards fluorine. Thus the bond distance of HF would be 0.92 ºA less than the expected bond distance (1.01 A) and hence very high energy is required to break it.

BOND ENERGY :

“ The energy required to break a bond between two atoms in a diatomic molecule is known as bond energy, it might be taken as energy released in the formation of a bond from free atoms and measured in Kj/mole”

Bond energy may be exothermic and endothermic depending upon bond formation or breaking respectively.

FACTORS INFLUENCING BOND ENERGY:

1. POLARITY OF MOLECULES:

 The presence of ionic character on the molecule shortens the bond length and atoms are more strongly bonded together, therefore energy needed to break the bond would be high.

2. BOND LENGTH:

              The distance between two nuclei in a diatomic molecule is called bond length. Shorter the bond length more firmly is the atoms held and stronger will be the bond. That is why bond energy will be high.

3. BOND ODER:     

            Bond energy increases with the increased order of bond.

DIPOLE MOMENT :

“The quantitative measurement of concentration of a molecule is known as dipole moment or it is the tendency of a molecule to orient in an electric field”

The dipole moment measures the concentration of positive and negative charges in different part of the molecule and is equal to the product of ionic charge and distance between the center of positive and negative charges.

Here q = charge on molecule

        d = distance between the center of positive and negative charges.

REPRESENTATION:

Dipole moment is represented by (→) along with (-) and the direction of arrow is towards high electronegative atom in the molecule.

MEASUREMENTS:

Dipole moment is expressed in Debye after the name of introducer but in S.I system the unit of dipole moment is C.m(coulomb meter).

The relation between debye and C.m is given below

                        1 Debye = 10^-18 esu   X   cm = 3.335 x 10^-30 sm

 SIGNIFICANCE:

Dipole moment is a measurement of degree of polarity of the molecule greater the value of dipole moment, more polar will be the molecule. does by knowing the value of dipole moment, one can predict.

i)          The strength and stability of bond .

ii)         Its melting and boiling point.

Iii)       Its solubility in water.

iv)        The geometry of molecule i.e it is linear or angular.

1-DIPOLE MOMENT OF DIATOMIC MOLECULE:

The dipole moment of homo diatomic molecule is zero because of unavailability of ionic character in the molecule.

  Cl-Cl                         H-H

   (U=0)                           (U=0)

The hetreodiatomic molecules posses dipole moment and its value depends upon extent of polarization of molecule.

2-DIPOLE MOMENT OF POLYATOMIC MOLECULE.

For polyatomic molecules dipole moment does not only depend up on the polarity of its bond but it also depends up on the geometry molecule.

Molecule of CO₂ is linear and the bond moment of each C= 0 bond is opposite and equal which cancels out the effect of each other making CO₂ molecule and non polar and its dipole moment is zero.

Similarly molecule of CS₂ is linear and has zero dipole moment.

H₂O is angular molecule the vector some of bond moment of two H---O bonds is equal but not in opposite direction and it is polar molecule having a specific dipole moment.

Molecule of CCl₄ has symmetrical structure .the four C-Cl bond are polar but vector sum of dipole moment of all these are cancelled to each other gives a zero dipole moment due to the symmetrical structure of molecule.

CO-ORDINATE COVALENT BOND:

This type of chemical bond was suggested by sidwick and defines as                                                                                                                                 “A covalent bond in which the shared pair of electron is donated by one atom is called coordinate covalent bond”

   In this covalent bond both the electrons are supplied entirely by one atom .it is a less equitable mode of partnership in which the contribution is one sided.

REASON OF COORDINATES BOND FORMATION :

    Element of group VA, VIA and VIIA remain have lone pair of electrons after utilizing their valences in the covalent bond. These lone pair electrons can form a new bond with anther atom having empty orbital, this is as coordinate bond.

REPRESENTATION:

The dative bond is represented by (→) and its direction is towards electron accepting group.

It may also be represented by short line (-) but with the inclusion of positive charge (+) for electron donor and negative charge (-) for electron accepter.

CHARACTERISTATICS OF COORDINATE COMPOUND

Coordinate compounds exhibit characterstatics similar to covalent compounds.

1.         They do not ionize in water and are poor conductors of     electricity.

2.         They are very sparingly soluble in water but dissolve in organic solvents.

3          Since a coordinate linkage is semi-polar, melting and boiling point are higher than   those of purely covalent compounds but lower than ionic compounds.

Example of coordinate compounds

 1. Ammoiumion(NH4⁺):

In ammonia molecule, the central atom linked to three H+ atoms and yet N has an unshared pair electrons. the H+ ion furnished by an acid has no electron to contribute and chain accept a pair of electrons loaned by N atom. Thus NH3 donates its unshared electrons to H+ forming ammonium ion.

2. Hydronium ion ( H₃O⁺):

The oxygen atom in water molecule is attached to two atoms by two covalent bonds. There are still two unshaired pairs of electrons with the O atom. O atom donates one of these pairs of electron to H+ ion and the hydronium ion is thus formed.

3. Nitromethane( CH₃NO₂):

The Lewis structure of nitromethane is shown below.here the N atom has five valence electons, three of which are used in forming a covalent bond with C atom and two covalent bonds will O atom. The N atom is still left with two unshared electrons which are denoted to another O atom.

SIGMA BOND:

   “A covalent bond which is formed by head to head overlapping of atomic orbital is called as sigma bond”

It is generally established between “S—S”,”S—Px”and Px—Px atomic orbital. sigma bond formation is based on parallel or linear overlapping of atomic orbital, therefore bond strength should be maximum and it needs high energy to break a sigma bond. However the strength of b/w S—S, S—Px and Px—Px are not exactly same.

The S—S overlap is not so effective due to its spherical charge distribution. P-orbital has directional charge distribution and longer lobes which cause more effective and deep overlapping causing short bond length. Thus S—S sigma bond is relatively weaker than s-p and p-p. The relative bond strength is given as.

Pi BOND:

“A covalent bond which is formed by lateral or side overlapping of half filled P- atomic orbital is called as Pi bond”

 The overlapping of atomic orbital takes place perpendicular to inter nuclear axis. In the formation of Pi bond only Py and Pz orbital take part through side way overlapping, the extent of overlapping is small and bond formed is weaker.

The sideway overlapping gives two molecular orbital, Pi bonding and Pi antibonding M.O. These Pi bonding Nd anti bonding M.O have two regions of electron density below and above the orbital plane.

VALENCE BOND THEORY (V.B.T.):

This theory was proposed by Heitler and London in 1927 and later explained by paulling.

This theory gives us a clear explanation about bond length, bond energy and the strength of covalent molecules. Main postulates of this are given below.

1-         A covalent bond is formed by the linear overlapping of half filled atomic orbital `  in which electron pair is in opposite spin.

2-         The pairing of electrons in the molecule should satisfy the Paul’s exclusion principle and the two electrons would have different values of spin quantum number.

3.         The electron pair may be localized by the two nuclei. The strength of bond, bond length and bond energy depend upon the extent of overlapping. Deeply overlapped orbital form relatively stronger bond.

MOLECULAR ORBITAL THEORY:

This theory was proposed by Hund and Huckle in 1930, it explains the bond order and magnetic property of covalent molecule. It consists of following postulates.

1.         The linear combination of atomic orbital gives two type of molecular orbital called as bonding and anti bonding molecular orbital.

2.         Bonding molecular orbital has low energy (high stability) and anti bonding molecular orbital has high energy (low stability) with atomic orbital.

3.         Electrons in molecular orbital are multicentered delocalized.

             Consider the molecule of H₂ molecule.

 HYBRIDIZATION:

                        The hypothetical process of mixing of different atomic orbital to produce the same number of equivalent orbital

 HAVING SAME SHAPE and energy is known as hybridization, the orbital so formed are called hybrid orbital” 

                        The concept of hybridization was introduced by pauling. The type of hybridization depends upon the number of mixing orbital i.e.SP³,SP²,SP,dSP²,d²SP³ etc.

1-        SP³ – HYBRIDIZATION:

It is a mixing of one s-orbital and three p-orbital to produce four sp3 hybrids orbital is known as sp³ or tetrahedral hybridization.

                        Each sp³ hybrid orbital possesses the character of s and p in the ratio of 1:3 these are directed at the corner of regular tetrahedron with an angle of 109⁰ to each other.

EXAMPLE : FORMATION OF METHANE (CH4):

                        The electronic configuration of carbon is given below.

                        C(z=6) =1s²,2s¹,2p_x¹,2p_y¹.

                        The 2s and three 2p orbital of carbon mixed together to get a set of four equivalent sp3 hybrid orbital which are located at the corner of regular tetrahedron and each sp3 orbital of carbon over laps with s orbital of hydrogen from four sigma bonds.

Other molecules which show sp3 hybridization are CCl_4, SiCl₄,SnCl₄ etc.

2-                 SP² – HYBRIDIZATION:

                        The mixing of one s and two p atomic orbital to produce three sp² hybrid orbital is referred to as sp² or trigonal hybridization.

                                      OR

                        These are the set of three hybrid orbital, which arise from the appropriate combination of one s, and two p orbital.    

                        These sp2 orbital are coplanar and directed towards the corners of equilateral with angle of 120º. Each sp² hybrid orbital has character of S and P in the ratio of 1:2 sp² hybrids orbital is identical in shape.

Example:

                        Formation of ethane (C₂H₄)

          Each carbon in ethene undergoes sp² hybridization which is coplanar overlap with each other. at an angle of 120º. Two-sp² orbital of each carbon overlap with two S orbital while the other sp2 orbital

        The  unhybrid  P_z orbital of the two-carbon atom overlap above and below the plane to form π bond.

                        Boron trifluoride:

                        It is also an example of sp2 hybridization. Boron utilizes its 2s,2p_x and 2p_y orbital for the formation of hybrid orbital. These three hybrids  orbital overlap with three p orbital from three different fluorine atoms to form three B—F bonds at an angle of 120º from each other.

Others examples of molecules showing sp2 hybridization are C₆H₆   CO₂      HCN etc

3-        SP—HYBRIDIZATION

                        The mixing of one S and one P atomic orbital is called SP or diagonal hybridization. The hybrids SP-orbital are collinear at an angle of 180º which provide a maximum separation. These SP orbital have character of S and P in the ratio of 1:1.

                        One S orbital can combine with one P orbital on the same atom to form two new and completely equivalent orbital called sp hybrid orbital.

Example:

                        Formation of ethyne (C₂H₂)

           Each carbon in ethyne undergoes SP hybridization, which are collinear at an angle of 180º. One sp1 orbital of each carbon overlap with one s orbital of hydrogen while the other sp orbital overlaps with each other.

           The unhybridized P_y and P_z orbital of the two carbon atoms overlap above and below the plane to form two π bonds. 

ELECTRON PAIR REPULSION THEORY:

                        This theory explains about the geometry of simple covalent molecules and the ions of non transitional elements. It is introduced by Sidgewick and Powell in 1940.it is based on the repulsion of electron pairs in the valence shell of central atom in a molecule.

                        Main postulates of this theory are given below.

i)          The central atom in a covalent molecule may have two types of electron pairs, the bond pairs, and lone pairs of electron. These are called as active set of electron pairs.

ii)         These bond and loan pair of electrons exert repulsive forces to  each other try to be as far apart as possible, hence they orient themselves in space in such a manner that forces of repulsion between them is minimized.

iii)        The force of repulsion between two bond pair and loan pair is    not the same. The order of repulsion is as follows.

In case of molecule with double or triple bond, the π electron pairs are not considered to be involved in the repulsion and hence called as “inactive set” of electron.

iv)        The shape of molecule depends upon total number of active set of electron pair around the central atom.