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

NEXT TOPIC                                  PREVIOUS TOPIC

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°

NEXT TOPIC            PREVIOUS TOPIC

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

NEXT TOPIC

trends in ionisation enthalpy of d-block elements part- 2

 

Ionisation enthalpy- M(g) -----> M⁺(g) + Δ_iH (ionisation enthalpy) 

·      Ionisation potential values increase in a period from left to right with increase in  nuclear charge. The increase, however, is not regular.

·      In vertical columns, i.e., groups, the ionisation potential decreases from first member to second member in most of the cases as expected,

·       however, the third member has higher value than second member. This is due to lanthanide contraction, i.e., the atomic radii of the elements of the same group of the second and third transition series are nearly the same but atomic numbers differ by 32. Thus, the outer electrons are held very firmly and the ionisation potential values are very high.   

·      On account of this, the members of 5d are inert under ordinary conditions, i.e., Pt, Au, Hg, etc., are noble metals.

·      Ionisation enthalpy depends upon

      a) Nuclear charge α I.E   α 1/size of atom     

      b) electronic configuration

  (order of ionisation enthalpy  à partially filled < half-filled < completely filled orbitals)

                                                       Fe                           Mn                     Zn

  Electronic configurationà3d⁶4s²                 3d⁵4s²           3d¹⁰4s²      

  Stability order                      less                 stable         most stable

·      I^st ionisation enthalpy of 5d series element are higher than those of 3d or 4d elements

 Ans it is due to lanthanoid contraction or the ineffective shielding of the nucleus by 4f electrons over effective nuclear charge.

·      Ist ionisation enthalpy of Mn is greater than Fe

Ans-due to presence of half-filled d-orbital in Mn(3d⁵4s²⁾ that is most stable than partially filled d-orbital present in Fe(3d⁶4s²)

·      Third ionisation energy of manganese is very high.

 Ans This is because the third electron has to be removed from a stable configuration, (Mn^{2 +} 3d⁵), i.e;, half-filled 3d-subshell.

·      The second ionisation potential values of Cu and Cr are sufficiently higher than those of neighbouring elements.

Ans This is because of the electronic configuration of Cu⁺ which is 3d¹⁰ (completely filled) and of Cr⁺ which is 3d⁵ (half-filled), i.e., for the second ionisation potential, the electron is to be removed from very stable configurations.

Click here to see next Trends in oxidation state of d-block elements Part-3

Click here to see d & f block elements slide presentation

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 

 

States of matter intermolecular forces Part -1

 

States of matter

There are two types of properties shown by individual or the collection of a large number of particles

A)        1)   Individual/atomic properties- are shown by single particle or atom of a matter

Examples- atomic size, atomic/ionic radii, ionisation enthalpy, electron gain enthalpy, electronegativity, electro positivity

B)      2). Bulk properties-

                           those properties which are associated with the collection of a large number of individual particles are called bulk properties.

Example - Melting, boiling, solubility, water having wetting properties

Water exist in three states of matter as

                                                                   Solid -     Ice

                                                                  liquid – water

                                                                  Gas - water vapour

Triple point- the state at which all three phases exists simultaneously is called triple point

i.e. there are three states of matter as         a) solid   b) liquid     c) gas

on the basis of shape, volume, rigidity, fluidity, interparticle spaces, interparticle force, K.E. and Diffusion  

Mainly, the states of matter depend upon –

                                                       1). Energies of molecules

                                                       2). On the manner in which the individual molecules bind/aggregate

Note -Chemical properties of a substance do not change with the change of its physical state but the rate of chemical reaction do depend upon the physical state.

          Therefore  ,it become necessary for a chemist to know the knowledge of state of matter and the physical laws which govern the behaviour of matter in different states

Intermolecular forces-

·       The forces of attraction and repulsion between interacting particles (atoms or molecules) are called intermolecular forces.

·       These exist all matters ( solid liquid gas)  and responsible for many structural feature and physical properties of matter.

·       this term does not involved the electrostatic/ionic force and covalent bonds

·       intermolecular forces are weaker forces.

·       The attrative intermolecular forces are known as Vander Waals forces, in honour of Dutch scientists Johannes who studied the effect of these molecular forces on the real gases

Intermolecular forces  arise due to the following type of interactions:-

1). Dispersion forces or London force or instantaneous dipole – induced dipole forces

      ·            This force of attraction was firstly proposed by the German physicist Fritz London.

      ·           Dispersion forces are weaker intermolecular forces

      ·          These forces exert among the atoms(noble gases) or molecules(H₂ O₂ N₂ CO2  which are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

       ·         Temporary dipoles can occur in non- polar molecules when the electrons that constantly orbit the nucleus occupy a similar location or unsymmetrical that means the charge cloud become more on one side than the other  by chance(momentarily).it is also called instantaneous dipole.

                     This instantaneous or transient dipole distorts the electron density of the other atom or molecule.

    i.e. Temporary dipoles can induce a dipole in neighbouring molecules, initiating an attraction called London force.

·       These forces are always attractive and interaction energy is proportional to 1

2. Dipole- induced dipole interaction·       This force of attraction was firstly proposed by the German physicist Fritz London.

·       Dispersion forces are weaker intermolecular forces

·       These forces exert among the atoms(noble gases) or molecules(H₂ O₂ N₂ CO2  which are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

·       Temporary dipoles can occur in non- polar molecules when the electrons that constantly orbit the nucleus occupy a similar location or unsymmetrical that means the charge cloud become more on one side than the other  by chance(momentarily).it is also called instantaneous dipole.

·       This instantaneous or transient dipole distorts the electron density of the other atom or molecule.

i.e. Temporary dipoles can induce a dipole in neighbouring molecules, initiating an attraction called London force.

·       These forces are always attractive and interaction energy is proportional to 1/r⁶

2. Dipole- diploe (Keesom) force-

     ·       This interaction was first studied by Kessom in 1912.

     ·       These forces occur in molecules which have permanent electric dipole such as HCl, NH3, H2O etc.

     ·       A polar molecule has separate centre of positive and negative charges and the ends of dipoles possess “partial charge” (δ=delta=partial)

    ·       The partial charges are always less than the unit electronics charge (1.6 X 10^-19C).

    ·       These forces arise due to interaction between oppositely charged ends of the polar molecules.

    ·       The positive end of one molecule attract the negative end of the other molecule and vice versa

     ·       Greater the dipole moment of the molecules, greater is the forces of attraction.

     ·       It is stronger than London force

      ·       Interaction energy is proportional to 1/r³ (between stationary polar molecules and is proportional to 1/r⁶

   Besides dipole-dipole interaction, polar molecule can interact by London forces also.as a result the net intermolecular forces in polar molecules increase.

4. Hydrogen Bond