Class 11 chemical bonding notes ppt part I.pptx

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1. Chemical bonding and M olecular structure Class XI Prepared by Dr. Tarang Tomar PGT chemistry

2. What is a chemical bond ? The attractive forces which holds various constituents such as atoms molecules and ions together in different chemical species is called a Chemical bond. Basically it is the same kind of electrostatic attraction that binds the electron of an atom to its positively charged nucleus to form a molecule. This process is accompanied by decrease in energy. Decrease in energy ∝ strength of the bond Therefore, molecules are more stable than atoms.

3. Lattice energy: Amount of energy released during the formation of 1 mole of ionic crystal from its constituent ions. Na + + Cl - NaCl L.E. = - ve Amount of energy required to dissociate 1 mole of ionic crystal into its constituent ions . NaCl Na + + Cl - L.E. = + ve

4. Questions based on Lattice energy Trick L.E ∝ charge L.E. ∝ 1/size Q1. Which one is having more lattice energy? NaF , MgF2, AlF3 Na2O, MgO , Al2O3 Li2O, Li3N NaCl , KCl NaF , NaCl , NaBr NaF , MgCl2

5. Cu + , Ag + , Au + , Zn 2+ , Cd 2+ , Hg 2+ Follow PNGC

6. Covalent bonds have several characteristics, including:    Low melting and boiling points: Covalent compounds have low melting and boiling points because of the relatively weak forces between their particles.    Poor conductors: Covalent compounds are poor conductors of electricity and heat in all states (solid, molten, or aqueous).    Insoluble in water: Covalent compounds are generally insoluble in water, but they can dissolve in organic solvents.    Soft solids, liquids, or gases: Covalent compounds can exist as soft solids, liquids, or gases.    Brittle: Covalent compounds are brittle solids Characteristics of Covalent compounds

7. Lewis Dot Structure

8. Kossel Lewis approach to chemical bonding Lewis postulated that atoms achieve a stable octet when linked via chemical bonds. In case of bonds formed from H2 , F2 etc. the bond formed from sharing of electrons between the atoms. In this case each atom attains a stable outer octet of electrons. Lewis symbols: I n the formation of a molecule only outer electrons or group valence of the electrons take part in chemical bonding and hence are called valence electrons. In electrons of the s, p, and f orbital's do not take part in chemical bonding. Only ‘d’ orbital's take part.

9. Lewis symbols Lewis symbols as shown are given below: Lewis symbols only show the group valence or electrons that take part in chemical bonding and hence it is called valence electrons. Significance of Lewis Symbols : The number of dots around the symbol represents the number of valence electrons. This number common or group valence of the element. The group valence of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.

10. Trick for Calculations:- Total electron (Q) = V.E of all atoms + (- ve charge) – (+ ve charge ) Bond Pair of electron (B.P e - ) = 2 x no. of bonds Lone Pair of electron = Q - B.P e - (i.e. (L.P e - ) or = bond or ≡ bond)

11. Draw the lewis dot structure of the following compounds:- 1. H 2 2. O 2 3. H 3 O + 4. NH 4 + 5. NO 3 - 6. O 3 7. CO 3 2- 8. NO 2 - 9. SO 4 2- 10. SO 2 11. SO 3 12. PO 4 3- 13. CO

12. # F ormal Charge (F.C.):- Formal charge is the charge assigned to an atom in a molecule Formula / Trick :- F.C. = V – L – Where, V = Valence electron on atom L = Lone pair on atom B = Bonding electron on atom  

13. Calculate the formal charge of each atom in a molecule NO 3 -1 SO 3 -2 NO 2 -1 COCl 2 CO BH 3 N 2 O 4

14. Limitations of the octet rule. The incomplete octet of the central atom In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially in the case of elements having less than 4 valence electrons. Eg: LiCl , BeH 2.

15. The odd electron molecules: In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO 2 , the octet rule is not satisfied for all the atoms. The expanded octet: Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet . For example:- PCl5, SF6, H2SO4, SO3 etc

16. Coordinate Bond :- It is a covalent bond in which the shared electron pair from one atom is known as coordinate bond. Necessary conditions for the formation of coordinate bond. a. Octet of donor atom should be complete and should have at least one lone pair of electron. b. Acceptor atom should have a deficiency of at least 1 pair of electron. For eg . c. Atom which provide electron pair for sharing is called donor and other atom which accepts electron pair is known as acceptor. This is why the bond is known as dative bond.

17. Sigma and pi bond :- Sigma and pi bond are two types of covalent bonds. It is a bond that is formed by the mutual sharing of electrons so as to complete their octet or duplet in the case of hydrogen, lithium and beryllium. 1. 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 head-on overlap or axial overlap. This can be formed by any one of the following types of combinations of atomic orbitals.

18. 2. Pi (  ) Bond During the formation of Pi bonds, atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis.

19. All single bonds are 𝞂- bonds. Multiple bonds contain one  𝞂 -bond, and the rest are π-bonds. π-bond is never formed alone. First, a 𝞂-bond is formed, and then the formation of the π- bond takes place. (Exception: C molecule contains both π-bonds) A sigma bond is always stronger than a  pi bond because the extent of overlapping of atomic orbitals along the internuclear axis is greater than sideways overlapping. The electron cloud of 𝞂-bond is symmetrical about the internuclear axis, while that of π-bond is not. Free relation about a 𝞂-bond is possible, but that about a π-bond is not possible. The  less pi bonds, the more stable the compound is. The  more the number of pi bonds, the compound is more reactive . Note:

20. Overlap of Atomic Orbitals The atomic orbitals of 2 atoms overlap once they are nearer to one another. The overlapping of atomic orbitals can have positive, negative, or zero overlaps, relying upon these characteristics. The figure beneath shows the different configurations of the s and p-orbitals that lead to positive, negative, and zero overlaps. Positive atomic orbital overlap: Whenever the two involved atomic orbitals phase is identical, positive overlap takes place. Bonds are created as a consequence of this overlap. Negative atomic orbital overlap: Negative overlap occurs whenever the phases of the involved atomic orbitals oppose one another. Bond formation doesn't take place in this instance. Zero overlaps of atomic orbital: Zero atomic orbital overlaps occur while two intriguing orbitals do not overlap with each other in an orbital. The orbital overlap diagram is shown in next page:-

21. Zero overlap Zero overlap

22. Hybridization Pauling introduced the concept of hybridisation . According to him 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. For example when one 2s and three 2p-orbitals of carbon hybridise , there is the formation of four new Sp 3 hybrid orbital

23. Limitations of Valence Bond Theory it fails to explain the shape and geometry of molecules. It fails to explain magnetic behavior of covalent molecules.

24. Valence Shell Electron Pair Repulsion Theory VSEPR Theory was suggested by Sidgwick and Powel[1940] It was developed by Gilllespe and Nyholm in 1957. Based on that in a polyatomic molecule the direction bonds around the central atom depends on the total number of Bonding &Non-bonding electron pairs in its valance shell.

25. The shape of the molecule is determined by repulsions between all of the electron present in the valance shell. Electron pairs in the valence shell of the central atom repel each other and align themselves to minimize this repulsion. Lone pair electrons takes up more space round the central atom than a bondpair . Lone pair attracted to one nucleus, but bond pair is shared by two nuclei. The minimum repulsions to the state minimum energy and maximum stability of the molecule. Postulates of VSEPR Theory

26. Repulsion strengths Lone pair-Lone pair  Lone pair-Bond pair  Bond pair-Bond pair Nyholmn and Gillespie stated that the electron pairs existing as l.p. causes greater repulsive interactions as compared to bonded pairs

27. Limitations of VSEPR theory  It fails to predict the shapes of isoelectronic species ( CH 4 &NH 4 + ) and transition metal compounds.  This model does not take relative sizes of substituents . Unable to explain atomic orbitals overlap.

28. Molecular Orbital Theory was proposed by Hund & Mulliken . With the help of MOT, we can explain and understand those things that VBT was unable to explain. (Exp. Paramagnetic nature of O 2  molecule. As per VBT, it should be Diamagnetic.) The atomic orbital loses its identity during molecule formation (by overlapping) and forms new orbitals called molecular orbitals. Molecular orbital formed by overlapping of atomic orbital of same energy. Electrons in a molecule occupy molecular orbitals in accordance with Aufbau principle, Pauli’s exclusion principle and Hund’s rule.  No. of molecular orbitals formed = No. of atomic orbitals involved in overlapping. Molecular Orbital Theory Salient features of Molecular orbital theory:-

29. According to this theory, all the atomic orbitals of the participating atoms gets disturbed when the concerned nuclei approach each other. Definition of Atomic orbitals:- “ The region in space around the nucleus of an atom when the probability of finding the electron density is maximum.” Definition of Molecular orbitals:- “ The region in space comprising the nuclei of the combining atom around which there is maximum probability of finding the electron density.”

30. Linear combination of atomic orbitals (L.C.A.O) Molecular orbitals are formed by the linear combination of the wave function of participating atomic orbitals They may combine either by addition or subtraction. Let ψ A and ψ B represents the wave functions of the two combining atomic orbitals A and B. Ψ = ψ A + ψ B Bonding Molecular orbitals Ψ * = ψ A - ψ B Antibonding Molecular orbitals

31. Ψ = ψ A + ψ B a. Combination by addition b . Combination by subtraction Ψ * = ψ A - ψ B Constructive interference Destructive interference

32. Shapes of molecular orbitals 1. Combination between s-atomic orbitals Node Note :- Node means nodal plane where the probability of finding electron density is nill .

33. 2. Combination between s and p-atomic orbitals

34. 3 . Combination between p x -atomic orbitals

35. 4. Combination between p y or p z -atomic orbitals

36. Energy Level Diagram of Homoatomic molecules Get repelled by magnetic fields Get attracted by magnetic fields

37. M.O. energy level diagram for homonuclear diatomic molecule. 1. Hydrogen molecule (H 2 ) Electronic Configuration of H (1): 1s 1 Electronic Configuration of H 2 (2): σ 1s 2 Energy A.O. of H A.O. of H M.O. of H 2 σ* 1s σ 1s

38. Bond order in H 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (2-0) = ½ x 2 = 1 Bond order in H 2 = 1 ……………i.e .( H-H ) Thus, the bond order in H 2 molecule is 1. It suggest that there is a single bond present between the two H-atoms in H 2 molecule. It is a of σ type. H 2 molecule is diamagnetic because all the electrons are paired .

39. 1a. Hydrogen ion (H 2 + ion) Electronic Configuration of H 2 (2): σ 1s 2 Electronic Configuration of H 2 + (1): σ 1s 1 Energy A.O. of H A.O. of H M.O. of H 2 σ* 1s σ 1s Electronic Configuration of H (1): 1s 1 Bond order in H 2 + ion = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (1-0 ) = ½

40. 2. Helium molecule (He 2 ) Electronic Configuration of He (2): 1s 2 Electronic Configuration of He 2 (4): σ 1s 2 , σ *1s 2 Energy A.O. of He A.O. of He M.O. of He 2 σ* 1s σ 1s

41. Bond order in He 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (2-2) = ½ x 0 = Bond order in He 2 = 0……………… i.e.( Molecule is unstable ) Thus, the bond order in He 2 molecule is 0. It suggest He 2 molecule is not stable . Hence He 2 molecule does not exist . Helium being inert gas element exist in atomic form only.

42. 3. Lithium molecule (Li 2 ) Electronic Configuration of Li (3): 1s 2 , 2s 1 Electronic Configuration of Li 2 (6): σ 1s 2 , σ *1s 2 , σ 2s 2 Energy A.O. of Li A.O. of Li M.O. of Li 2 σ 1s σ * 1s σ 2s σ * 2s 1s 1s 2s 2s Bond order in Li 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (4-2) = ½ x 2 = 1 Thus, the bond order in Li 2 molecule is 1 . It suggest that there is a single bond present between the two Li-atoms in Li 2 molecule. It is a of σ type. The Li 2 molecule is diamagnetic due to the presence of paired electrons . Bond order in Li 2 = 1……… i.e .( Molecule is stable )

43. 4. Beryllium molecule (Be 2 ) Electronic Configuration of Be (4): 1s 2 , 2s 2 Electronic Configuration of Be 2 (8): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 Energy A.O. of Be A.O. of Be M.O. of Be 2 σ 1s σ * 1s σ 2s σ * 2s 1s 1s 2s 2s Bond order in Be 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (4-4) = ½ x 0 = Bond order in Be 2 = 0……………… i.e.( Molecule is unstable ) Thus , the bond order in Be 2 molecule is 0. It suggest Be 2 molecule is not stable . Hence Be 2 molecule does not exist . Beryllium being inert gas element exist in atomic form only.

44. 5. Boron molecule (B 2 ) Electronic Configuration of B (5): 1s 2 ,2s 2 ,2px 1 Electronic Configuration of B 2 (10): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 Energy A.O. of B A.O. of B M.O. of B 2 σ 2s σ * 2s 2s 2s 2px 2py 2pz 2px 2py 2pz σ *2px π* 2py π* 2py σ 2px π 2py π 2py

45. Bond order in B 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (4-2) = ½ x 2 = 1 Bond order in B 2 = 1……………… i.e.( B-B ) Thus, the bond order in B 2 molecule is 1 . It suggest that there is a single bond present between the two B-atoms in B 2 molecule. It is a of σ type. The B 2 molecule is diamagnetic due to the presence of paired electrons .

46. 6. Carbon molecule (C 2 ) Electronic Configuration of C (6): 1s 2 ,2s 2 ,2px 1 ,2py 1 Electronic Configuration of C 2 (12): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 1 , π 2pz 1 Energy A.O. of C A.O. of C M.O. of C 2 σ 2s σ * 2s 2s 2s 2px 2py 2pz 2px 2py 2pz σ *2px π* 2py π* 2py σ 2px π 2py π 2py

47. Bond order in C 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (6-2) = ½ x 4 = 2 Bond order in C 2 = 2……………… i.e.( C=C ) Thus, the bond order in C 2 molecule is 2 . It suggest that there are two bonds present between the two C-atoms in C 2 molecule. Out of two bonds one is σ bond & another is bond. The C 2 molecule is paramagnetic due to the presence of two unpaired electrons . σ π π

48. 7. Nitrogen molecule (N 2 ) Electronic Configuration of N (7): 1s 2 ,2s 2 ,2px 1 ,2py 1 ,2pz 1 Electronic Configuration of N 2 (14): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 Energy A.O. of N A.O. of N M.O. of N 2 σ 2s σ * 2s 2s 2s 2px 2py 2pz 2px 2py 2pz σ *2px π* 2py π* 2py σ 2px π 2py π 2py

49. Bond order in N 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-2) = ½ x 6 = 3 Thus, the bond order in N 2 molecule is 3 . It suggest that there are three bonds present between the two N-atoms in N 2 molecule. Out of three bonds, one is σ bond & there are two bond. The N 2 molecule is diamagnetic due to the presence of paired electrons . Bond order in N 2 = 3……………… i.e.( N= N ) σ π π π

50. 8. Oxygen molecule (O 2 ) Electronic Configuration of O (8): 1s 2 ,2s 2 ,2px 2 ,2py 1 ,2pz 1 Electronic Configuration of O 2 (16): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 1 , π *2pz 1 , Energy A.O. of O A.O. of O M.O. of O 2 σ 2s σ * 2s 2s 2s 2px 2py 2pz 2px 2py 2pz σ *2px π* 2py π* 2py σ 2px π 2py π 2py

51. Bond order in O 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-4) = ½ x 4 = 2 Bond order in O 2 = 2 (stable)……………… i.e.( O O ) Thus, the bond order in O 2 molecule is 2 . It suggest that there are two bonds present between the two O-atoms in O 2 molecule. Out of two bonds one is σ bond & one is π bond . The O 2 molecule is paramagnetic due to the presence of two unpaired electrons .

52. 8a. Oxygen molecule (O 2 + ion) Electronic Configuration of O (8): 1s 2 ,2s 2 ,2px 2 ,2py 1 ,2pz 1 Electronic Configuration of O 2 (16): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 1 , π *2pz 1 Electronic Configuration of O 2 + (15): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 1 Bond order in O 2 + ion = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-3) = ½ x 5 = 2.5 Bond order in O 2 + = 2.5………… i.e.( O= O ) Thus, the bond order in O 2 + ion is 2.5 . In O 2 + ion there is one σ bond, one bond & one three electron bond. … π ½ σ π

53. 8b. Oxygen molecule (O 2 - ion Superoxide ion ) Electronic Configuration of O (8): 1s 2 ,2s 2 ,2px 2 ,2py 1 ,2pz 1 Electronic Configuration of O 2 (16): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 1 , π *2pz 1 Electronic Configuration of O 2 - (17): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 2 , π *2pz 1 Bond order in O 2 - ion = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-5) = ½ x 3 = 1.5 Bond order in O 2 - = 1.5………… i.e.( O - O ) Thus, the bond order in O 2 - ion is 1.5 . In O 2 - ion there is one σ bond & one three electron bond. … σ

54. 8c. Oxygen molecule (O 2 - - ion peroxide ion ) Electronic Configuration of O (8): 1s 2 ,2s 2 ,2px 2 ,2py 1 ,2pz 1 Electronic Configuration of O 2 (16): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 1 , π *2pz 1 Electronic Configuration of O 2 - - (18): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 2 , π *2pz 2 Bond order in O 2 - ion = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-6) = ½ x 2 = 1 Bond order in O 2 - - = 1………… i.e.( O - O ) Thus, the bond order in O 2 - - ion is 1 . In O 2 - - ion there is one σ bond. σ Bond Order O 2 + > O 2 > O 2 - > O 2 - - 2.5 2 1.5 1

55. 9. Fluorine molecule (F 2 ) Electronic Configuration of F (9): 1s 2 ,2s 2 ,2px 2 ,2py 2 ,2pz 1 Electronic Configuration of F 2 (18): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 2 , π *2pz 2 , Energy A.O. of F A.O. of F M.O. of F 2 σ 2s σ * 2s 2s 2s 2px 2py 2pz 2px 2py 2pz σ *2px π* 2py π* 2py σ 2px π 2py π 2py

56. Bond order in F 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-6) = ½ x 2 = 1 Bond order in F 2 = 1……………… i.e.( F-F ) Thus, the bond order in F 2 molecule is 1 . It suggest that there is a single bond present between the two F-atoms in F 2 molecule. It is a of σ type. The F 2 molecule is diamagnetic due to the presence of paired electrons .

57. 10. Neon molecule (Ne 2 ) Electronic Configuration of F(9 ): 1s 2 ,2s 2 ,2px 2 ,2py 2 ,2pz 2 Electronic Configuration of F 2 (18): σ 1s 2 , σ *1s 2 , σ 2s 2 , σ *2s 2 , σ 2px 2 , π 2py 2 , π 2pz 2 , π *2py 2 , π *2pz 2 , σ *2px 2 Energy A.O. of Ne A.O. of Ne M.O. of Ne 2 σ 2s σ * 2s 2s 2s 2px 2py 2pz 2px 2py 2pz σ *2px π* 2py π* 2py σ 2px π 2py π 2py

58. Bond order in Ne 2 Molecule = ½ (Number of electron in BMO - Number of electron in ABMO ) = ½ (8-8) = ½ x 0 = Bond order in Ne 2 = 0……………… i.e.( Molecule is unstable ) Thus, the bond order in Ne 2 molecule is 0. It suggest Ne 2 molecule is not stable . Hence Ne 2 molecule does not exist . Neon being inert gas element exist in atomic form only.

59. For example:- 1. Hydrogen Bond in Water (H 2 O) A highly electronegative oxygen atom is connected to a hydrogen atom in the water molecule. The shared pair of electrons are attracted to the oxygen atoms more, and this end of the molecule becomes negative, while the hydrogen atoms become positive.

60. 2. Hydrogen Bond in Hydrogen Fluoride (HF ) A stronger-than-average hydrogen bond is created by hydrofluoric acid and is known as a symmetric hydrogen bond. Formic acid can also make this type of bond.

61. 3. Hydrogen Bond in Ammonia (NH 3 ) Between the hydrogen in one molecule and the nitrogen in another, hydrogen bonds are formed. Since each nitrogen has a single electron pair, the bond that develops in the case of ammonia is relatively weak. Methylamine also has this form of hydrogen bonding with nitrogen.

62. 4. Hydrogen Bond in Alcohol and Carboxylic Acid  A type of chemical molecule with a -OH group is alcohol. In most cases, hydrogen bonding is easily generated if any molecule containing the hydrogen atom is immediately coupled to either oxygen or nitrogen. Hydrogen Bond in Alcohol

63. Hydrogen Bond in Carboxylic Acid   Strength of Hydrogen Bond:  The hydrogen bond is a relatively weak one. Hydrogen bonds have a strength that is halfway between weak van der Waals forces and strong covalent bonds. The attraction of the shared pair of electrons, and hence the atom’s electronegativity, determines the hydrogen bond’s dissociation energy.

64. Volatility –  The boiling point of compounds incorporating hydrogen bonding between distinct molecules is greater, hence they are less volatile. Solubility –  Because of the hydrogen bonding that can occur between water and the alcohol molecule, lower alcohols are soluble in water. Lower density of ice than water –  In the case of solid ice, hydrogen bonding causes water molecules to form a cage-like structure. In fact, each water molecule is tetrahedrally connected to four other water molecules. In the solid state, the molecules are not as tightly packed as they are in the liquid state. This case-like structure collapses as ice melts, bringing the molecules closer together. As a result, the volume of water reduces while the density increases for the same quantity of water. As a result, at 273 K, ice has a lower density than water. Ice floats because of this. Viscosity and Surface Tension –  Hydrogen bonding is found in compounds that have an associated molecule. As a result, their flow becomes more complicated. They have high surface tension and higher viscosity. Properties of Hydrogen Bonding

65. Types of Hydrogen Bonding 1. Intermolecular Hydrogen Bonding   Intermolecular hydrogen bonding occurs when hydrogen bonds are formed between molecules of the same or distinct substances. Hydrogen bonding in water, alcohol, and ammonia, for example.

66. 2. Intramolecular Hydrogen Bonding Intramolecular hydrogen bonding refers to hydrogen bonding that occurs within a single molecule. It occurs in compounds with two groups, one of which has a hydrogen atom linked to an electronegative atom and the other of which has a highly electronegative atom linked to a less electronegative atom of the other group. The link is created between the more electronegative atoms of one group and the hydrogen atoms of the other group. For e g . o-nitrophenol, o-hydroxy benzoic acid etc.