Chemical Bonding And Molecular Structures

Prepared By:-Swastik Mishra Upasana NathRahul Kanungo Pratik Patnaik

1st ShiftKendriya Vidyalaya No-1, BHubaneswar

Kossel Lewis Approach

Kossel Lewis Approach To Chemical Bonding

• In 1916, Kossel and Lewis were the first to become independently successful in giving a satisfactory explanation about the formation chemical bond in terms of electrons.

• They were the first to provide some logical explanation of valence which was based on the inertness of noble gases.

Lewis approach

• He pictured the atom in terms of a positively charged ‘Kernel’ and the outer shell which could accommodate a maximum of eight electrons.

• He assumed that these eight electrons occupy the corners of a cube which surround the ‘Kernel’.

• Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds.

• In sodium and chlorine, this can happen by transfer of electron from sodium to chlorine.

• In case of other molecules like CL2 the bond is formed by sharing of a pair of electrons between atoms.

Lewis symbolIn formation of a molecule, only the

outer shell electron take part in chemical combination and they are called valence electrons.

The inner shells do not get involved in the combination process.

G.N. Lewis, an American chemist introduced simple notations to represent valence shell electrons in an atom and these are known as Lewis symbols.

Lewis Dot Structures for elements known at his time

Significance of Lewis symbol

• The number of dots represent the number of valence electrons.

• This number of valence electrons help to calculate the common or group valence of that element.

• The group valence of the element is generally either equal to the number of dots in Lewis symbol or 8 minus the number of dots.

• Other important facts given by Kossel are-• In periodic table, the highly electronegative

halogens and the highly electropositive alkali metals are separated by noble gases.

• The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of electron by the respective atoms.

• The negative and positive ions thus formed attain stable noble gas electronic configurations.

• The negative and positive ions are stabilized by electrostatic attraction.

Example• Na Na+ + e-

[Ne] 3s1 [Ne]

• Cl + e- Cl-[Ne} 3s2 3p5 [Ne] 3s2 3p6 or [Ar]

=> Na+ + Cl- NaCl

Octet rule

• According to electronic theory of chemical bonding, atoms can combine either by transfer of valence electrons from one atom to another or by sharing of valence in order to have an octet in their valence shells. this is known as octet rule.

Recall: Electrons are divided between

core and valence electrons.

ATOM core valenceNa 1s2 2s2 2p6 3s1 [Ne] 3s1

Br [Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5

Covalent Bonding

Covalent bond is the sharing of the VALENCE ELECTRONS of each atom in a bond

Br Br

Covalent bond

• The bond formed by atoms when they share their valence electrons to gain octet is known as covalent bond.

• When two atoms share one electron pair they are said to be joined by a single covalent bond.

• If two atoms share two pairs of electrons, the covalent bond between them is called a double bond.

• When combining atoms share three electron pairs , a triple bond is formed.

Covalent bond diagrams

H2O

NH3

CH4

Lewis representation of simple molecules(the Lewis structures)

• The Lewis dot structures provide a picture of bonding in molecules and ions in terms of shared pairs of electrons and the octet rule.

• The Lewis dot structures can be written by adopting the following steps.– 1)the total number of electrons required for

writing the structures are obtained by adding the valence electrons of the combining atoms.

– 2) for anions, each negative charge would mean addition of one electron. For cations, each positive charge would result in subtraction of one electron from the total number of valence electrons.

– 3)one must know the chemical symbols of the combining atoms and their skeletal structures.

– 4)in general, the least electronegative atom occupies the central position.

– 5)after accounting for shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as lone pairs.

Formal charge

• A formal charge (FC) is the charge assigned to an atom in a molecule, assuming that electrons in a chemical bond are shared equally between atoms, regardless of relative electro negativity.

• The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in Lewis structure.

• Formal charge = (total no. of valence electrons in the

free atom)-(total no. of non bonding electrons)-

(1/2(total no. of bonding electrons)

• Formal charge on ‘1’ =6-2-1/2(6)=+1• Formal charge on ‘2’ =6-4-1/2(4)=0• Formal charge on ‘3’ =6-6-1/2(2)=-1

• Formal charges do not indicate real charge separation within the molecule. These charges help in keeping track of the valence electrons in the molecule.

• Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species.

• Generally the lowest energy structure is the one with the smallest formal charges on the atoms. The formal charge is a factor based on pure covalent view of bonding in which electron pairs are shared equally by neighboring atoms.

Limitations of the octet rule• It is not universal. There are

exceptions to it.1)the incomplete octet of the central atom

In some compounds, the number of electrons

surrounding the central atom is less than eight. Examples- LiCl, BeH2

and BCl3

2) Odd-electron molecule in molecules with an odd number of

electrons like nitric oxide, NO and nitrogen dioxide, the octet rule is not satisfied for all atoms.

3)The Expanded Octet Elements in and beyond the third

period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as expanded octet.

Examples-PF5, SF6, H2SO4

4) Octet rule is based on the chemical inertness of noble gases. However, some noble gases also combine with oxygen and fluorine to form a number of compounds like XeF2, KrF2, XeOF2 etc.

5) This theory does not account for the shape of the molecules.

6) It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.

Bond Parameters

Bond parameters• Bond length bond length is defined as the equilibrium

distance between the nuclei of two bonded atoms in a molecule.

Bond lengths are measured by spectroscopic, X-ray diffraction techniques etc.

Each atom of the bonded pair contribute to the bond length.

• In case of a covalent bond, the contribution from each atom is called the covalent radius of that atom.

• The covalent radius is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation.

• The van der Waals radius represents the overall size of the atom which includes its valence shell in a non bonded situation.

• Bond angle It is defined as the angle between the

orbitals containing bonding electron pairs around the central atom in a molecule or complex ion.

Bond angle is expressed in degree.

• Bond enthalpy It is defined as the amount of energy required to break one

mole of bonds of a particular type between two atoms in a gaseous state.

The unit of bond enthalpy is kJ per mol.Example-

The H-H bond enthalpy in hydrogen molecule is 435.8 kJ/mol.

H2 -------> H(g)+H(g) ΔaH=435.8 kJ/ mol

O2 (O=O) (g) ------> O(g)+O(g) ΔaH=498 kJ/mol

• The larger the bond dissociation enthalpy, stronger will be the bond in the molecule.

HCl---H + Cl • In case of polyatomic molecules, the

measurement of bond strength in the following way- let us take the example of H2O molecule.H2O---H+OH; ΔaH1=502 kJ/mol

OH---H+O; ΔaH2=427 kJ/mol In this case, the average bond enthalpy is taken. average bond enthalpy=502+427=464.5 kJ/mol

2

Bond order• In Lewis description of covalent bond, the

bond order is given by the number of bonds between the two atoms in a molecule.

• For example, the bond order of h2 which has a single shared pair of electrons is 1

• The bond order of O2 with 2 shared pairs of electrons is 2

• The bond order of N2 with 3 shared pairs of electron is 3.

• For N2, the bond order is 3 and its bond enthalpy is 946 kJ/mol and it is one of the highest for diatomic molecule.

• Isoelectronic molecules and ions have identical bond orders.

example- N2 and CO have bond order 3.

A general correlation useful for understanding the stabilities of molecules is that: with increase in bond order, bond enthalpy increases and bond length decreases.

Resonance

Resonance structures• A single Lewis structure may be inadequate

for representation of a molecule.• According to the concept of resonance,

whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, position of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.

• For example, O3 molecule can be represented by two structures as shown in the figure.

• For O3 these two structures constitute the canonical structures or resonance structures and their hybrid shown by the third figure represents the structure of o3 more accurately. this is called resonance hybrid.

• In general it may be stated that– Resonance stabilizes the molecule as

the energy of the resonance hybrid is less than the energy of any single canonical structure

– Resonance averages the bond characteristics as a whole.Resonance Energy

It is the difference between the actual bond energy of the molecule and that of the most stable of the resonating structures (having least energy).

Polarity Of Bonds

 Polarity of BondsThe existence of a hundred percent ionic or covalent bond represents

an ideal situation. In reality no bond or a compound is either completely covalent or ionic. Even in case of covalent bond between two hydrogen atoms, there is some ionic character.

When covalent bond is formed between two similar atoms, for example in H2, O2, Cl2, N2 or F2, the shared pair of electrons is equally attracted by the two atoms. As a result electron pair is situated exactly between the two identical nuclei. The bond so formed is called non polar covalent bond. Contrary to this in case of a heteronuclear molecule like HF, the shared electron pair between the two atoms gets displaced more towards fluorine since the electronegativity of fluorine (Unit 3) is far greater than that of hydrogen. The resultant covalent bond is a polar covalent bond.

As a result of polarisation, the molecule possesses the dipole moment (depicted below) which can be defined as the product of the magnitude of the charge and the distance between the centres of positive and negative charge. It is usually designated by a Greek letter ‘μ;’. Mathematically, it is expressed as follows :

Dipole moment (μ) = charge (Q) x distance of separation (r)Dipole moment is usually expressed in Debye units (D).The

conversion factor is1 D = 3.33564 x 10 �-30 C mwhere C is coulomb and m is meter.

• Further dipole moment is a vector quantity and is depicted by a small arrow with tail on the positive centre and head pointing towards the negative centre. For example the dipole moment of HF may be represented as :

• The shift in electron density is symbolised by crossed arrow ( ) above the Lewis structure to indicate the direction of the shift.

• In case of polyatomic molecules the dipole moment not only depend upon the individual dipole moments of bonds known as bond dipoles but also on the spatial arrangement of various bonds in the molecule. In such case, the dipole moment of a molecule is the vector sum of the dipole moments of various bonds. For example in H2O molecule, which has a bent structure, the two O-H bonds are oriented at an angle of 104.50. Net dipole moment of 6.17 x 10 �-30 C m (1D = 3.33564 x 10 �-30 C m) is the resultant of the dipole moments of two O-H bonds.

• Peter Debye, the Dutch chemist received Nobel prize in 1936 for his work on X-ray diffraction and dipole moments. The magnitude of the dipole moment is given in Debye units in order to honour him.

• Net Dipole moment, μ = 1.85 D = 1.85 x 3.33564 x 10 �-30 C m = 6.17 1030 C m�

• The dipole moment in case of BeF2 is zero. This is because the two equal bond dipoles point in opposite directions and cancel the effect of each other.

• In tetra-atomic molecule, for example in BF3, the dipole moment is zero although the B – F bonds are oriented at an angle of 120° to one another, the three bond moments give a net sum of zero as the resultant of any two is equal and opposite to the third.

• In case of NH3 and NF3 molecule. Both the molecules have pyramidal shape with a lone pair of electronson nitrogen atom. Although fluorine is more electronegative than nitrogen, the resultant dipole moment of NH3 ( 4.90 x 10–3 the orbital dipole due to lone pair is in the same direction as the resultant dipole moment of the N – H bonds, whereas in NF3the orbital dipole is in the direction opposite to the resultant dipole moment 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.

The partial covalent character of ionic bonds was discussed by Fajan’s in terms of the following rules:

• The smaller the size of the cation and the larger the size of the anion, the greater the covalent character of an ionic bond.• The greater the charge on the cation, the greater the covalent character of the ionic bond.• For cations of the same size and charge, the one, with electronic configuration (n-1)dnnso, typical of transition metals, is more polarising than the one with a noble gas configuration, ns2 np6, typical of alkali and alkaline earth metal cations. The cation polarises the anion, pulling the electronic charge toward itself and thereby increasing the electronic charge between the two. This is precisely what happens in a covalent bond, i.e., build-up of electron charge density between the nuclei. The polarising power of the cation, the polarisability of the anion and the extent of distortion (polarisation) of anion are the factors, which determine the per cent covalent character of the ionic bond.

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory

• Sidgwick and Powell in 1940, proposed a simple theory based on the repulsive interactions of the electron pairs in the valence shell of the atoms. It was further developed and redefined by Nyholm and Gillespie (1957).

• The main postulates of VSEPR theory are as follows:

• • The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom.

• • Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.

• • These pairs of electrons tend to occupy such positions in space that minimise repulsion and thus maximise distance between them.

• • The valence shell is taken as a sphere with the electron pairs localising on the spherical surface at maximum distance from one another.

• • 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.

• • Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure.

• The repulsive interaction of electron pairs decrease in the order:

• Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp)

• Nyholm and Gillespie (1957) refined the VSEPR model by explaining the important difference between the lone pairs and bonding pairs of electrons. While the lone pairs are localised on the central atom, each bonded pair is shared between two atoms. As a result, the lone pair electrons in a molecule occupy more space as compared to the bonding pairs of electrons. This results in greater repulsion between lone pairs of electrons as compared to the lone pair – bond pair and bond pair – bond pair repulsions. These repulsion effects result in deviations from idealised shapes and alterations in bond angles in molecules.

• For the prediction of geometrical shapes of molecules with the help of VSEPR theory, it is convenient to divide molecules into two categories as (i) molecules in which the central atom has no lone pair and (ii) molecules in which the central atom has one or more lone pairs.

• The VSEPR Theory is able to predict geometry of a large number of molecules, especially the compounds of p-block elements  accurately. It is also quite successful in determining the geometry quite-accurately even when the energy difference between possible structures is very small. The theoretical basis of the VSEPR theory regarding the effects of electron pair repulsions on molecular shapes is not clear and continues to be a subject of doubt and discussion.

•  Shape (geometry) of Some Simple Molecules/Ions with Central Ions having One or more Lone Pairs of Electron(E).

• Shapes of Molecules containing Bond Pair and Lone Pair

ORBITAL OVERLAP CONCEPT

ORBITAL OVERLAP CONCEPT

• 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.

• 3.Zero overlap: If overlap of orbitals present in 2 different planes takes place , then it is called as zero overlap. E.g.. Px overlaps with Py (in real situation, overlapping does not takes place.)

TYPES OF OVERLAPS RESULTING IN SIGMA BOND

FORMATION

TYPES OF OVERLAPS RESULTING IN PI BOND FORMATION

Due to lateral/sideways overlap of P-P orbitals present in the same

plane, Pi bond is formed.

COMPARISION 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.

PI BONDS• It is a weak bond.• Electron cloud is

asymmetrical.

• 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.

HYBRIDIZATIONThe 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.

DIFFERENT TYPES OF HYBRIDISATIONS:

Sp HybridizationNO. of hybrid orbitals produced=2Structure = LINEARBond angle = 180 degree …..e.g. BeF₂

sp HYBRIDIDSATION

• 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 ( lone pairs = 0)• Shape may be:• Square pyramidal…..e.g. BrF₅ ( lone pairs = 1)• Square planar…….e.g. XeF4₄ (lone pairs = 2)• SP3d3 Hybridization:• No. of hybrid orbitals = 7• Arrangement = PENTAGONAL BIPYRAMIDAL• BOND ANGLES:• Axial with equatorial = 90 degree.• Equatorial to equatorial = 72 degree.

• SP3d2 Hybridization

• Sp3d3 Hybridization

RULES REGARDING HYBRIDIZATION

• 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.

Molecular Orbital Theory

MOT

• CONDITIONS FOR COMBINATION OF ATOMIC ORBITALS

• 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, 2Pxorbital of one atom can combine with 2Px orbital of the other atom but not with 2Pz orbital .

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

Designations of Molecular OrbitalsJust 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.

MOLECULAR ORBITALSThe 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 two2 s - molecular orbitals designated as σ 2s and σ * 2s

Orbital Diagram Of O2 And O

O2 O

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.

DETERMINATION OF SHAPEOF MOLECULES

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.

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.

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

Examples

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).

BONDING IN HOMONUCLEAR DIATOMIC MOLECULES

INTRODUCTION1.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.

H2 MOLECULE

Bonding and Anti-bonding molecular orbitals in H2

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 (σ1s). 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

Bond order in H2

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

Hydrogen Bond

HYDROGEN BONDIn 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 attractive force which binds the atom of one molecule with an electronegative atom (such as fluorine, oxygen or nitrogen) of another molecule, generally of the same substance , is known as hydrogen bond .

• 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:

• 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.

• Types of hydrogen bonding

• Hydrogen bonding may be classified into two types :

• Intermolecular hydrogen bonding 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.

• Intramolecular hydrogen bondingThis 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.

• 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.

• Some Unique Properties of Water

• 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.

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