Chemical Bonding and Molecular Structure : Notes and Study Materials -pdf
- Concepts of Chemical Bonding and Molecular Structure
- Chemical Bonding and Molecular Structure Master File
- Chemical Bonding and Molecular Structure Revision Notes
- Chemical Bonding and Molecular Structure MindMap
- NCERT Solution Chemical Bonding and Molecular Structure
- NCERT Exemplar Solution Chemical Bonding and Molecular Structure
- Chemical Bonding and Molecular Structure: Solved Example 1
- Chemical Bonding and Molecular Structure: Solved Example 2
- Chemical Bonding and Molecular Structure : Practice Paper 1
- Chemical Bonding and Molecular Structure : Practice Paper 2
- Chemical Bonding and Molecular Structure : Practice Paper 3
Subtopics of Class 11 Chemistry Chapter 4: Chemical Bonding and Molecular Structure
- Octet Rule
- Covalent Bond
- Lewis Representation Of Simple Molecules (The Lewis Structures)
- Formal Charge
- Limitations Of The Octet Rule
- Ionic Or Electrovalent Bond
- Lattice Enthalpy
- Bond Parameters
- Bond Length
- Bond Angle
- Bond Enthalpy
- Bond Order
- Resonance Structures
- Polarity Of Bonds
- The Valence Shell Electron Pair Repulsion (VSEPR) Theory
- Valence Bond Theory
- Orbital Overlap Concept
- Directional Properties Of Bonds
- Overlapping Of Atomic Orbitals
- Types Of Overlapping And Nature Of Covalent Bonds
- The Strength Of Sigma And Pi Bonds
- Types Of Hybridisation
- Other Examples Of Sp3, Sp2 And Sp Hybridisation
- The Hybridisation Of Elements Involving D Orbitals
- Molecular Orbital Theory
- Formation Of Molecular Orbitals Linear Combination Of Atomic Orbitals (Lcao)
- Conditions For The Combination Of Atomic Orbitals
- Types Of Molecular Orbitals
- Energy Level Diagram For Molecular Orbitals
- Electronic Configuration And Molecular Behaviour
- Bonding In Some Homonuclear Diatomic Molecules
- Hydrogen Bonding
- Cause Of Formation Of Hydrogen Bond
- Types Of H-bonds.
Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4
• Chemical Bond
The force that holds different atoms in a molecule is called chemical bond.
• Octet Rule
Atoms of different elements take part in chemical combination in order to complete their octet or to attain the noble gas configuration.
• Valence Electrons
It is the outermost shell electron which takes part in chemical combination.
• Facts Stated by Kossel in Relation to Chemical Bonding
— In the periodic table, the highly electronegative halogens and the highly electro-positive alkali metals are separated by noble gases.
— Formation of an anion and cation by the halogens and alkali metals are formed by gain of electron and loss of electron respectively.
— Both the negative and positive ions acquire the noble gas configuration.
— The negative and positive ions are stabilized by electrostatic attraction Example,
• Modes of Chemical Combination
— By the transfer of electrons: The chemical bond which formed by the complete transfer of one or more electrons from one atom to another is termed as electrovalent bond or ionic bond.
— By sharing of electrons: The bond which is formed by the equal sharing of electrons between one or two atoms is called covalent bond. In these bonds electrons are contributed by both.
— Co-ordinate bond: When the electrons are contributed by one atom and shared by both, the bond is formed and it is known as dative bond or co-ordinate bond.
• Ionic or Electrovalent Bond
Ionic or Electrovalent bond is formed by the complete transfer of electrons from one atom to another. Generally, it is formed between metals and non-metals. We can say that it is the electrostatic force of attraction which holds the oppositely charged ions together.
The compounds which is formed by ionic or electrovalent bond is known as electrovalent compounds. For Example, ,
(i) NaCl is an electrovalent compound. Formation of NaCl is given below:
Na+ ion has the configuration of Ne while Cl– ion represents the configuration of Ar.
(ii) Formation of magnesium oxide from magnesium and oxygen.
Electrovalency: Electrovalency is the number of electrons lost or gained during the formation of an ionic bond or electrovalent bond.
• Factors Affecting the Formation of Ionic Bond
(i) Ionization enthalpy: As we know that ionization enthalpy of any element is the amount of energy required to remove an electron from outermost shell of an isolated gaseous atom to convert it into cation.
Hence, lesser the ionization enthalpy, easier will be the formation of a cation and have greater chance to form an ionic bond. Due to this reason alkali metals have more tendency to form an ionic bond.
For example, in formation of Na+ ion I.E = 496 kJ/mole
While in case of magnesium, it is 743 kJ/mole. That’s why the formation of positive ion for sodium is easier than that of magnesium.
Therefore, we can conclude that lower the ionization enthalpy, greater the chances of ionic bond formation.
(ii) Electron gain enthalpy (Electron affinities): It is defined as the energy released when an isolated gaseous atom takes up an electron to form anion. Greater the negative electron gain enthalpy, easier will be the formation of anion. Consequently, the probability of formation of ionic bond increases.
For example. Halogens possess high electron affinity. So, the formation of anion is very common in halogens.
(iii) Lattice energy or enthalpy: It is defined as the amount of energy required to separate 1 mole of ionic compound into separate oppositely charged ions.
Lattice energy of an ionic compound depends upon following factors:
(i) Size of the ions: Smaller the size, greater will be the lattice energy.
(ii) Charge on the ions: Greater the magnitude of charge, greater the interionic attraction and hence higher the lattice energy.
• General Characteristics of ionic Compounds
(i) Physical’State: They generally exist as crystalline solids, known as crystal lattice. Ionic compounds do not exist as single molecules like other gaseous molecules e.g., H2 , N2 , 02 , Cl2 etc.
(ii) Melting and boiling points: Since ionic compounds contain high interionic force between them, they generally have high melting and boiling points.
(iii) Solubility: They are soluble in polar solvents such as water but do not dissolve in organic solvents like benzene, CCl4etc.
(iv) Electrical conductivity: In solid state they are poor conductors of electricity but in molten state or when dissolved in water, they conduct electricity.
(v) Ionic reactions: Ionic compounds produce ions in the solution which gives very fast reaction with oppositely charged ions.
• Covalent Bond—Lewis-Langmuir Concept
When the bond is formed between two or more atoms by mutual contribution and sharing of electrons, it is known as covalent bond.
If the combining atoms are same the covalent molecule is known as homoatomic. If they are different, they are known as heteroatomic molecule.
• Lewis Representation of Simple Molecules (the Lewis Structures)
The Lewis dot Structure can be written through the following steps:
(i) Calculate the total number of valence electrons of the combining atoms.
(ii) Each anion means addition of one electron and each cation means removal of one electron. This gives the total number of electrons to be distributed.
(iii) By knowing the chemical symbols of the combining atoms.
(iv) After placing shared pairs of electrons for single bond, the remaining electrons may account for either multiple bonds or as lone pairs. It is to be noted that octet of each atom should be completed.
• Formal Charge
In polyatomic ions, the net charge is the charge on the ion as a whole and not by particular atom. However, charges can be assigned to individual atoms or ions. These are called formal charges.
It can be expressed as
• Limitations of the Octet Rule
(i) The incomplete octet of the central atoms: In some covalent compounds central atom has less than eight electrons, i.e., it has an incomplete octet. For example,
Li, Be and B have 1, 2, and 3 valence electrons only.
(ii) Odd-electron molecules: There are certain molecules which have odd number of electrons the octet rule is not applied for all the atoms.
(iii) The expanded Octet: In many compounds there are more than eight valence electrons around the central atom. It is termed as expanded octet. For Example,
• Other Drawbacks of Octet Theory
(i) Some noble gases, also combine with oxygen and fluorine to form a number of compounds like XeF2 , XeOF2 etc.
(ii) This theory does not account for the shape of the molecule.
(iii) It does not give any idea about the energy of The molecule and relative stability.
• Bond Length
It is defined as the equilibrium distance between the centres of the nuclei of the two bonded atoms. It is expressed in terms of A. Experimentally, it can be defined by X-ray diffraction or electron diffraction method.
• Bond Angle
It is defined as -the angle between the lines representing the orbitals containing the bonding – electrons.
It helps us in determining the shape. It can be expressed in degree. Bond angle can be experimentally determined by spectroscopic methods.
• Bond Enthalpy
It is defined as the amount of energy required to break one mole of bonds of a particular type to separate them into gaseous atoms.
Bond Enthalpy is also known as bond dissociation enthalpy or simple bond enthalpy. Unit of bond enthalpy = kJ mol-1
Greater the bond enthalpy, stronger is the bond. For e.g., the H—H bond enthalpy in hydrogen is 435.8 kJ mol-1.
The magnitude of bond enthalpy is also related to bond multiplicity. Greater the bond multiplicity, more will be the bond enthalpy. For e.g., bond enthalpy of C —C bond is 347 kJ mol-1 while that of C = C bond is 610 kJ mol-1.
In polyatomic molecules, the term mean or average bond enthalpy is used.
• Bond Order
According to Lewis, in a covalent bond, the bond order is given by the number of bonds between two atoms in a molecule. For example,
Bond order of H2 (H —H) =1
Bond order of 02 (O = O) =2
Bond order of N2 (N = N) =3
Isoelectronic molecules and ions have identical bond orders. For example, F2 and O22- have bond order = 1. N2, CO and NO+ have bond order = 3. With the increase in bond order, bond enthalpy increases and bond length decreases. For example,
• Resonance Structures
There are many molecules whose behaviour cannot be explained by a single-Lew is structure, Tor example, Lewis structure of Ozone represented as follows:
Thus, according to the concept of resonance, whenever a single Lewis structure cannot explain all the properties of the molecule, the molecule is then supposed to have many structures with similar energy. Positions of nuclei, bonding and nonbonding pairs of electrons are taken as the canonical structure of the hybrid which describes the molecule accurately. For 03, the two structures shown above are canonical structures and the III structure represents the structure of 03 more accurately. This is also called resonance hybrid.
Some resonating structures of some more molecules and ions are shown as follows:
• Polarity of Bonds
Polar and Non-Polar Covalent bonds
Non-Polar Covalent bonds: When the atoms joined by covalent bond are the same like; H2, 02, Cl2, the shared pair of electrons is equally attracted by two atoms and thus the shared electron pair is equidistant to both of them.
Alternatively, we can say that it lies exactly in the centre of the bonding atoms. As a result, no poles are developed and the bond is called as non-polar covalent bond. The corresponding molecules are known as non-polar molecules.
Polar bond: When covalent bonds formed between different atoms of different electronegativity, shared electron pair between two atoms gets displaced towards highly electronegative atoms.
For Example, in HCl molecule, since electronegativity of chlorine is high as compared to hydrogen thus, electron pair is displaced more towards chlorine atom, thus chlorine will acquire a partial negative charge (δ–) and hydrogen atom have a partial positive charge (δ+) with the magnitude of charge same as on chlorination. Such covalent bond is called polar covalent bond.
• Dipole Moment
Due to polarity, polar molecules are also known as dipole molecules and they possess dipole moment. Dipole moment is defined as the product of magnitude of the positive or negative charge and the distance between the charges.
• Applications of Dipole Moment
(i) For determining the polarity of the molecules.
(ii) In finding the shapes of the molecules.
For example, the molecules with zero dipole moment will be linear or symmetrical. Those molecules which have unsymmetrical shapes will be either bent or angular.
(e.g., NH3with μ = 1.47 D).
(iii) In calculating the percentage ionic character of polar bonds.
• The Valence Shell Electron Pair Repulsion (VSEPR) Theory
Sidgwick and Powell in 1940, proposed a simple theory based on repulsive character of electron pairs in the valence shell of the atoms. It was further developed by Nyholm and Gillespie (1957).
Main Postulates are the following:
(i) The exact shape of molecule depends upon the number of electron pairs (bonded or non bonded) around the central atoms.
(ii) The electron pairs have a tendency to repel each other since they exist around the central atom and the electron clouds are negatively charged.
(iii) Electron pairs try to take such position which can minimize the rupulsion between them.
(iv) The valence shell is taken as a sphere with the electron pairs placed at maximum distance.
(v) A multiple bond is treated as if it is a single electron pair and the electron pairs which constitute the bond as single pairs.
• Valence Bond Theory
Valence bond theory was introduced by Heitler and London (1927) and developed by Pauling and others. It is based on the concept of atomic orbitals and the electronic configuration of the atoms.
Let us consider the formation of hydrogen molecule based on valence-bond theory.
Let two hydrogen atoms A and B having their nuclei NA and NB and electrons present in them are eA and eB .
As these two atoms come closer new attractive and repulsive forces begin to operate.
(i) The nucleus of one atom is attracted towards its own electron and the electron of the other and vice versa.
(ii) Repulsive forces arise between the electrons of two atoms and nuclei of two atoms. Attractive forces tend to bring the two atoms closer whereas repulsive forces tend to push them apart.
• Orbital Overlap Concept
According to orbital overlap concept, covalent bond formed between atoms results in the overlap of orbitals belonging to the atoms having opposite spins of electrons. Formation of hydrogen molecule as a result of overlap of the two atomic orbitals of hydrogen atoms is shown in the figures that follows:
Stability of a Molecular orbital depends upon the extent of the overlap of the atomic orbitals.
• Types of Orbital Overlap
Depending upon the type of overlapping, the covalent bonds are of two types, known as sigma (σ ) and pi (π) bonds.
(i) Sigma (σ bond): Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis.
The axial overlap involving these orbitals is of three types:
• s-s overlapping: In this case, there is overlap of two half-filled s-orbitals along the internuclear axis as shown below:
• s-p overlapping: This type of overlapping occurs between half-filled s-orbitals of one atom and half filled p-orbitals of another atoms.
• p-p overlapping: This type of overlapping takes place between half filled p-orbitals of the two approaching atoms.
(ii) pi (π bond): π bond is formed by the atomic orbitals when they overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis.The orbital formed is due to lateral overlapping or side wise overlapping.
• Strength of Sigma and pf Bonds
Sigma bond (σ bond) is formed by the axial overlapping of the atomic orbitals while the π-bond is formed by side wise overlapping. Since axial overlapping is greater as compared to side wise. Thus, the sigma bond is said to be stronger bond in comparison to a π-bond.
Distinction between sigma and n bonds
Hybridisation is 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.
Salient Features of Hybridisation:
(i) Orbitals with almost equal energy take part in the hybridisation.
(ii) Number of hybrid orbitals produced is equal to the number of atomic orbitals mixed,
(iii) Geometry of a covalent molecule can be indicated by the type of hybridisation.
(iv) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.
Conditions necessary for hybridisation:
(i) Orbitals of valence shell take part in the hybridisation.
(ii) Orbitals involved in hybridisation should have almost equal energy.
(iii) Promotion of electron is not necessary condition prior to hybridisation.
(iv) In some cases filled orbitals of valence shell also take part in hybridisation.
Types of Hybridisation:
(i) sp hybridisation: When one s and one p-orbital hybridise to form two equivalent orbitals, the orbital is known as sp hybrid orbital, and the type of hybridisation is called sp hybridisation.
Each of the hybrid orbitals formed has 50% s-characer and 50%, p-character. This type of hybridisation is also known as diagonal hybridisation.
(ii) sp2 hybridisation: In this type, one s and two p-orbitals hybridise to form three equivalent sp2 hybridised orbitals.
All the three hybrid orbitals remain in the same plane making an angle of 120°. Example. A few compounds in which sp2 hybridisation takes place are BF3, BH3, BCl3 carbon compounds containing double bond etc.
(iii) sp3 hybridisation: In this type, one s and three p-orbitals in the valence shell of an atom get hybridised to form four equivalent hybrid orbitals. There is 25% s-character and 75% p-character in each sp3 hybrid orbital. The four sp3 orbitals are directed towards four corners of the tetrahedron.
The angle between sp3 hybrid orbitals is 109.5°.
A compound in which sp3 hybridisation occurs is, (CH4). The structures of NH2 and H20 molecules can also be explained with the help of sp3 hybridisation.
• Formation of Molecular Orbitals: Linear Combination of Atomic Orbitals (LCAO)
The formation of molecular orbitals can be explained by the linear combination of atomic orbitals. Combination takes place either by addition or by subtraction of wave function as shown below.
The molecular orbital formed by addition of atomic orbitals is called bonding molecular orbital while molecular orbital formed by subtraction of atomic orbitals is called antibonding molecular orbital.
Conditions for the combination of atomic orbitals:
(1) The combining atomic orbitals must have almost equal energy.
(2) The combining atomic orbitals must have same symmetry about the molecular axis.
(3) The combining atomic orbitals must overlap to the maximum extent.
• Types of Molecular Orbitals
Sigma (σ) Molecular Orbitals: They are symmetrical around the bond-axis.
pi (π) Molecular Orbitals: They are not symmetrical, because of the presence of positive lobes above and negative lobes below the molecular plane.
• Electronic configuration and Molecular Behaviour
The distribution of electrons among various molecular orbitals is called electronic configuration of the molecule.
• Stability of Molecules
• Bond Order
Bond order is defined as half of the difference between the number of electrons present in bonding and antibonding molecular orbitals.
Bond order (B.O.) = 1/2 [Nb-Na]
The bond order may be a whole number, a fraction or even zero.
It may also be positive or negative.
Nature of the bond: Integral bond order value for single double and triple bond will be 1, 2 and 3 respectively.
Bond-Length: Bond order is inversely proportional to bond-length. Thus, greater the bond order, smaller will be the bond-length.
Magnetic Nature: If all the molecular orbitals have paired electrons, the substance is diamagnetic. If one or more molecular orbitals have unpaired electrons, it is paramagnetic e.g., 02 molecule.
• Bonding in Some Homonuclear (Diatomic) Molecules
(1) Hydrogen molecule (H2): It is formed by the combination of two hydrogen atoms. Each hydrogen atom has one electron in Is orbital, so, the electronic configuration of hydrogen molecule is
This indicates that two hydrogen atoms are bonded by a single covalent bond. Bond dissociation energy of hydrogen has been found = 438 kJ/mole. Bond-Length = 74 pm
No unpaired electron is present therefore,, it is diamagnetic.
(2) Helium molecule (He2): Each helium atom contains 2 electrons, thus in He2 molecule there would be 4 electrons.
The electrons will be accommodated in σ1s and σ*1s molecular orbitals:
• Hydrogen Bonding
When highly electronegative elements like nitrogen, oxygen, flourine are attached to hydrogen to form covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. Thus, partial positive charge develops on hydrogen atom which forms a bond with the other electronegative atom. This bond is known as hydrogen bond and it is weaker than the covalent bond. For example, in HF molecule, hydrogen bond exists between hydrogen atom of one molecule and fluorine atom of another molecule.
It can be depicted as
• Types of H-Bonds
(i) Intermolecular hydrogen bond (ii) Intramolecular hydrogen bond.
(i) Intermolecular hydrogen bond: It is formed between two different molecules of the same or different compounds. For Example, in HF molecules, water molecules etc.
(ii) Intramolecular hydrogen bond: In this type, hydrogen atom is in between the two highly electronegative F, N, O atoms present within the same molecule. For example, in o-nitrophenol, the hydrogen is in between the two oxygen atoms.
CBSE Class 11 Chemistry Chapter-4 Important Questions
1 Marks Questions
1.Define a chemical bond.
Ans. The attractive force which holds various constituents (atoms, ions etc.) together in different chemical species is called a chemical bond.
2.Give the main feature of Lewis approach of chemical bonding.
Ans. Lewis postulated that atoms achieve the stable octet when they are linked by chemical bonds. He assured that atoms are positively charged centre and the outer shell that could accommodate a maximum of eight electrons. These electrons occupy the corners of a cube which surrounds the centre. Lewis introduced simple notations to represent valence electrons in an atom called Lewis symbol
3.Write electron dot structure (Lewis structure) of Na, Ca, B, Br, Xe, As, Ge, N3-.
4.Give the octet rule in short.
Ans.The atoms tend to adjust the arrangement of their electrons in such a way that they (except H and He) achieve eight electrons in their outermost shell. This is known as the octet rule.
5.Define an ionic bonding. [?]
Ans.An ionic bond (or electrovalent bond) is formed by a complete transfer of one or more of outer most electrons from the atom of a metal to that of a non – metal.
6.Which one of the following has the highest bond order?N2, N2+ or N2–.
Ans. N2 has the highest bond order.
7.Define bond order.
Ans. Bond order is defined as number of bonds between two atoms in a molecule.
8.What type of bond is formed when atoms have high difference of electornegativity?
Ans. Electrovalent or ionic bond.
9.Define dipole moment.
Ans. Dipole moment is defined as the product of the magnitude of the charge and the distance between the centers of positive and negative charge.
10.Give the mathematical expression of dipole moment.
Ans. Mathematically dipole moment is expressed as dipole moment (M) = charge (Q) x distance of separation (r). Dipole moment is usually expressed in Debye units (D).
11.Why is dipole moment of CO2, BF3, CCl4 is zero?
Ans. Because there molecules have symmetrical shapes and thus the dipoles gets cancelled and the net dipole moment is zero.
12.Why is BF3 non – polar?
Ans. Because BF3 has symmetrical shape, the net dipole moment is zero and thus it is non – polar.
13. Write the resonating structure of O3 molecule.
14.What is sigma bond?
Ans.A covalent bond formed due to the overlap of orbitals of the two atoms along the line going the two nuclei (orbital axis) is called sigma (s) bond.
15.What is pi – bond?
Ans. A covalent bond formed between the two atoms due to the sideways overlap of their p – orbitals is called a pi () bond.
16.How many s – and – bond are there in a molecule of C2H4 (ethene )?
Ans. In a molecule of ethene, there are 5 s – bonds (one between C-C , and four between C-H and one – bond.
17.How many s – and – bonds are there in a molecule of CH2 = CH – CH = CH2 ?
Ans. There are 9 s – bonds (three between C – C and 6 between C – H) and 2 – bonds.
18.What type of bond exists in multiple bond (double / triple)?
Ans. pi (p) – bond is always present in molecules containing multiple bond.
19.What type of bond are formed due to orbital overlap?
Ans. Covalent bonds are formed due to the overlap of certain orbitals that are oriented favourably in the space.
20.How do covalent bonds form due to orbital overlapping?
Ans. According to orbital overlap concept, the formation of a covalent bond between two atoms results by pairing of electrons present in the valence shell having opposite spins.
Ans. Hybridisation is 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.
22.State the hybrid orbitals associated with B in BCl3 and C in C2H4
Ans. (i) hybridization (ii)hybridization.
23.What is the state of hybridization of carbon atoms in diamond and graphite?
Ans.In Diamond it is
In graphite it is
24.What type of hybridisation takes place in (i) p in PCL5 and (ii) S in S F6?
25.Define bonding molecular orbital.
Ans. The molecular orbital formed by the addition of atomic orbitals is called bonding molecular orbital.
26.Define antibonding molecular orbital.
Ans. The molecular orbital formed by the subtraction of atomic orbitals is called antibonding molecular orbital.
27.Explain diagrammatically the formation of molecular orbital by LCAO.
Ans. The molecular orbital formed by subtraction of atomic orbital is called antibonding molecular orbital.
28.Which one may exhibit paramagnetism?
Ans.would exhibit paramagnetism because it contains one unpaired electron in its Mo configuration.
29.Why are bonding molecular orbitals more stable than antibonding molecular orbitals?
Ans. Bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital.
30.Define bond order.
Ans. Bond order (b.o) is defined as one half the difference between the number of electrons present in the bonding and the antibonding orbitals i.e;
Bond order (b.o) =
If molecule is stable and
If molecule is unstable.
31.Define hydrogen bonding
Ans. Hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule.
32.What are the types of H-bonding? Which of them is stronger?
Ans. (i)Inter-molecular H-bonding
(ii) Intra molecular H-bonding. Inter molecular H-bonding is stronger than intra-molecular H-bonding.
33. has higher boiling point than . Give reason.
Ans.In, there is hydrogen bonding whereas in PH3 there is no hydrogen bonding.
34.Define electrovalent bond.
Ans. The bond formed, as a result of the electrostatic attraction between the positive and negative ions are termed as the electrovalent bond.
2 Marks Questions
1.Give the main feature of Kossel’s explanation of chemical bonding.
Ans. Kossel in relation to chemical bonding drew attention to the following facts –
(i) In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases.
(ii) In the formation of a negative ion from a halogen atom and a positive ion from an alkali metal, atom is associated with a gain and loss of an electron by the respective atoms.
(iii) The negative and positive ions so formed attain stable noble gas electronic configurations. The noble gases have particularly eight electrons, ns2 np6.
The –ve and +ve ions are stabilized by electrostatic attraction.
2.How can you explain the formation of NaCl according to kossel concept?
Ans. The formation of NaCl from sodium and chlorine can be explained as
Na ® Na+ + e–
[Ne] 3s1 ® [Ne]
Cl + e– ® Cl–
[Ne] 3s2 3p5 . [Ne] 3s2 3p6 or [Ar]
Na+ + Cl– ® Na+ Cl– or NaCl.
3.Write the significance of octet rule.
Ans. Octet rule signifies –
(i) It is useful for understanding the structures of most of the organic compounds.
It mainly applies to the second period elements of the periodic table.
4.Write the Lewis structure for CO molecule
Ans. (i) The outer (valence) shell configurations of carbon and oxygen atoms are
Carbon : (6) – 1s2 2s2 2p2
Oxygen : (8) – 1s2 2s2 2p4.
The valence electrons (4 + 6 = 10)
But it does not complete octet, thus multiple bond is exhibited.
(ii) N (2s2 2p3), O (2s2 2p4)
5 + (2 x 6) + 1 = 18 electrons.
5.Give the Lewis dot structure of HNO3
Ans. HNO3 ®
6.What changes are observed in atoms undergoing ionic bonding?
Ans. Due to the electron transfer the following changes occurs –
(i) Both the atoms acquire stable noble gas configuration.
(ii) The atom that loses electrons becomes +vely charged called cation whereas that gains electrons becomes –vely charged called anion.
(iii) Cation and anion are held together by the coulombic forces of attraction to form an ionic bond.
7.Mention the factors that influence the formation of an Ionic bond.
Ans.Ionic bond formation mainly depends upon three factors –
(i) Low ionization energy – elements with low ionization enthalpy have greater tendency to form an ionic bonds.
(ii) High electron gain enthalpy – high negative value of electron gain enthalpy favours ionic bond.
(iii) Lattice energy – high lattice energy value favours ionic bond formation.
8.Give reason why H2+ ions are more stable than H2– though they have the same bond order.
Ans.In H2– ion, one electron is present in anti bonding orbital due to which destabilizing effect is more and thus the stability is less than that of H2+ ion.
9.How would the bond lengths vary in the following species? C2, C2– C22-.
Ans.The order of bond lengths in C2 , C2– and C22- is C2 > C2– > C22-.
10.Out of covalent and hydrogen bonds, which is stronger.
Ans. Covalent bond.
11.Define covalent radius.
Ans. The covalent radius 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.
12.Why NH3 has high dipole moment than NF3 though both are pyramidal?
Ans. In case of NH3 the orbital dipole due to lone pair is in the same direction as the resultant dipole moment of the N-H bonds, whereas in NF3 the orbital dipole is in the direction opposite to the resultant dipole moment of the three N-F
bonds. The orbital dipole become of lone pair decreases, which results in the low dipole moment.
13.Draw the resonating structure of NO3–
14.On which factor does dipole moment depend in case of polyatomic molecules.
Ans.The dipole moment of the polyatomic molecule depends on individual dipole moments of bonds and also on the spatial arrangement of various bonds in the molecule.
15.Dipole moment of Be F2 is zero. Give reason.
Ans. In BeF2 the dipole moment is zero because the two equal bond dipoles point in opposite directions and cancel the effect of each other.
16.Bond dipoles in Be F2
Give the main features of VSEPR Theory.
Ans.The main postulates of VSEPR theory are as follows :
(i)The shape of a molecule depends upon the number of valence shell electron pairs around the central atom.
(ii)Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
(iii)There pairs of electrons tends to occupy such position in space that minimize repulsion and thus maximize distance between them.
(iv)The valence shell is taken as a sphere with the electron pairs localizing on the sphere at maximum distance from one another.
(v)A multiple bond is treated as it is a single electron pair and two or three electron pairs of a multiple bond is treated as super pair.
(vi)When two or more resonance structures can represent a molecule, the VSEPR nodal is applicable to any such structure.
17.What’s difference between lone pair and bonded pair of electrons?
Ans. Lone pair electrons do not take part in bond formation whereas bond pair electrons take part in bond formation.
18.CO2 is linear whereas SO2 is bend – shaped. Give reason.
Ans. In CO2, the bond electron are furtherest away from each other forming1800 angle. Thus, CO2 is linear.
In SO2, the number of bonding pairs is 4 where it has an lone pair of electron which does not participate in bond formation thereby repulsive strain is experienced.
19.Why does H2O have bent structure?
Ans. In water molecule, there are two bonding pairs and two lone pairs of electrons. The shape should have been tetrahedral if there were all bp but two lp are present. Thus the shaped is distorted to an angular shape. Because lp – lp repulsion is more than lp – bp repulsion.
20.For the molecule,
Why is structure (b) more stable than structure (a)?
Ans. In (a) the lp is present at axial position so there are three lp – bp repulsions at 900 . Whereas in (b) the lp is in an equatorial position are there are two lp – bp repulsions. Hence, arrangement (b) is more stable than (a).
21.How would you attribute the structure of PH3 molecule using VSEPR model?
Ans. Phosphorus atom has 5 electrons in its outermost orbit. H – atoms contribute one electron each to make in all 8 electron around P – atom. Thus 4 pairs of electrons would be distributed in a tetrahedral manner around the central atom. Three pairs from three P – H bonds while the fourth pair remains unused. Due to repulsion between the bp and lp, the shape is not of tetrahedral but trigonal pyramidal molecule.
22.In SF4 molecule, the lp electrons occupies an equatorial position in the trigonal bipyramidal arrangement to an axial position. Give reason.
Ans. In SF4 molecule, the lp electrons occupies an equatorial position because, lp – bp repulsion is minimum.
23.How is VBT different from Lewis concept?
Ans.In Lewis concept, bond formation is explained in terms of sharing of electron pairs and the Octet rule whereas in VBT bond formation is described in terms of hybridization and overlap of the orbitals.
24.S – orbital does not show any preference for direction. Why?
Ans. S – Orbital does not show any preference for direction because it is spherically symmetrical.
25.Why is s– bond stronger than – bond?
Ans. Orbitals can overlap to a greater extent in a s – bond due to axial orientation, so s – bond is strong. Whereas, in a pi – bond sideways overlapping is not to an appreciable extent due to the presence of s – bond which restricts the distance between the involved atoms.
26.What are the different types of s – bond formation?
Ans. s – bond can be formed by any of the following types of combinations of atoms orbitals.
(a) S – S – overlapping : In this case, there is a overlap of two half – filled S – orbitals along the inter nuclear axis.
(b) S- P overlapping : This type of over lapping occurs between half – filled s-orbitals of one atom and half-filled p-orbitals of another atom.
(c) P – P overlapping : This type of overlap takes place between half-filled p-orbitals of the two approaching atoms.
27.What is zero over lap?
Ans. The unsymmetrical overlap of orbitals results in zero overlap i-e; between px-s and px-py orbital
28. the features of hybridisation.
Ans. The main features of hybridization are
(i) The number of hybrid orbitals is equal to number of the atomic orbitals that get hybridized.
(ii) The hybridized orbitals are always equivalent in energy and shape.
(iii) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.
(iv) The hybrid orbitals orient in a manner to minimize repulsion resulting in a particular geometrical shape.
29.What are the important consolations for hybridisation?
Ans. (i) The orbitals present in the valence shell of the atom are hybridised.
(ii)The orbitals undergoing hybridization should have almost the same energy.
(iii)It is not essential that electrons get promoted prior to hybridization.
It is necessary that only half filled orbitals participate in hybridisation even filled orbitlals can take part.
30.Describe the shape of sp, sp2 and sp3 hybrid orbital?
Ans. (i) Sp-hybrid orbital is oriented to an angle .
(ii)-hybrid orbital lie in a plane and is directed towards the corners of equilateral triangle making an angle of .
(iii) -hybrid orbitals are directed towards the four corners of tetrahedron making an angle of
31.Ethylene is a planar molecule whereas acetylene is a linear molecule. Give reason.
Ans. In case of ethylene, hybridization where the four hydrogen atoms are placed in four corners of a plane sharing
Whereas acetylene shows sp hybridization and shares an angle of and thus it is linear.
32.In H2O, H2S, H2Se, H2Te, the bond angle decreases though all have the same bent shape. Why?
Ans. In all the four cases, the molecules undergo hybridization forming four hybrid orbitals, two of which are occupied by lp of electrons and two by bp electrons. Thus they are expected to have angle but this does not happen. In case of molecule, as oxygen is small in size and has high electronegativity value, the bp are closer due to which it is subjected to larger repulsion (bp-bp). In case of as S atom is larger than O, bp-bp repulsion is less as compared to and it is true for as well.
33.Out of p-orbital and sp-hybrid orbital which has greater directional character and Why?
Ans. Sp-hybrid orbital has greater directional character than p-orbital. Because in case of p-orbitals, the two lobes are equal in size and equal electron density is distributed whereas in Sp-hybrid orbital, electron density is greater on one side.
34.does not exist. Explain in terms of LCAO.
Ans. The electronic configuration of helium atom is. Each helium atom contains 2 electrons, therefore, in molecule there would be 4 electrons. These electrons will accommodated in molecular orbitals leading to electronic configuration :
molecule is there unstable and does not exist.
35.Dipole moment is a scalar or a vector quantity?
Ans.Dipole moment is a vector quantity and is depicted by a small arrow with tail on the +ve centre and head pointing towards the negative centre.