Trends in oxidation state of d-block elements Part-3

 Oxidation States (oxidation number)-describes loss/gain of electron by an atom

·      Stability of oxidation state of an element is defined by –

i)      Electronic configuration

ii)      Sum of the ionisation enthalpies

iii)    SEP value

iv)    Hydration enthalpy

·      Sc(+3) and Zn(+2) exhibit only one oxidation state

·      Except scandium, (which has +3 oxidation state) for the elements of first transition series +2 oxidation state is the most common. This state arises due to loss of 4s-electron

·      In 3d series highest oxidation state is +7 (Mn)

·      In d-block series highest oxidation state is +8 (Os, Ru)

·      All transition elements, except first and last member of the series, exhibit a number of oxidation state

·      Generally, within the transition series, the highest oxidation state increases with increase of atomic number, reaching to a maximum in the middle and then starts decreasing.

·      The variable oxidation states of a transition metal is due to the involvement of (n-1)d and, outer ns electrons in bonding as the energies of ns and (n-1)d subshells are nearly equal.

·      In a group of d-block elements, the higher states are more stable for heavier elements.

For example, in group 6, Mo(vi) and W(vi) are more stable than Cr(vi). For example, dichromate having Cr(vi) is a strong. Oxidising agent in acidic medium while MoO3and W03 are stable. oxides.

·      Higher oxidation states are exhibited when ns and (n-1)d-electrons take part in bonding.

·      higher oxidation states are found in compound with fluorine and oxygen because fluorine and oxygen are most electronegative in nature

       Higher oxidation states in oxides are normally more stable than fluorides due to capability of oxygen to form multiple bonds.

·      In p-block lower oxidation states of heavier elements are more stable while in d-block heavier element, higher oxidation state are more stable. Due to inert pair effect found in p-block elements.

            For example, in group 6, Mo(vi) and W(vi) are more stable than Cr(vi). For example, dichromate having Cr(vi) is a strong. Oxidising agent in acidic medium while MoO3and W03 are stable. oxides.

·      Some of the transition metals form compounds in zero oxidation state or lower oxidation state In transition element when a complex compound has ligands capable of π-acceptor character in addition to the σ-bonding or ligand having ability of forming synergic bond/back bonding

Example – Ni(CO)4, Fe(CO)₅

Problems-1. (on the basis of electronic configuration)

a). Ti⁴⁺ (3d^o4s⁰) is more stable than Ti^{3 +} (3d¹4s^o) 

Ans - due to completely filled orbitals in Ti4+ which is most stable than partially filled orbital

b).  Mn²⁺⁽3d⁵4s⁰) is more stable than Mn^{3 +} (3d⁴ 4s⁰)

Ans - due to half-filled orbitals which   is most stable than partially filled orbital

 

 c). Fe³⁺(3d⁵4s⁰) is more stable than Fe^{2 +} (3d⁶ 4s⁰) 

Ans- due to half-filled orbitals which   is most stable than partially filled orbital

Problem-2 (on the basis of sum of ionisation enthalpy)

a). Ni²⁺is more stable than Pt²⁺ while Pt⁴⁺ is more stable than Ni⁴⁺

  Ans- due to sum of IE₁ + IE₂ for Ni²⁺ is lesser than IE₁ + IE₂ for Pt²⁺ while IE₁ + IE₂+ IE₃ + IE₄ for Pt⁴⁺ is smaller than Ni⁴⁺ .

Problem -3 ( on the basis of hydration enthalpy)

a).Cu⁺ is not stable in aqueous solution than Cu²⁺ why?

Ans – since energy is required to remove one electron from Cu⁺.large hydration enery evolve during hydration of Cu⁺ that compensates it .

Problem-4 Name the element which do not show variable oxidation state – Sc

·      Problem -5 which of the 3d- series of transition metal exhibit variable oxidation state and why ?  Ans--Mn due to the involvement of (n-1)d and, outer ns electrons in bonding as the energies of ns and (n-1)d subshells are nearly equal.

Click here to see next The Inner Transition f-Block Elements

Click here to see d & f block elements slide presentation

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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