1. Chemical Bonding-I Covalent Bonds
Presented By:
Dr.Vatsala Soni

2. Covalent Bond
Formed by mutual sharing electron pairs between the combining atoms of the same or different elements.
The atoms involved constitute equal number of electron for sharing in order to acquire configuration of nearest noble gas.
The pair of electrons which constitute the bond between atoms is called a bonding pair of electrons.
If there is one bonding pair of electron between the two atoms, the bond formed is single covalent bond. Eg. H2, F2 etc.
And if two or three pairs of electron between two atoms, constitute a double (eg. O2) or triple (eg. N2) bond respectively.

3. Theory of covalent bond formation
Lewis Model (A classical treatment)
Valence Bond Theory (VBT)
(by Linus Pauling)
Molecular Orbital Theory (MOT)
(By Robert Mulliken)

4. Lewis Model
A covalent bond is formed by sharing of valence electrons between atoms of non-metals.
Reasons : 1. They have small/no difference in E.N.
2. They can achieve stable electronic configurations (octet /duplet structures)
3. The total energy of the system is lowered when electrons are shared between the two nuclei.

5. Electron Sharing in Covalent Bonds
Formation of Covalent Bonds
Electron Sharing in Covalent Bonds
Attraction between oppositely charged nuclei and shared electrons ( _____________ in nature) H
Shared electrons
electrostatic e-
The shared electron pair spends most of the time between the two nuclei.
Overlapping of atomic orbitals  covalent bond formation

6. Covalent Bonds in Elements
Formation of Covalent Bonds
Covalent Bonds in Elements
Hydrogen molecule
Dot and cross diagram

7. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)

8. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

9. Bonding in Molecules Covalent Bonding
The term covalent implies sharing of electrons between atoms.
Valence electrons and valence shell orbitals
Only valence electrons are used for bonding: ns, np, nd
“Core” electrons are held too tightly (too low in energy)
Filled nd orbitals are considered core electrons
Valence state electron configurations and Promotion Energies
- The promotion energy is the energy required to promote electrons from the ground state to a “valence state”, which is one type of excited state configuration that is used for bonding. C*
E.g. C
2s2
2p2
2s1
2p3
ground state
valence state

10. Localized Bonding Models
Localized implies that electrons are confined to a particular bond or atom.
The Lewis approach to bonding
Pairs of electrons are localized in bonds or as non-bonding “lone pairs” on atoms. Each bond is formed by a pair of electrons shared by two atoms.
G.N. Lewis
Octet rule: most main group atoms will tend to end up with an ns2 np6 electron configuration.
This is mostly true for the molecules of organic chemistry not necessarily for inorganic compounds. ns np

11. Rules for drawing Lewis diagrams
a. Pick the central atom.
- Atoms that are present only once in the formula, especially heavy elements and metals, tend to be at the center of the structure.
- Oxygen is often terminal and hydrogen almost always is.
- Often the formula is written with the central atom first.
(Sometimes there may be more than one central atom.)
Write out the valence shell electron configurations for the neutral central atom and the "terminal" atoms in their ground states.
c. If there is a negative charge distribute it among the terminal atoms in the first instance. Bear in mind that all the terminal atoms must make at least one covalent bond with the central atom, so do not create any noble gas configurations on them. Positive charge is best initially assigned by removing electrons from the central atom.
The total number of unpaired electrons on the terminal atoms will have to match the number of unpaired electrons on the central atom to account for the bonds and leave no unpaired electrons. If this is not the case, once the first three steps have been carried out, there are two strategies available:
Move electrons between the central atom and the terminal atoms as necessary. Make sure you keep track of the formal charges because you must be specific about their location. Enclosing a Lewis structure in brackets with the charge outside is not acceptable.
f. If and only if the central atom comes from the second period or below (Na onwards, n=3 and up), electrons can be placed into the nd subshell. (Whether the d orbitals play a significant role in bonding in main group compounds is debatable, but they do help to predict correct structure without invoking canonical structures with unreasonable charge separations.)

12. Typical Lewis structural types:
Molecules that conform to the “Octet Rule”: saturated molecules
NH3
CH4 2s 2p 2s 2p N C
ground state C*
valence state 1s 1s 1s
3 H
4 H 1s 1s 1s 1s
These are typical of the molecules of organic chemistry.

13. Molecules that conform to the “Octet Rule”: unsaturated molecules.
ClNO
NO3- 2s 2p 2s 2p N N N+ O Cl O 2s 2p 3s 3p 2s 2p O- 2s 2p O- 2s 2p

14. Resonance
Resonance implies that there is more than one possible way to distribute the valence electrons in a Lewis structure. For an adequate description, each “canonical” structure must be drawn.
If different equivalent resonance structures are possible, the molecule tends to be more stable than one would otherwise expect. This is a quantum mechanical effect that we will talk about later.
I expect you to be able to:
Draw Lewis structures (including resonance structures when necessary), determine bond orders, determine and place formal charges.
Less favourable canonical structure

15. Molecules that don’t conform to the “Octet Rule”:
Electron-deficient molecules
Expanded valence shell molecules
ClF3
BH3 3s 3p 2s 2p Cl 3d B
Cl* B* F
3 H 1s 1s 1s 2s 2p F 2s 2p F 2s 2p
“Hypervalent molecules”
“Lewis acids”

16. Resonance structures
Some molecules cannot be adequately represented by a single Lewis structure. For example, two valid Lewis structures can be drawn for the anion (HCONH)¯. One structure has a negatively charged N atom and a C-O double bond; the other has a negatively charged O atom and a C-N double bond.
These structures are called resonance structures or resonance forms. A double headed arrow is used to separate the two resonance structures. 20

17. Introduction to Resonance Theory
Regarding the two resonance forms of (HCONH)¯ shown below, it should be noted that:
Neither resonance structure is an accurate representation for (HCONH)¯. The true structure is a composite of both resonance forms and is called a resonance hybrid.
The hybrid shows characteristics of both structures.
Resonance allows certain electron pairs to be delocalized over two or more atoms, and this delocalization adds stability.
A molecule with two or more resonance forms is said to be resonance stabilized.

18. Introduction to Resonance Theory
The following basic principles of resonance theory should be kept in mind:
Resonance structures are not real. An individual resonance structure does not accurately represent the structure of a molecule or ion. Only the hybrid does.
Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another.
Resonance structures are not isomers. Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons.

19. Drawing Resonance Structures
Rule [1]: Two resonance structures differ in the position of multiple bonds and nonbonded electrons. The placement of atoms and single bonds always stays the same.
Rule [2]: Two resonance structures must have the same number of unpaired electrons.

20. Drawing Resonance Structures
Rule [3]: Resonance structures must be valid Lewis structures. Hydrogen must have two electrons and no second-row element can have more than eight.

21. trigonal bipyramidal*
Valence Shell Electron Pair Repulsion Theory
A basic geometry can be assigned to each non-terminal atom based on the number of “objects” attached to it. Objects include bonded atoms (single, double, triple, partial bonds) and “lone pairs” of electrons.
VSEPR theory lets us predict the shape of a molecule based on the electron configurations of the constituent atoms. It is based on maximizing the distance between points on a spherical surface.
Number of
Objects 2 3 4 5 6
Geometry
linear
trigonal planar
tetrahedral
trigonal bipyramidal*
Octahedral

22. (t.b.p. or square pyramidal)
The geometry around an atom is described by the general formula:
AXmEn
Where X is a bonded atom, E is a lone pair and (m+n) is the number of objects (sometimes called the steric number, SN) around the central atom A.
Number of
Objects 2 3 4 5 6
Geometry
linear
trigonal planar
tetrahedral
trigonal bipyramidal
Octahedral
Formula (Shape)
AX2
AX3
(trig. planar)
AX2E
(bent)
AX4
(tetrahedral)
AX3E
(pyramidal)
AX2E2
AX5
(t.b.p. or square pyramidal)
AX4E
(seesaw)
AX3E2
(T-shaped)
AX2E3
(linear)
AX6
(octahedral)
AX5E
(square pyramidal)
AX4E2
(square planar)
AX3E3

23. Less common geometries
Number of
Objects 7 8
Geometry
pentagonal
bipyramidal
square
anti-prismatic
Xe- F F F F F
Xe is described as AX5E2 and has a pentagonal planar shape derived from the pentagonal bipyramidal geometry.
XeF5-
NMe4+

24. Refinement of VSEPR theory predicted geometries
The relative steric demand of objects is different and amount of repulsion caused by the object will alter the arrangement of the atoms around the central atom.
Lone pair of electrons
CH4
109.5°
Multiple bond polarized toward central atom
NH3
Increasing steric demand
106.6°
Normal single bond
OH2
Long single bond
polarized away from central atom
104.5°

25. Valence Bond Theory
Valence bond theory (VBT) is a localized quantum mechanical approach to describe the bonding in molecules. VBT provides a mathematical justification for the Lewis interpretation of electron pairs making bonds between atoms. VBT asserts that electron pairs occupy directed orbitals localized on a particular atom. The directionality of the orbitals is determined by the geometry around the atom which is obtained from the predictions of VSEPR theory.
In VBT, a bond will be formed if there is overlap of appropriate orbitals on two atoms and these orbitals are populated by a maximum of two electrons.
 bonds: have a node on the inter-nuclear axis and the sign of the lobes changes across the axis.
 bonds: symmetric about the internuclear axis

26. Valence Bond Theory
Detailed valence bond theory treatment of bonding in H2.
VBT considers the interactions between separate atoms as they are brought together to form molecules.
HA 1s1
HB 1s1
electron
A (1)
B (2)
Atomic wavefunction on atom B
1 = A(1) B(2)
Quantum mechanics demands that electrons can be interchangeable so we must use a linear combination of 1 and 2.
2 = A(2) B(1)
+ = N (1 + 2) (bonding, H-H)
3 = A(1) A(2) (ionic H- H+)
- = N (1 - 2) (anti-bonding)
4 = B(1) B(2) (ionic H+ H-)
molecule = N [1 + 2] + (C [3 + 4])
molecule = N [covalent + (C ionic)]
N is a normalizing coefficient
C is a coefficient related to the amount of ionic character

27. Valence Bond Theory
Valence bond theory treatment of bonding in H2 and F2 – the way it is generally used. F 2s 2p
HA 1s1
HB 1s1 F
A a
B b 2s 2p
Z axis
2pz
2pz
This gives a 1s-1s  bond between the two H atoms.
This gives a 2p-2p  bond between the two F atoms.
For VBT treatment of bonding, people generally ignore the anti-bonding combinations and the ionic contributions.

28. Valence bond theory treatment of bonding in O2
Z axis O 2s 2p
2pz
2pz
This gives a 2p-2p  bond between the two O atoms.
Lewis structure
Z axis
2py
2py
(the choice of 2py is arbitrary)
Double bond:  bond +  bond
Triple bond:  bond + 2  bond
This gives a 2p-2p  bond between the two O atoms. In VBT,  bonds are predicted to be weaker than  bonds because there is less overlap.
The Lewis approach and VBT predict that O2 is diamagnetic – this is wrong!

29. Directionality
The bonding in diatomic molecules is adequately described by combinations of “pure” atomic orbitals on each atom. The only direction that exists in such molecules is the inter-nuclear axis and the geometry of each atom is undefined in terms of VSEPR theory (both atoms are terminal). This is not the case with polyatomic molecules and the orientation of orbitals is important for an accurate description of the bonding and the molecular geometry.
Examine the predicted bonding in ammonia using “pure” atomic orbitals: 2s 2p N
The 2p orbitals on N are oriented along the X, Y, and Z axes so we would predict that the angles between the 2p-1s  bonds in NH3 would be 90°. We know that this is not the case. 1s 1s 1s
3 H
106.6°

30. Orthogonality and Normalization
Two properties of acceptable orbitals (wavefunctions) that we have not yet considered are that they must be orthogonal to every other orbital and they must be normalized. These conditions are related to the probability of finding an electron in a given space.
Orthogonal means that the integral of the product of an orbital with any other orbital is equal to 0, i.e.:
where n  m and dt means that the integral is taken over “all of space” (everywhere).
Normal means that the integral of the product of an orbital with itself is equal to 1, i.e.:
This means that we must find normalization coefficients that satisfy these conditions. Note that the atomic orbitals () we use can be considered to be both orthogonal and normal or “orthonormal”.

31. Example of the orthogonality of 1 and 2
Thus our hybrid sp orbitals are orthogonal to each other, as required.

32. A 2s orbital superimposed on a 2px orbital
Hybridization
The problem of accounting for the true geometry of molecules and the directionality of orbitals is handled using the concept of hybrid orbitals. Hybrid orbitals are mixtures of atomic orbitals and are treated mathematically as linear combinations of the appropriate s, p and d atomic orbitals.
Linear sp hybrid orbitals
A 2s orbital superimposed on a 2px orbital
The two resultant sp hybrid orbitals that are directed along the X-axis (in this case)
The 1/2 are normalization coefficients.

33. Valence bond theory treatment of a linear molecule: the bonding in BeH2
The promotion energy can be considered a part of the energy required to form hybrid orbitals. 2s 2p Be
Be* sp 2p
Be* (sp)
2 H 1s 1s
The overlap of the hybrid orbitals on Be with the 1s orbitals on the H atoms gives two Be-H (sp)-1s  bonds oriented 180° from each other. This agrees with the VSEPR theory prediction.

34. Valence bond theory treatment of a trigonal planar molecule: the bonding in BH32s 2p B B*
sp2 2p
B* (sp2)
This gives three sp2 orbitals that are oriented 120° apart in the xy plane – be careful: the choice of axes in this example determines the set of coefficients.

35. Valence bond theory treatment of a trigonal planar molecule: the bonding in BH3
sp2 2p B*
3 H 1s 1s 1s
The overlap of the sp2 hybrid orbitals on B with the 1s orbitals on the H atoms gives three B-H (sp2)-1s  bonds oriented 120° from each other. This agrees with the VSEPR theory prediction.

36. Valence bond theory treatment of a tetrahedral molecule: the bonding in CH42s 2p C C*
sp3
C* (sp3)
This gives four sp3 orbitals that are oriented in a tetrahedral fashion.

37. Valence bond theory treatment of a tetrahedral molecule: the bonding in CH42s 2p C C*
sp3
C* (sp3)
4 H 1s 1s 1s 1s
The overlap of the sp3 hybrid orbitals on C with the 1s orbitals on the H atoms gives four C-H (sp3)-1s  bonds oriented ° from each other. This provides the tetrahedral geometry predicted by VSEPR theory.

38. Valence bond theory treatment of a trigonal bipyramidal molecule: the bonding in PF53s 3p
PF5 has an VSEPR theory AX5 geometry so we need hybrid orbitals suitable for bonds to 5 atoms. ns and np combinations can only provide four, so we need to use nd orbitals (if they are available). P 3d P*
P* (sp3d) 3d 3s
3pz
3py
3px
3dz2
sp3dz2
The appropriate mixture to form a trigonal bipyramidal arrangement of hybrids involves all the ns and np orbitals as well as the ndz2 orbital.

39. Valence bond theory treatment of a trigonal bipyramidal molecule
The orbitals are treated in two different sets.
These coefficients are exactly the same as the result for the trigonal planar molecules because they are derived from the same orbitals (sp2)
These coefficients are similar to those for the sp hybrids because they are formed from a combination of two orbitals (pd).
Remember that d orbitals are more diffuse than s or p orbitals so VBT predicts that the bonds formed by hybrids involving d orbitals will be longer than those formed by s and p hybrids.

40. Valence bond theory treatment of a trigonal bipyramidal molecule: the bonding in PF5
P* (sp3d) 3d F 2s 2p F F 2s 2p 2s 2p F 2s 2p F 2s 2p
The overlap of the sp3d hybrid orbitals on P with the 2p orbitals on the F atoms gives five P-F (sp3d)-2p  bonds in two sets: the two axial bonds along the z-axis (180° from each other) and three equatorial bonds in the xy plane (120° from each other and 90° from each axial bond). This means that the 5 bonds are not equivalent!

41. An alternative, and maybe more reasonable, version of VBT treatment of a trigonal bipyramidal molecule:
The d orbitals are too high in energy to mix effectively with the s and p orbitals, so the trigonal bipyramidal molecule is actually composed of an equatorial set of trigonal (sp2) hybrids and the axial bonds come from an MO interaction between the two ligand orbitals and the pz orbital on the central atom.

42. The square pyramidal AX5 geometry requires mixing with a different d orbital than in the trigonal bipyramidal case.
Sb(C6H5)5
d orbitals
You should consider what orbital(s) would be useful for such a geometry and we will see a way to figure it out unambiguously when we examine the symmetry of molecules.

43. Valence bond theory treatment of an octahedral molecule:
the bonding in SF6 3s 3p S 3d S*
S* (sp3d2) 3d F F F F F F 3s
3pz
3py
3px
3dz2
3dx2-y2
sp3d2
The overlap of the sp3d2 hybrid orbitals on S with the 2p orbitals on the F atoms gives six S-F (sp3d2)-2p  bonds 90° from each other that are equivalent. You can figure out the normalization coefficients. As in the case of the TBP, there is also an MO approach that does not require d orbitals.

44. Valence bond theory treatment of p-bonding: the bonding in ClNO2s 2p N
sp2 2p
There are three “objects” around N so the geometry is trigonal planar. The shape is given by AX2E (angular or bent).
N*(sp2)    Cl O 3s 3p 2s 2p
A drawing of the VBT p bond in ClNO.
The overlap of the sp2 hybrid orbitals on N with the 3p orbital on Cl and the 2p orbital on O give the two  bonds and it is the overlap of the “left over” p orbital on N with the appropriate orbital on O that forms the (2p-2p) p bond between the two atoms.

45. Valence bond theory treatment of p-bonding: the bonding in the nitrate anion2s 2p N N+
sp2 2p
There are three “objects” around N so the geometry is trigonal planar. The shape is given by AX3 (trigonal planar).
N+*(sp2)  O-    2s 2p O- 2s 2p O
VBT gives only one of the canonical structures at a time. 2s 2p
The overlap of the sp2 hybrid orbitals on N with the the 2p orbitals on the O give the three (sp2-2p)  bonds and it is the overlap of the “left over” p orbital on N with the appropriate orbital on the uncharged O atom that forms the (2p-2p) p bond.

46. Valence bond theory treatment of p-bonding: the bonding in ethene2s 2p
Each C
Each C*
There are three “objects” around each C so the geometry is trigonal planar at each carbon. The shape is given by AX3 for each carbon.
sp2 2p
C*(sp2)  2p
sp2     
C*(sp2)
4 H 1s 1s 1s 1s
The overlap of the sp2 hybrid orbitals on C with the the 1s orbitals on each H give the four terminal (sp2-1s)  bonds. The double bond between the C atoms is formed by a (sp2- sp2)  bond and the (2p-2p) p bond.

47. Valence bond theory treatment of p-bonding: the bonding in SOCl2
sp3 3d
There are four “objects” around S so the geometry is tetrahedral and the shape is given by AX3E (pyramidal).
S*(sp3)     O Cl Cl 2s 2p
The overlap of the sp3 hybrid orbitals on S with the 3p orbitals on Cl and the 2p orbital on O give the three  bonds and, because the lone pair is located in the final sp3 hybrid, it is the overlap of the “left over” d orbital on S with an appropriate p orbital on O that forms the (3d-2p) p bond in the molecule.

48. Valence bond theory treatment of bonding: a hypervalent molecule, ClF33s 3p Cl 3d
Cl*
There are five “objects” around Cl so the geometry is trigonal bipyramidal and the shape is given by AX3E2 (T-shaped). Consider this: Why are such molecules T-shaped instead of pyramidal?
Cl* (sp3d) 3d F F F
The overlap of the sp3d hybrid orbitals on Cl with the 2p orbitals on the F atoms gives three P-F (sp3d)-2p  bonds in two sets: the two axial bonds along the z-axis (less than 180° from each other because of the repulsion from the lone pairs) and the one equatorial bond halfway between the other Cl bonds. Again, the bond lengths will not be the same because there is more d contribution to the axial hybrid orbitals.

49. Summary of Valence Bond Theory
Write an acceptable Lewis structure for the molecule.
Determine the number of VSEPR objects around all central atoms and determine the geometry around the atom.
Construct hybrid orbitals suitable for the predicted bonding.
Link orbitals together to make bonds.
Describe the bonding. Include the names of the orbitals involved in each bond. Draw pictures of the bonds formed by the overlap of these orbitals.
Two objects around Be, so AX2 (linear)
Two orbitals pointing 180° from each other needed, so use two sp hybrids 1s sp
Two (sp-1s) Be-H  bonds.

50. Limitations of Valence Bond Theory:
It fails to explain the tetravalency of carbon.
This theory does not discuss about energies of electrons.
The assumptions about the electrons being localized to specific locations.

51. Molecular Orbital Theory
• For example, when two hydrogen atoms bond, a σ1s (bonding) molecular orbital is formed as well as a σ1s* (antibonding) molecular orbital.
• The following slide illustrates the relative energies of the molecular orbitals compared to the original atomic orbitals.
• Because the energy of the two electrons is lower than the energy of the individual atoms, the molecule is stable.

53. In order to participate in MOs, atomic orbitals must overlap in space
In order to participate in MOs, atomic orbitals must overlap in space. (Therefore, only valence orbitals of atoms contribute significantly to MOs.)
Factors that determine orbital interaction: ! energy difference between the interacting orbitals ! magnitude of their overlap For the interaction to be strong, the energies of the two orbitals must be approximately equal and the overlap must be large

54. Bond Order
• The term bond order refers to the number of bonds that exist between two atoms. • The bond order of a diatomic molecule is defined as one-half the difference between the number of electrons in bonding orbitals, nb, and the number of electrons in antibonding orbitals, na.
bond order = 1/2 (na - nb )

56. Electron Configurations of Diatomic Molecules of the Second Period
1. Homonuclear diatomic molecules such as Li2 utilize only F orbitals.
For filled K shell bonding and antibonding orbitals use KK designation.
2. Be2 = KK(F2s)2 (F2s*)2
Bond order = 1/2 (2-2) = 0 So Be2 is unstable
3. For remaining elements, molecular orbitals must also be formed using p orbitals.

59. The overlap of “p” orbitals results in two sets of B orbitals (two bonding and two antibonding) and one set of F orbitals (one bonding and one antibonding).
The next slide illustrates the relative energies of these molecular orbitals for homonuclear diatomic molecules composed of 2nd period elements.
Note that the relative energies of the F and B bonding orbitals changes between N2 and O2.
For Li2 through C2, interactions between the F orbitals formed from the 2s and 2p atomic orbitals cause an increase in the energy of the F 2p MO’s.

64. Diamagnetism paired electrons repelled from induced magnetic field much weaker than paramagnetism
Paramagnetism unpaired electrons attracted to induced magnetic field much stronger than diamagnetism

67. O2 has 12 valence electrons. Its MO configuration is:
O2 = KK(σ2s) 2 (σ2s*)2 (σ2p) 2 (π2p) 2 (π2p) 2 (π2p*)(π2p*) with one electron in each of the π2p* orbitals spin aligned (Hund’s rule) and, bond order for O2 = 1/2(8-4) = 2

77. Outcomes of MO Model
As bond order increases, bond energy increases and bond length decreases.
Bond order is not absolutely associated with a particular bond energy.
N2 has a triple bond, and a correspondingly high bond energy.
O2 is paramagnetic. This is predicted by the MO model, not by the LE model, which predicts diamagnetism.
Combining LE and MO Models • σ bonds can be described as being localized. • π bonding must be treated as being delocalized.

78. THANK YOU