Chemical Bonding And Molecular Structures

O R B I T A L

O V E R L A P C O N C E P T

• If orbitals of 2 atoms are mixed with each other partially during bond formation, then the phenomenon is called as overlapping of orbitals.
• TYPES OF ORBITAL OVERLAPS:
• 1.Positive overlap: If the symmetry of both the atomic orbitals is the same, then it is called as positive overlap i.e. if the symmetry of the overlapping orbitals is either positive, or negative, then it is a type of positive overlap.

• 2.Negative overlap: If the symmetry of the atomic orbitals is not the same, i.e., one is positive and the other negative, then it is a type of negative overlap.

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• 3.Zero overlap: If overlap of orbitals present in 2 different planes takes place

, then it is called as zero overlap. E.g.. P_x overlaps with P_y (in real

situation, overlapping does not takes place.)

T Y P E S O F

O V E R L A P S

R E S U L T I N G

I N S I G M A

B O N D F O R M A T I O N

O V E R L A P S

R E S U L T I N G I N P I B O N D F O R M A T I O N

Due to lateral/sideways overlap of P-P orbitals present in the same plane, Pi bond is formed.

C O M P A R I

SIGMA BONDS

• It is a strong bond.
• Electron cloud is symmetrical along the inter- nuclear axis.
• There can be free rotations of atoms around this bond.
• These are less reactive.
• Shape of the molecule is determined by these bonds.
• Sigma bonds have independent existence.

S I O N

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PI BONDS

• It is a weak bond.
• Electron cloud is asymmetrical.

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• Free rotation is not possible around this bond.
• These are more reactive.
• These bonds do not affect the shape of the molecule.
• Pi bond always exists with a sigma bond.

H Y B R I D I Z A T I O N

The intermixing of different atomic orbitals of approximately equal energy levels to produce hybrid orbitals before bond formation is called as HYBRIDIZATION.

Here arrangement of hybrid orbitals are such that there is minimum repulsion in

between the hybrid orbitals.

No. of orbitals mixed=No. of hybrid orbitals produced.

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DIFFERENT TYPES OF HYBRIDISATIONS:

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Sp Hybridization

NO. of hybrid orbitals produced=2 Structure = LINEAR

Bond angle = 180 degree …..e.g. BeF₂

s p

H Y B R I D I D S A T I O N

• SP² Hybridization:
• No. of hybrid orbitals produced = 3
• Arrangement of these orbitals = TRIGONAL PLANAR
• Bond angle=120 degree……..ex. BF₃, etc.
• SP³ Hybridization:
• No. of hybrid orbitals = 4
• Arrangement= TETRAHEDRAL
• BOND ANGLE…..
• IN CH₄ = 109.28 degree
• IN NH₃ = 107.3 degree.
• Here, in methane, sigma bond is formed between H and C atom due to overlapping of sp3 orbital of H atom with S orbital of H atom.
• Structure of NH₃ – TRIANGULR PYRAMIDAL.

•

SP²

Hybridization

•

SP³

Hybridization

• SP³d Hybridization:
• NO. of hybrid orbitals produced = 5
• Structure = TRIGONAL BIPYRAMIDAL
• BOND ANGLES:
• Equatorial = 120 degree.
• Axial = 90 degree.
• If lone pair of electron is present at central atom, its position is always equatorial.Ex.PCl5
• Length of axial bonds is longer than that of equatorial bonds because of

minimum repulsion.

• IT may have different shapes according to no. of lone pairs it has:
• 1 lone pair – seesaw shape………e.g. SF₄
• 2 lone pairs – bent T shape………E.g. BrF₃
• 3 Lone pairs – Linear shape

• SP³d² Hybridization:
• No. of hybrid orbitals produced = 6
• Arrangement = OCTAHEDRAL……e.g. SF₆
• Shape may be:
• Square pyramidal…..e.g. BrF₅
• Square planar…….e.g. XeF4₄
• SP³d³ Hybridization:
• No. of hybrid orbitals = 7
• Arrangement = PENTAGONAL BIPYRAMIDAL
• BOND ANGLES:
• Axial with equatorial = 90 degree.
• Equatorial to equatorial = 72 degree.

( lone pairs = 0)

( lone pairs = 1) (lone pairs = 2)

• SP3d2 Hybridization

• Sp3d3 Hybridization

R U L E S R E G A R D I N G H Y B R I D I Z A T I O N

• Only orbitals of approximately same energy levels can take part.
• No. of orbitals mixed = No. of hybrid orbitals produced.
• Most hybrid orbitals are similar but not always identical in shape. They may differ from one another in their orientation in space.
• The electron waves in hybrid orbitals repel each other and this tend to the farthest apart.
• Hybrid orbitals can form only sigma bonds.
• Depending on the number and the nature of the orbitals undergoing hybridization, various types of hybrid orbitals directing towards the corners of specified geometrical figures come into existence.

• C O N D I T I O N S F O R C O M B I N A T I O N O F A T O M I C O R B I T A L S
• For atomic orbitals to combine, resulting in the formation of molecular orbitals , the main conditions are :
• The combining atomic orbitals should have almost the same energies. For example, in the case of diatomic molecules, 1s-orbital of one atom can combine with 1s- orbital of the other atom, but 1s-orbital of one atom cannot combine with 2s-orbital of the other atom.
• The extent of overlap between the atomic orbitals of the two atoms should be large.
• The combining atomic orbitals should have the same symmetry about the molecular axis. For

example, 2P_xorbital of one atom can combine with 2P_x orbital of the other atom but not with 2P_z orbital .

• Note : It may be noted that Z-axis is taken as the inter-nuclear axis according to modern conventions.

D e s i

M o l e c u l a r O r b i t a l s

Just as atomic orbitals are designated as

s, p, d, f etc molecular orbitals of diatomic molecules are named σ (sigma) ,

π (pi) , δ (delta) etc.

M O L E C U L A R O R B I T A L S

The molecular orbitals which are cylindrically symmetrical

around inter-nuclear axis are called σ - molecular orbitals. The molecular orbital formed by the addition of 1s orbitals is designated as σ 1s and the molecular orbital formed by subtraction of 1s orbitals is designated as σ * 1s .

Similarly combination of 2s orbital results in the

formation of two

2 s - molecular orbitals designated as σ 2s and σ * 2s

g n a t i

o n s o f

1. Determine the number of electrons in the molecule. We get the number of electrons per atom from their atomic number on the periodic table. (Remember to determine the total number of electrons, not just the valence electrons.)
2. Fill the molecular orbitals from bottom to top until all the electrons are added. Describe the electrons with arrows. Put two arrows in each molecular orbital, with the first arrow pointing up and the second pointing down.
3. Orbitals of equal energy are half filled with parallel spin before they begin to pair up.

Stability of the molecule with bond order.

Bond order = 1/2 (#e- in bonding MO's - #e- in antibonding MO's)

We use bond orders to predict the stability of molecules :-

• If the bond order for a molecule is equal to zero, the molecule is unstable.
• A bond order of greater than zero suggests a stable

molecule.

• The higher the bond order is, the more stable the bond.

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We can use the molecular orbital diagram to predict whether the molecule is paramagnetic or diamagnetic. If all the electrons are paired, the molecule is diamagnetic. If one or more electrons are unpaired, the molecule is paramagnetic.

1. The molecular orbital diagram for a diatomic hydrogen molecule, H₂, is

• The bond order is 1. Bond Order = 1/2(2 - 0) = 1
• The bond order above zero suggests that H₂is stable.
• Because there are no unpaired electrons, H₂ is diamagnetic.

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2. The molecular orbital diagram for a diatomic helium molecule, He₂, shows the following.

• The bond order is 0 for He₂. Bond Order = 1/2(2 - 2) = 0
• The zero bond order for He₂suggests that He₂is unstable.
• If He₂did form, it would be diamagnetic

3. The molecular orbital diagram for a diatomic oxygen molecule, O₂, is

• O₂has a bond order of 2. Bond Order = 1/2(10 - 6) = 2
• The bond order of two suggests that the oxygen molecule is stable.
• The two unpaired electrons show that O₂is paramagnetic

Diatomic molecules are molecules composed only of two atoms, of either the same or different chemical elements. The prefix di- is of Greek origin, meaning two. Common diatomic molecules are hydrogen (H₂), nitrogen (N₂), oxygen (O₂), and carbon monoxide (CO). Seven elements exist as homonuclear diatomic molecules at room temperature: H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂. Many elements and chemical compounds aside from these form diatomic molecules when evaporated. The noble gases do not form diatomic molecules: this can be explained using molecular orbital theory (see molecular orbital diagram).

INTRODUCTION

1. Two H atoms in their ground state configuration come together and form a single bond. The bond formation stabilizes both atoms and, therefore, is lower in energy than the atomic orbitals. This is also observed in Valence Bond Theory, which implies that each H atom in H₂shares its electron with one another, so that both can achieve the stable configuration of He.
2. On top of that, MO Theory allows one to compute the amount of energy released from a bond formation and a distance between two bonded atoms as well as predict the magnetic property of a molecule (or a substance). For H₂, the bond strength is -432 kJ/Mol, and the bond length is 74 angstrom (or 74 pm). H₂ is a diamagnetic molecule because the electrons paired up; therefore, it is not attracted by a magnetic field.

B o n d i n g

m o l e c u l a r o r

a n d

b i t a l

A n t i -b o n d i n g

s i n H ₂

1. Each H atom has a 1s atomic orbital. When two H atoms come to a proper proximity, their 1s orbitals interact and produce two molecular orbitals: a bonding MO and an anti-bonding MO.
2. If the electrons are in phase, they have a constructive interference. This results in a bonding sigma MO (σ_1s). This MO has an increased probability of finding electrons in the bonding region.

Figure 2: Schematic representation of the bonding molecular orbital σ(1s)

If the electrons are out of phase, they have a destructive interference. This results in an anti- bonding sigma MO (σ*_1s). This MO has a decreased probability of finding electrons in the bonding region. (Valence Bond Theory does not explain this phenomenon.)

Figure 3:Schematic representation of antibonding molecular orbital σ*(1s) Note that there is a nodal plane in the anti-bonding

B o n d o r d e r i n H 2

Bond order = 1/2 (#e- in bonding MO - #e- in antibonding MO)

For H₂, bond order = 1/2 (2-0) = 1, which means H₂ has only one bond. The antibonding orbital is empty. Thus, H₂ is a stable molecule.

Again, in the MO, there is no unpaired electron, so H₂ is diamagnetic

H Y D R O G E N B O N D

In compounds of hydrogen with strongly electronegative elements, such as fluorine, oxygen and

nitrogen, electron pair shared between the two atoms lie far away from the hydrogen atom. As a result, the hydrogen atom becomes highly electropositive with respect to the other atom. This phenomenon of charge separation in the case of hydrogen fluoride is represented as . Such a molecule is said to be polar .

The molecule behaves as a dipole because one end carries a positive charge and the other end a negative charge. The electrostatic force of attraction between such molecules should be very strong. This is

because the positive end of one molecule is attracted by the negative end of the other molecule . Thus, two or molecules may associate together to form larger cluster of molecules. This is illustrated below for the association of several molecules of hydrogen fluoride.

• The cluster of HF molecules may be described as (HF)_n.
• It may be noted that hydrogen atom is bonded to fluorine atom by a covalent bond in one molecule and by electrostatic force or by hydrogen bond to the fluorine atom in the adjacent molecule . Hydrogen atom is thus seen to act as a bridge between the two fluorine atoms.

• The hydrogen bond is represented by a dotted line. The solid lines represent the original(covalent ) bond present in the molecule.
• Chlorine, bromine and iodine are not as highly electronegative as fluorine and therefore, the shared pair of electrons in the case of HCl , HBr and HI do not lie as far away from hydrogen as in the case of HF. The tendency to form hydrogen bond in these cases is therefore less.
• Water molecule, because of its bent structure, is also a dipole, oxygen end carrying a negative charge and hydrogen end carrying a positive charge. Hydrogen bond taking place in this case as well, as represented below:

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• The cluster of water molecules may be described as (H₂O)_n

• The nature of hydrogen bond
• The hydrogen bond is a class in itself. It arises from electrostatic forces between positive end (pole) of one molecule and the negative end(pole) of the other molecule generally of the same substance. The strength of hydrogen bond has been has been found to vary between 10 - 40 kJ mol⁻¹ (i.e., 6.02 x

10²³ bonds) while that of a covalent bond has been found to be of the order of 400 kJ mol⁻¹ . Thus a hydrogen bond is very much weaker than a covalent

bond. Consequently, the length of hydrogen bond is bigger than the length of a covalent bond.

• In the case of hydrogen fluoride, for instance, while the length of the covalent bond between F and H atoms is 100 pm, the length of hydrogen bond between F and H atoms of neighbouring molecules is 155 pm.

• T y p e s o f h y d r o g e n b o n d i n g
• Hydrogen bonding may be classified into two types :
• I n t e r m o l e c u l a r

h y d r o g e n b o n d i n g

This type of hydrogen bonding involves electrostatic forces of attraction between hydrogen and

electronegative element of two different molecules of the substance. Hydrogen bonding in molecules of HF, NH₃ , H₂O etc. are examples of intermolecular hydrogen bonding.

• I n t r a m o l e c u l a r

h y d r o g e n b o n d i n g

This type of bonding involves electrostatic forces of attraction between hydrogen and electronegative element both present in the same molecule of the substance. Examples o-nitrophenol and salicylaldehyde.

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• p-Nitrophenol , on account of large distance between two groups , does not show any intramolecular hydrogen bonding. On the other hand, it shows the usual inter molecular hydrogen bonding , as illustrated below:

• As a result of intermolecular hydrogen bonding, the para derivative undergoes association, resulting in an increase in molar mass and hence an increase in boiling point. In ortho derivative, on account

of intramolecular hydrogen bonding , no such association is possible. Consequently, the ortho derivative is more volatile than the para derivative. Thus, while ortho nitrophenol is readily volatile in steam , para nitrophenol is completely non-volatile.

The two derivatives can thus be separated from each other by steam distillation.

• Density in solid state(ice) is less than that in liquid state . This is some what unusual because in most substances density in solid is more than that in liquid state.
• Water contracts when heated between 0°C and 4°C . This is again unusual because most substances expand when heated in all temperature ranges.
• Both these peculiar features are due to hydrogen bonding, as discussed below :

• In ice, hydrogen bonding between H₂O molecules is more extensive than in liquid water. A substance in solid state has a definite structure and the molecules are more rigidly fixed relative to one another than in the liquid state. In ice, the H₂O molecules are tetrahedrally oriented with respect to one another.At the same time

, each oxygen atom is surrounded tetrahedrally by four hydrogen atoms, two of these are bonded covalently and the other two by hydrogen bonds.The tetrahedral open cage-like crystal structure of ice. The central oxygen atom A is surrounded tetrahedrally by the oxygen atoms marked

1,2, 3 and 4.The hydrogen bonds are weaker and therefore, longer than covalent bonds. This arrangement gives rise to an open cage-like structure , as shown in the Fig. There are evidently a number of ‘holes' or open spaces.

• These holes are formed because the hydrogen bonds holding the H₂O molecules in ice are directed in certain definite angles . In liquid water such hydrogen bonds are fewer in number. Therefore, as ice melts, a large number of hydrogen bonds are broken. The molecules, therefore, move into the ‘hole' or open spaces and come closer to one another than they were in the solid state. This results in a sharp increase in density . The density of liquid water is, therefore higher than that of ice.
• As liquid water is heated from 0°C to 4°C, hydrogen bonds continue to be broken and the molecules come closer and closer together. This leads

to contraction. However, there is some expansion of

water also due to rise in temperature as in other liquids. It appears that up to 4°C, the former effect predominates and hence the volume increases as the temperature rises.

• It can be easily realised that without hydrogen bonding , water would have existed as a gas like hydrogen sulphide. In that case no life would have been possible on this globe.
• Hydrogen bonding also exists in all living organisms, whether of animal or of vegetable kingdom. Thus, it exists in various

tissues, organs, blood, skin and bones in animal life. It plays an important role in determining structure of proteins which are so essential for life.

• Hydrogen bonding plays an important role in making wood fibres more rigid and thus makes it an article of great utility. The cotton, silk or synthetic fibres owe their rigidity and tensile strength to hydrogen bonding. Thus hydrogen bonding is of vital importance for our clothing as well. Most of our food materials also consists of hydrogen bonded molecules. Sugars and carbohydrates , for example, have many -OH groups. The oxygen of one such group in one molecule is bonded with -OH group of another molecule through hydrogen bonding. Hydrogen bonding is thus a phenomenon of great importance in every day life.