Case Study Questions Class 11 Chemistry Chapter 4 Chemical Bonding and Molecular Structure
CBSE Class 11 Case Study Questions Chemistry Chemical Bonding and Molecular Structure. Important Case Study Questions for Class 11 Board Exam Students. Here we have arranged some Important Case Base Questions for students who are searching for Paragraph Based Questions Chemical Bonding and Molecular Structure.
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CBSE Case Study Questions Class 11 Chemistry Chemical Bonding and Molecular Structure
Case I
The attractive force which holds variousconstituents (atoms, ions, etc.) together in differentchemical species is called a chemical bond. In order to explain the formation of chemicalbond in terms of electrons, a number ofattempts were made, but it was only in 1916when Kössel and Lewis succeededindependently in giving a satisfactoryexplanation. They were the first to providesome logical explanation of valence which wasbased on the inertness of noble gases. Lewis postulated that atoms achieve thestable octet when they are linked bychemical bonds. In the formation of amolecule, only the outer shell electrons takepart in chemical combination and they areknown as valence electrons. The inner shellelectrons are well protected and are generallynot involved in the combination process.G.N. Lewis, an American chemist introducedsimple notations to represent valenceelectrons in an atom. These notations arecalled Lewis symbols. For example, the Lewissymbols for the elements of second period areas under:
The bond formed, as a result of theelectrostatic attraction between thepositive and negative ions was termed as the electrovalent bond. The electrovalenceis thus equal to the number of unitcharge(s) on the ion.
Kössel and Lewis in 1916 developed animportant theory of chemical combinationbetween atoms known as electronic theoryof chemical bonding. According to this,atoms can combine either by transfer ofvalence electrons from one atom to another(gaining or losing) or by sharing of valenceelectrons in order to have an octet in theirvalence shells. This is known as octet rule. when two atoms share oneelectron pair they are said to be joined bya single covalent bond. In many compoundswe have multiple bonds between atoms. Theformation of multiple bonds envisagessharing of more than one electron pairbetween two atoms. If two atoms share twopairs of electrons, the covalent bondbetween them is called a double bond. Forexample, in the carbon dioxide molecule, wehave two double bonds between the carbonand oxygen atoms. Similarly in ethenemolecule the two carbon atoms are joined bya double bond. The Lewis dot structures provide a pictureof bonding in molecules and ions in termsof the shared pairs of electrons and theoctet rule. The Lewis dotstructures can be written by adopting thefollowing steps:
- The total number of electrons required forwriting the structures are obtained byadding the valence electrons of thecombining atoms. For example, in the CH4molecule there are eight valence electronsavailable for bonding.
- For anions, each negative charge wouldmean addition of one electron. Forcations, each positive charge would result in subtraction of one electron from the totalnumber of valence electrons. For example,for the CO32– ion, the two negative chargesindicate that there are two additionalelectrons than those provided by theneutral atoms.
- Knowing the chemical symbols of thecombining atoms and having knowledgeof the skeletal structure of the compound, it is easyto distribute the total number of electronsas bonding shared pairs between theatoms in proportion to the total bonds.
- In general the least electronegative atomoccupies the central position in themolecule/ion. For example in the NF3 andCO32–, nitrogen and carbon are the centralatoms whereas fluorine and oxygenoccupy the terminal positions.
- After accounting for the shared pairs ofelectrons for single bonds, the remainingelectron pairs are either utilized for multiplebonding or remain as the lone pairs. Thebasic requirement being that each bondedatom gets an octet of electrons.
[A] MCQ
1) … postulated that atoms achieve the stable octet when they are linked by chemical bonds.
(a) Lewis
(b) Debye
(c) Charles
(d) Sidgwick
Ans- a) Lewis
2) … in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding.
(a) Kössel
(b) Lewis
(c) Both a) & b)
(d) Sidgwick
Ans- c) Both a) & b)
3) In the formation of a molecule, only the outer shell electrons take part in chemical combination and they are known as …
(a) Backscattered electrons
(b) valence electrons
(c) Primary electrons
(d) Secondary electrons
Ans- b) valence electrons
4) In the CH4 molecule there are … valence electrons available for bonding.
(a) 4
(b) 6
(c) 8
(d) 10
Ans- c) 8
5) The type of bond between atoms in a molecule of CO2 is ….
(a) Ionic bond
(b) Metallic bond
(c) Hydrogen bond
(d) covalent bond.
Ans- d) covalent bond.
[B] Short Answers
1) Describe – Lewis Symbol
Ans- G.N. Lewis, an American chemist introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols. For example, the Lewis symbols for the elements of second period are as under
2) Describe – electrovalent bond.
Ans- The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the electrovalent bond. The electrovalence is thus equal to the number of unit charge(s) on the ion.2) Describe – electrovalent bond.
3) What is octate rule ?
Ans- Atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.
[C] Long Answers
1) How to write Lewis Dot structure ?
Ans-The Lewis dot structures can be written by adopting the following steps:
- The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms.
- For anions, each negative charge would mean addition of one electron. For cations, each positive charge would result in subtraction of one electron from the total number of valence electrons.
- Knowing the chemical symbols of the combining atoms and having knowledge of the skeletal structure of the compound , it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.
- In general the least electronegative atom occupies the central position in the molecule/ion.
- After accounting for the shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs. The basic requirement being that each bonded atom gets an octet of electrons.
Case –II
Lewis dot structures, in general, do notrepresent the actual shapes of the molecules.In case of polyatomic ions, the net charge ispossessed by the ion as a whole and not by aparticular atom. It is, however, feasible toassign a formal charge on each atom. Theformal charge of an atom in a polyatomicmolecule or ion may be defined as thedifference between the number of valenceelectrons of that atom in an isolated or freestate and the number of electrons assignedto that atom in the Lewis structure. It isexpressed as :Generally the lowest energystructure is the one with the smallestformal charges on the atoms. The formalcharge is a factor based on a pure covalentview of bonding in which electron pairsare shared equally by neighbouring atoms. The octet rule, though useful, is not universal.It is quite useful for understanding thestructures of most of the organic compoundsand it applies mainly to the second periodelements of the periodic table. There are threetypes of exceptions to the octet rule.
- The incomplete octet of the central atom
- Odd-electron molecules
- The expanded octetFrom the Kössel and Lewis treatment of theformation of an ionic bond, it follows that theformation of ionic compounds wouldprimarily depend upon:
- The ease of formation of the positive andnegative ions from the respective neutralatoms;
- The arrangement of the positive andnegative ions in the solid, that is, thelattice of the crystalline compound.
The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJ mol–1. This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+ (g) and one mole of Cl– (g) to an infinite distance.Bond length is defined as the equilibriumdistance between the nuclei of two bondedatoms in a molecule. Bond lengths aremeasured by spectroscopic, X-ray diffractionand electron-diffraction techniques . The covalent radius is measuredapproximately as the radius of an atom’score which is in contact with the core ofan adjacent atom in a bonded situation. The vander Waals radius represents the overall sizeof the atom which includes its valence shellin a nonbonded situation. Bond Angle is defined as the angle between the orbitalscontaining bonding electron pairs around thecentral atom in a molecule/complex ion. Bondangle is expressed in degree which can beexperimentally determined by spectroscopicmethods. It gives some idea regarding thedistribution of orbitals around the centralatom in a molecule/complex ion and hence ithelps us in determining its shape. Forexample H–O–H bond angle in water can berepresented as under :
Bond Enthalpy It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. The unit of bond enthalpy is kJ mol–1. For example, the H–H bond enthalpy in hydrogen molecule is 435.8 kJ mol–1 . H2 (g) → H(g) + H(g); ∆aH = 435.8 kJ mol–1. Bond OrderIn the Lewis description of covalent bond,the Bond Order is given by the number ofbonds between the two atoms in amolecule. The bond order, for example in H2(with a single shared electron pair), in O2(with two shared electron pairs) and in N2(with three shared electron pairs) is 1,2,3respectively. A general correlation useful forunderstanding the stablities of moleculesis that: with increase in bond order, bondenthalpy increases and bond lengthdecreases. The concept of resonance was introducedto deal with the type of difficulty experiencedin the depiction of accurate structures ofmolecules like O3. According to the conceptof resonance, whenever a single Lewis structure cannot describe a moleculeaccurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.
Thus for O3, the two structures shown above constitute the canonical structures or resonance structures and their hybrid i.e., theIII structure represents the structure of O3more accurately. This is also called resonance hybrid. Resonance is represented by a double headed arrow. In general, it may be stated that
- Resonance stabilizes the molecule as the energy of the resonance hybrid is lessthan the energy of any single cannonical structure; and,
- Resonance averages the bond characteristics as a whole. Thus the energy of theO3resonancehybrid is lower than either of the two cannonical froms I and II .
[A] MCQ
1) Which of the following technique use to measure bond length?
(a) spectroscopic techniques
(b) X-ray diffraction
(c) electron-diffraction techniques
(d) All the above
Ans- d) All the above
2) The unit of bond enthalpy is …
(a) kJ mol–1
(b) Cal mol-1
(c) Cal mol
(d) kJ mol
Ans- a) kJ mol–1
3) With increase in bond order, bond enthalpy … and bond length ….
(a) decreases , decreases
(b) increases , decreases
(c) increases , increases
(d) decreases , increases
Ans- b) increases , decreases
4) The …. is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation.
(a) ionic radius
(b) Metallic radius
(c) covalent radius
(d) None of above
Ans- c) covalent radius
5) … is given by the number of bonds between the two atoms in a molecule.
(a) Bond Order
(b) Bond size
(c) Bond enthalpy
(d) Bond angle
Ans- a) Bond order
[B] Short Answers
1) What are exceptions to octate rule ?
Ans- There are three types of exceptions to the octet rule.
- The incomplete octet of the central atom
- Odd-electron molecules
- The expanded octet From the Kössel and Lewis treatment of the formation of an ionic bond, it follows that the formation of ionic compounds would primarily depend upon:
- The ease of formation of the positive and negative ions from the respective neutral atoms;
- The arrangement of the positive and negative ions in the solid, that is, the lattice of the crystalline compound.
2) Explain the Lattice Enthalpy of an ionic solid with suitable example .
Ans- The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJ mol–1. This means that 788 kJ of energy is required to separate one mole of solid NaCl into one mole of Na+ (g) and one mole of Cl– (g) to an infinite distance.
3) What is bond enthalpy ? Give an example
Ans- Bond Enthalpy It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. The unit of bond enthalpy is kJ mol–1. For example, the H – H bond enthalpy in hydrogen molecule is 435.8 kJ mol–1 . H2 (g) → H(g) + H(g); ∆aH = 435.8 kJ mol–1
[C] Long Answers
1) Explain the given term – formal charge .
Ans- Lewis dot structures, in general, do not represent the actual shapes of the molecules. In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom. It is, however, feasible to assign a formal charge on each atom. The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It is expressed as : Generally the lowest energy structure is the one with the smallest formal charges on the atoms. The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighbouring atoms.
2) What is bond angle ?
Ans – Bond Angle is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. Bond angle is expressed in degree which can be experimentally determined by spectroscopic methods. It gives some idea regarding the distribution of orbitals around the central atom in a molecule/complex ion and hence it helps us in determining its shape. For example H–O–H bond angle in water can be represented as under :
3) Explain the structure of O3 in terms of resonance.
Ans- The concept of resonance was introduced to deal with the type of difficulty experienced in the depiction of accurate structures of molecules like O3 . According to the concept of resonance, whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.
Thus for O3 , the two structures shown above constitute the canonical structures or resonance structures and their hybrid i.e., the III structure represents the structure of O3 more accurately. This is also called resonance hybrid.
Case –III
When covalent bond is formed betweentwo similar atoms, for example in H2,O2, Cl2,N2Or F2, the shared pair of electrons is equally Attracted by the two atoms. As a result electronPair is situated exactly between the twoIdentical nuclei. The bond so formed is calledNonpolar covalent bond. As a result of polarisation, the moleculePossesses the dipole moment which can be defined as the productOf the magnitude of the charge and theDistance between the centres of positive andNegative charge. It is usually designated by aGreek letter ‘µ’. Mathematically, it is expressedAs follows :Dipole moment (µ) = charge (Q) × distance ofSeparationDipole moment is usually expressed inDebye units (D). The conversion factor is1 D = 3.33564×10–30 C mWhere C is coulomb and m is meter. Just as all the covalent bonds haveSome partial ionic character, the ionicBonds also have partial covalentCharacter. The partial covalent character of ionic bonds was discussed by Fajans in terms of the following rules:
- The smaller the size of the cation and theLarger the size of the anion, the greater theCovalent character of an ionic bond.
- The greater the charge on the cation, theGreater the covalent character of the ionic bond.
- For cations of the same size and charge,The one, with electronic configuration(n-1)d0ns0, typical of transition metals, isMore polarising than the one with a nobleGas configuration, ns2 np6, typical of alkali and alkaline earth metal cations.
Sidgwick and Powell in 1940, proposed a simple theoryBased on the repulsive interactions of theElectron pairs in the valence shell of the atoms.It was further developed and redefined byNyholm and Gillespie (1957).The main postulates of VSEPR theory areAs follows:
- The shape of a molecule depends uponThe number of valence shell electron pairs(bonded or nonbonded) around the centralAtom.
- Pairs of electrons in the valence shell repelone another since their electron clouds arenegatively charged.
- These pairs of electrons tend to occupySuch positions in space that minimiseRepulsion and thus maximise distanceBetween them.
- The valence shell is taken as a sphere withThe electron pairs localising on theSpherical surface at maximum distanceFrom one another.
- A multiple bond is treated as if it is a singleElectron pair and the two or three electronPairs of a multiple bond are treated as aSingle super pair.
- Where two or more resonance structuresCan represent a molecule, the VSEPRModel is applicable to any such structure.
The arrangement of electron pairs and the atoms around the central atom can be : linear,Trigonal planar, tetrahedral, trigonal-Bipyramidal and octahedral . Valence bond theory was introduced byHeitler and London (1927) and developedFurther by Pauling and others. A discussionOf the valence bond theory is based on the knowledge of atomic orbitals, electronicConfigurations of elements.partialmerging of atomic orbitals is called overlappingof atomic orbitals which results in the pairingof electrons. The extent of overlap decides thestrength of a covalent bond . according toorbital overlap concept, the formation of acovalent bond between two atoms results bypairing of electrons present in the valence shellhaving opposite spins. When orbitals of two atoms come close to formbond, their overlap may be positive, negativeor zero depending upon the sign anddirection of orientation of amplitude of orbitalwave function in space. Positive andnegative sign on boundary surface diagramsin the show the sign (phase) of orbitalwave function and are not related to charge.Orbitals forming bond should have same sign(phase) and orientation in space. This is calledpositive overlap. The criterion of overlap, as the main factorfor the formation of covalent bonds appliesuniformly to the homonuclear/heteronucleardiatomic molecules and polyatomic molecules.
[A] MCQ
1) Dipole moment is usually expressed in….
(a) Debye
(b) Centimeter
(c) Columbs
(d) Ergs
Ans- a) Debye
2) 1 D = …
(a) 33564×10–28 C m
(b) 33564×10–30 C m
(c) 33564×10–32 C m
(d) 33564×10–34 C m
Ans- b) 3.33564×10–30 C m
3) Valence bond theory was introduced by ….
(a) Pauling and lewis
(b) Nyholm and Gillespie
(c) Heitler and London
(d) Sidgwick and Powell
Ans- c) Sidgwick and Powell
4) Pair is situated exactly between the two Identical nuclei the bond so formed is called …. covalent bond.
(a) Unipolar
(b) Bipolar
(c) Polar
(d) nonpolar
Ans- d) nonpolar
5) Pairs of electrons in the valence shell … one another since their electron clouds are negatively charged.
(a) Attract
(b) Repel
(c) Both a) & b)
(d) None if above
Ans- b) repel
[B] Short Answers
1) What is meant by non-polar covalent bond ?
Ans-When covalent bond is formed between two similar atoms, the shared pair of electrons is equally attracted by the two atoms. as a result electron Pair is situated exactly between the two Identical nuclei. The bond so formed is called Nonpolar covalent bond.
2) Write the fajans rule for the partial covalent character
of ionic bonds .
Ans- The partial covalent character of ionic bonds was discussed by Fajans in terms of the following rules:
- The smaller the size of the cation and the Larger the size of the anion, the greater the Covalent character of an ionic bond.
- The greater the charge on the cation, the Greater the covalent character of the ionic bond.
- For cations of the same size and charge, The one, with electronic configuration (n-1)d0ns0 , typical of transition metals, is More polarising than the one with a noble Gas configuration, ns2 np6 , typical of alkali and alkaline earth metal cations.
3) What is overlapping of atomic orbitals ? What is positive overlapping ?
Ans- Partial merging of atomic orbitals is called overlapping of atomic orbitals which results in the pairing of electrons. Orbitals forming bond have same sign (phase) and orientation in space. This is called positive overlap.
[C] Long Answers
1) Give the postulates of VSEPR theory.
Ans- The postulates of VSEPR theory are as follows:
- The shape of a molecule depends upon The number of valence shell electron pairs (bonded or nonbonded) around the central Atom.
- Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
- These pairs of electrons tend to occupy Such positions in space that minimise Repulsion and thus maximise distance Between them.
- The valence shell is taken as a sphere with The electron pairs localising on the Spherical surface at maximum distance From one another.
- A multiple bond is treated as if it is a single Electron pair and the two or three electron Pairs of a multiple bond are treated as a Single super pair.
- Where two or more resonance structures Can represent a molecule, the VSEPR Model is applicable to any such structure.
2) Explain the term – Dipole moment.
Ans-The dipole moment which can be defined as the product Of the magnitude of the charge and the Distance between the centres of positive and Negative charge. It is usually designated by a Greek letter ‘µ’.Mathematically, it is expressed As follows: Dipole moment (µ) = charge (Q) × distance of Separation Dipole moment is usually expressed in Debye units (D). The conversion factor is 1D = 3.33564×10–30C m Where C is coulomb and m is meter.
Case –IV
The covalent bond may be classified into twotypes depending upon the types ofoverlapping:(i) Sigma(σ) bond, and (ii) pi(π) bond
i) Sigma(σ) bond : This type of covalent bondis formed by the end to end (head-on)overlap of bonding orbitals along theinternuclear axis. This is called as headon overlap or axial overlap. This can beformed by any one of the following typesof combinations of atomic orbitals.
s-s overlapping : In this case, there isoverlap of two half filled s-orbitals alongthe internuclear axis.
s-p overlapping: This type of overlapoccurs between half filled s-orbitals of oneatom and half filled p-orbitals of anotheratom.
p–p overlapping : This type of overlaptakes place between half filled p-orbitalsof the two approaching atoms.
(ii) pi(π ) bond : In the formation of π bondthe atomic orbitals overlap in such a waythat their axes remain parallel to each otherand perpendicular to the internuclear axis.The orbitals formed due to sidewiseoverlapping consists of two saucer type charged clouds above and below the planeof the participating atoms.
Basically the strength of a bond depends uponthe extent of overlapping. In case of sigma bond,the overlapping of orbitals takes place to alarger extent. Hence, it is stronger as comparedto the pi bond where the extent of overlappingoccurs to a smaller extent. Further, it isimportant to note that in the formation ofmultiple bonds between two atoms of amolecule, pi bond(s) is formed in addition to asigma bond. In order to explain the characteristicgeometrical shapes of polyatomic moleculeslike CH4,NH3 and H2O etc., Pauling introducedthe concept of hybridisation. According to himthe atomic orbitals combine to form new set ofequivalent orbitals known as hybrid orbitals.Unlike pure orbitals, the hybrid orbitals areused in bond formation. The phenomenon isknown as hybridisation which can be definedas the process of intermixing of the orbitals ofslightly different energies so as to redistributetheir energies, resulting in the formation of newset of orbitals of equivalent energies and shape.For example when one 2s and three 2p-orbitalsof carbon hybridise, there is the formation offour new sp3 hybrid orbitals. Salient features of hybridisation: The mainfeatures of hybridisation are as under :
- The number of hybrid orbitals is equal tothe number of the atomic orbitals that gethybridised.
- The hybridised orbitals are alwaysequivalent in energy and shape.
- The hybrid orbitals are more effective informing stable bonds than the pure atomicorbitals.
- These hybrid orbitals are directed in spacein some preferred direction to haveminimum repulsion between electronpairs and thus a stable arrangement.Therefore, the type of hybridisationindicates the geometry of the molecules. Important conditions for hybridisation
- The orbitals present in the valence shell of the atom are hybridised.
- The orbitals undergoing hybridisation should have almost equal energy.
- Promotion of electron is not essential condition prior to hybridisation.
- It is not necessary that only half filled orbitals participate in hybridisation.
some cases, even filled orbitals of valence shell take part in hybridisation.
There are various types of hybridisationinvolving s, p and d orbitals. The differenttypes of hybridisation are as under:
(i) sp hybridisation: This type ofhybridisation involves the mixing of one s andone p orbital resulting in the formation of twoequivalent sp hybrid orbitals. The suitableorbitals for sp hybridisation are s and pz, ifthe hybrid orbitals are to lie along the z-axis. Example of molecule having sphybridisationBeCl2: The ground state electronicconfiguration of Be is 1s22s2. In the exited stateone of the 2s-electrons is promoted to vacant 2p orbital to account for its bivalency.One 2s and one 2p-orbital gets hybridised toform two sp hybridised orbitals.
(ii) sp2 hybridisation : In this hybridisationthere is involvement of one s and twop-orbitals in order to form three equivalent sp2hybridised orbitals. For example, in BCl3molecule, the ground state electronicconfiguration of central boron atom is1s22s22p1. In the excited state, one of the 2selectrons is promoted to vacant 2p orbital as a result boron has three unpaired electrons.These three orbitals (one 2s and two 2p)hybridise to form three sp2 hybrid orbitals.
(iii) sp3 hybridisation: This type ofhybridisation can be explained by taking theexample of CH4 molecule in which there ismixing of one s-orbital and three p-orbitals ofthe valence shell to form four sp3 hybrid orbitalof equivalent energies and shape. There is 25%s-character and 75% p-character in each sp3hybrid orbital. The four sp3 hybrid orbitals soformed are directed towards the four cornersof the tetrahedron. The angle between sp3hybrid orbital is 109.5° .
[A] MCQ
1) … introduced the concept of hybridisation.
(a) Pauling
(b) Lewis
(c) Nyholm
(d) Gillespie
Ans- a) Pauling
2) Which of the following is an example of sp3 hybridization .?
(a) BeCl2
(b) Ch4
(c) BCl3
(d) C2H4
Ans- b) CH4
3) The angle between sp3 hybrid orbital is ….
(a) 5°
(b) 9°
(c) 5°
(d) 120°
Ans- c) 109.5°
4) A sigma bond is formed by the overlapping of …
(a) s−s,
(b) s−p
(c) p−p
(d) All the above
Ans- d) All the above
5) when one 2s and three 2p-orbitals of carbon hybridise, there is the formation of four new … hybrid orbitals.
(a) sp3
(b) sp2
(c) sp
(d) None of above
Ans- a) sp3
[B] Short Answers
1) What are important conditions of hybridisation?
Ans- Important conditions for hybridisation
The orbitals present in the valence shell of the atom are hybridised.
The orbitals undergoing hybridisation should have almost equal energy.
Promotion of electron is not essential condition prior to hybridisation.
It is not necessary that only half filled orbitals participate in hybridisation.
some cases, even filled orbitals of valence shell take part in hybridisation.
2) Explain the term – Hybridisation.
Ans- The atomic orbitals combine to form new set of equivalent orbitals known as hybrid orbitals. Unlike pure orbitals, the hybrid orbitals are used in bond formation. The phenomenon is known as hybridisation which can be defined as the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape.
3) What are the main features of hybridisation ?
Ans- The main features of hybridisation are as under :
- The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.
- The hybridised orbitals are always equivalent in energy and shape.
- The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.
- These hybrid orbitals are directed in space in some preferred direction to have minimum repulsion between electron pairs and thus a stable arrangement. Therefore, the type of hybridisation indicates the geometry of the molecules.
[C] Long Answers
1) Which are different types of hybridization ?
Ans- There are various types of hybridisation involving s, p and d orbitals. The different types of hybridisation are as under:
1) Sp hybridisation: This type of hybridisation involves the mixing of one s and one p orbital resulting in the formation of two equivalent sp hybrid orbitals. The suitable orbitals for sp hybridisation are s and pz , if the hybrid orbitals are to lie along the z-axis. Example of molecule having sp hybridisation BeCl2 : The ground state electronic configuration of Be is1s22s2 . In the exited state one of the 2s-electrons is promoted to vacant 2p orbital to account for its bivalency. One 2s and one 2p-orbital gets hybridised to form two sp hybridised orbitals.
2) sp2 hybridisation : In this hybridisation there is involvement of one s and two p-orbitals in order to form three equivalent sp2 hybridised orbitals. For example, in BCl3 molecule, the ground state electronic configuration of central boron atom is 1s 1s22s22p1 . In the excited state, one of the 2s electrons is promoted to vacant 2p orbital as a result boron has three unpaired electrons. These three orbitals (one 2s and two 2p) hybridise to form three sp2 hybrid orbitals.
3) Sp3 hybridisation: This type of hybridisation can be explained by taking the example of CH4 molecule in which there is mixing of one s-orbital and three p-orbitals of the valence shell to form four sp3 hybrid orbital of equivalent energies and shape. There is 25% s-character and 75% p-character in each sp3 hybrid orbital. The four sp3 hybrid orbitals so formed are directed towards the four corners of the tetrahedron. The angle between sp3 hybrid orbital is 109.5°.
2) Give the classification of the covalent bond depending upon the types of overlapping.
Ans-The covalent bond may be classified into two types depending upon the types of overlapping: (i) Sigma(σ) bond, and (ii) pi(π) bond
- Sigma(σ) bond : This type of covalent bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. This is called as head on overlap or axial overlap. This can be formed by any one of the following types of combinations of atomic orbitals.
s-s overlapping : In this case, there is overlap of two half filled s-orbitals along. the internuclear axis.
s-p overlapping: This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom.
p–p overlapping : This type of overlap takes place between half filled p-orbitals of the two approaching atoms.
(ii) pi(π ) bond : In the formation of π bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbitals formed due to sidewise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms.
Case Study Questions Class 11 Chemistry Chemical Bonding and Molecular Structure PDF