Download Study notes on Chemical Bonding and more Study notes Chemistry in PDF only on Docsity!
Atoms of different elements excepting noble gases donot have
complete octet so they combine with other atoms to form chemical bond.
The force which holds the atoms or ions together within the molecule is
called a chemical bond and the process of their combination is called
Chemical Bonding. It depends on the valency of atoms.
Cause and Modes of chemical combination
Chemical bonding takes place due toacquire a state of minimum
energy and maximum stability and toconvert atoms into molecule to
acquire stable configuration of the nearest noble gas. We divide atoms into
three classes,
(1) Electropositive elements which give up one or more electrons easily. They have low ionisation potentials.
(2) Electronegative elements, which can gain electrons. They have higher value of electronegativity.
(3) Elements which have little tendency to lose or gain electrons. Different types of bonds are formed from these types of atoms. Atoms involved Type
A +B Electrovalent
B +B Covalent
A +A Metallic
Electrons deficient molecule or ion (Lewis acid) and electrons rich molecule or ion (Lewis base)
Coordinate
H and electronegative element (F,
N,O)
Hydrogen
Electrovalent bond
An electrovalent bond is formed when a metal atom transfers one or more electrons to a non-metal atom.
Na Cl Na Cl or NaCl
Some other examples are:MgCl 2 ,CaCl 2 ,MgO,Na 2 S,CaH 2 ,AlF 3 ,NaH,
KH, K 2 O ,KI,RbCl,NaBr, CaH 2 etc.
(1) Conditions for formation of electrovalent bond (i) The atom which changes into cation (+ ive ion) should possess 1,
2 or 3 valency electrons. The other atom which changes into anion (– ve
ion) should possess 5, 6 or 7 electrons in the valency shell. (ii) A high difference of electronegativity (about 2) of the two atoms
is necessary for the formation of an electrovalent bond.Electrovalent bond
is not possible between similar atoms.
(iii)There must be overall decrease in energy i.e., energy must be
released. For this an atom should have low value of Ionisation potential and
the other atom should have high value of electron affinity.
(iv)Higher the lattice energy, greater will be the case of forming an ionic
compound. The amount of energy released when free ions combine together to
form one mole of a crystal is called lattice energy (U). Lattice energy
r r
K
r ^ r is internuclear distance. The energy changes involved in the formation of ionic compounds from their constituent elements can be studied with the help of a thermochemical cycle called Born Haber cycle.
According to Hess's law of constant heat summation, heat of formation of an ionic solid is net resultant of the above changes.
H (^) f H Subl. H diss. IE EA U 2
Na ( s ) () 2
Cl 2 g
Na ( g )
Na ( g ) Cl ( g )
H sub
1/2 Hdiss. Cl ( g)
H f
Na+ Cl– (s)
– U
(Born Haber Cycle)
(Lattice energy)
Chemical Bonding
Chapter
(2) Characteristics of electrovalent compounds (i) Electrovalent compounds are generally crystalline is nature. The constituent ions are arranged in a regular way in their lattice.
(ii) Electrovalent compounds possess high melting and boiling points. Order of melting and boiling points in halides of sodium and oxides of IInd group elements is as,
NaF NaCl NaBr NaI , MgO CaO BaO (iii) Electrovalent compounds are hard and brittle in nature. (iv) Electrovalent solids do not conduct electricity. While electrovalent compounds in the molten state or in solution conduct electricity.
(v) Electrovalent compounds are fairly soluble in polar solvents and insoluble in non-polar solvents.
(vi) The electrovalent bonds are non-rigid and non-directional. Thus these compound do not show space isomerism e.g. geometrical or optical isomerism.
(vii) Electrovalent compounds furnish ions in solution.The chemical
reaction of these compounds are known as ionic reactions, which are fast.
3 (Precipitate)
KCl AgNO 3 AgCl KNO
(viii) Electrovalent compounds show isomorphism. (ix) Cooling curve of an ionic compound is not smooth, it has two break points corresponding to time of solidification.
(x) Ionic compounds show variable electrovalency due to unstability of core and inert pair effect.
Covalent bond
Covalent bond was first proposed by Lewis in 1916. The bond
formed between the two atoms by mutual sharing of electrons so as to
complete their octets or duplets (in case of elements having only one shell)
is called covalent bond or covalent linkage. A covalent bond between two
similar atoms is non-polar covalent bond while it is polar between two different atom having different electronegativities. Covalent bond may be single, double or a triple bond. We explain covalent bond formation by Lewis octet rule.
Chlorine atom has seven electrons in the valency shell. In the formation of chlorine molecule, each chlorine atom contributes one electron and the pair of electrons is shared between two atoms. both the atoms acquire stable configuration of argon.
( 2 , 8 , 8 ) ( 2 , 8 , 8 )
( 2 , 8 , 7 )
**
**
( 2 , 8 , 7 )
Cl ^ Cl Cl Cl^ or Cl^ Cl
Some other examples are : H 2 S , NH 3 , HCN , PCl 3 , PH 3 , C 2 H 2 , H 2 , C 2 H 4 , SnCl 4 , FeCl 3 , BH 3 , graphite, BeCl 2 etc. (1) Conditions for formation of covalent bond (i) The combining atoms should be short by 1, 2 or 3 electrons in the valency shell in comparison to stable noble gas configuration.
(ii) Electronegativity difference between the two atoms should be zero or very small.
(iii) The approach of the atoms towards one another should be accompanied by decrease of energy.
(2) Characteristics of covalent compounds (i) These exist as gases or liquids under the normal conditions of temperature and pressure. Some covalent compounds exist as soft solids.
(ii) Diamond, Carborandum (SiC), Silica (SiO 2 ),AlN etc. have giant
three dimensional network structures; therefore have exceptionally high melting points otherwise these compounds have relatively low melting and boiling points. (iii) In general covalent substances are bad conductor of electricity.
Polar covalent compounds likeHCl in solution conduct electricity. Graphite
can conduct electricity in solid state since electrons can pass from one layer to the other. (iv) These compounds are generally insoluble in polar solvent like water but soluble in non-polar solvents like benzene etc. some covalent compounds like alcohol, dissolve in water due to hydrogen bonding. (v) The covalent bond is rigid and directional. These compounds, thus show isomerism (structural and space). (vi) Covalent substances show molecular reactions. The reaction rates are usually low.
(vii)The number of electrons contributed by an atom of the element
for sharing with other atoms is called covalency of the element. Covalency
= 8 – [Number of the group to which element belongs].The variable
covalency of an element is equal to the total number of unpaired electrons
in s,p and d-orbitals of its valency shell.
The element such asP,S,Cl,Br,I have vacantd-orbitals in their
valency shell. These elements show variable covalency by increasing the number of unpaired electrons under excited conditions. The electrons from
paired orbitals get excited to vacantd-orbitals of the same shell.
Four elements, H, N, O and F do not possess d-orbitals in their
valency shell. Thus, such an excitation is not possible and variable valency is
not shown by these elements. This is reason that NCl 3 exists while NCl 5 does
not.
(3) The Lewis theory : The tendency of atoms to achieve eight electrons in their outermost shell is known as lewis octet rule. Lewis symbol for the representative elements are given in the following table, 1 2 13 14 15 16 17 Group IA IIA IIIA IVA VA VIA VIIA Lewis symbol
X X
X
X
X
X
X
(4) Failure of octet rule : There are several stable molecules known
in which the octet rule is violatedi.e., atoms in these molecules have
number of electrons in the valency shell either short of octet or more than octet. BeF 2 , BF 3 , AlH 3 are electron- deficients (Octet incomplete) hence are Lewis acid. In PCl (^) 5 , P has 10 electrons in valency shell while in SF (^) 6 , S has 12 electrons in valence shell. Sugden introduced singlet linkage in which one electron is donated (Instead of one pair of electrons) to the electron deficient atom so that octet rule is not violated. This singlet is represented as (⇁). Thus, PCl 5 and SF 6 have structures as,
(5) Construction of structures for molecules and poly atomic ions : The following method is applicable to species in which the octet rule is not violated. (i) Determine the total number of valence electrons in all the atoms
present, including the net charge on the species (n 1 ).
(ii) Determinen 2 = [2 × (number ofH atoms) + 8 × (number of
other atoms)].
P
Cl Cl
Cl Cl
Cl
S
F F
F F
F F
Square planar See saw
Zero Non zero
XeF 4 SF 4 , TeCl 4 AX 5 Trigonal bipyramidal Square pyramidal
Zero Non zero
PCl 5 BrCl 5 AX 6 Octahedral Distorted octahedral
Zero Non zero
SF 6
XeF 6
AX 7 Pentagonal bipyramidal^ Zero IF 7 (2) Every ionic compound having some percentage of covalent character according to Fajan's rule. The percentage of ionic character in compound having some covalent character can be calculated by the following equation.
The % ionic character 100 Theoretical
Observed
(3) The trans isomer usually possesses either zero dipole moment or
very low value in comparison tocis–form
H C Cl
H C Cl
Cl C H
H C Cl
Fajan’s rule The magnitude of polarization or increased covalent character depends upon a number of factors. These factors are,
(1) Small size of cation : Smaller size of cation greater is its
polarizing power i.e. greater will be the covalent nature of the bond.
(2) Large size of anion :Larger the size of anion greater is its
polarizing power i.e. greater will be the covalent nature of the bond.
(3) Large charge on either of the two ions : As the charge on the
ion increases, the electrostatic attraction of the cation for the outer
electrons of the anion also increases with the result its ability for forming
the covalent bond increases.
(4) Electronic configuration of the cation : For the two ions of the same size and charge, one with a pseudo noble gas configuration (i.e. 18 electrons in the outermost shell) will be more polarizing than a cation with noble gas configuration (i.e., 8 electron in outer most shell).
Valence bond theory or VBT
It was developed by Heitler and London in 1927 and modified by Pauling and Slater in 1931.
(1) To form a covalent bond, two atoms must come close to each other so that orbitals of one overlaps with the other.
(2) Orbitals having unpaired electrons of anti spin overlaps with each other.
(3) After overlapping a new localized bond orbital is formed which has maximum probability of finding electrons.
(4) Covalent bond is formed due to electrostatic attraction between radii and the accumulated electrons cloud and by attraction between spins of anti spin electrons.
(5)Greater is the overlapping, lesser will be the bond length, more
will be attraction and more will be bond energy and the stability of bond
will also be high.
(6) The extent of overlapping depends upon: Nature of orbitals involved in overlapping, and nature of overlapping.
(7) More closer the valence shells are to the nucleus, more will be the overlapping and the bond energy will also be high.
(8) Between two sub shells of same energy level, the sub shell more directionally concentrated shows more overlapping. Bond energy :
2 s 2 s < 2 s 2 p < 2 p 2 p
(9) s -orbitals are spherically symmetrical and thus show only head
on overlapping. On the other hand, p -orbitals are directionally
concentrated and thus show either head on overlapping or lateral
overlapping.Overlapping of different type gives sigma ( ) and pi ( ) bond.
Sigma () bond Pi () bond
It results from the end to end
overlapping of two s-orbitals or
two p-orbitals or one s and one p-
orbital.
It result from the sidewise (lateral)
overlapping of two p-orbitals.
Stronger Less strong
Bond energy 80 kcals Bond energy 65 kcals
More stable Less stable Less reactive More reactive Can exist independently (^) Always exist along with a -bond The electron cloud is symmetrical about the internuclear axis.
The electron cloud is above and below the plane of internuclear axis.
Hybridization The concept of hybridization was introduced by Pauling and Slater. Hybridization is defined as the intermixing of dissimilar orbitals of the same atom but having slightly different energies to form same number of new orbitals of equal energies and identical shapes. The new orbitals so formed are known as hybrid orbitals. Characteristics of hybridization (1) Only orbitals of almost similar energies and belonging to the same atom or ion undergoes hybridization. (2) Hybridization takes place only in orbitals, electrons are not involved in it. (3) The number of hybrid orbitals produced is equal to the number of pure orbitals, mixed during hybridization. (4) In the excited state, the number of unpaired electrons must correspond to the oxidation state of the central atom in the molecule. (5) Both half filled orbitals or fully filled orbitals of equivalent energy can involve in hybridization. (6) Hybrid orbitals form only sigma bonds. (7) Orbitals involved in bond formation do not participate in hybridization. (8) Hybridization never takes place in an isolated atom but it occurs only at the time of bond formation. (9) The hybrid orbitals are distributed in space as apart as possible resulting in a definite geometry of molecule. (10) Hybridized orbitals provide efficient overlapping than overlapping by pure s, p and d-orbitals. (11) Hybridized orbitals possess lower energy.
How to determine type of hybridization : The structure of any
molecule can be predicted on the basis of hybridization which in turn can be known by the following general formulation,
H V M C A
WhereH = Number of orbitals involved in hybridization viz. 2, 3, 4,
5, 6 and 7, hence nature of hybridization will besp,sp^2 ,sp^3 ,sp^3 d,sp^3 d^2 ,sp^3 d^3
respectively.
V = Number electrons in valence shell of the central atom,
M = Number of monovalent atom
C = Charge on cation,
A = Charge on anion
Resonance
The phenomenon of resonance was put forward by Heisenberg to explain the properties of certain molecules.
In case of certain molecules, a single Lewis structure cannot explain all the properties of the molecule. The molecule is then supposed to have many structures, each of which can explain most of the properties of the molecule but none can explain all the properties of the molecule. The actual structure is in between of all these contributing structures and is called resonance hybrid and the different individual structures are called resonating structures or canonical forms. This phenomenon is called resonance.
To illustrate this, consider a molecule of ozone O 3. Its structure
can be written as
( a ) ( b ) ( c )
O O
O
O O
O
O O
O
As a resonance hybrid of above two structures (a) and (b. For
simplicity, ozone may be represented by structure (c), which shows the
resonance hybrid having equal bonds between single and double.
Resonance is shown by benzene, toluene, O 3 , allenes (>C = C =
C<), CO, CO 2 , CO 3 , SO 3 , NO, NO 2 while it is not shown by H 2 O 2 , H 2 O, NH 3 ,
CH 4 , SiO 2.
As a result of resonance, the bond lengths of single and double bond
in a molecule become equal e.g.O– O bond lengths in ozone orC– O bond
lengths in CO 32 – ion.
The resonance hybrid has lower energy and hence greater stability
than any of the contributing structures.
Greater is the number of canonical forms especially with nearly
same energy, greater is the stability of the molecule.
Difference between the energy of resonance hybrid and that of the most stable of the resonating structures (having least energy) is called resonance energy. Thus,
Resonance energy = Energy of resonance hybrid – Energy of the
most stable of resonating structure.
In the case of molecules or ions having resonance, the bond order changes and is calculated as follows,
Totalno.ofresonatingstructures
Totalno.of bondsbetweentwoatomsinall thestructures Bond order
In benzene
doublebond inglebond Bond order
s
In carbonate ion
O O
C
O
O O
C
O
O O
C
O
/ \ \
/ \
// \
Bond order
Bond characteristics (1) Bond length “The average distance between the centre of the nuclei of the two bonded atoms is called bond length”.
It is expressed in terms of Angstrom (1 Å = 10 ^10 m) or picometer
(1pm = 10 ^12 m).
In an ionic compound, the bond length is the sum of their ionic radii ( d r r ) and in a covalent compound, it is the sum of their
covalent radii (e.g., forHCl, d rH rCl ).
Factors affecting bond length (i) The bond length increases with increase in the size of the atoms.
For example, bond length of H X are in the order,
HI HBr HCl HF.
(ii) The bond length decreases with the multiplicity of the bond. Thus, bond length of carbon–carbon bonds are in the order,
C C C C C – C.
(iii) As an s-orbital is smaller in size, greater the s-character shorter is the hybrid orbital and hence shorter is the bond length.
For example, sp^3 C – H sp^2 C – H spC – H
(iv) Polar bond length is usually smaller than the theoretical non- polar bond length. (2) Bond energy “The amount of energy required to break one mole of bonds of a particular type so as to separate them into gaseous atoms is called bond dissociation energy or simply bond energy”. Greater is the bond energy, stronger is the bond. Bond energy is usually expressed in kJ mol –^1.
Factors affecting bond energy (i) Greater the size of the atom, greater is the bond length and less is the bond dissociation energy i.e. less is the bond strength. (ii) For the bond between the two similar atoms, greater is the multiplicity of the bond, greater is the bond dissociation energy. (iii) Greater the number of lone pairs of electrons present on the bonded atoms, greater is the repulsion between the atoms and hence less is the bond dissociation energy. (iv) The bond energy increases as the hybrid orbitals have greater
amount ofs orbital contribution. Thus,bond energy decreases in the
following order, sp sp^2 sp^3
(v)Greater the electronegativity difference, greater is the bond
polarity and hence greater will be the bond strength i.e., bond energy,
H F H Cl H Br H I ,
(5) The number of molecular orbitals formed is equal to the number of combining atomic orbitals.
(6) When two atomic orbitals combine, they form two new orbitals called bonding molecular orbital and antibonding molecular orbital.
(7) The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital.
(8) The bonding molecular orbitals are represented by , etc,
whereas the corresponding antibonding molecular orbitals are represented
by ^ , etc.
(9) The shapes of the molecular orbitals formed depend upon the type of combining atomic orbitals.
(10) The filling of molecular orbitals in a molecule takes place in accordance with Aufbau principle, Pauli's exclusion principle and Hund's rule. The general order of increasing energy among the molecular orbitals formed by the elements of second period and hydrogen and their general electronic configurations are given below.
(11) Electrons are filled in the increasing energy of theMO which is
in order
(a) 1 s ,* 1 s , 2 s ,* 2 s , 2 p x , 2 py * 2 p y , * 2 px
2 pz , * 2 pz
Increasing energy (for electrons > 14)
(b) 1 s ,^1 s , 2 s , 2 s , 2 py , 2 px ,
2 pz
Increasing energy (for electrons 14) (12) Number of bonds between two atoms is called bond order and is given by
Bond order B^ A
N N
where NB number of electrons in bondingMO.
NA number of electrons in antibondingMO.
For a stable molecule/ion, N (^) B NA
(13)Bond order Stability of molecule Dissociation energy
Bondlength
(14) If all the electrons in a molecule are paired then the substance is a diamagnetic on the other hand if there are unpaired electrons in the molecule, then the substance is paramagnetic. More the number of unpaired electron in the molecule greater is the paramagnetism of the substance.
Hydrogen bonding
In 1920, Latimer and Rodebush introduced the idea of “hydrogen bond”.
For the formation ofH-bonding the molecule should contain an
atom of high electronegativity such asF,O orN bonded to hydrogen atom
and the size of the electronegative atom should be quite small.
Types of hydrogen bonding (1) Intermolecular hydrogen bond : Intermolecular hydrogen bond is formed between two different molecules of the same or different substances.
(i) Hydrogen bond between the molecules of hydrogen fluoride.
(ii) Hydrogen bond in alcohol and water molecules
(2) Intramolecular hydrogen bond (Chelation) Intramolecular hydrogen bond is formed between the hydrogen atom
and the highly electronegative atom(F,O or N) present in the same
molecule. Intramolecular hydrogen bond results in the cyclisation of the molecules and prevents their association. Consequently, the effect of intramolecular hydrogen bond on the physical properties is negligible. For example : Intramolecular hydrogen bonds are present in
molecules such aso-nitrophenol,o-nitrobenzoic acid, etc.
O
H
O
N
O
Ortho nitrophenol
O
C
O
O
N
O
Ortho nitrobenzoic acid
H
H | C O H O
Salicyldehyde (o-hydroxy benzaldehyde)
*(2 py)
*(2 pz)
*(2 px)
(2 py)^ ^ (2^ px)
(2 pz)
2 p 2 p
*(2 s)
(2 s)
2 s 2 s
*(1 s)
(1 s)
1 s 1 s
Atomic orbitals
Molecular orbitals
Atomic orbitals
Molecular orbital energy level diagram (Applicable for elements with Z > 7)
Increasing energy
*(2 py)
*(2 pz)
*(2 px)
(2 pz)
(2 py)
2 p 2 p
*(2 s)
(2 s)
2 s (^2) s
*(1 s)
(1 s)
1 s (^1) s
Atomic orbitals
Molecular orbitals
Atomic orbitals
Molecular orbital energy level diagram obtained by the overlap of 2 s and 2 pz atomic orbitals after mixing (Applicable for elements with Z < 7)
Increasing energy
(2 px)
The extent of both intramolecular and intermolecular hydrogen
bonding depends on temperature.
Effects of hydrogen bonding
Hydrogen bond helps in explaining the abnormal physical properties in several cases. Some of the properties affected by H-bond are given below,
(1) Dissociation : In aqueous solution, hydrogen fluoride dissociates
and gives the difluoride ion ( HF 2 )instead of fluoride ion ( F ). This is
due toH-bonding inHF. This explains the existence of KHF 2 .H-bond
formed is usually longer than the covalent bond present in the molecule
(e.g. inH 2 O, O–H bond= 0.99Å butH-bond= 1.77Å).
(2) Association : The molecules of carboxylic acids exist as dimers because of the hydrogen bonding. The molecular masses of such compounds are found to be double than those calculated from their simple formulae. For example, molecular mass of acetic acid is found to be 120.
(3) High melting and boiling point : The compounds having hydrogen bonding show abnormally high melting and boiling points.
The high melting points and boiling points of the compounds ( H (^) 2 O , HF and NH 3 )containing hydrogen bonds is due to the fact that
some extra energy is needed to break these bonds.
(4) Solubility : The compound which can form hydrogen bonds with
the covalent molecules are soluble in such solvents. For example,lower
alcohols are soluble in water because of the hydrogen bonding which can
take place between water and alcohol molecules as shown below,
2 5 2 5
CH H C H
H O H O H O
The intermolecular hydrogen bonding increases solubility of the compound in water while, the intramolecular hydrogen bonding decreases.
(5) As the compounds involving hydrogen bonding between different molecules (intermolecular hydrogen bonding) have higher boiling points, so they are less volatile.
(6) The substances which contain hydrogen bonding have higher viscosity and high surface tension.
(7) Explanation of lower density of ice than water and maximum
density of water at 277K : In case ofsolid ice, the hydrogen bonding gives
rise to a cage like structure of water molecules as shown in following figure.
As a matter of fact,each water molecule is linked tetrahedrally to four other
water molecules.Due to this structure ice has lower density than water at
273 K. That is why ice floats on water. On heating, the hydrogen bonds
start collapsing, obviously the molecules are not so closely packed as they are in the liquid state and thus the molecules start coming together resulting in the decrease of volume and hence increase of density. This goes
on upto 277K.After 277K, the increase in volume due to expansion of the
liquid water becomes much more than the decrease in volume due to
breaking of H-bonds. Thus, after 277 K , there is net increase of
volume on heating which means decrease in density. Hence density of water
is maximum 277 K.
O N = O
Due to chelation, – OH group is not available to form hydrogen bond with water hence it is sparingly soluble in water.
O
H
O
N
O
o- Nitrophenol
- OH group available to form hydrogen bond with water, hence it is completely soluble in water.
H – O …… H – O – H
p- Nitrophenol
H
H
H
H
O
H
H
O
H
H
H^ H
H
O
H
O
H
O
H
H
O
H
H
O
H
O
O
0.90 Å (99 pm) 1.77 Å (177 pm)
Vacant spaces
Cage like structure of H 2 O in the ice
