Bond Parameters

Bond parameters refer to the characterization of covalent bond on the

basis of various parameters like bond length, bond angle, bond

enthalpy. In this chapter, we will learn more about the concept of bond

parameters.

Introduction

Different atoms combine all together in order to become stable. By

forming bonds, this combination takes place. There are different types

of bond, namely, ionic or electrovalent bond, covalent bond, and

coordinate bond. This, in turn, shows that every bond has some feature

associated with it.

Here are some of the different features or characteristics of bonds

which can be termed as bond parameters. A particular covalent bond

is characterized by certain parameters:

● Bond Length

● Bond Angle

● The Bond Enthalpy

● Bond Order

Bond Length

The bond length refers to the distance between the centers of the

nuclei of two bonded atoms in an equilibrium position. The stronger

the force of attraction in between the bonding atoms, the smaller is the

length of the bond. However, the bigger the atom size, the longer the

bond length.

Browse more Topics under Chemical Bonding And Molecular Structure

● Covalent Compounds

● Fundamentals of Chemical Bonding

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Resonance Structures

● Valence Bond Theory

● VSEPR Theory

It is measured by spectroscopic, X-ray diffraction and electron

diffraction technique. Each atom of the bonded pair contributes to the

bond length. In case of a covalent bond, the contribution by each atom

is the covalent radius of that atom.

The certain factors upon which the bond length is dependent are –

● Bond Multiplicity: The bond length decreases with an increase

in bond multiplicity.

● Size of an Atom: Bond length is directly proportional to the

size of an atom. The bond length increases with the increase in

the size of the atoms.

The stronger the force of attraction between the bonding atom, the

smaller the bond length. However, the bigger the size of an atom, the

longer will be the bond length. Also, it is to be noted that in the case

of a covalent bond, the contribution by each atom is referred to as the

covalent radius of that atom.

Bond Angle

Bond angle refers to the angle between the two bonds i.e. the angle

between two orbitals that contains a pair of bonding electron around

the central atom in a complex molecule or an ion. This angle is usually

measured in degrees, further calculated using the spectroscopic

method.

This gives a clear idea about the distribution of bonded electron pairs

around the atoms and helps in determination of the shape of the

molecules. It also gives an idea about the bonded electron pairs

distribution around the atoms and determining the shape of the

molecules.

Bond Enthalpy

The amount of energy which is needed in order to break one mole of

the bond of a particular type between two atoms in a gaseous state is

referred to as the Bond Enthalpies. Bond enthalpy is directly

proportional to the strength of the bond between the molecules.

In case of polyatomic molecules, the two bonds of the same type can

have different bond enthalpy. For e.g.: Two O-H bonds of water

molecule have different bond enthalpy. Due to differences in bond

enthalpy, polyatomic molecules have average bond enthalpy.

Factors affecting the bond Enthalpy:

● Atomic Size

● Electronegativity

● Extent of overlapping

● Bond Order

Bond Order

As per the Lewis description of covalent bonds, the bond order is the

number of bonds that forms in between the two atoms in a molecule.

The Isoelectronic molecules or ions have the same bond order.

For example, the two isoelectronicc molecules, F2 and O22- are

isoelectronic molecules and so have the same bond order of 1. The

greater the order of the bond, there is an increase in bond enthalpy and

a decrease in the length of the bond.

The bond order in H2 wherein one electron pair is shared is one, in O2

where two electron pairs shared is two and in N2 in which three

electron pairs are shared is three.

● H – H Bond order = 1

● O = O Bond order = 2

● N ≡ N Bond order = 3

● C ≡ O Bond order = 3

Isoelectronic species have the same bond order. For Example, F2, O22-

(18 electrons) have bond order 1 N2, CO and NO+ (14 electrons) have

bond order = 3

Important Points Regarding the Bond Order

● The Isoelectronic species, those species that have the same

number of electrons, have equal bond orders. Considering, for

example, N2, NO+ and CO have a total of 14 electrons and all

of them have the equal bond order of 3.

● The greater the order of the bond, the greater is the stability of

molecules.

● The greater the order of the bond, the shorter is the length of

the bond.

Solved Example For You

Question: Determine the bond order for nitrate ( NO3− ) ion.

Answer :

1. The Lewis structure of nitrate ion is:

2. Total number of bonds in the molecule = 4

3. Total number of bond groups between individual atoms = 3

4. Hence, the bond order is 4/3 = 1.33

Covalent Compounds

You have now a brief idea of why different elements behave

differently. But do you a know a major part of it is because of the

“nature” of the bonds in the compounds. Just like you and your best

friends have a number of differences due to the “inner” qualities, so is

the case with ionic and covalent compounds. In this chapter, we will

learn more about the concept of covalent compounds, look at their

properties and more.

What is a Covalent Compound?

Covalent compounds are the ones having strong intra-molecular

bonds. This is because the atoms within the covalent molecules are

very tightly held together. Each molecule is indeed quite separate and

the force of attraction between the individual molecules in a covalent

compound tends to be weak.

We require very little energy in separating the molecules. This is

because of the attractive forces between the molecules with the

absence of overall electric charge. Covalent compounds are usually

gaseous molecules at room temperature and pressure. They might also

be liquids with low relatively low boiling points.

These characteristics could be attributed to their weak intermolecular

forces which hold these atoms together. However, we also have a lot

of solid covalent compounds. They have low melting points.

However, it is interesting to note that a small number of these have a

completely different structure. They form huge structures where a

huge number of atoms are held together. This is possible due to the

presence of shared electrons.

These giant molecular structures are basically lattices made up of

molecules which are held together by covalent bonds structure. These

covalent bonds are very strong. They also tend to be very hard with

high melting points which are different from most of the covalent

compounds. The example of this kind of covalent compounds includes

diamond and graphite of carbon atom network. They also include

silica of silicon and oxygen atoms network.

Browse more Topics under Chemical Bonding And Molecular Structure

● Bond Parameters

● Fundamentals of Chemical Bonding

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Resonance Structures

● Valence Bond Theory

● VSEPR Theory

General Properties of Covalent Compounds

● Covalent compounds usually have low melting points. An

exception to this include molecules of silica and diamonds that

have a high melting point.

● These compounds have low boiling points. This can be

attributed to their weak force of attraction between the various

bonded atoms. Van Der Waals forces bind these atoms.

● These compounds are usually gases and liquids with low

boiling and melting points.

● The solid covalent compounds have soft structures like

graphite. This is because of the presence of a cloud of electrons

in between each layer of carbons atoms.

● These compounds are non-conductors of electrical charge. The

absence of charged ions is the main reason behind this. An

exception to this is graphite, where we see a cloud of electrons.

These make graphite a good conductor.

● They are bad conductors of heat also. Their molecules lack free

electrons and that obstructs the flow of heat energy.

● Covalent compounds do not possess polar characteristics as a

general property. Therefore, these compounds are insoluble in

water. Water molecules are not absolutely neutral and have a

slight negative charge on the oxygen atom and slight positive

charges on the hydrogen atoms and since covalent compounds

are made up of neutral molecules or molecules with slight

charges and hence are not attracted to water molecules

strongly.

(Source: Google)

Physical And Chemical Properties

● The liquid covalent compounds evaporate. This means the

molecules of liquids and solids loses from their surface into the

air.

● These compounds have very less affinity between their

molecules.

● Various covalent compounds have their own characteristically

shaped molecules. Their bonds are directed at pre-set angles.

● Some compounds especially medicines are soluble in water.

The rest are soluble in oil.

● Most of the covalent compounds are non-polar or have very

little tendency to split completely to form ions and hence never

conduct electricity.

● At normal temperature and pressure, we will find these

compounds as either liquids or gases. But, there are solids as

well and they have higher molecular weights.

● The covalent compounds crystals are of two types: One that has

weak van der Waal force holding these together like in Iodine.

These are easily fusible and volatile The other having a large

network of atoms setting up the macromolecules.

● These compounds are soluble in organic solvents like ether and

benzene.

● Covalent bonds are directional in nature. Therefore, they

exhibit the phenomenon of isomerism.

● Covalent compounds majorly have a very slow rate of

reactions, unlike the various ionic compounds.

Solved Examples for You

Question: Why are covalent compounds not soluble in water?

Answer: Water molecules are not absolutely neutral. These molecules

have a slight negative charge on the oxygen atom and slight positive

charges on the hydrogen atoms. On the other hand, we know that the

covalent compounds are made up of neutral molecules or molecules

with slight charges. It is for this reason that these compounds are not

attracted to water molecules strongly.

Fundamentals of Chemical Bonding

So, why do the reactions occur as they occur? Can you mix just

hydrogen and oxygenin a gas chamber and get the water that you

drink? No! That is why we are going to look at the fundamentals of

chemical bonding in this chapter. We will first understand what

chemical bonding actually is. Then we will look at its types and more.

What is Chemical Bonding?

Why do different elements or substances undergo a transformation

under a set of condition? As students of chemistry, we must know

that. The answer to this interesting question lies in the chapter of

fundamentals of chemical bonding.

Each and everything in this universe tries to become stable. This

happens only by losing energy. In an interaction between two types of

matter, one exerts a force on another. The energies between the two

types of matter depend on the nature of the force between them. If the

force is attractive, the energy decreases. On the other hand, if it’s

repulsive in nature then the energy increases.

In cases where the force is attractive, the two atoms get bound

together. Such a force is what we call a chemical bond. Thus, we can

define chemical bonding as follows: “The attractive force which holds

various constituents (atom, ions, etc.) together and stabilizes them by

the overall loss of energy is chemical bonding.”

Browse more Topics under Chemical Bonding And Molecular Structure

● Bond Parameters

● Covalent Compounds

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Resonance Structures

● Valence Bond Theory

● VSEPR Theory

Learn the Tricks In understanding the Chemical Bonding

Fundamentals of Chemical Bonding

● How does it work?

Like many other people, you might also question why do some atoms

combine with only specific atoms? How do we know which pair of

atoms will combine and which will not? Well, there are reasons

behind this. Let us explore in the below segment.

● Why do atoms combine?

Atoms combine together to lose their energy. This would make them

stable.

● Why do certain atoms combine while others do not?

This is mainly because a compound forms only when there is an

attractive force leading to the lowering of energy. On the other hand,

in case of a repulsive force, we find an increase in overall energy of

the system. Thus, we do not see the formation of any compounds.

● How do we know which pair of atoms will combine and which

will not?

To answer this question, we will have to combine our knowledge of

the periodic table, the elements’ electronic configuration and the

atomic structure.

Types of Chemical Bonding

Atoms gain stability by majorly four types of bonding methods. They

are:

● Ionic bond

● Covalent bond

● Hydrogen bond

● Polar bond

1) Ionic Bond

Kössel and Lewis were the first scientists to explain the formation of

chemical bonds successfully. They used the concept of inertness of

noble gases to explain the fundamentals of chemical bonding.

According to their theory of ionic and chemical bonding, ionic

bonding involves the transfer of electrons between atoms. In such

cases, one atom loses an electron and the other atom gains an electron.

When an electron transfer occurs, one atom has a negative charge

making it the anion. On the other hand, the other atom has a positive

charge making it the cation. The ionic bond is majorly strong because

of the concept of “opposite charges attract.”

2) Covalent Bond

Covalent bonding is the most common method of bonding that we

witness in compounds containing carbon. These are basically the

organic compounds. A Covalent bond signifies the sharing of

electrons between atoms.

In such cases, we know that there is a formation of a new orbit by the

shared pair of electrons. This orbit extends around the nuclei of both

the atoms and creates a new molecule.

Polar bonds and hydrogen bonds are actually secondary types of

covalent bonding.

3) Polar Bond

A covalent bond can be of two types:

● Polar Bond

● Non-polar Bond

In a Polar bond, we witness the electrons to be shared unequally.

These tend to be closer to one atom than the other. Due to this uneven

distance between the electron and atom, a charge difference is created

in the different parts of the atom.

Because of this, one end of the molecule will be slightly positively

charged and the other end is slightly negatively charged. Water is an

example of a polar molecule.

4) Hydrogen Bond

A hydrogen bond is a weaker form of bonding as a contrast to ionic

and covalent bonding. Hydrogen bonding is a type of polar covalent

bonding (O-H bonding). In this case, the hydrogen has a slightly

positive charge. This means that electrons are pulled more towards the

other element.

Hydrogen bonds are responsible for many important characteristics of

the substances around us. Some of these instances include the

structure of DNA, proteins and the properties of water.

Solved Example for You

Q: What are London dispersion forces?

Ans: London dispersion forces are those that occur due to the

temporary imbalances of charge occurring inside an atom. The

concentration of charge of the atoms undergoes constant changes as

these electrons are always in motion. This creates a temporary shift in

the overall charge distribution of the atom.

When this particle happened to come in contact with another, the

temporary imbalance of charge will result in an attraction of positive

and negative charges. These are the London dispersion forces.

Hybridisation

The formation of bonds is no less than the act of courtship. Atoms

come closer, attract to each other and gradually lose a little part of

themselves to the other atoms. In chemistry, the study of bonding, that

is, Hybridization is of prime importance. What happens to the atoms

during bonding? What happens to the atomic orbitals? The answer lies

in the concept of Hybridisation. Let us see!

Introducing Hybridisation

All elements around us, behave in strange yet surprising ways. The

electronic configuration of these elements, along with their properties,

is a unique concept to study and observe. Owing to the uniqueness of

such properties and uses of an element, we are able to derive many

practical applications of such elements.

When it comes to the elements around us, we can observe a variety of

physical properties that these elements display. The study of

hybridization and how it allows the combination of various molecules

in an interesting way is a very important study in science.

Understanding the properties of hybridisation lets us dive into the

realms of science in a way that is hard to grasp in one go but excellent

to study once we get to know more about it. Let us get to know more

about the process of hybridization, which will help us understand the

properties of different elements.

Browse more Topics under Chemical Bonding And Molecular Structure

● Bond Parameters

● Covalent Compounds

● Fundamentals of Chemical Bonding

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Resonance Structures

● Valence Bond Theory

● VSEPR Theory

What is Hybridization?

Scientist Pauling introduced the revolutionary concept of

hybridization in the year 1931. He described it as the redistribution of

the energy of orbitals of individual atoms to give new orbitals of

equivalent energy and named the process as hybridisation. In this

process, the new orbitals come into existence and named as the hybrid

orbitals.

Rules for Observing the Type of Hybridisation

The following rules are observed to understand the type of

hybridisation in a compound or an ion.

● Calculate the total number of valence electrons.

● Calculate the number of duplex or octet OR

● Number of lone pairs of electrons

● Number of used orbital = Number of duplex or octet + Number

of lone pairs of electrons

● If there is no lone pair of electrons then the geometry of

orbitals and molecule is different.

Types of Hybridisation

The following are the types of hybridisation:

1) sp – Hybridisation

In such hybridisation one s- and one p-orbital are mixed to form two

sp – hybrid orbitals, having a linear structure with bond angle 180

degrees. For example in the formation of BeCl2, first be atom comes

in excited state 2s12p1, then hybridized to form two sp – hybrid

orbitals. These hybrid orbitals overlap with the two p-orbitals of two

chlorine atoms to form BeCl2

2) sp2 – Hybridisation

In such hybridisation one s- and to p-orbitals are mixed form three

sp2– hybrid orbitals, having a planar triangular structure with bond

angle 120 degrees.

3) sp3 – Hybridisation

In such hybridisation one s- and three p-orbitals are mixed to form

four sp3– hybrid orbitals having a tetrahedral structure with bond angle

109 degrees 28′, that is, 109.5 degrees.

Studying the Formation of Various Molecules

1) Methane

4 equivalent C-H σ bonds can be made by the interactions of C-sp3

with an H-1s

2) Ethane

6 C-H sigma(σ) bonds are made by the interaction of C-sp3 with H-1s

orbitals and 1 C-C σ bond is made by the interaction of C-sp3 with

another C-sp3 orbital.

3) Formation of NH3 and H2O molecules

In NH2 molecule nitrogen atom is sp3-hybridised and one hybrid

orbital contains two electrons. Now three 1s- orbitals of three

hydrogen atoms overlap with three sp3 hybrid orbitals to form NH3

molecule. The angle between H-N-H should be 109.50 but due to the

presence of one occupied sp3-hybrid orbital the angle decreases to

107.80. Hence, the bond angle in NH3 molecule is 107.80.

4) Formation of C2H4 and C2H2 Molecules

In C2H4 molecule carbon atoms are sp2-hybridised and one 2p-orbital

remains out to hybridisation. This forms p-bond while sp2 –hybrid

orbitals form sigma- bonds.

5) Formation of NH3 and H2O Molecules by sp2 hybridization

In H2O molecule, the oxygen atom is sp3 – hybridized and has two

occupied orbitals. Thus, the bond angle in the water molecule is

105.50.

A Solved Question for You

Q: Discuss the rules of hybridisation. Are they important to the study

of the concept as a whole?

Ans: Yes, the rules of hybridisation are very important to be studied

before diving into the subject of hybridisation. Hence, these rules are

essential to the understanding of the concepts of the topic. The

following are the rules related to hybridisation:

● Orbitals of only a central atom would undergo hybridisation.

● The orbitals of almost the same energy level combine to form

hybrid orbitals.

● The numbers of atomic orbitals mixed together are always

equal to the number of hybrid orbitals.

● During hybridisation, the mixing of a number of orbitals is as

per requirement.

● The hybrid orbitals scattered in space and tend to the farthest

apart.

● Hybrid bonds are stronger than the non-hybridised bonds.

When you once use an orbital to build a hybrid orbital it is no longer

available to hold electrons in its ‘pure’ form. You can hybridize the s

– and p – orbitals in three ways.

Hydrogen Bond

How many bonds have you learnt up till now? We are not talking

about the relationship bond. Here we are talking about the bonds in

chemistry. You must have heard about the covalent bond, ionic bond,

etc, right? Now we will see the hydrogen bond in detail in this section.

Hearing this for the first time? Or just have a vague idea about it? No

worries! We will clear all your doubts. Keep reading below.

What is a Hydrogen Bond?

The hydrogen bond is an interaction involving a hydrogen atom

located between a pair of other atoms having a high affinity for

electrons such as nitrogen, oxygen or fluorine. It is an electrostatic

attraction between two polar groups. This occurs when a hydrogen

atom (H) is bound to a highly electronegative atom such as nitrogen

(N), oxygen (O), or fluorine (F), and experiences the electrostatic field

of another highly electronegative atom nearby.

The electrons constituting the covalent bond are shifted toward the

more electronegative atom. this happens when there is covalent bond

formation between highly electronegative elements and the hydrogen

atom. This leads to the development of a partial positive charge on the

hydrogen atom which helps in the bond formation with the

electronegative atoms of the other molecules. This is hydrogen bond

which is comparatively weaker than the covale

Browse more Topics Under Chemical Bonding And Molecular Structure

● Bond Parameters

● Covalent Compounds

● Fundamentals of Chemical Bonding

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Resonance Structures

● Valence Bond Theory

● VSEPR Theory

Mechanism of Hydrogen Bond Formation

The electrons carry with them a negative charge, so wherever the

electrons move they give the negative charge. This results in unequal

sharing of electrons. In a molecule, hydrogen bonds are formed, when

the hydrogen atom covalently linked to a highly electronegative atom

like oxygen, nitrogen, or fluorine, experiences the electrostatic field of

another highly electronegative atom of the nearby molecule.

One atom of the pair (the donor), generally a fluorine, nitrogen or

oxygen atom, is covalently bonded to a hydrogen atom (-FH, -NH or

–OH), whose electrons it shares unequally. Its high electron affinity

gives hydrogen a slight positive charge. The other atom of the pair,

typically F, N, or O has an unshared pair of the electron; hence it has

slight negative charge. Mainly through electrostatic attractions, the

donor atom shares its hydrogen with the acceptor atom hence forming

a hydrogen bond.

The small sizes of nitrogen, oxygen, and fluorine are essential to

hydrogen bonding because it makes those atoms electronegative that

their covalently bonded hydrogen is highly positive. Another reason is

that it allows the lone pair on the other oxygen, nitrogen, or fluorine to

come close to the hydrogen.

Hydrogen Bonding in HF Molecule

In HF molecule there is a hydrogen bond between the hydrogen atom

of one molecule and the fluorine atom of another molecule.

– – H+– F– – – – H+– F–– – – H+– F–– –

In this case, the hydrogen bond acts as a bridge between two atoms,

where one atom is held by a covalent bond and the other atom is held

by a hydrogen bond. In the structure above, the dotted line (– – –)

depicts the hydrogen bond and the solid line depicts the covalent

bond.

The shared pair of electrons move away from the hydrogen atom

toward the electronegative atom as the hydrogen atom is bonded to a

highly electronegative element. Hydrogen atom becomes

electropositive with respect to the electronegative element. This

results in the development of positive charge over hydrogen atom and

partial negative charge over the electronegative element.

This further leads to the formation of a polar molecule with an

electrostatic force of attraction. The magnitude of H-bond depends on

the physical state of the compounds. It reaches a maximum value in

solid state and minimum in a gaseous state.

● Intermolecular hydrogen bonding occurs between different

molecules of same or different compounds. Whereas

● Intramolecular hydrogen bonding occurs when hydrogen atom

lies in between the two electronegative elements present in the

same molecule.

Hydrogen Bonding in Water

Hydrogen bonds account for some important qualities of water. Even

though a hydrogen bond is only 5% as strong as a covalent bond, it’s

enough to stabilize water molecules.

● Hydrogen bonding causes water to remain liquid over a wide

temperature range.

● As it takes extra energy to break hydrogen bonds, water has an

unusually high heat of vaporization. Water has a much higher

boiling point than other hydrides.

There are many important consequences of the effects of hydrogen

bonding between water molecules:

● Hydrogen bonding makes ice less dense than liquid water, so

ice floats on water.

● The effect of hydrogen bonding on the heat of vaporization

helps make perspiration an effective means of lowering

temperature for animals.

● The effect on heat capacity means water protects against

extreme temperature shifts near large water bodies or humid

environments. Water helps regulate temperature on a global

scale.

Solved Examples for You

Question: On what parameters the energy of hydrogen bond depends?

Answer: The energy of hydrogen atom depends on the nature of donor

and acceptor atoms that is their geometry, bond, and environment. The

energy can be as high as 40kcal/mol

Ionic or Electrovalent Compounds

Do you know why some compounds show strong conductivity, while

some are pretty slow at it? If you were to melt such substances, you

will find that they have a sharp melting point. Why does it happen?

These substances are Ionic or Electrovalent compounds. In this

chapter, we will have a closer look at these ionic or electrovalent

compounds. Let’s begin.

Ionic or Electrovalent Bond

There are primarily three ways in which two atoms combine together

to lose energy and to become stable. One of the ways is by donating or

accepting electrons so as to complete their octet configuration. The

bond formed by this kind of combination is an ionic bond or

electrovalent bond.

An Ionic bond is the bond formed by the complete transfer of valence

electron so as to attain stability. This type of bonding leads to the

formation of two oppositely charged ions. These include the positive

ion, cations and negative ions, anions. The presence of two oppositely charged ions results

in strong attractive force between them. This force is the ionic or electrovalent bond.

Properties of an Ionic Bond

Due to the presence of a strong force of attraction between cations and

anions in ionic bonded molecules, we observe the following

properties:

● The ionic bonds are the strongest of all the bonds.

● The ionic bonds have a charge separation. So, they are the most

reactive of all the bonds in the proper medium.

● The ionic bonded molecules have high melting and boiling

point.

● The ionic bonded molecules in their aqueous solutions or in the

molten state are good conductors of electricity. This is due to

the presence of ions which acts as charge carriers.

Examples of Ionic Bonds

The following table shows the elements and the ions formed by them

when they lose or gain an e‑.

Element

Electronic config. Reaction Formed ion

Na(11) 2,8,1 Na → Na+ + e– …………………Reaction 1 Na+

Ca(20) 2,8,8,2 Ca → Ca2+ + 2e–……………….. Reaction 2 Ca2+

Cl(17) 2,8,7 Cl + e–→ Cl– ………………….…. Reaction 3 Cl–

O(8) 2,6 O + 2e–→ O2-……………………..Reaction 4 O2-

Now, when Na reacts with Cl, reaction 1 and reaction 3 will take

place. The resultant compound will be NaCl. When Na reacts with O,

reaction 1 and reaction 4 will take place. The resultant compound will

be Na2. In the case when Ca reacts with Cl, reaction 2 and reaction 3

will take place. The resultant compound will be CaCl2.

When Ca reacts with O, reaction 2 and reaction 4 will take place and

the resultant compound will be CaO. We, now have some information

about ionic or electrovalent bonds. Let us now look at what

electrovalent compounds are and what their characteristics are.

Electrovalent Compounds

The compounds which contain ionic or electrovalent bonds are

Electrovalent or Ionic Compounds. Mainly electrovalent compounds

are formed due to the reaction between highly electropositive and

highly electronegative atoms.

Characteristics of Electrovalent Compounds

● Crystal Structure: In the solid state of electrovalent compounds,

anions and cations are arranged in a regular manner. This is a

crystal in which anions are surrounded by a definite number of

cations and vice-versa.

● Physical Nature: Ionic or electrovalent compounds are

generally hard. Their hardness increases with increasing ionic

charge and the decreasing distance between ions.

● Solubility: Positive ion of ionic compound attaches to the

negative part of a polar solvent and negative ion of ionic

compound attach with the positive part of the polar solvent.

Therefore, ionic or electrovalent compounds are soluble in

polar solvents like water and insoluble in non-polar solvents

like benzene, ether, alcohol.

● Melting Point and Boiling Point: Electrovalent or ionic

compounds have high melting and boiling points because they

need a large amount of energy to break strong ionic bonds.

Solved Example for You

Q: The state in which HCl is a bad conductor of electricity is

_________ but, the state in which HCl is a good conductor of

electricity is __________ .

A) Solid, Anhydrous B) Aqueous, Solid C) Anhydrous, Solid

D) Anhydrous, Aqueous

Solution: D) Well any substance that can give rise to ions inside a

solution or has free charges within itself is a conductor. As anhydrous

HCl doesn’t have any free ions or charges, it will not conduct. While

on the other hand, Aqueous HCl has Hydrogen and Chloride ions

present in the solution, so it will conduct electricity.

Molecular Orbital Theory

We know that atoms bond. That results in the diversity of matter

around us. But what rules does the atomic or molecular bonding obey?

Are there any rules at all? How do you think the molecules are

arranged in an element? For that, we need to know the Molecular

Orbital Theory. Let us begin!

Molecular Orbital Theory

The Valence Bond Theory fails to answer certain questions like why

He2 molecule does not exist and why O2 is paramagnetic. Therefore in

1932 F. Hood and R.S. Mulliken came up with Molecular Orbital

Theory to explain questions like the ones above.

Learn VSEPR Theory to know the geometrical arrangement of various

molecules.

According to the Molecular Orbital Theory, individual atoms combine

to form molecular orbitals. Thus the electrons of an atom are present

in various atomic orbitals and are associated with several nuclei.

We know that we can consider electrons as either particle or wave

nature. Therefore, we can describe an electron in an atom as

occupying an atomic orbital, or by a wave function Ψ. These are

solutions to the Schrodinger wave equation. Electrons in a molecule

occupy molecular orbitals. We can obtain the wave function of a

molecular orbital by the following methods.

● Linear Combination of Atomic Orbitals (LCAO)

● United Atom Method

Linear Combination of Atomic Orbitals (LCAO)

As per this method, the formation of orbitals is because of Linear

Combination (addition or subtraction) of atomic orbitals which

combine to form the molecule. Consider two atoms A and B which

have atomic orbitals described by the wave functions ΨA and ΨB.

If the electron cloud of these two atoms overlaps, then we can obtain

the wave function for the molecule by a linear combination of the

atomic orbitals ΨA and ΨB. The below equation forms two molecular

orbitals.

ΨMO = ΨA + ΨB

Bonding Molecular Orbitals

When the addition of wave function takes place, the type of molecular

orbitals formed are Bonding Molecular Orbitals. We can represent

them by ΨMO = ΨA + ΨB. They have lower energy than atomic orbitals

involved.

Anti-Bonding Molecular Orbitals

When molecular orbital forms by the subtraction of wave function, the

type of molecular orbitals formed are antibonding Molecular Orbitals.

We can represent them as ΨMO = ΨA – ΨB. They have higher energy

than atomic orbitals. Therefore, the combination of two atomic

orbitals results in the formation of two molecular orbitals. They are

the bonding molecular orbital (BMO) and the anti-bonding molecular

orbital (ABMO).

Learn Polarity of Molecules and factors on which Polarity depends.

Relative Energies of Molecular Orbitals

● Bonding Molecular Orbitals (BMO) – Energy of Bonding

Molecular Orbitals is less than that of Anti Bonding Molecular

Orbitals. This is because of the increase in the attraction of both

the nuclei for both the electron (of the combining atom).

● Anti-Bonding Molecular Orbitals (ABMO) – Energy of Anti

Bonding Molecular Orbitals is higher than Bonding Molecular

Orbitals. This is because the electron tries to move away from

the nuclei and are in a repulsive state.

What happens to an atom and atomic orbital during bonding? Learn

Hybridisationto know.

Understand Schrodinger’s Wave Equation

Rules for Filling of Molecular Orbitals

We have to follow certain rules while filling up molecular orbitals

with electrons in order to write correct molecular configurations. They

are

● Aufbau Principle – This principle states that those molecular

orbitals which have the lowest energy are filled first.

● Pauli’s Exclusion Principle – According to this principle, each

molecular orbital can accommodate a maximum of two

electrons having opposite spins.

Read the General Properties of Covalent Compounds here.

Solved Examples for You

Question: Write the Hund’s Rule.

Answer: The Hund’s rule states that in two molecular orbitals of the

same energy, the pairing of electrons will occur when each orbital of

same energy consist of one electron.

Polarity of Bonds

What do you mean by Polarity of Bonds? Sounds difficult? Well, it is

not! Have you played the “tug-of-war” game? Similar to the game of

“tug-of-war”, in chemistry when two atoms share a pair of electrons,

they try to pull it towards themselves. This gives rise to the concept of

bond polarity. However, before we get into the details of the chapter,

let us first know what polarity is.

Polarity of Bonds

Polarity refers to the physical properties of compounds such as boiling

point, melting points and their solubilities. The polarity of bonds is

caused due to the interaction of the bonds between molecules and

atoms with different electronegativities.

Consider an electromotive force (EMF) or an electric potential, acting

between two points. Here, the points or poles have a greater number of

electrons than the other. The pole that has more electrons possesses a

negative polarity whereas the other end possesses a positive polarity.

Polarity in Chemistry is nothing but the concept of the separation of

an electric charge leading a molecule to have a positive and negative

end. Consider the below example:

In an H-F bond, the fluorine atom is more electronegative than that of

the Hydrogen atom. The electrons eventually spend more time at the

Fluorine atom. Hence, this F atom slightly becomes negative whereas

the Hydrogen atom tends to become slightly positive.

Browse more Topics under Chemical Bonding And Molecular Structure

● Bond Parameters

● Covalent Compounds

● Fundamentals of Chemical Bonding

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Resonance Structures

● Valence Bond Theory

● VSEPR Theory

Definition of Polarity

“A state or a condition of an atom or a molecule having positive and

also negative charges, especially in case of magnetic or an electrical

poles.”

Polarity Of Molecules

The bond or the molecular polarities are related to the

electronegativities of the atoms or the molecules. A molecule can

basically be either polar molecule, non-polar molecule or an ionic

molecule.

Polar Molecules

A polar molecule usually forms when the one end of the molecule is

said to possess a number of positive charges and whereas the opposite

end of the molecule has negative charges. Thus, they end up creating

an electrical pole. In a molecule having a polar bond, the centre of the

negative charge will be on one side. Whereas the centre of positive

charge will be on the different side. The entire molecule will be a

polar molecule.

Non- Polar Molecules

A molecule which does not have the charges present at the end due to

the reason that electrons are finely distributed and those which

symmetrically cancel out each other are the non- polar molecules. In a

solution, we cannot mix a polar molecule with the non-polar molecule.

For example, consider water and oil. In this solution, water is the polar

molecule. On the other hand, oil behaves as a non-polar molecule.

These two molecules do not form a solution. This is because they

cannot ever be mixed up.

Examples of Polar and Nonpolar Molecules

A molecule may be polar or Non-polar. A non-polar molecule has the

structure of its atoms lined up in a way that the orbital electrons in the

outer region cancel out the electronegativity. In general,

pyramid-shaped and V-shaped molecules are said to be polar.

Whereas the Linear molecules said to be non-polar in nature.

Water is said to be a polar molecule due to the difference in the

electronegativities between the oxygen atom and the hydrogen.

Oxygen is a highly electronegative atom when compared to hydrogen.

Fats, petrol, oil, gasoline are said to be non-polar molecules as they do

not dissolve in water and nonpolar are insoluble in water. Glucose is

another such example of a polar molecule. It is based on the

arrangement of the oxygen and hydrogen atoms in it.

Factors on which the Polarity of Bonds Depends

1) Relative Electronegativity of Participating Atoms

Since the bond polarity involves pulling of electrons towards itself,

hence a more electronegative element will be able to attract the

electrons more towards itself. As a result, the electrons will definitely

move towards the more electronegative element. The amount of their

shifting will depend upon the relative electronegativity of the

participating atoms.

2) The Spatial Arrangement of Various Bonds in the Atom

The shared pair of electrons also experience pulling force from the

other bonded and non-bonded pair of electrons. This results in

different bond polarity between same participating atoms that are

present in different molecules. For e.g. Bond Polarity of O-H bond in

a water molecule and acetic acid molecule is different. This is due to

the different spatial arrangement of various bonds in the molecule.

Solved Example for You

Q: The electronegativity of C,H,O,N and S are 2.5, 2.1, 3.5, 3.0 and

2.5 respectively. Which of the following bond is most polar?

A) O – H B) S – H C) N – H D) C – H

Solution: A) If the difference in the electronegativity between two or

more atoms is more, the bond between them is more polar. For the

given atoms, we can see that:

● O – H = 3.5 – 2.1 = 1.4

● S – H = 3.5 – 2.5 = 1

● N – H = 3.0 – 2.1 = 0.9

● C – H = 2.5 – 2.1 = 0.4 .

Therefore, the O-H bond is the most polar among the given bonds.

Resonance Structures

Do you know how to represent compounds through Lewis dot

method? Can you represent benzene that way? Oh, you can NOT!

Don’t worry! That is where the concept of resonance structures comes

into play. But, what is it? In this chapter, we will read more about

resonance structure and how we find that. But, before we proceed to

that, let us first look at what resonance effect is all about.

What is Resonance Effect?

We cannot predict the properties of many organic compounds with the

help of single Lewis dot structure. For example, let’s consider the case

of benzene. Going by the Lewis dot method, we would end up

predicting Benzene to have three C-C bonds and three C=C bonds.

But, the actual property deviates from this prediction.

Thus, we define resonance structures for defining properties of these

compounds. The resonance structures (canonical structures) are

actually hypothetical. This is because they do not represent any real

molecule individually. They contribute to the actual structure in

proportionately according to their stability.

The energy of actual structure of the molecule (the resonance hybrid)

is lower than that of any of the canonical structures.The resonance

energy increases with the number of important contributing structures.

The number of unpaired electrons is the same in the resonance

structures and so also are the positions of nuclei.

Browse more Topics under Chemical Bonding And Molecular Structure

● Bond Parameters

● Covalent Compounds

● Fundamentals of Chemical Bonding

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Valence Bond Theory

● VSEPR Theory

Stability

The stability of resonance increases with:

● Number of covalent bonds

● Number of atoms with an octet of electrons (except hydrogen

which has a duplex)

● Separation of opposite charges,

● Dispersal of charge

● A negative charge if any on a more electronegative atom, a

positive charge if any on the more electropositive atom,

increases the stability of the atom.

More on Resonance Effect

Resonance is the phenomenon which causes a polarity to be produced

in the molecule. This could happen either by the interaction of two

π-bonds or between a π-bond and lone pair of electrons present on an

adjacent atom. The delocalisation of π-electrons is what causes this

effect. We can classify the resonance effect into two main categories,

as described below.

1) Positive Resonance Effect (+R effect)

In the positive resonance effect, we notice that the transfer of electrons

takes place away from an atom or substituent group attached to the

conjugated system (presence of alternate single and double bonds in

an open-chain or cyclic system) due to resonance. Some of the

substituent groups which attribute to positive resonance effect are –

COOH, –CHO, >C=O, – CN, –NO2, etc.

2) Negative Resonance Effect (-R effect)

In this effect, we see that the transfer of electrons is towards the atom

or substituent group attached to the conjugated system (presence of

alternate single and double bonds in an open-chain or cyclic system)

due to resonance. Examples of the substituent groups that attribute to

negative resonance effect include – COOH, –CHO, >C=O, – CN,

–NO2, etc.

Let us now look at resonance structures in more specific details. We

will also look at some of the examples of the same.

Resonance Structures

From few experiments, it was observed that the bond parameters of

some molecules were not same as what was calculated on the basis of

different bond theories. Thus, to explain these differences, the theory

of resonance came into the picture which suggested that whenever a

single Lewis structure is insufficient to describe a molecule correctly,

then multiple Lewis structures can be superimposed over each other to

describe the molecule leading to hybrid structures with similar energy,

position of nuclei, bonding and non-bonding pairs of electrons.

Definition

Resonance structures are the multiple Lewis structures of similar

energy, the position of nuclei, bonding and the non-bonding pair of

electrons that can accurately describe a molecule. They are taken as

canonical structures of the hybrid molecules formed by the

superimposition of multiple Lewis structures. The hybrid molecule

alone can accurately describe the molecule.

An Indicative example – Resonance Structure of SO3

The above figure shows the different canonical structures of SO3.

They all are similar in energy, the position of nuclei, bonding and

non-bonding pairs. The three O-S bond have the same bond length.

This was actually not equivalent to the length of the double bond

between O and S. So, it requires resonance structures to describe it

correctly.

Solved Example for You

Q: Write a note on the characteristics of resonance.

Ans:

● Every structure is associated with a certain quantity of energy,

which determines the stability of a molecule or ion.

● A resonance hybrid is one particular structure that is an

intermediary structure between the contributing structures. The

total quantity of potential energy, however, is lesser than the

intermediate. This way, the molecule is a hybrid molecule.

● Resonance averages the bond characteristics as a whole.

● The canonical forms don’t exist in reality actually. They are

only discussed to make the study of molecules easier for us.

The superimposition of canonical forms leads to the formation

of hybrid molecules. This procedure accurately describes the

molecule.

● The concept of resonating structures or canonical structures

came into being so as to account for the different bond

parameters found in the molecule than suggested by their

Lewis structures.

Valence Bond Theory

Nothing is perfect! Haven’t you heard it too many times in your life?

Yeah, and it’s true! This belief applies to chemistry as well. If you

thought that the Lewis theory explained all about compounds and

molecules, you are wrong! It failed to explain many concepts and that

is why we have the Valence Bond Theory. Here, we will read more

about the valence bond theory and also look at its limitations. Yes,

even this theory isn’t perfect guys! Let’s learn why.

Why a Need for Valence Bond Theory Arose?

The theory given by Lewis explained the structure of molecules.

However, it failed to explain the chemical bond formation. Similarly,

VSEPR theory explained the shape of simple molecules. But, it’s

application was very limited. It also failed to explain the geometry of

complex molecules. Hence, scientists had to introduce the theory of

valence bonds to answer and overcome these limitations.

Valence Bond Theory

Heitler and London introduced this theory. This is primarily based on

the concepts of atomic orbitals, electronic configuration of elements,

the overlapping of atomic orbitals, hybridization of atomic orbitals.

The overlapping of atomic orbitals results in the formation of a

chemical bond. The electrons are localized in the bond region due to

overlapping.

Valence bond theory describes the electronic structure of molecules.

The theory says that electrons fill the atomic orbitals of an atom

within a molecule. It also states that the nucleus of one atom is

attracted to the electrons of another atom. Now, we move on and look

at the various postulates of the valence bond theory.

Postulates of Valence Bond Theory

● The overlapping of two half-filled valence orbitals of two

different atoms results in the formation of the covalent bond.

The overlapping causes the electron density between two

bonded atoms to increase. This gives the property of stability to

the molecule.

● In case the atomic orbitals possess more than one unpaired

electron, more than one bond can be formed and electrons

paired in the valence shell cannot take part in such a bond

formation.

● A covalent bond is directional. Such a bond is also parallel to

the region of overlapping atomic orbitals.

● Based on the pattern of overlapping, there are two types of

covalent bonds: sigma bond and a pi bond. The covalent bond

formed by sidewise overlapping of atomic orbitals is known as

pi bond whereas the bond formed by overlapping of atomic

orbital along the inter nucleus axis is known as a sigma bond.

Source: Quora

Limitations of Valence Bond Theory

As we pointed out earlier, nothing is perfect! In a similar way, the

Valence Bond theory is also not perfect. It has its own set of

limitations. They are:

● It fails to explain the tetravalency of carbon.

● This theory does not discuss the electrons’ energies.

● The assumptions are about the electrons being localized to

specific locations.

Solved Examples for You

Question: Based on the overlapping of orbitals, how many types of

covalent bonds are formed and what are they?

Answer: Based on the overlapping of orbitals, two types of covalent

bonds are formed. These are known as sigma(σ) and pi(π) bonds.

● Sigma bonds are formed by the end-to-end overlap of atomic

orbitals along the inter-nuclear axis known as a head-on or

axial overlap. End-on overlapping is of three types, they are s-s

overlapping, s-p overlapping and p-p overlapping.

● A pi bond is formed when atomic orbitals overlap in a specific

way that their axes remain parallel to each other and

perpendicular to the internuclear axis.

VSEPR Theory

We have already covered the Lewis structure and the Valence Bond

Theory. Did you get all the answers to your queries? No. We do not

know anything about the geometrical arrangement of the various

molecules. the VSEPR theory comes to our rescue! In this chapter, we

will know more about the arrangement of molecules by the VSEPR

theory.

What is the VSEPR Theory?

The Valence Shell Electron Pair Repulsion Model is often abbreviated

as VSEPR (pronounced “vesper”). It is basically a model to predict

the geometry of molecules. Specifically, VSEPR models look at the

bonding and molecular geometry of organic molecules and polyatomic

ions. It is useful for nearly all compounds that have a central atom that

is not a metal.

Browse more Topics under Chemical Bonding And Molecular Structure

● ond Parameters

● Covalent Compounds

● Fundamentals of Chemical Bonding

● Hybridisation

● Hydrogen Bonding

● Ionic or Electrovalent Compounds

● Molecular Orbital Theory

● Polarity of Bonds

● Resonance Structures

● Valence Bond Theory

Importance of VSEPR Models

● Lewis structures only tell the number and types of bonds

between atoms, as they are limited to two dimensions. The

VSEPR model predicts the 3-D shape of molecules and ions

but is ineffective in providing any specific information

regarding the bond length or the bond itself.

● VSEPR models are based on the concept that electrons around

a central atom will configure themselves to minimize repulsion,

and that dictates the geometry of the molecule.

● It can predict the shape of nearly all compounds that have a

central atom, as long as the central atom is not a metal. Each

shape has a name and an idealized bond angle associated with

it.

The following terms are commonly used in discussing the shapes of

molecules.

● Bond Angle: This is the angle between a bonded atom, the

central atom, and another bonded atom.

● Lone Pair: This refers to a pair of valence electrons that are not

shared with another atom.

● Molecular Geometry: This is the 3-D arrangement of bonded

atoms in a polyatomic ion or molecule.

● Electron Pair Geometry: This is the 3-D arrangement of

electron pairs around the central atom of a polyatomic ion or

molecule.

The main difference between molecular geometry and electron pair

geometry is that molecular geometry does not include unpaired

electrons, whereas electron pair geometry includes both bonded atoms

and unpaired electrons. If there are no unpaired electrons in the

compound being assessed, the molecular and electron pair geometries

will be the same.

Molecular Geometry

Steps to Using VSEPR

● Draw a Lewis structure for the ion or molecule in question.

● Determine the number of electron groups around the central

atom. Each lone pair of electrons counts as a single group.

Each bond counts as a single group, even if it is a double or

triple bond. Find the corresponding electron geometry from the

table.

● Determine the number of lone pairs and the number of bonding

pairs around the central atom, and use that to find the molecular

geometry.

VSEPR Notation

VSEPR notation gives a general formula for classifying chemical

species based on the number of electron pairs around a central atom.

Note, however, that not all species have the same molecular geometry.

For example, carbon dioxide and sulfur dioxide are both species, but

one is linear and the other is bent.

Sometimes, the notation is expanded to include lone pair electrons.

This can get confusing because water can be referred to as a species

depending on the conventions the author or text chooses. In general,

● A is used to represent the central atom.

● B or X is used to represent the number of atoms bonded to the

central atom.

● E represents the number of lone pairs on the central atom

(ignore lone pairs on bonded atoms).

Again, this theory is also not void of any limitations. We will now

discuss the common limitations of the VSEPR theory.

Limitations of the VSEPR theory

● The VSEPR model is not a theory. It does not explain or

attempt to explain any observations or predictions. Rather, it is

an algorithm that accurately predicts the structures of a large

number of compounds.

● VSEPR is simple and useful but does not work for all chemical

species.

● First, the idealized bond angles do not always match the

measured values. For example, VSEPR predicts that and will

have the same bond angles, but structural studies have shown

the bonds in the two molecules are different by 12 degrees.

● VSEPR also predicts that group-2 halides such as will be linear

when they are actually bent. Quantum mechanics and atomic

orbitals can give more sophisticated predictions when VSEPR

is inadequate.

Solved Example for You

Q: On the basis of VSEPR theory explain the structure of NH3

molecule.

Ans: In ammonia, N is the central atom. Nitrogen is a group 15

element and therefore has 5 electrons in its outmost shell. Three

electrons of N are bonded with hydrogen and the rest two which do

not take part in bonding form the lone pair. The outer shell then has a

share in eight electrons, that is, three pairs bonded and one lone pair.

These four pairs of electrons give rise to a tetrahedral structure where

three positions are occupied by H atoms and fourth position by the

lone pair. This shape may either be described as tetrahedral or

pyramidal. The presence of lone pair causes slight distortion from

109°28’ to 107°48’.