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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’.