Structure and bonding concept


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Introduction to structures and models of bonding:

Introduction to structures and models of bonding Introduction to structures and models of bonding 1 Mr.Dhiraj S.Nikam ( Asst.Prof .) Bharati vidyapeeth’s College of Pharmacy , Navi Mumbai

Niels Bohr’s Model (1913):

Niels Bohr’s Model (1913) Electrons orbit the nucleus in circular paths of fixed energy ( energy levels). 2

Niel Bohr’s:

Niel Bohr’s Electrons can jump from one energy level to another energy level. Electrons absorb or emit light energy when they jump from one energy level to another. 3


A quantum of energy is the amount of energy required to move an electron from one energy level to another. Excited State and Ground State: Ground state: the lowest possible energy level an electron be at. Excited state: an energy level higher than the ground state. Quantum 4

Quantum Mechanics overview::

Quantum Mechanics overview: Electrons have discrete energies, not because they are in shells but because they can only have certain wavelengths Line spectra are not due to electrons jumping from shell to shell (as in Bohr’s model)… Instead they’re due to electrons transforming from one wavelength (waveform) to another Each electron is a wave that can be described by a series of “quantum numbers” There are four quantum numbers: n, l , m l , m s 5

Quantum Numbers: 4 values used to identify or index quantized properties of bound electrons.:

Quantum Numbers: 4 values used to identify or index quantized properties of bound electrons. Principal Quantum Number Orbital Quantum Number Magnetic Quantum Number Spin Quantum Number 6

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What are Quantum Numbers? Quantum number are a set of four values that define the energy state of an electron in an atom. Quantum number values are designated as n, l, m and s (s is often written as m s ) “n” is called the principal quantum number and ranges from 1, 2, 3, etc. (also refers to the energy level or shell “l” represents the orbital type and depends on n. It ranges from 0 through n – 1. It often called the azimuthal quantum number “m” depends on l. It ranges from – l thru 0 to + l. It defines the orbital orientation in space and is call the magnetic quantum number . “S” is the spin number and is either + ½ or – ½ 7

Principal Quantum Number-n:

Principal Quantum Number-n Principal shell in which electrons belong. Referred as major energy level. Describes size of electron cloud. Has positive integer value n-1,2,3…..Infinity. n = 1, 2, 3, 4, ...... “ shells ” (n = K, L, M, N, ......) Electron with n=1 has lowest energy bound firmly to nucleus. Max # electrons = 2n 2 1 2 3 8

Azimuthal Quantum number: l:

Azimuthal Quantum number: l Def.: Spatial distribution of electron cloud about the nucleus and describes angular momentum of electron. Defines shape of orbital. Has all Integer value 0 to n-1 ,each value refers to energy sublevel or sub-shell. Total level of such possible sublevels in each principal level equals to principal quantum number. n=1: l= 0 (1s) n=2: l= 0(2s), l=1 (2p) n= 3, l= 0 (3s),l=1(3p),l=2(3d) 9

l : azimuthal quantum number:

l : azimutha l quantum number If n can be thought of as shells, l can be thought of as “subshells” dividing each shell into subsections … (l = 0  n - 1) n = 1 l = 0 (s) n = 2 l = 0 (s) l = 1 (p) n = 3 l = 0 (s) l = 1 (p) l = 2 (d) 10

Magnetic quantum number: m:

Magnetic quantum number: m Based on the observation (Zeeman, 1897) that single lines on line spectra split into new lines near a strong magnet. It is also called as orientation quantum number, it gives orientation and distribution of electron cloud. It has positive and negative integral values from +l to –l through 0. Formula: (2l+1) For l=1 m= +1,0,-1 designated as Px,Py,Pz . 11

The Spin Quantum Number, ms (Pauli, 1925):

The Spin Quantum Number, m s (Pauli, 1925) Based on the observation that magnets could further split lines in line spectra, and that some elements exhibit paramagnetism ‘ m s ’ relates to the ‘spin’ of an electron A llowable values: m s = - ½ or + ½ i.e. for any possible set of n , l , and m l values, there are two possible m s values when two electrons of opposite spin are paired, there is no magnetism observed; an unparied electron is weakly magnetic 12

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Spin quantum number effects: Every orbital can hold up to two electrons. Consequence of the Pauli Exclusion Principle. The two electrons are designated as having one spin up  and one spin down   Spin describes the direction of the electron’s magnetic fields. 13

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Atomic orbital::

Atomic orbital: The Electron Cloud: The electron cloud represents positions where there is probability of finding an electron. The higher the electron density, the higher the probability that an electron may be found in that region. 15

Shape of Atomic Orbitals::

Shape of Atomic Orbitals: Orbital Depictions a ) Electron density plot b ) Orbital density picture c ) An electron contour drawing Each atomic energy level can be associated with a specific three-dimensional atomic orbital 16

Shape of Atomic Orbitals::

Shape of Atomic Orbitals: Orbital size- Orbital becomes larger as value of n increases. In any particular atom, all orbitals with the same principle quantum number are similar in size 17

Orbital shapes::

Orbital shapes: Atomic orbital shapes are surfaces that surround 90% of the total probability of where its electrons are Look at l = 0, the s-orbitals Basic shape of an s-orbital is spherical centered on the nucleus Basic shape is same for same l values Nodes = areas of zero probability Number of nodes changes for larger n 18

Orbital shapes::

Orbital shapes: p-orbitals There are no 1p orbitals (n = 1, l = 0 only) 2p orbitals (n = 2, l = 1) have 2 lobes with a node at the nucleus There are three different p-orbitals ( l = 1, m l = -1, 0, 1) 2p x lies along the x-axis 2p y lies along the y-axis 2p z lies along the z-axis All three 2p orbitals have the same energy = degenerate 3p, 4p, 5p, etc… have the same shape and number 19

Orbital shapes::

Orbital shapes: d-orbitals There are no 1d or 2d orbitals (d needs l = 2, so n = 3) 5 degenerate d-orbitals (m l = -2, -1, 0, 1, 2) 4 of the d-orbitals have 4 lob e s which lie in planes on or between the xyz axes: 3d xy , 3d xz , 3d yz , 3d x 2 -y 2 1 is composed of 2 lobes and a torus-shaped area: 3d z 2 The 4d orbitals etc…are the same shape, only larger 20

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f-orbitals n = 4, l = 3, m l = -3, -2, -1, 0, 1, 2, 3 7 f-orbitals in the fourth shell are degenerate The f-orbital are only used for the lanthanides and actinides and are complex shapes. We won’t use them. 21

Electron Configurations: :

Electron Configurations: Important principles: aufbau principle : (building up): Lower energy orbitals are populated first.Only 2 electrons can be in each orbital. When electrons are in same orbital they must be spin paired.(Pauli principle). 5 6 7 8 22


Hund’s rule tells us that the electrons will fill the p orbitals by placing electrons in each orbital singly and with same spin until half-filled. Then the electrons will pair to finish the p orbitals. 23

Hybridization Theory:

Hybridization Theory In biology, the word "hybrid" usually refers to a genetic mixing. Hybridization: process of combining two or more atomic orbitals to create new orbitals(hybrids) that explain the geometry of the compound. When determining the types of hybrid orbitals that an atom has, you must to first determine the number of groups around that atom much like VSEPR theory For basic hybridization theory, atoms will generally be surrounded by two, three, or four groups of electrons. 24

Valence bond theory::

Valence bond theory: The theory we will describes the placement of electrons into bonding orbitals located around the individual atoms from which they originated. Central atoms mix or hybridize their valence atomic orbitals to form new bonding orbitals called hybrid orbitals. 25

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Orbital shapes, Individual (“isolated”) Atoms 26

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COVALENT BOND FORMATION (VB THEORY) In order for a covalent bond to form between two atoms, overlap must occur between the orbitals containing the valence electrons. The best overlap occurs when two orbitals are allowed to meet “head on” in a straight line. When this occurs, the atomic orbitals merge to form a single bonding orbital and a “single bond” is formed, called a sigma ( ) bond . 28

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In order for “best overlap” to occur, valence electrons need to be re-oriented and electron clouds reshaped to allow optimum contact . To form as many bonds as possible from the available valence electrons, sometimes separation of electron must also occur. We describe the transformation process as “ orbital hybridization ” and we focus on the central atom in the species... 30

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s p -Hybridization of Be in BeCl 2 Atomic Be: 1s 2 2s 2 Valence e’s Hybrid sp orbitals: 1 part s , 1 part p 31

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Sp2 -Hybridization of B in BF 3 Atomic B : 1s 2 2s 2 2p 1 Valence e’s Hybrid sp 2 orbitals: 1 part s , 2 parts p 33

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Sp3- Hybridization of C in C H 4 Atomic C : 1s 2 2s 2 2p 2 Valence e’s Hybrid sp 3 orbitals: 1 part s , 3 parts p 35

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Sp3d -Hybridization of P in P F 5 P : 1s 2 2s 2 2p 6 3s 2 3p 3 37

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Lewis Structures: :

Lewis Structures: In 1916 G. N. Lewis proposed that atoms combine in order to achieve a more stable electron configuration. Maximum stability results when an atom is isoelectronic with a noble gas. An electron pair that is shared between two atoms constitutes a covalent bond. 39

Rules for Lewis Structures:

Rules for Lewis Structures 1. Make certain that the bond is a covalent bond then set up the skeleton structure as follows: The atom with the lowest electronegativity will tend to go in middle. Place all the other atoms around this central atom. Attach these atoms to the central atom in reasonable fashion with single bonds. 2.Sum valence electrons. 3. Complete octets of peripheral atoms. 4. Place leftover e - on central atom. 5. If necessary use multiple bonds to fill the center atom's octet. 40

Covalent Bonding in H2:

Covalent Bonding in H 2 H . H . Two hydrogen atoms, each with 1 electron, can share those electrons in a covalent bond. H : H Sharing the electron pair gives each hydrogen an electron configuration analogous to helium.(Duet Rule) 41

Covalent Bonding in F2:

Covalent Bonding in F 2 Two fluorine atoms, each with 7 valence electrons, can share those electrons in a covalent bond. Sharing the electron pair gives each fluorine an electron configuration analogous to neon. F : F : : .. .. .. .. F : .. . : F : .. . : 42

The Octet Rule:

The Octet Rule F : F : : .. .. .. .. In forming compounds, atoms gain, lose, or share electrons to give a stable electron configuration characterized by 8 valence electrons. 43

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Example C . . . . F : .. . Combine carbon (4 valence electrons) and four fluorines (7 valence electrons each) to write a Lewis structure for CF 4 . : F : .. .. C : F : .. .. : F : .. .. : F : .. .. The octet rule is satisfied for carbon and each fluorine. : 44

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Example It is common practice to represent a covalent bond by a line. We can rewrite : F : .. .. C : F : .. .. : F : .. .. : F : .. .. .. C F F F F .. .. .. .. : : : : : : .. as 45

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C : O .. O .. : : N C H Inorganic examples C : : : O .. : O .. : : : : : N : C : H Carbon dioxide Hydrogen cyanide 46

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Organic examples Ethylene Acetylene : : : C : C : H H C C H H C : : C .. H : : .. H H H C C H H H H 47

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Writing Lewis Structures for Molecules with One Central Atom SOLUTION: PROBLEM: Write a Lewis structure for CCl 2 F 2 , one of the compounds responsible for the depletion of stratospheric ozone. PLAN: Follow the steps outlined previously Step 1: Carbon has the lowest EN and is the central atom. The other atoms are placed around it. C Cl Cl F F 48

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SOLUTION: Steps 2-4: C has 4 valence e - , Cl and F each have 7. The sum is 4 + 4(7) = 32 valence e - . C Cl Cl F F Make bonds and fill in remaining valence electrons placing 8e - around each atom. : : : : : : : : : : : : 49

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Writing Lewis Structures for Molecules with Multiple Bonds. PLAN: PROBLEM: Write Lewis structures for the following: Nitrogen (N 2 ), the most abundant atmospheric gas If an atom does not have an octet, Step 5 which follows the other steps in Lewis structure construction must be done. If a central atom does not have 8e - , an octet, then an e - can be moved in to form a multiple bond. 50

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SOLUTION: N 2 has 2(5) = 10 valence e - . N . : N : . . . N . N : : . N : N : 51

Formal Charge::

Formal Charge: Formal charge of atom =Valence electrons-Electron count) Electron Count = (no of unshared electrons + 1/2 the shared electrons) 52

Nitric acid::

Nitric acid: : .. H O O O N : : .. .. 53

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. Hydrogen shares 2 electrons with oxygen. Assign 1 electron to H and 1 to O. A neutral hydrogen atom has 1 electron. Therefore, the formal charge of H in nitric acid is 0. Formal charge of H 54

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. Oxygen has 4 electrons in covalent bonds. Assign 2 of these 4 electrons to O. Oxygen has 2 unshared pairs. Assign all 4 of these electrons to O. Therefore, the total number of electrons assigned to O is 2 + 4 = 6. Formal charge of O 55

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. Electron count of O is 6. A neutral oxygen has 6 electrons. Therefore, the formal charge of oxygen is 0. Formal charge of O 56

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons). A neutral oxygen has 6 electrons. Therefore, the formal charge of oxygen is 0. Formal charge of O 57

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons). A neutral oxygen has 6 electrons. Therefore, the formal charge of oxygen is -1. Formal charge of O 58

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. Electron count of N is 4 (half of 8 electrons in covalent bonds). A neutral nitrogen has 5 electrons. Therefore, the formal charge of N is +1. Formal charge of N – 59

Nitric acid:

Nitric acid .. : .. H O O O N : : .. .. A Lewis structure is not complete unless formal charges (if any) are shown. Formal charges – + 60

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– "Electron Counts" and Formal Charges in NH 4 + and BF 4 - 1 4 N H H H H + 7 4 .. B F F F F .. .. .. .. : : : : : : .. 61

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For O A # valence e - = 6 # nonbonding e - = 4 # bonding e - = 4 X 1/2 = 2 Formal charge = 0 For O B # valence e - = 6 # nonbonding e - = 2 # bonding e - = 6 X 1/2 = 3 Formal charge = +1 For O C # valence e - = 6 # nonbonding e - = 6 # bonding e - = 2 X 1/2 = 1 Formal charge = -1 62

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P : 5 – 0 – ½ (8) = +1 O: 6 – 6 – ½ (2) = –1 Cl : 7 – 6 – ½ (2) = 0 63

Draw Lewis structure and find out formal charges on Following::

Draw Lewis structure and find out formal charges on Following: Nitromethane : CH 3 NO 2 Dimethylsulfoxide :(CH 3 ) 2 SO 64

Molecular Structure::

Molecular Structure: Three dimensional arrangement of the atoms in a molecule. 65

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Linear structure – atoms in a line Carbon dioxide 66

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Trigonal planar – atoms in a triangle Boron trifluoride 67

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68 Tetrahedral structure Methane

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Six pairs of electrons around the central atom are based on the Octahedron structure. AB 6 : e.g. SF 6 The central atom can be visualized as being at the centre of an octahedron, with the six electrons pointing to the six vertices – all bond angles are 90 ° Octahedral Square Pyramidal E.g. BrF 5 Square Planar E.g. XeF 4 69

VSEPR Model:

VSEPR Model VSEPR: Valence Shell Electron-Pair Repulsion. The structure around a given atom is determined principally by minimizing electron pair repulsions. BeCl 2 180 ° Linear 70

Three Pairs of Electrons:

Three Pairs of Electrons BF 3 120 ° Trigonal planar 71

Four Pairs of Electrons::

Four Pairs of Electrons: CH 4 109.5 ° Tetrahedral 72

Steps for Predicting Molecular Structure Using the VSEPR Model:

Steps for Predicting Molecular Structure Using the VSEPR Model Draw the Lewis structure for the molecule. Count the electron pairs and arrange them in the way that minimizes repulsion (put the pairs as far apart as possible). Determine the positions of the atoms from the way electron pairs are shared (how electrons are shared between the central atom and surrounding atoms). Determine the name of the molecular structure from positions of the atoms. 73

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VSEPR: Valence shell electron pair repulsion:

VSEPR: Valence s hell electron pair repulsion VSEPR Rule states that all groups emenating from an atom-whether single, double or triple or lone pairs will be in spatial positions that are far apart form one a nother as possible. The most stable arrangement of groups attached to a central atom is the one that has the maximum separation of electron pairs ( bonded or nonbonded ). 75

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But When some other atoms are attached to carbon there is certain deviation from their ideal values. The lone pair on Ammonia is larger than bonding pair, so it causes contraction in H-N-H angle. VSEPR: Valence shell electron pair repulsion 76

VSEPR: Valence shell electron pair repulsion :

If it is related with space repulsion between two groups, steric repulsion arises from butressing of filled orbitals that cannot participate in bonding. Order of repulsion: Lone pair-Lone pair> Lone pair-bond pair> Bond pair-bond pair VSEPR: Valence shell electron pair repulsion 77

Polar Covalent Bond::

Polar Covalent Bond: Covalent bonds can have ionic character Polar covalent bonds Bonding electrons attracted more strongly by one atom than by the other. Electron distribution between atoms is not symmetrical. 78


Electronegativity Electronegativity (EN): intrinsic ability of an atom to attract the shared electrons in a covalent bond. Differences in EN produce bond polarity. Scales of electronegativity 1.Pauling Scale 2.Mulliken Scale 79

Electronegativity scales:

Electronegativity scales Pauling Scale: Based upon bond dissociation energies. Mulliken Scale: Based upon following. Ionization potential: Energy required to remove electron from an atom or molecule.(number reflects the affinity of an atom for electron it already has) Electron affinity: Electron affinity is the amount of energy released or required to attach another electron to an atom or molecule. (This number reflects affinity of an atom for additional electron) 80

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Factor influencing electronegativity:

Factor influencing electronegativity Energy of orbital: The atom with lowest energy valence orbitals i.e F is most electronegative. 82

Electrostatic Potential Surfaces:

Electrostatic Potential Surfaces Surfaces are based on the van der Waals surface – allow us to visualize the charge distribution and the size and shape of the molecule. Areas shaded in red represent negative electrostatic potential and those shaded in blue represent positive electrostatic potential . Areas shaded in green are essentially neutral. 83

Inductive Effect:

Inductive Effect Phenomenon of withdrawing electrons through sigma bond to the more electronegative atom. Permanent effect in saturated carbon chain compounds. Group attached to carbon chain should have tendency to release or withdraw electrons . Types of inductive effect + I effect effect –electron donating groups e.g., CH 3 , C 2 H 5 – I effect effect –electron withdrawing groups e.g., - NO 2 , –CN 84

Features of Inductive Effect:

Chloroacetic acid is a stronger acid than acetic acid because.. Features of Inductive Effect K a = 1.4 × 10 -3 K a = 1.75 × 10 -5 85

Bond Dipoles::

Bond Dipoles: A bond dipole is the local moment associated with a polar Covalent bond. Electrostatic potential surfaces also reveal bond dipoles.  = Q  r, in debyes (D), 1 D = 3.336  10  30 coulomb meter 86

Molecular Dipole:

Molecular Dipole A molecular dipole is present in a molecule whenever the center of positive charge in the molecule is not coincident with the center of negative charge. Molecular dipole is vector sum of individual bond dipoles. It is based on magnitude of charge separation, and the intensity of electric field around the molecule. 0.0 D CCl4 and CH3CN 4.0 D 87

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Molecular Quadrupole Moment::

A Quadrupole is simple two dipoles aligned in such way that there is no net dipole. Benzene is common example which undergo large permanent quadrupole moment and charge distribution. In Benzene C is more elctronegative than Hydrogen and C undergoes SP2 Hybidization . Cyclohexane has negligible quadrupole moment indicating SP3 Hybidization Molecular Quadrupole Moment: 89

Hybridization Effects::

Hybridization Effects: Electronegativity of C and H depends upon hybridization effect. More the S character More will be electronegativity, because S orbital have substantial density of electrons at the nucleus. Sp >Sp2>SP3 More the S character shorter will be the bond length. 90


Resonance Some molecules are have structures that cannot be shown with a single representation. In these cases we draw structures that contribute to the final structure but which differ in the position of the  bond(s) or lone pair(s ). Such a structure is delocalized and is represented by resonance forms The resonance forms are connected by a double-headed arrow. 91

Resonance Hybrids called as Real Structures.:

Resonance Hybrids called as Real Structures. A structure with resonance forms does not alternate between the forms. Instead, it is a hybrid of the two resonance forms, so the structure is called a resonance hybrid For example, benzene (C 6 H 6 ) has two resonance forms with alternating double and single bonds In the resonance hybrid, the actual structure, all its C-C bonds are equivalent, midway between double and single 92

Rules for Resonance Forms:

Rules for Resonance Forms Individual resonance forms are imaginary - the real structure is a hybrid (only by knowing the contributors can you visualize the actual structure) Resonance forms differ only in the placement of their  or nonbonding (lone pairs) electrons. Different resonance forms of a substance don’t have to be equivalent. Resonance forms must be valid Lewis structures : the octet rule applies. The resonance hybrid is more stable than any individual resonance form would be. 93


Examples Electrons are spread over large number of atoms i.e described by single resonance structure. Energy of stabilization imparted by resonance is Resonance energy or delocalization energy . 94

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Only Electrons are move around the position of nuclei never change. Making resonance structure acceptable- 1.Nobel gas configuration 2.Maximum number of covalent bond. 3.Placement of negative charge on electronegative atom. 4.Minimum like charges. 95

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Cyclopentadine :

Cyclopentadine 99

Benzyl bromide:

Benzyl bromide 100

Basic Bonding Concepts: Bond Radius:

Basic Bonding Concepts: Bond Radius Covalent radius - is half the distance between two identical atoms bonded together. Ionic radius is the size of an atom as determined in the crystal lattice of various salts. Radius depends on the charge of the ion. Van der walls radius of an atom is the effective size of electron cloud when in a covalent bond, as preceived by an atom to which it is not attached. 101


Polarizability Def : Ability of electron cloud to distort in response to an electric field. Upon distortion the dipole is typically induced in the molecule, adding to any permanent dipole already present. Electronegativity plays important role in polarizability . Atom that holds on to their electrons are more polarizable. C>N>O>F Alkanes are generally most polarizable molecules. Water is very polar molecule, but not polarizable. E.g - Methane-2.6 Benzene -10.32 Water-1.45 102

Comparison C-I and C-Cl:

Comparison C-I and C- Cl C-I bond is more reactive than C- Cl ? Ans - - Although C-I bond is not very polar, but when it goes in SN2 reaction, the anionic nucleophile induces a large dipole in C-I which makes I more polarizable. -Thus large polarizability makes lower elctronegativity , hence C-I is more reactive than C- Cl 103

Bond Order:

Bond Order Bond order= Nb -Na/2 Only MO s formed from valence orbitals are considered for determining bond order. Molecule is stable if Nb >Na, unstable if Na>Nb. Information given by bond order 1.Stability of molecule or ion- Nb >Na 2.Bond dissociation energy: Greater the bond order greater is bond dissociation energy. 3.Bond length: Bond order is inversely proportional to bond length. 4.Magentic properties: Greater the number of unparied electrons ,more will be paramagnetic character. If no unpaired electrons species will be diamagnetic. 104

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