Chemical bonding -HYDROGEN BONDING

 

HYDROGEN BONDING

·     ·       In 1920, Latimer and Rodebush introduced the idea of "hydrogen bond" to explain the nature of association in liquid state of substances like water, hydrogen fluoride, ammonia, formic acid, etc.

     ·       Hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) 

     ·       It exerts in polar molecule or molecule having polar group.

     ·       This bond is represented by dotted line (---------) while solid represents the covalent bond.

     ·       The magnitude of the hydrogen bonding depends on the physical state of the compound.

     ·        order of  increasing strength of hydrogen bond in various physical state as

        gas < Liquid< solid  

        Type of hydrogen bonding-

         a)       Intermolecular hydrogen bonding: This type of bonding results between the positive and negative ends of different molecules of the same or different substances. Example-

   a) H^δ+ — F^δ--------H^δ+ — F^δ+--------H^δ+ — F^δ-

   b) H^δ+ — F^δ--------H^δ+ — O —H^δ+-------H^δ+ — F^δ       

        b).    Intramolecular hydrogen bonding: This type of bonding results between hydrogen and an electronegative element both present in the same molecule.

Importance of Hydrogen bond-

   ·      It helpful to explain that H2O exits in liquid state while H2S exists in gaseous state since in H2O, oxygen is more electronegative element than Sulphur in H2S therefore oxygen has ability to form H-bond with other water molecules. so it exists.

   ·      Ortho-nitrophenol has low boiling point than para-nitrophenol since P-nitrophenol form intermolecular hydrogen bond which is stronger than intra molecular hydrogen bond which is formed by ortho-nitrophenol.

     ·      Solubility of compound in water can be explained on the basis of h-bonding. Example alcohol is soluble in water but oil not. Because alcohol forms H-bond with water but oil does not form.

Bond parameter- Bond length, bond angle and bond energy

 Bond parameter-

 Bond Length-

 The average distance between the nucleus of two atoms is known as bond length, normally it is represented in Å or pm  

Factors Affecting Bond Length

(a) electronegativity: - Bond length proportional to EN

                                             Therefore               H—F < H—Cl < H—Br < H—I

(b) Bond order or number of bonds: - Bond length α 1/Number of bonds or bond order

                                                                     ex.        C—C,      >      C = C  >    C = C

                                                 bond length          154pm             134pm      120pm

Bond Angle- The minimum angle between any two adjacent bonds is known as bond angle. It is represented in degree (°), min (') and second (")

Factors affecting the bond angle –

a). Hybridization— bond angle is proportional to  % 's' character

                                   BeCl2     > BCl3           > CCl4

          bond Angle      180°          120°          109.28'

Hybridizations-           sp               sp2              sp3

% s- character            50%           33.33%       25%

b). Lone pair —When hybridization is same, lone pair are different

      Bond angle is inversely proportional to No. of lone pair Example-

                               CH₄                        NH₃                     H₂O

Hybridization     sp³                      sp³                       sp³

Bond angle        109.28'                107°                     104.5°         

No of lone pair    0                       1                            2

c) nature of Central Atom

                 Bond angle is proportional to the Electronegativity of central atom

Example -

                       NH3         >       PH3       > AsH3

 Bond angle  107°                   93°            91°  

 - Electronegativity decreases from N to As , so,  Bond angle will decrease

d). Nature of Side atom bonded to a central atom

       Bond angle is proportional to size of the side atom  but is inversely proportional to electronegativity of bonded atom

  Example-

    PF₃ < PCl₃ < PBr₃ < PI₃   ---->   Reason - (Electronegativity of side atom decrease)

    OF₂ < Cl₂O < Br₂O ---->   Reason - (Electronegativity of side atom decrease)

     SF₂ < SCl₂ < SBr₂ ---->   Reason - (Electronegativity of side atom decrease

HYBRIDISATION

 

HYBRIDISATION

   ·      It is introduced by Pauling and Slater, to explain equivalent nature of covalent bonds in a molecule.

   ·      Mixing of different shape and approximate equal energy atomic orbitals, and redistribution of energy to form new orbitals, of same shape & same energy. These new orbitals are called hybrid orbitals and the phenomenon is called hybridization.

 Characteristic of Hybridization

(1) Hybridization is a mixing of orbitals and not electrons. Therefore, in hybridization full filled, half-filled and empty orbitals may take part.

(2) Number of the hybrid orbitals formed is always be equal  to number of atomic orbitals which  participates in the process of hybridization.

(3) The number of hybrid orbitals on central atom of a molecule or ion

 (4) One element can represent many hybridization state depending on experimental conditions for example, C showing sp, sp² and sp³ hybridization in its compounds.

(5) Hybrid orbitals are designated as sp, sp², sp³ etc.

 (6) The order of repulsion between lp – lp > lp – bp > bp – bp

(7) The directional properties in hybrid orbital is more than atomic orbitals. Therefore, hybrid orbitals form stronger sigma bond. The directional property of different hybrid orbitals will be in following order. sp < sp2 < sp3 < sp3d < sp3d 2 < sp3d3 since it depends upon the directional nature of orbitals.

Types of Hybridization

(A) sp hybridization

 (a) In this hybridization one s  & one p – orbital of an atom are mixed to give two new hybrid orbitals which are equivalent in shape & energy known as sp hybrid orbitals.

 (b) These two sp hybrid orbitals are arrange in straight line & at bond angle 180°.

(B) sp² Hybridization:

(a) In this hybridization one s & two p orbitals are mixed to give three new sp2 hybrid orbitals which are in same shape & equivalent energies.

 (b) These three sp2 hybrid orbitals are at angle of 120° & giving trigonal planar shape.

(C) sp³ Hybridization:

 (I) In this hybridization one  ‘ s’ orbital and three ‘p’ orbitals of an atom of a molecule or ion, are mixed to give four new hybrid orbitals called as sp3 hybrid orbitals.

 (II) The angle between hybrid orbitals will be 109° 28'

(D) sp³d Hybridization:

(I) In this hybridization one s orbital, three p orbitals and one d orbital are mixed to give five new hybrid orbitals which are equivalent in shape and energy called as sp3d hybrid orbitals.

(II) Out of these five hybrid orbitals, three hybrid orbitals are at 120° angle and two hybrid orbitals are perpendicular to the plane of three hybrid orbitals that is trigonal planar, the shape of molecule becomes trigonal bipyramidal. For example, PF5 showing sp3d hybridisation

(C) sp³d²Hybridisation:

 (I) In this hybridization, one s-orbital, three p-orbitals & two d-orbitals (dz2 , dx2–y2 ) are mixed to give six new hybrid orbitals known as sp3d 2 hybrid orbitals.

(II) The geometry of molecule obtained from above six hybrid orbitals will be symmetrical octahedral or square bipyramidal.

 (III) The angle between all hybrid orbitals will be 90°. Example: SF₆, AlF6 ^–3, PF₆ ^–, ICl5, XeF4, XeOF_4, ICl₄ ^–

(F) sp³d³ Hybridization:

(I) In this hybridization, one s-orbital, three p-orbitals & three d-orbitals are mixed to give seven new hybrid orbitals known as sp³d ³ hybrid orbitals.

(II) These seven sp³d ³ orbitals are configurated in pentagonal bipyramidal shape.

(III) Five bond angles are of 72° and 10 bond angles of 90°. (IV) The following examples showing sp3d 3 hybridization –IF₇ & XeF₆

_{explaination is coming soon} 

Valence bond Theory (VBT)

 VALENCE BOND THEORY (VBT) :

 It was presented by Heitler & London to explain how a covalent bond is formed. It was extended by Pauling & Slater.

 It is based on the knowledge of atomic orbitals, electronic configuration of elements, the overlap of atomic orbitals, hybridization of atomic orbitals

The main points of theory are –

 (a) To form a covalent bond overlapping occurs between half-filled valence shell orbitals of the two atoms.

(b) Resulting bond acquires a pair of electrons with opposite spins to get stability.

(c) Orbitals come closer to each other from the direction in which there is maximum overlapping

(d) So covalent bond has directional character.

(e) Strength of covalent bond α extent of overlapping.

(f) Extent of overlapping depends on two factors.

(i) Nature of orbitals – 

p, d and f are directional orbitals -------> more overlapping

s-orbital -----> non directional – less overlapping

(ii) Nature of overlapping –

 Co-axial overlapping - extent of overlapping more.

Collateral overlapping - extent of overlapping less Order of strength of Co - axial overlapping –

p - p > s - p > s - s                                                                                      

Two types of bonds are formed on account of overlapping.

(A) Sigma bond                                               

(B) Pi- bond

(A) Sigma bond

a) Bond formed between two atoms by the overlapping of half-filled orbitals along their axis (end to end overlap) is called sigma bond.

(b) sigma bond is directional.

(c) sigma bond do not take part in resonance.

(d) Free rotation is possible about a single s bond.

(e) Maximum overlapping is possible between electron clouds and hence it is strong bond.

(f) There can be only one sigma bond between two atoms

Sigma bonds are formed by four types of overlapping

 (i) s - s overlapping – Two half-filled s-orbitals overlap along the internuclear axis. Ex. H2 molecule

(ii) s - p overlapping (Formation of HF) – When half fill s-orbital of one atom overlap with half-filled p- orbital of other atom.

(iii) p - p overlapping – (Coaxial) – It involves the coaxial overlapping between half filled p-orbitals of two different or same atoms.

Pi(π)-Bond

(a) The bond formed by sidewise (lateral) overlapping are known as pi bonds.

(b) Lateral overlapping is only partial, so formed are weaker and hence more reactive than sigma bonds (Repulsion between nucleus is more as orbitals have to come much close to each other for pi- bonds formation) Example – Formation of O2

(c) Free rotation about a pi bond is not possible.

(d)  π bond is weaker than sigma bond (Bond energy difference is 63.5 KJ or 15 K cal/mole)

(e) π bonds are less directional, so do not determine the shape of a molecule

F) it is former by overlapping of unhybridized orbitals

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VSEPR theory - VALENCE SHELL ELECTRON PAIR REPULSION THEORY

 VALENCE SHELL ELECTRON PAIR REPULSION THEORY (VSEPRT)

It was described by Sidgwick and Powel in 1940 and further developed by Gillespie and Nyholm in 1957

    ·       Shape of molecule depends upon the no of bonded or non-bonded electron pairs around  the central atom

    ·       These electron pairs having negative charge repel each other and tend to occupy such position in space that minimize repulsion and thus maximize distance between them

    ·        If the central atom possesses only bonded pairs of electrons along with identical atoms then shape of the compound is symmetrical and according to Sidgwick & Powel.

Eg.        

CO2 — 180° — linear

 BF3 — 120° — triangular

CH4 — 109° 28' — tetrahedral

PCl5 - 120° and 90°   - Trigonal bipyramidal

 If the central atom possesses bonded pair of electrons as well as lone pair of electrons, then shape of the molecule will be unsymmetrical ie. the original bond angle will be disturbed due to repulsion between lone pair of electrons.

 Similarly, on having different type of side atoms, molecule becomes unsymmetrical due to unequal force of repulsion between e– .

  Order of repulsion is -  

  lp – lp > lp – bp > bp – bp

 Bond angle α     1/ no of lone pair

By increasing no. of lone pair of electrons, bond angle is decreased approx. by 2.5°.

 eg.:-        

CH4       NH3      H2O --------> sp3

109°      107°       105°              hybridization

d). The VSEPR model considers double and triple bonds to have slightly greater repulsive effects than single bonds because of the repulsive effect of a electrons.

For example, the H3C-C-CH3 angle in (CH3)2C=CH2 is smaller and H3C-C=CH2 angle is larger than the trigonal 120°

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Formal charge

 Formal charge- defined as the difference between the no of valence electrons of that atom in an isolated of free state and the no of electron assigned to that atom in the lewis structure

= no. of valence electron in the free atom – no of non-bonding electrons (lone pair)- (no of bonding or shared electrons)/2

Charge on the molecule or ion = sum of all the formal charges

Significance-

·       It do not indicate real charge separation within the molecules. Indicating the charge on the atoms in the lewis structure helps in keeping track of the valence electrons in the molecule

·       Helpful in selection in lowest energy structure from a no of possible lewis structure for a given species

KOSSEL - LEWIS APPROACH TO CHEMICAL BONDING - Octet rule

KOSSEL - LEWIS APPROACH TO CHEMICAL BONDING (Octet rule)

     ·       Every atom has a tendency to complete its octet in outermost shell

     ·       H has the tendency to complete its duplet.

     ·       To acquire inert gas configuration atoms loose or gain electron or share electron.

      ·       The tendency of atoms to achieve eight electrons in their outer most shell is known as Lewis octet rule.

Exception of Octet Rule

(a) Incomplete octet molecules: - or (electron deficient molecules): -

·    Compound in which octet is not complete in outer most orbit of central atom.

 Example - 

Halides of IIIA groups, BF3, AlCl3, BCl3

hydride of III A/13th group etc.

Other examples – BeCl2 (4e),

(b) Expansion of octet or (electron efficient molecules)

Compound in which central atom has more than 8e– in outermost orbits.

 Example - In PCl5 , PF5,SF6 and IF7 the central atom P, S and contain 10, 12, and 14 electrons respectively.

(c) Odd electron molecules: - Central atom in a molecule has an unpaired electron or odd no of electrons in their outer most shell. e.g. NO, NO2, ClO2 etc.

d). It is based upon chemical inertness of noble gases. However, Xe and Kr also combine with other to form compound likes XeF₂, KrF₂, XeOF₂ etc.

E) could not explain relative stability of molecules

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Lewis symbol and rule of Lewis dot structure -why does atom combine with other atom?

 Chemical bond- 

·       A force that acts between two or more atoms to hold them together as a stable molecule. 

·       This process accompanied by decrease in energy.  

·       Decrease in potential energy (P.E.) is proportional to Strength of the bond.  

·       Therefore, molecules are more stable than atoms. 

Cause of Chemical Combination  

(A) Tendency to acquire minimum energy  

(a) When two atoms approach to each other. Nucleus of one atom attracts the electron of another atom.  

(b) If net result is attraction, the total energy of the molecule decreases and a chemical bond form. 

 (d) So, Attraction is inversely proportional 1/energy and proportional to Stability.  

(e) Bond formation is an exothermic process during formation of a molecule by combining of molecules 

(B) Tendency to acquire noble gas configuration:  

(a) Atom combines to acquire noble gas configuration.  

(b) Only outermost electrons or valence electrons i.e. ns, np and (n-1) d shells electrons participate in bond formation or chemical reaction. 

(c) Inert gas elements do not participate in bond formation, since they have stable electronic configuration as 1s² or ns²np⁶) 

Lewis Symbols  

A Lewis symbol is a convenient way to represent the valence electrons, which are shown as dots placed on the sides, top, or bottom of the symbol for the element.  

Na

Lewis Structures  

·       The Lewis structure is a representation of a molecule that shows how the valence  electrons are arranged among the atoms in the molecule. 

·       These representations are named after G. N. Lewis, who conceived the idea while lecturing to a class of general chemistry students in 1902. 

·       Lewis structures are based on achieving a noble gas electron configuration by atoms 

·       Writing Lewis structures is a trial-and-error process. 

·       The following steps are followed in constructing dot formulae for molecules and polyatomic ions:  

      (i) Write a symmetrical 'skeleton' for the molecules and poly atomic ions. 

                   a. The least electronegative element is usually taken as the central                                                                   element except Hydrogen. 

                   b. Oxygen atoms do not bond to each other except in O₂, 0 ₃, peroxide and superoxides. 

    c.       Hydrogen actually bonds to an oxygen atom and not to the central atom in ternary acids (oxyacids) ex-Nitrous acid HN02, has the skeleton H-O-N=O 

(ii) Calculate the number of electrons available in the valency shell of all the atoms (sum of the valence electrons for a molecule.) this is expressed by A  

         For negatively charged ions add to the total number of electrons equal to the charge on the anion and for positively charged ions, subtract the number of electrons equal to the charge on the cations 

Example – 

1). Sum of valence electrons for PO₄^3-  = 1 x 5 (for P atom) + 4 x 6 (for 0 atoms) + 3 (for charge)   

                      = 5 + 24 + 3  

                      = 32 electrons.

  

  2). Sum of valence electrons for NH₄⁺ ion  = 1 x 5 (for nitrogen atom) + 4 x 1 (for H                                      atoms) 1 (for positive charge)   

                                                                      = 5+4-1 

                                                                      =  8 electrons. 

  iii). Calculate the total number of electrons needed by all atoms to achieve noble gas                                       configuration.     

                                      This number is represented by N. 

                                         For example: N for  H2S04 

                                          N =2x2+8xl+8x4 =4+8+32 = 44 electrons. 

  iv).  Calculate the total number of electrons shared. This is represented by S which is equal to

                 N - A. For example: S for H2SO4 .

                                                S = N-A = 44- 32 =12 electrons. 

   v) Place the shared pair of electrons into the skeleton, using double and triple bonds only when                   necessary. 

  vi) Place the additional unshared (lone) pairs of electrons to fill the octet to every atom except                        hydrogen which can have only 2 electrons as the total comes equal to A