UNIT 4 ( PAGE 2)
HANNAY AND SMITH EQUATION
Hanny and Smith proposed the following equation for calculating the percentage ionic character :
Percentage ionic = 16 (cA - cB) + 3.5 (cA - cB)2
character
Both the equations , give only approximate values.
DIPOLE MOMENTS
The molecules containing polar bonds are said to be polar and possesses an electric dipole. The magnitude of the charge displacement in a polar covalent bond is measured through a quantity called the dipole moment m. Dipole moment is the product of the magnitude of charges(d) and the distance separating them(d). (The symbol d suggests a small magnitude of charge , less than the charge on an electron).
m = d d
d is of the order of 10-10 esu and ‘d’ is of the order of
10-8 cm. Dipole moment is of the order of 10-18 esu . cm and the unit is known as Debye (D).
In SI units the dipole moment is expressed in Coulomb . meter ( C . m ) In CGS units it is Debye i.e.,
10-18 esu. cm
1 D = 10-18 esu. cm
The unit is still used and its conversion factor to SI units is :
1 D = 3.33564 x 10-30 C. m
Dipole moment has a direction and is indicated by drawing an arrow with its tail on the positive end and arrow head on the negative. Thus HF is shown as :
The molecular dipole of a molecule having one polar bond is same as that of individual bonds (e.g. HF), while in the case of molecules containing more than one polar bonds, the molecular dipole moment is the sum of the dipole moments of all the individual bonds. Whether a molecule as a whole possesses a dipole or not depends upon the relative orientations of the bond dipoles.
Applications of Dipole moment Measurements
1. Distinction between polar and non-polar molecules
Molecules having some dipole moments are polar, while those having no dipole moments are non-polar. For example HF, NH3, H2O etc. are polar as they have significant dipole moments. On the other hand CO2, BeF2, H2, N2 etc. are non-polar molecules as their dipole moments is zero.
2. Degree of polarity : The dipole moment gives an idea of degree of polarity. Higher the dipole moments, higher is the polarity.
3. Shapes of molecules
In the carbon monoxide molecule, CO, the oxygen atom has a more electronegativity than does the carbon atom and so there is a slight displacement of electrons toward the oxygen atom. The CO molecule is slightly polar . In the representation the cross-base arrow (®) points to the atom that attracts the electrons more strongly.( Not shown here is the fact that the carbon-to-oxygen bond is a multiple covalent bond).
Even though the difference in electron negativity between C and O produces a bond dipole moment in each carbon-to-oxygen bond in CO2, the CO2 molecule is non-polar What this means is the effects of two O atoms in attracting electrons from the C atom cancel each other out. We say that the resultant dipole moment is zero. And this happens because the O atoms lie along the same straight line through the central C atom.
Because of the difference in electronegativities of H and O , there is a bond dipole moment of 1.51 D in the O-H bond. If water were a linear molecule, the bond moments would be in opposite directions and there would be no result dipole moment. But the measured dipole moment of water is 1.84 D. The molecule is polar and so cannot be linear. The two bond moments combine to yield a resultant dipole moment of 1.84 D for a particular bond angle of 104.50.
Measurement of dipole moments can assist us in describing molecular shapes.
We would expect the bond dipole moment of C-Cl bond to be large. This is because the Cl atom is considerably more electronegative than the C atom. The experimentally measured dipole moment is large, 2.05 D. Since a molecule of CCl4 has four C- Cl bonds , we might expect an even larger resultant dipole moment for the entire molecule. In fact , when we measure the resultant dipole moment for CCl4 we find it to be zero. CCl4 is a nonpolar molecule. The only way to account for this fact is that if the Cl atoms are arranged around the C atoms in some symmetrical fashion in which the individual bond dipole moments cancel each other.
CCl4 : a non-polar molecule CHCl3 : a polar molecule
m = 0 m = 1.92 D
(a) Symmetrical distribution of the four carbon-to-chlorine bonds in CCl4 causes a cancellation of all bond dipole moments ; there is no resultant dipole moment in the molecule. CCl4 is nonpolar.
(b) The carbon-to-hydrogen bond has a dipole moment of essentially zero because the electronegativities of C and H are quite similar. The three carbon-to-chlorine bond dipole moments cause chlorine end of the molecule to develop a slight negative charge. The hydrogen end carries a slight positive charge. CHCl3 is polar.
In CCl4 there are four pairs of electrons, all bond pairs, distributed around the central C atom. The electron pair orientation and shape of the molecule are both tetrahedral. From the above Fig , we see that a symmetrical tetrahedral orientation of the C- Cl bonds in fact just we need to account for the lack of dipole moment
The molecule CHCl3, in which there is an imbalance in dipole moments, is indeed a polar molecule. The cases of CCl4 and CHCl3 are compared in the above Fig.
Both NH3 and NF3 molecules are pyramidal shape with a lone pair of electrons on nitrogen atom. Although fluorine is more electronegative than nitrogen, the resultant dipole moment of NH3 ( 4.90 x 10-30 C. m) is greater than that of NF3 (0.80 x 10-30 C. m). This is because , in the case of NH3 the orbital dipole moment due to lone pair is in the same direction as the resultant dipole moment of the N-H bonds, whereas in NF3 the orbital dipole is in the direction opposite to the resultant dipole of the three N-F bonds . The orbital dipole because of lone pair decreases the effect of the resultant N-F bond moments, which results in the low dipole moment of NF3 as represented below:
Ionic Bond as Extreme Case of Polar Covalent Bond
A covalent bond develops a polar character due to differences in the electronegativities of the atoms forming the bond. The more electronegative atom acquires a partial negative charge and less electronegative atom acquires a partial positive charge. When electronegativity difference of two atoms becomes very large, the bond pair almost lies in complete possession of more electronegative atom. In this situation , it is said that the transference of electron has taken place from an atom with lower electronegativity to the atom with higher electronegativity. The tranference of electron has resulted in the formation of ionic species with opposite charges which are held by electrostatic force of attraction called ionic bond. Thus, ionic bond is an extreme case of polar covalent bond.
Distinction between Ionic bond and Covalent bonds
The main points of distinction between ionic and covalent bonds have been given in the following TABLE.
Ionic bond | Covalent bond |
1. This bond is formed by transference of electron from one atom to another. 2. The bond usually exists between metal and non-metal atoms 3. The bond is non-directional. 4. It is always polar 5. It holds charged particles of opposite charges 6. No multiplicity is observed | 1. The bond is formed by mutual sharing of electrons between the atoms. 2. The bond usually exists between non-metal atoms. 3. The bond has directional character. 4. It may or may not be polar 5. It holds neutral particles. 6. The bond can be single, double or triple bond. |
Problems
07. The inter-nuclear separation of KCl molecule in the vapour state is 267 pm. Assuming complete transfer of electron from potassium to chlorine atom. Calculate the dipole moment of KCl molecule. What will be the direction of this dipole ?
08. The bond length in HCl molecule is 127.6 pm. The observed dipole moment is 1.03 D. What is the percentage ionic character in HCl molecule ?
09. Arrange the bonds in order of increasing order of ionic character in the molecules :
LiF , K2O , N2 , SO2 and ClF3
10. Arrange the following in the order of increasing ionic character :
C-H , F-H , Br-H , Na-I , K-F and Li-Cl
11. Sketch the bond moments and resultant dipole moment in the following molecules :
H2O , PCl3 , NH3 , NF3
12. Explain why BeH2 molecule has zero dipole moment although the Be-H bonds are polar.
13. Dipole moment of hydrogen halides decreases from H-F to H -I.
14. Which out of OCS and CS2 as higher dipole moment and why ?
15. Which out of NH3 and NF3 has higher dipole moment and why ?
16. Although both CO2 and H2O are triatomic molecules, the shape of H2O molecules, the shapes of water molecule is bent while that of CO2 is linear. Explain this on the basis of dipole moment.
GEOMETRY OF MOLECULES
A molecule consists of two or more atoms liked together through certain types of attractive force. In polyatomic molecule, the constituent atoms get arranged in a definite geometric fashion around the central atom. This definite relative arrangement of the bonded atoms in a molecule is known as the geometry or shape of the molecule. The shapes acquired by different molecules are of different types. The molecule may be linear, triangular , square-planar, pyramidal, triangular bipyramidal , octahedral etc., The shape of the molecule contributes to a considerable extent to the physical and chemical properties of the compound. For example water molecule is angular in shape. Had water molecule been linear, its properties would have been much different.
There are various theories and experimental technique to find the shape of the molecule of a particular substance. Among these , the Valence Shell Electron Pair Repulsion Theory (VSEPR theory) is a simple tool to have an idea of the shape of a particular molecule.
THE VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY
This theory provides a simple method to predict the shapes of covalent molecules. This approach was developed by Gillespie and Nyholm. The theory is primarily based upon the fact that in a polyatomic molecule, the direction of bonds around the central atom depends upon the total number of electron pairs(bonding as well as non-bonding) in its valence shell. The electron pairs place themselves as far apart as possible in space so as to have minimum repulsive interactions between them. The minimum repulsion corresponds to the state of minimum energy and maximum stability of the molecule. The main postulates of the theory are :
(i) The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around central atom.
(ii) Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
(iii) These pairs of electrons tend to occupy such positions in space that minimize repulsion and thus maximize distance between them.
(iv) The valence shell is taken as a sphere with the electron pairs localizing on the spherical surface at maximum distance from one another.
(v) A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair.
(vi) Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure.
(vii) The repulsive interactions between two lone pairs(lp) are different from those between two bond pairs (bp) or those between a lone pair and a bond pair. The repulsive interactions between various electron pairs decreases in the order as :
ℓ p - lp > ℓ p - bp > bp – bp
Geometry of molecules in which central atom has no lone pairs of electrons.
Number of electron pairs | Arrangement of electron pairs | Shape | Examples |
2 AB2 |  | Linear | BeF2, BeCl2 |
3 AB3 |  | Triogonal planar | BF3, BCl3,AlCl3 |
4 AB4 |  | Tetrahedral | CH4, SiH4 SiF4, NH4+ |
5 AB5 |  | Trigonal bipyramidal | PF5, PCl5 SbCl5 |
6 AB6 |  | Octahedral | SF6, TeF6 |
7 AB7 |  | Pentagonal bipramidal | IF7 |
Prediction of shapes of molecules by VSEPR
1. If there are two electron pairs around a central atom in the molecule, the only way to keep them farthest apart is to have them directed at an angle of 1800. Thus the molecule adopts a linear geometry. Example : BeCl2.
2. If there are three electron pairs around the central atom in a molecule, in order to have them farthest apart, the bond angle should be 1200. The geometry of the molecule is trigonal planar. Example : B F3
3. If there are four electron pairs around the central atom in a molecule, the only way to keep them farthest apart is to have them directed towards the corners of a regular tetrahedron. The bond angle in this case is 109.50. Examples : Molecules like CCl4, SiH4, NH4+, BF4- etc have four electron pairs in the valence shell of their respective central atom and are tetrahedral in shape.
4. When there are five electron pairs around the central atom , the molecule adopts triogonal bipyramidal geometry. In this case all the bond angles are not equal. Three electron pairs are in the same plane at angle of 1200 to another. These are called equatorial bonds. Of the remaining two , one lies above and other lies below the plane of equatorial bonds making an angle of 900 with the plane. These are called axial bonds. Thus, in this arrangements , three bonds are 1200 each and two are 900 each.
Example : PCl5 molecule. The central atom , phosphorus has five electron pairs. The arrangement which keeps the five pairs as far apart as possible is trigonal bipyramidal arrangement. Thus PCl5 molecule has a trigonal bipyramidal shape.
5. If there are six electron pairs around the central atom, the molecule has octahedral geometry. In this case all the bond angles are equivalent and 900 each.
Example : SF6 molecule. The central atom has six electron pairs. SF6 molecule is octahedral because the repulsion between the six pairs around the sulphur atom is minimised. 6. The molecule having seven electron pairs around the central atom adopts pentagonal bipyramidal geometry. Five bonds are in the same plane at an angle of 720 to one another, while the remaining two (one above and one below the plane) are 900 each.
Example : IF7 molecule , there are seven electron pairs around the central iodine atom. In order to have minimum repulsion between bonded electrons, the molecule adopts pentagonal bipyramidal geometry.
REGULAR AND IRREGULAR GEOMETRY
If the central atom is surrounded by only orbitals containing shared pairs of electrons, i.e., bond pairs and there are no orbitals containing lone pairs in the valence shell, the molecule has a regular geometry. There is no distortion in shape. If however, the central atom is surrounded also by one or more orbitals containing lone pairs of electrons in the valence shell, the geometry of the molecule is distorted to some extent. This is because the repulsive interactions between lone pair-lone pair , lone pair -bond pair and bond pair-bond pair are different. The bond pair between two atoms is under the influence of two nuclei and most of the electron cloud is oriented between them. On the other hand the lone pair of electron is under the influence of only one nucleus and therefore its electron cloud is spread out and tends to occupy more space. For example, in the case of ammonia, there are three bond pairs and one lone pair. Since the lone pairs occupies more space around the central atom, it can interact more and therefore it will repel the electron pairs in the neighbouring orbitals more strongly whereas the repulsive interactions between the bond pairs will be less. Thus, it may be concluded that the lone pair-lone pair repulsion, which in turn is greater than lone pair-bond-pair repulsion, which in turn is greater than bond pair-bond pair repulsion, i.e.,
Lone pair - lone pair repulsion > lone pair-bond pair repulsion > bond pair - bond pair repulsion
Thus in the presence of one or more orbitals with lone pairs has the effect of altering the bond angles to a smallest extent. The molecule now do not retain any regular geometry.
In sulphur dioxide (type AB2E) there are three electron pairs on the S atom, the overall arrangement is trigonal planar. However, because one of the three electron pairs is a lp, the SO2 molecule has a ‘’bent’’ shape and because of ℓ p - ℓ p and bp-bp repulsions the angle OSO gets reduced to 119.5° from the value of 120°.
Ammonia molecule
Ammonia has four electron pairs around the nitrogen atom, though one electron pair is non-bonding electron type. Its repulsion is more than the bonding electron pairs. The geometry is distorted from tetrahedral one and is called pyramidal. The bond angle is 106045’. PCl3, NF3 and H3O+ are of the same shape.
Water molecule
Water molecule has four electron pairs around the oxygen atom, though there are two lone pairs or non-bonding electron pairs. Since their repulsive force is higher than bonding electron pairs, the tetrahedral geometry is distorted and water is predicted to be angular with H-O-H bond angle 1040 28’. F2O, NH2- and SnCl2 are having the same shape.
Angular or bent shape of water molecule
The central sulphur atom has five electron pairs whose arrangement should be trigonal bipyramidal (TABLE) . However one of the electron pairs is a lone pair. SF4 can , therefore have structure (a) or (b) shown below :
In arrangement (a) the lone pair is in axial position which leads to three lp-bp repulsions at 90°. In structure (b) the lone pair is in an equatorial position and there are only two ℓ p – b p repulsions. Hence arrangement (b) is the favoured structure. The shape shown in structure (b) is variously described as : a distorted tetrahedron, a folded square or seesaw.
Chlorine trifluoride ClF3
The ClF3 molecule is isoelectronic with SF4. It is expected to have a basic trigonal bi-pyramidal (tbp) arrangement of five electron pairs. However there are two lps and three bps. Three arrangements are possible. Calculations based on lp-lp, lp-bp and bp-bp repulsions for three arrangements show that arrangement (a) is most stable and ClF3 , therefore has a T-shaped structure.
Shapes(geometry) of some molecules /ions with central ions having one or more lone pairs of electrons
Molecule Type | Number of bonding pairs | Number of lone pairs | Arrangement of electron pairs | Shape Geometry | Examples |
AB2E | 2 | 1 |  Trigonal planar | Bent | SO2, O3 |
AB3E | 3 | 1 |  | pyramidal | NH3 |
AB2E2 | 2 | 2 |  | Bent | H2O |
AB4E | 4 | 1 |  | See saw | SF4 |
AB3E2 | 3 | 2 |  | T-shaped | ClF3 |
AB5E | 5 | 1 |  | Square pyramidal | BrF5 |
AB4E2 | 4 | 2 |  | Square planar | XeF4 |
Problems
21. Although geometries of NH3 and H2O molecules are distorted tetrahedral , bond angle in water is less than that of ammonia. Discuss.
22. Apart from tetrahedral geometry , another possible geometry for CH4 , is square planar with four H atoms at the corners of a square and carbon atom at its centre . Explain why CH4 is not square planar.
MODERN CONCEPT OF COVALENT BOND
All the mechanical systems in the universe tend to lower their potential energy. Lower the energy of the system, the greater is its stability. In a similar manner , the formation of bonds between the atoms occurs if it is accompanied by a decrease of energy. The bond formation takes place , if the attractive forces are stronger than the repulsive forces, so that it is accompanied by a net decrease of energy.
Consider the formation of hydrogen molecule from two atoms of hydrogen. Isolated hydrogen atoms possess no attractive or repulsive forces between them.
Formation of a molecule by combination of two atoms
a. atoms at a large distance, no interaction.
b. Atoms coming closer, interaction begins
c. Bonded atoms.
However, when two hydrogen atoms approach each other, the interaction between them starts taking place, due to operation of the following new forces.
(i) Forces of attraction between the nucleus of one atom and electrons of the other, and
(ii) Forces of repulsion between :
(a) the nuclei of two atoms and
(b) electron of one atom and electron of the other atom.
Fig 1
The forces acting between two atoms when they come closer.
NA and NB represent the nuclei of the two atoms while eA and eB represent their electrons respectively.
When two atoms approach each other ( Fig above), the magnitude of attractive forces is greater than that of repulsive forces, the energy of system begins to decrease(Fig 2)
As the two atoms moves closer , the decrease in energy continues, till a certain minimum value is attained. At this minimum value, the bond formation takes place and the distance (r0) between the two nuclei (corresponding to formation of a stable covalent bond between the two hydrogen atom) called the bond length or bond distance.
If the two nuclei approach more nearer than r0, the repulsive forces become greater than the attractive force. Owing to this, the potential energy of the system increases, thereby the molecule become unstable
Fig 2 Variation of potential energy as the two hydrogen atoms approach each other.
Experiments have shown that the value of r0 is 0.74 Å or 74 pm. The energy corresponding to minimum in the curve is called bond energy (435 kJ) If this much of energy is supplied from out side , the bond will break and the molecule will dissociate into atoms.
VALANCE BOND THEORY
The classical idea of Lewis and others that electron pairing takes place during the formation of a covalent bond , leads to the development of an important theory known as Valence Bond Theory. This theory was suggested by Heitler and London
HEITLER-LONDON THEORY
According to Heitler and London
(i) Atoms having one or more unpaired electrons can only form covalent bonds.
(ii) The maximum number of covalence formed by an atom is equal to the number of its unpaired electrons.
(iii) A covalent bond is formed only when the electrons of combining atoms get paired up by the mutual neutralisation of their spins. Therefore, the unpaired electrons of the combining atoms must be of opposite spins.
(iv) The pair of electrons is localised between the two bonded atoms.
(v) After bond formation ,electrons are indistinguishable, i.e., it is not possible to ascertain which electron belongs to which atom.
(vi) Each atom in a molecule tends to acquire a closed shell structure.
For example, the formation of the H2 molecule by combination of two hydrogen atoms can be explained as follows.
When the two hydrogen atoms approach each other, then at a certain internuclear distance there may be two possibilities. If the two atoms possess electrons of opposite spin, they attract each other and fuse together to form H2 molecule (Fig )
Fig Formation of H2 molecule
In this case , potential energy of the system decreases and a minimum in P.E diagram (curve I Fig ) . However if the electron of the combining atoms are of parallel spin, the two atoms repel each other and potential energy of the system increases ( curve II Fig ). In this case no bond is formed.
Fig Potential energy diagram for H2 molecule
PAULING SLATER THEORY
Linus Pauling and J.C Slater improved Heitler London theory to account for the directional characteristics in covalent bond and gave a new form to Valence Bond Theory. The main postulates of the theory are as follows :
(i) A covalent bond is formed when the atomic orbitals of the two combining atoms , possessing unpaired electrons of opposite spin , mutually overlap together. During the formation of bond, spins of two electrons get normally neutralised. If atomic orbitals have parallel spins, repulsion takes place and no bond is formed.
(ii) The strength of the covalent bond thus formed depends upon the extent of overlapping of two atomic orbitals. Higher the overlapping, greater is the strength of covalent bond formed.
The extent of overlapping between two orbitals having wave functions yA and yB is given by the following integral S, known as overlap integral
S = òyA yB dt
Where dt is small volume element.
When S is positive, there is an increase in electron density between two nuclei, resulting in the formation of a covalent bond. If S is negative, electron density decreases between the two nuclei. The nuclei thus repel each other and no bond is formed. When S is zero, no overlapping takes place and there is neither attraction nor repulsion between the two nuclei.
(iii) The atomic orbitals of only unpaired electrons (bonding electrons) take part in overlapping. The atomic orbitals of other electrons (non bonding electrons) do not overlap.
(iv) If an atom possesses more than one atomic orbital containing unpaired electrons, multiple bonds may be formed.
(v) Combining atoms retain their identity in the molecule formed by their combination.
(vi) The bond formed by a given set of orbitals lies in the direction of their maximum overlapping . The covalent bonds are thus directional in nature.
(vii) When two atomic orbitals come closer and overlap, the maximum electron density lies in the region of overlap. This subsequently develops an attraction between the two nuclei and accounts for the bond formation.
(viii) An axial overlap between the two orbitals of same energy content always forms a stronger bond.
(ix) A spherically symmetric orbital does not show any directional preference but non spherical orbitals have a tendency to form a bond in the direction of maximum electron density i.e., along their axes. s-orbitals being spherically symmetric overlap to the same extent in all directions whereas p- , d- and f-orbitals have a tendency to overlap along their axes where electron density is maximum.
OVERLAPPING OF ATOMIC ORBITALS
When two atoms come close to each other , there is overlapping of atomic orbitals. This overlap may be positive, negative or zero depending up on the properties of overlapping of atomic orbitals. The various arrangements of s and p orbitals resulting in positive , negative or zero are depicted in Fig.
Positive overlap
Negative overlap
Zero overlap
Positive , negative and zero overlaps of s and p atomic orbitals
The criterion of overlap , as the main factor for the formation of covalent bonds applies uniformly to homonuclear/hetronuclear diatomic molecules and polyatomic molecules. In the case of polyatomic molecules like CH4 , NH3 and H2O the VB theory has to account for their characteristic shapes. The shapes of CH4 , NH3 and H2O molecules are tetrahedral, pyramidal and bent respectively.
Consider the CH4 (methane) molecule. The electronic configuration of carbon in the ground state is [He]2s22p2 which in the excited state becomes [He]2s12Px12Py12Pz1. The energy required for this excitation is compensated by the release of energy due to overlap between the orbitals of carbon and the hydrogen. The four atomic orbitals of carbon , each of which an unpaired electron can overlap with 1s orbitals of the four H atoms which are also singly occupied. This will result in the formation of four C-H bond. It will , however, be observed that while the three p orbitals of carbon are at 90° to one another, the HCH angle for these will also be 90°. That is three C-H bond bonds will be oriented at 90° to one another. The 2s orbital of carbon and 1s orbital of H are spherically symmetrical and they can overlap in any direction. Therefore the direction of the fourth
C-H bond cannot be ascertained. This description does not fit in with the tetrahedral HCH angles of 109.5°. Clearly, it follows that simple atomic orbital overlap does not account for the directional characteristics of bonds in CH4. Using similar procedure, and arguments, it can be seen that in the case of NH3 and H2O molecules, the HNH and HOH angles should be 90°. This is in disagreement with the bond angles of 107° and 104.5° in NH3 and H2O molecules respectively.
Types of Overlapping and Nature of Covalent Bonds
The covalent bond may be classified into two types depending up on the types of overlapping :
i) Sigma (s) bond
ii) Pi (p) bond
Sigma(s) bond
This type of covalent bond is formed by end to end (head on) overlap of bonding orbitals along internuclear axis. This is called head on overlap or axial overlap. This can be formed by any one of the following types of combinations of atomic orbitals.
· s-s overlapping : In this case , there is overlap of two half-filled s-orbitals along the internuclear axis as shown below :
· s-p Overlapping : This type of overlap occurs between half-filled s-orbitals of one atom and half-filled p-orbitals of another atom.
· p-p overlapping : This type of overlap takes place between half-filled p-orbitals of two approaching atoms.
Pi (p) bond
In the formation of p-bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbitals formed due to side wise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms.
The characteristics of s and p bonds have been compared in TABLE.
Comparison of sigma and pi bonds
Sigma bond | Pi bond |
1. It is formed by the end to end overlap of two s- , one s- and one p-orbitals and two p-orbitals along their internuclear axis. It also results from end to end overlap of hybrid orbitals. 2. It determines the direction and extent of internuclear distance. 3. It is a strong bond due to greater overlap of orbitals 4. The molecular sigma orbital consists of a single electron cloud symmetrically around the internuclear axis. 5. There can be free rotation of atoms around the s-bond. | 1. It is formed by side wise overlap of two p-orbitals. 2. It has no primary effect on the direction of bond but shortens the internuclear distance 3. It is a weak bond due to poor overlap of orbitals. It is formed when sigma bond already exists between atoms. 4. The molecular pi-orbital consists of two electron clouds , one above and one below the plane of the bonded atoms. 5. As the electron cloud overlap is above and below the plane of atoms, free rotation is not possible around pi-bond. |
Strength of s and p-bonds : The strength of a bond depends upon the extent of overlapping of half-filled atomic orbitals. The extent of overlapping between two atoms is always greater when there is end to end overlapping of orbitals than when there is lateral ovelapping of orbitals. Hence a s-bond is stronger than a p-bond.
Characteristics of Covalent Compounds
The covalent compounds have the following characteristic properties.
1. State of existence : The covalent compounds are formed by the mutual sharing of electrons. There is no transfer of charge(electrons) from one atom to another. Hence , there is no ion formation. The compound exist as molecules and not as ions.
2. Conduction of electricity : Since covalent compounds do not contain ions, they do not conduct electricity in the fused or dissolved state.
3. Melting and boiling points : Most of the covalent compounds are gases, or liquids with low boiling points. Some are solids also which have low melting points. The low melting and boiling points are due to the fact that, the intermolecular forces in covalent compounds are generally weak and are easily overcome even at lower temperatures.
4. Solubility : Covalent compounds are generally insoluble or less soluble in water or other polar solvents. However, they are easily soluble in non-polar solvents.
5. Directional character of bond : The covalent bond is rigid and directional and hence there is a possibility of position and stereo isomerism in such compounds.
The various postulates of orbital concept regarding the formation of covalent bonds are :
(i) Covalent bonds are formed by the overlapping of half-filled atomic orbitals present in the valence shell of the atoms participating in bonding.
(ii) The orbitals undergoing overlapping must have electrons with opposite spins.
(iii) Overlapping of atomic orbitals results in the decrease of energy and formation of covalent bond.
(iv) The strength of the covalent bond depends upon the extent of overlapping. The greater the overlapping, more is the energy released and consequently, stronger will be the covalent bond.
The above treatment of the formation of covalent bond involving the overlap of the half-filled atomic orbitals is based upon the wave mechanical model is called Valence
According to orbital concept, the covalence of the element is defined as the number of half-filled orbitals present in the valence shell of its atoms.
Representation of orbitals and Overlapping of orbitals
Problems
23. Discuss the orbital model of the following molecules:
(i) Fluorine (ii) Hydrogen fluoride (iii) Nitrogen
(iv) Oxygen (v) Water
(i) Fluorine molecule F2 : Fluorine molecule is formed by overlap of two half-filled p-orbitals of two fluorine atoms as shown below :
(2) Hydrogen fluoride molecule, HF : Hydrogen fluoride molecule is formed by the overlap of 1s-orbitals of hydrogen atom , with one of the 2p-orbitals of fluorine atom which is half-filled as shown below :
3. Nitrogen molecule N2 : Nitrogen molecule is formed when three half-filled p-orbitals of the nitrogen atom overlap with three-half filled p-orbitals of nitrogen atom.
4. Oxygen molecule O2 : Oxygen molecule is formed when two half-filled 2p-orbitals of each oxygen atom overlap with the two 2p-orbitals, of the other oxygen atom to form a double bond.
4. Water molecule, H2O : Water molecule is formed when two half-filled 2p-orbitals of oxygen atom overlap with two half-filled 1s-orbitals of two hydrogen atoms.
Non-formation of He2 molecule
Helium atom(1s2) has 2 electrons in its 1s-orbitals. The new attractive and repulsive forces which come into existence when two helium atoms approach each other is shown in Fig (a) and (b).
In this case it has been found that, the new repulsive forces are dominant over, the attractive forces during the approach of two helium atoms. As a result of larger magnitude of repulsive forces, the energy of the system increases which leads to instability. Since energy of the separate helium atoms is smaller than that of the system when they are close to each other, they prefer to stay separate and do not form molecule He2.
Explanation on the basis of orbital overlap
The 1s orbitals of each helium atom has 2 electrons. These orbitals cannot overlap because, it would go against Pauli’s exclusion principle. Here helium atoms do not form bonds to constitute He2 molecule.
CO-ORDINATE BOND
Lewis showed that a different type of covalent bonds can be formed when both the electrons for sharing between the atoms are contributed by one atom only. This type of bond is known as co-ordinate bond or dative bond. The bond formed is indistinguishable from a normal covalent bond because all electrons are alike irrespective of their source.
A co-ordinate bond is established between two atoms, one of which has a complete octet and at the same time possesses at least one pair of unshared electrons while the other is short of two electrons. The former atom contributes one such pair of electrons for mutual sharing between two atoms as a result of which the second atom also completes its octet and acquires a stable noble gas configuration. The atom which contributes the pair of electrons is known as the donor and the atom which accepts these electrons is called the acceptor. The free pair of electrons is known as the lone pair.
In the formation of hydronium ion, for example, the oxygen atom in water molecule is the donor whereas the hydrogen ion is the acceptor. The co-ordinate covalent bond is represented by an arrow pointing from the donor to the acceptor as shown below in the formation of hydronium ion.
It is important to note that the co-ordinate covalent bond once formed is indistinguishable from a covalent bond.
According to the orbital overlap concept of covalency , a cordinate covalent bond results when an orbital containing a lone pair of electrons in one atom overlaps with an empty orbital present in the other atom. This may be illustrated by taking into consideration the formation of hydronium ion.
Formation of hydronium ion H3O+ : The central atom , oxygen , in water molecule has four orbitals. Two of these orbitals contain bond pairs of electrons, while the other two (shades ones) contain lone pairs of electrons, as shown in Fig.
Formation of hydronium ion.
H2O molecule H3O+ ion
A hydrogen atom loses its solitary electron(1s electron) to form a hydrogen ion. The H+ ion, , therefore has an empty s-orbital. The empty s-orbital of hydrogen ion overlaps with one of the p-orbitals of oxygen atom containing a lone pair of electrons forming hydronium ion, H3O+, as illustrated in the above figure.
Characteristics of Co-ordinate Covalent compounds
The properties of co-ordinate covalent compounds are very much similar to those of covalent compounds in many respects.
1. The electrons in co-ordinate covalent compounds are held firmly by the nuclei and therefore, they do not form ions in water.
2. The co-ordinate compounds are sparingly soluble in water. A number of them , however, are largely soluble in organic solvents.
3. The co-ordinate covalent bond is also rigid and directional and therefore, the structure of such compounds holds out the possibilities for space isomerism just as in covalent compounds.
4. A co-ordinate covalent bond results in semi-polarity in the molecule because the shared electron pair is contributed by one of the atoms only. The co-ordinate covalent compounds, therefore lie in between electrovalent and covalent compounds in their volatility.