{"id":22456,"date":"2022-01-30T13:56:28","date_gmt":"2022-01-30T08:26:28","guid":{"rendered":"https:\/\/infinitylearn.com\/surge\/?p=22456"},"modified":"2024-01-17T15:49:52","modified_gmt":"2024-01-17T10:19:52","slug":"chemical-bonding-and-molecular-structure-class-11-notes-chemistry-chapter-4","status":"publish","type":"post","link":"https:\/\/infinitylearn.com\/surge\/study-material-notes\/class-11\/chemistry\/chemical-bonding-and-molecular-structure-class-11-notes-chemistry-chapter-4\/","title":{"rendered":"Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4"},"content":{"rendered":"

Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4<\/h2>\n

\u2022 Chemical Bond<\/strong>
\nThe force that holds different atoms in a molecule is called chemical bond.
\n\u2022 Octet Rule<\/strong>
\nAtoms of different elements take part in chemical combination in order to complete their octet or to attain the noble gas configuration.
\n\u2022 Valence Electrons<\/strong>
\nIt is the outermost shell electron which takes part in chemical combination.
\n\u2022 Facts Stated by Kossel in Relation to Chemical Bonding<\/strong>
\n\u2014 In the periodic table, the highly electronegative halogens and the highly electro-positive alkali metals are separated by noble gases.
\n\u2014 Formation of an anion and cation by the halogens and alkali metals are formed by gain of electron and loss of electron respectively.
\n\u2014 Both the negative and positive ions acquire the noble gas configuration.
\n\u2014 The negative and positive ions are stabilized by electrostatic attraction Example,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-1\"
\n\u2022 Modes of Chemical Combination<\/strong>
\n\u2014 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.
\n\u2014 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.
\n\u2014 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.
\n\u2022 Ionic or Electrovalent Bond<\/strong>
\nIonic 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.
\nThe compounds which is formed by ionic or electrovalent bond is known as electrovalent compounds. For Example, ,
\n(i) NaCl is an electrovalent compound. Formation of NaCl is given below:
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-2\"
\nNa+<\/sup> ion has the configuration of Ne while Cl\u2013<\/sup> ion represents the configuration of Ar.
\n(ii) Formation of magnesium oxide from magnesium and oxygen.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-3\"
\nElectrovalency: Electrovalency is the number of electrons lost or gained during the formation of an ionic bond or electrovalent bond.
\n\u2022 Factors Affecting the Formation of Ionic Bond<\/strong>
\n(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.
\nHence, 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.
\nFor example, in formation of Na+<\/sup> ion I.E = 496 kJ\/mole
\nWhile in case of magnesium, it is 743 kJ\/mole. That\u2019s why the formation of positive ion for sodium is easier than that of magnesium.
\nTherefore, we can conclude that lower the ionization enthalpy, greater the chances of ionic bond formation.
\n(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.
\nFor example. Halogens possess high electron affinity. So, the formation of anion is very common in halogens.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-4\"
\n(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.
\nLattice energy of an ionic compound depends upon following factors:<\/strong>
\n(i) Size of the ions: Smaller the size, greater will be the lattice energy.
\n(ii) Charge on the ions: Greater the magnitude of charge, greater the interionic attraction and hence higher the lattice energy.
\n\u2022 General Characteristics of ionic Compounds<\/strong>
\n(i) Physical\u2019State: 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 <\/sub>, N2 <\/sub>, 02 <\/sub>, Cl2 <\/sub> etc.
\n(ii) Melting and boiling points: Since ionic compounds contain high interionic force between them, they generally have high melting and boiling points.
\n(iii) Solubility: They are soluble in polar solvents such as water but do not dissolve in organic solvents like benzene, CCl4<\/sub>etc.
\n(iv) Electrical conductivity: In solid state they are poor conductors of electricity but in molten state or when dissolved in water, they conduct electricity.
\n(v) Ionic reactions: Ionic compounds produce ions in the solution which gives very fast reaction with oppositely charged ions.
\nFor example,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-5\"
\n\u2022 Covalent Bond\u2014Lewis-Langmuir Concept<\/strong>
\nWhen the bond is formed between two or more atoms by mutual contribution and sharing of electrons, it is known as covalent bond.
\nIf the combining atoms are same the covalent molecule is known as homoatomic. If they are different, they are known as heteroatomic molecule.
\nFor Example,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-6\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-7\"
\n\u2022 Lewis Representation of Simple Molecules (the Lewis Structures)<\/strong>
\nThe Lewis dot Structure can be written through the following steps:
\n(i) Calculate the total number of valence electrons of the combining atoms.
\n(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.
\n(iii) By knowing the chemical symbols of the combining atoms.
\n(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.<\/p>\n

\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-9\"
\n\u2022 Formal Charge<\/strong>
\nIn 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.
\nIt can be expressed as
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-10\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-11\"
\n\u2022 Limitations of the Octet Rule<\/strong>
\n(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,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-12\"
\nLi, Be and B have 1, 2, and 3 valence electrons only.
\n(ii) Odd-electron molecules: There are certain molecules which have odd number of electrons the octet rule is not applied for all the atoms.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-13\"
\n(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,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-14\"
\n\u2022 Other Drawbacks of Octet Theory<\/strong>
\n(i) Some noble gases, also combine with oxygen and fluorine to form a number of compounds like XeF2 <\/sub>, XeOF2 <\/sub> etc.
\n(ii) This theory does not account for the shape of the molecule.
\n(iii) It does not give any idea about the energy of The molecule and relative stability.
\n\u2022 Bond Length<\/strong>
\nIt 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.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-15\"
\n\u2022 Bond Angle<\/strong>
\nIt is defined as -the angle between the lines representing the orbitals containing the bonding \u2013 electrons.
\nIt helps us in determining the shape. It can be expressed in degree. Bond angle can be experimentally determined by spectroscopic methods.
\n\u2022 Bond Enthalpy<\/strong>
\nIt is defined as the amount of energy required to break one mole of bonds of a particular type to separate them into gaseous atoms.
\nBond Enthalpy is also known as bond dissociation enthalpy or simple bond enthalpy. Unit of bond enthalpy = kJ mol-1<\/sup>
\nGreater the bond enthalpy, stronger is the bond. For e.g., the H\u2014H bond enthalpy in hydrogen is 435.8 kJ mol-1<\/sup>.
\nThe 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 \u2014C bond is 347 kJ mol-1<\/sup> while that of C = C bond is 610 kJ mol-1<\/sup>.
\nIn polyatomic molecules, the term mean or average bond enthalpy is used.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-16\"
\n\u2022 Bond Order<\/strong>
\nAccording to Lewis, in a covalent bond, the bond order is given by the number of bonds between two atoms in a molecule. For example,
\nBond order of H2<\/sub> (H \u2014H) =1
\nBond order of 02<\/sub> (O = O) =2
\nBond order of N2<\/sub> (N = N) =3
\nIsoelectronic molecules and ions have identical bond orders. For example, F2<\/sub> and O2<\/sub>2-<\/sup> have bond order = 1. N2<\/sub>, CO and NO+ have bond order = 3. With the increase in bond order, bond enthalpy increases and bond length decreases. For example,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-17\"
\n\u2022 Resonance Structures<\/strong>
\nThere are many molecules whose behaviour cannot be explained by a single-Lew is structure, Tor example, Lewis structure of Ozone represented as follows:
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-18\"
\nThus, 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<\/sub>, the two structures shown above are canonical structures and the III structure represents the structure of 03<\/sub> more accurately. This is also called resonance hybrid.
\nSome resonating structures of some more molecules and ions are shown as follows:
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-19\"
\n\u2022 Polarity of Bonds<\/strong>
\nPolar and Non-Polar Covalent bonds
\nNon-Polar Covalent bonds: When the atoms joined by covalent bond are the same like; H2<\/sub>, 02<\/sub>, Cl2<\/sub>, the shared pair of electrons is equally attracted by two atoms and thus the shared electron pair is equidistant to both of them.
\nAlternatively, 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.
\nFor Example,
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-20\"
\nPolar bond: When covalent bonds formed between different atoms of different electronegativity, shared electron pair between two atoms gets displaced towards highly electronegative atoms.
\nFor 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 (\u03b4\u2013<\/sup>) and hydrogen atom have a partial positive charge (\u03b4+<\/sup>) with the magnitude of charge same as on chlorination. Such covalent bond is called polar covalent bond.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-21\"
\n\u2022 Dipole Moment<\/strong>
\nDue 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.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-22\"
\n\u2022 Applications of Dipole Moment<\/strong>
\n(i) For determining the polarity of the molecules.
\n(ii) In finding the shapes of the molecules.
\nFor example, the molecules with zero dipole moment will be linear or symmetrical. Those molecules which have unsymmetrical shapes will be either bent or angular.
\n(e.g., NH3<\/sub>with \u03bc = 1.47 D).
\n(iii) In calculating the percentage ionic character of polar bonds.
\n\u2022 The Valence Shell Electron Pair Repulsion (VSEPR) Theory<\/strong>
\nSidgwick 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).
\nMain Postulates are the following:
\n(i) The exact shape of molecule depends upon the number of electron pairs (bonded or non bonded) around the central atoms.
\n(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.
\n(iii) Electron pairs try to take such position which can minimize the rupulsion between them.
\n(iv) The valence shell is taken as a sphere with the electron pairs placed at maximum distance.
\n(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.
\n\u2022 Valence Bond Theory<\/strong>
\nValence 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.
\nLet us consider the formation of hydrogen molecule based on valence-bond theory.
\nLet two hydrogen atoms A and B having their nuclei NA <\/sub> and NB <\/sub> and electrons present in them are eA<\/sub> and eB <\/sub>.
\nAs these two atoms come closer new attractive and repulsive forces begin to operate.
\n(i) The nucleus of one atom is attracted towards its own electron and the electron of the other and vice versa.
\n(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.
\n\u2022 Orbital Overlap Concept<\/strong>
\nAccording 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:
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-23\"
\nStability of a Molecular orbital depends upon the extent of the overlap of the atomic orbitals.
\n\u2022 Types of Orbital Overlap<\/strong>
\nDepending upon the type of overlapping, the covalent bonds are of two types, known as sigma (\u03c3 ) and pi (\u03c0) bonds.
\n(i) Sigma (\u03c3 bond): Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis.
\nThe axial overlap involving these orbitals is of three types:
\n\u2022 s-s overlapping:<\/strong> In this case, there is overlap of two half-filled s-orbitals along the internuclear axis as shown below:
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-24\"
\n\u2022 s-p overlapping<\/strong>: This type of overlapping occurs between half-filled s-orbitals of one atom and half filled p-orbitals of another atoms.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-25\"
\n\u2022 p-p overlapping:<\/strong> This type of overlapping takes place between half filled p-orbitals of the two approaching atoms.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-26\"
\n (ii) pi (\u03c0 bond):<\/strong> \u03c0 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.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-27\"
\n\u2022 Strength of Sigma and pf Bonds<\/strong>
\nSigma bond (\u03c3 bond) is formed by the axial overlapping of the atomic orbitals while the \u03c0-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 \u03c0-bond.
\nDistinction between sigma and n bonds
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-28\"
\n\u2022 Hybridisation<\/strong>
\nHybridisation 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.
\nSalient Features of Hybridisation:<\/strong>
\n(i) Orbitals with almost equal energy take part in the hybridisation.
\n(ii) Number of hybrid orbitals produced is equal to the number of atomic orbitals mixed,
\n(iii) Geometry of a covalent molecule can be indicated by the type of hybridisation.
\n(iv) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.
\nConditions necessary for hybridisation:<\/strong>
\n(i) Orbitals of valence shell take part in the hybridisation.
\n(ii) Orbitals involved in hybridisation should have almost equal energy.
\n(iii) Promotion of electron is not necessary condition prior to hybridisation.
\n(iv) In some cases filled orbitals of valence shell also take part in hybridisation.
\nTypes of Hybridisation:<\/strong>
\n(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.
\nEach of the hybrid orbitals formed has 50% s-characer and 50%, p-character. This type of hybridisation is also known as diagonal hybridisation.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-29\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-30\"
\n(ii) sp2<\/sup> hybridisation: In this type, one s and two p-orbitals hybridise to form three equivalent sp2<\/sup> hybridised orbitals.
\nAll the three hybrid orbitals remain in the same plane making an angle of 120\u00b0. Example. A few compounds in which sp2<\/sup> hybridisation takes place are BF3<\/sub>, BH3<\/sub>, BCl3<\/sub> carbon compounds containing double bond etc.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-31\"
\n(iii) sp3<\/sup> 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<\/sup> hybrid orbital. The four sp3<\/sup> orbitals are directed towards four corners of the tetrahedron.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-32\"
\nThe angle between sp3<\/sup> hybrid orbitals is 109.5\u00b0.
\nA compound in which sp3<\/sup> hybridisation occurs is, (CH4<\/sub>). The structures of NH2<\/sub> and H2<\/sub>0 molecules can also be explained with the help of sp3<\/sup> hybridisation.
\n\u2022 Formation of Molecular Orbitals:<\/strong> Linear Combination of Atomic Orbitals (LCAO)
\nThe 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.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-33\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-34\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-35\"
\nThe 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.
\nConditions for the combination of atomic orbitals:
\n(1) The combining atomic orbitals must have almost equal energy.
\n(2) The combining atomic orbitals must have same symmetry about the molecular axis.
\n(3) The combining atomic orbitals must overlap to the maximum extent.
\n\u2022 Types of Molecular Orbitals<\/strong>
\nSigma (\u03c3) Molecular Orbitals: They are symmetrical around the bond-axis.
\npi (\u03c0) Molecular Orbitals: They are not symmetrical, because of the presence of positive lobes above and negative lobes below the molecular plane.
\n\u2022 Electronic configuration and Molecular Behaviour<\/strong>
\nThe distribution of electrons among various molecular orbitals is called electronic configuration of the molecule.
\n\u2022 Stability of Molecules<\/strong>
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-36\"
\n\u2022 Bond Order<\/strong>
\nBond order is defined as half of the difference between the number of electrons present in bonding and antibonding molecular orbitals.
\nBond order (B.O.) = 1\/2 [Nb<\/sub>-Na<\/sub>]\nThe bond order may be a whole number, a fraction or even zero.
\nIt may also be positive or negative.
\nNature of the bond:<\/strong> Integral bond order value for single double and triple bond will be 1, 2 and 3 respectively.
\nBond-Length:<\/strong> Bond order is inversely proportional to bond-length. Thus, greater the bond order, smaller will be the bond-length.
\nMagnetic Nature:<\/strong> 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<\/sub> molecule.
\n\u2022 Bonding in Some Homonuclear (Diatomic) Molecules<\/strong>
\n(1) Hydrogen molecule (H2<\/sub>): 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
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-37\"
\nThis 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
\nNo unpaired electron is present therefore,, it is diamagnetic.
\n(2) Helium molecule (He2<\/sub>): Each helium atom contains 2 electrons, thus in He2<\/sub> molecule there would be 4 electrons.
\nThe electrons will be accommodated in \u03c31s and \u03c3*1s molecular orbitals:
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-38\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-43\"
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-40\"
\n\u2022 Hydrogen Bonding<\/strong>
\nWhen 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.
\nIt can be depicted as
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-41\"
\n\u2022 Types of H-Bonds<\/strong>
\n(i) Intermolecular hydrogen bond (ii) Intramolecular hydrogen bond.
\n(i) Intermolecular hydrogen bond:<\/strong> It is formed between two different molecules of the same or different compounds. For Example, in HF molecules, water molecules etc.
\n(ii) Intramolecular hydrogen bond:<\/strong> 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.
\n\"chemical-bonding-and-molecular-structure-cbse-notes-for-class-11-chemistry-42\"<\/p>\n

<\/h4>\n","protected":false},"excerpt":{"rendered":"

Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4 \u2022 Chemical Bond The force that holds different atoms […]<\/p>\n","protected":false},"author":7,"featured_media":0,"comment_status":"closed","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"_yoast_wpseo_focuskw":"","_yoast_wpseo_title":"","_yoast_wpseo_metadesc":"\"Get Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4 on Infinity Learn.\"","custom_permalink":"study-material-notes\/class-11\/chemistry\/chemical-bonding-and-molecular-structure-class-11-notes-chemistry-chapter-4\/"},"categories":[92,21],"tags":[],"table_tags":[],"acf":[],"yoast_head":"\nChemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4 - Infinity Learn by Sri Chaitanya<\/title>\n<meta name=\"description\" content=\""Get Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4 on Infinity Learn."\" \/>\n<meta name=\"robots\" content=\"index, follow, max-snippet:-1, max-image-preview:large, max-video-preview:-1\" \/>\n<link rel=\"canonical\" href=\"https:\/\/infinitylearn.com\/surge\/study-material-notes\/class-11\/chemistry\/chemical-bonding-and-molecular-structure-class-11-notes-chemistry-chapter-4\/\" \/>\n<meta property=\"og:locale\" content=\"en_US\" \/>\n<meta property=\"og:type\" content=\"article\" \/>\n<meta property=\"og:title\" content=\"Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4 - 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