Chemical bonding and molecular structure

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This slides describes the different concepts of topic that come under chemical bonding and molecular structure

This slides describes the different concepts of topic that come under chemical bonding and molecular structure

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  • 1. • Know about valence electrons • Learn various types of bond and bond parameters • Lewis Structure • Understand polar and covalent characters of covalent and ionic bonds • Concept of hybridisation • Study of VSEPR and Molecular Orbital Theory
  • 2. 1. Atomicity of a gas: The number of atoms present in the molecule of a gas is called its atomicity. 2. Bond dipole moment ( µ ).:A covalent bond between two atoms of different elements is called a polar covalent bond . A polar bond is partly covalent bond and partly ionic. The percentage of ionicity in a covalent bond is called percentage ionic character in that bond . The ionic character in a bond is expressed in terms of bond dipole moment ( µ ). 3. BORN-HABER CYCLE: This thermochemical cycle was devised by Born and Haber in 1919. It relates the lattice energy of a crystalline substance to other thermochemical data. The Born-Haber cycle is the application of Hess's law to the enthalpy of formation of an ionic solid at 298 K.
  • 3. 4. Chemical bond : The chemical force which keeps the atoms in any molecule together is commonly described as a chemical bond. 5. Chemical compounds :Compounds are generally called chemical compounds because they are formed due to the chemical combination of the combining element. 6. Covalent bond : The Bond formed by Mutual sharing of electrons between the combining atoms of the same or different elements is called covalent bonds. 7. Double covalent bond :The bond formed between two atoms due to the sharing of two electron-pairs is called a double covalent bond or simply a double bond. It is denoted by two small horizontal lines (=) drawn between the two atoms, e.g., O = O, O = C = O etc. 8. Electronegative or nonmetallic character :The tendency of an element to accept electrons to form an anion is called its non metallic or electronegative character. 9. ELECTRONEGATIVITY :The relative tendency of an atom in a molecule to attract a shared pair of electrons towards itself is termed its electronegativity.
  • 4. 10. Electronic configuration:The distribution of electrons amongst various energy levels of a atom is termed its electronic configuration 11. HYBRIDISATION :The process of mixing of the atomic orbitals to form new hybrid orbitals is called hybridisation. 12. Hybrid orbitals :According to the concept of hybridisation, certain atomic orbitals of nearly the same energy undergo mixing to produce equal number of new orbitals. The new orbitals so obtained are called hybrid orbitals. 13. Hybridisation in carbon :Carbon shows sp 3 hybridisation in alkanes, sp 2 hybridisation in alkenes and sp hybridisation in alkynes. 14. HYDROGEN BOND :The bond between the hydrogen atom of one molecule and a more electronegative atom of the same or another molecule is called hydrogen bond. 15. Ionic (or Electrovalent) bond : An ionic (or electrovalent) bond is formed by a complete transfer of one or more electrons from the atom of a metal to that of a non-metal. 16. LATTICE ENERGY (Lattice Enthalpy) :The strength of binding forces in solids is described by the term lattice enthalpy ( ∆ L H ) (earlier the term lattice energy was used). The molar enthalpy change accompanying the complete separation of the constituent particles that composed of the solid (such as ions for ionic solids and molecules for molecular solids) under standard conditions is called lattice enthalpy ( ∆ H° ). The lattice enthalpy is a positive quantity.
  • 5. 17. Lewis Formula (or Electronic Formula) of a Compound :The formula showing the mode of electron-sharing between different atoms in the molecule of a compound is called its electronic formula or Lewis formula. 18. Metallic crystals : In metallic crystals, the valence electrons of all the atoms form a pool of mobile electrons. The nuclei with their inner electrons (called Kernels) are embedded into this pool of free electrons. Thus, the constituent particles in a metallic crystal are the positive kernels in a pool of electrons. 19. NON-POLAR COVALENT BOND :When a covalent bond is formed between two atoms of the same element, the electrons are shared equally between the two atoms. In other words, the shared electron-pair will lie exactly midway between the two atoms. The resulting molecule will be electrically symmetrical, i.e ., centre of the negative charge coincides with the centre of the positive charge. This type of covalent bond is described as a non-polar covalent bond. The bonds in the molecules H 2 , O 2 , Cl 2 etc., are non-polar covalent bonds. 20. OCTET RULE : According to this theory, 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 . 21. Pi ( π ) Bond : A covalent bond formed between the two atoms due to the sideways overlap of their p -orbitals is called a pi ( π ) bond.
  • 6. 22. POLAR COVALENT BOND :When a covalent bond is formed between two atoms of different elements, the bonding pair of electrons does not lie exactly midway between the two atoms. In fact, it lies more towards the atom which has more affinity for electrons. The atom with higher affinity for electrons, thus, develops a slight negative charge, and the atom with lesser affinity for electrons a slight positive charge. Such molecules are called polar molecules. The covalent bond between two unlike atoms which differ in their affinities for electrons is said to be a polar covalent bond. 23. RESONANCE : When a molecule is represented by a number of electronic structures such that none of them can exactly describe all the properties of the molecule, but each structure has a contribution to it, then the molecule is termed as a resonance hybrid of all these structures. Such structures are called resonance structures and such a phenomenon is called resonance. 24. Resonance hybrid :When a molecule is represented by a number of electronic structures such that none of them can exactly describe all the properties of the molecule, but each structure has a contribution to it, then the molecule is termed as a resonance hybrid of all these structures.
  • 7. 25. Sigma ( σ ) Bond : A covalent bond formed due to the overlap of orbitals of the two atoms along the line joining the two nuclei (orbital axis) is called sigma ( σ ) bond. 26. Single Covalent Bond : A covalent bond formed by mutual sharing of one pair of electrons is called a single covalent bond, or simply a single bond. A single covalent bond is represented by a small line (−) between the two atoms. 27. Triple covalent bond : Bond formed due to the sharing of three electron- pairs is called a triple covalent bond or simply a triple bond. 28. Valence electrons : Valence is one of the most important chemical property of the elements. The chemical behaviour of an element depends upon the number of electrons in the outermost shell of its atom. The electrons present in the outermost shell are called valence electrons. The electrons in the outermost shell are called valence electrons because the electrons in the outermost shell determine the valence of an element . 29. Valency : The combining capacity of an atom of an element is described in terms of its valency. It may be defined as, The number of hydrogen or chlorine or double the number of oxygen atoms which combine with one atom of the element is termed its valency. It may also be defined as, The number of electrons which an atom loses or gains or shares with other
  • 8. • The electrons that are in the highest(outermost) energy level • That level is also called the valence shell of the atom they are held most loosely • The number of valence electrons in an atom determines: • The properties of the atom • The way that atom will bond chemically •As a rule, the fewer electrons in the valence shell, the more reactive the element
  • 9. • When an atom has eight electrons in the valence shell, it is stable • Atoms usually react in a way that makes each atom more stable • There are two ways this can happen: •The number of valence electrons increases to eight •Loosely held valence electrons are given up
  • 10. • Atomic number = number of Electrons • Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells. • Electron shells determine how an atom behaves when it encounters other atoms
  • 11. • Electrons are placed in shells according to rules: • The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.
  • 12. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons C would like to N would like to O would like to Gain 4 electrons Gain 3 electrons Gain 2 electrons
  • 13. Group # Group Name # of valence electrons 1 Alkali Metals 1 2 Alkaline Earth Metals 2 3-12 Transition Metals 1 or 2 13 Boron Group 3 14 Carbon Group 4 15 Nitrogen Group 5 16 Oxygen Group 6 17 Halogens 7 18 Noble Gases 8
  • 14. •The octet rule is a simple chemical rule of thumb •Octet Rule says atoms with 8 electrons in their outer shell areatoms with 8 electrons in their outer shell are stablestable •Atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electron configuration as a noble gas • The rule applies to the main-group elements, especially carbon, nitrogen, oxygen, the halogens, and also to metals such as sodium or magnesium •In simple terms, molecules or ions tend to be most stable when the outermost electron shells of their constituent atoms contain 8 electrons
  • 15. Doesn’t allow for • H, He or Li [stable with 2 e- in their outer shells] - Duet Rule • Transition elements - 18 electron rule • BF3 which only has 6 e- in its outer shell
  • 16. • named after Gilbert N. Lewis, who introduced it in his 1916 • also known as Lewis dot diagrams, electron dot diagrams, "Lewis Dot formula" Lewis dot structures, and electron dot structures) • are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. • can be drawn for any covalently bonded molecule, as well as coordination compounds.
  • 17. Symbols of atoms with dots to represent the valence-shell electrons 1 2 13 14 15 16 17 18 H• He: • • • • • • • • • • • Li• Be• • B • • C • • N • • O • : F • :Ne : • • • • • • • • • • • • • • • • • • • Na• Mg• • Al• • Si• • P• • S• :Cl • :Ar : • • • • • • • •
  • 18. • A bond results from the attraction of nuclei for electrons – All atoms trying to achieve a stable octet • IN OTHER WORDS – the p+ in one nucleus are attracted to the e- of another atom • Electronegativity
  • 19. • It is an exothermic process Energy released E N E R G Y Reactants Products
  • 20. • Endothermic reaction – energy must be put into the bond in order to break it E N E R G Y Reactants Products Energy Absorbed
  • 21. • Strong, STABLE bonds require lots of energy to be formed or broken • weak bonds require little E
  • 22. Double bondDouble bond Single bondSingle bond TripleTriple bondbond AcrylonitrileAcrylonitrile
  • 23. Fractional bond orders in resonance structures. Consider NO2 - bondsO—N2 bondsNOinpairs-e3 =orderBond Bond order = Total # of e- pairs used for a type of bond Total # of bonds of that type The N—O bond order = 1.5The N—O bond order = 1.5 O O O O N •• •• •• •• •• •••••• •• •• •• •• •• N
  • 24. Bond order is proportional to two important bond properties: (a) bond strength (b) bond length 745 kJ745 kJ 414 kJ414 kJ 110 pm110 pm 123 pm123 pm
  • 25. • Bond length is the distance between the nuclei of two bonded atoms.
  • 26. Bond length depends on bond order. Bond distances measuredBond distances measured using CAChe software. Inusing CAChe software. In Angstrom units where 1 A =Angstrom units where 1 A = 1010-2-2 pm.pm. Bond distances measuredBond distances measured using CAChe software. Inusing CAChe software. In Angstrom units where 1 A =Angstrom units where 1 A = 1010-2-2 pm.pm.
  • 27. Using Bond Energies Estimate the energy of the reaction H—H + Cl—Cl ----> 2 H—Cl Net energy = H∆ rxn = = energy required to break bonds - energy evolved when bonds are made H—H = 436 kJ/molH—H = 436 kJ/mol Cl—Cl = 242 kJ/molCl—Cl = 242 kJ/mol H—Cl = 432 kJ/molH—Cl = 432 kJ/mol H—H = 436 kJ/molH—H = 436 kJ/mol Cl—Cl = 242 kJ/molCl—Cl = 242 kJ/mol H—Cl = 432 kJ/molH—Cl = 432 kJ/mol
  • 28. Estimate the energy of the reaction H—H + Cl—Cl ----> 2 H—ClH—H = 436 kJ/molH—H = 436 kJ/mol Cl—Cl = 242 kJ/molCl—Cl = 242 kJ/mol H—Cl = 432 kJ/molH—Cl = 432 kJ/mol H—H = 436 kJ/molH—H = 436 kJ/mol Cl—Cl = 242 kJ/molCl—Cl = 242 kJ/mol H—Cl = 432 kJ/molH—Cl = 432 kJ/mol Sum of H-H + Cl-Cl bond energies = 436 kJ +Sum of H-H + Cl-Cl bond energies = 436 kJ + 242 kJ = +678 kJ242 kJ = +678 kJ Using Bond Energies 2 mol H-Cl bond energies = 864 kJ2 mol H-Cl bond energies = 864 kJ Net = ∆H = +678 kJ - 864 kJ = -186 kJNet = ∆H = +678 kJ - 864 kJ = -186 kJ
  • 29. Why do ionic compounds dissolve inWhy do ionic compounds dissolve in water?water? Boiling point = 100 ˚CBoiling point = 100 ˚C Boiling point = -161 ˚CBoiling point = -161 ˚C Why do water and methaneWhy do water and methane differ so much in theirdiffer so much in their boiling points?boiling points?
  • 30. HCl isHCl is POLARPOLAR because itbecause it has a positive end and ahas a positive end and a negative end.negative end. Cl has a greater share inCl has a greater share in bonding electrons thanbonding electrons than does H.does H. Cl has slight negative chargeCl has slight negative charge (-(-δδ)) and H has slightand H has slight positive chargepositive charge (+(+ δδ)) H Cl •• •• +δ -δ ••H Cl •• •• +δ -δ ••
  • 31. Due to the bond polarity, the H—Cl bondDue to the bond polarity, the H—Cl bond energy is GREATER than expectedenergy is GREATER than expected for a “pure” covalent bond.for a “pure” covalent bond. BONDBOND ENERGYENERGY ““pure” bondpure” bond 339 kJ/mol calc’d339 kJ/mol calc’d real bondreal bond 432 kJ/mol measured432 kJ/mol measured BONDBOND ENERGYENERGY ““pure” bondpure” bond 339 kJ/mol calc’d339 kJ/mol calc’d real bondreal bond 432 kJ/mol measured432 kJ/mol measured Difference = 92 kJ. This difference is proportional to the difference inDifference = 92 kJ. This difference is proportional to the difference in ELECTRONEGATIVITYELECTRONEGATIVITY,, χχ.. Difference = 92 kJ. This difference is proportional to the difference inDifference = 92 kJ. This difference is proportional to the difference in ELECTRONEGATIVITYELECTRONEGATIVITY,, χχ.. H Cl •• •• +δ -δ ••
  • 32. • Metallic bonding – Occurs between like atoms of a metal in the free state – Valence e- are mobile (move freely among all metal atoms) – Positive ions in a sea of electrons • Metallic characteristics – High mp temps, ductile, malleable, shiny – Hard substances – Good conductors of heat and electricity as (s) and (l)
  • 33. • electrons are transferred between valence shells of atoms • ionic compounds are made of ions • ionic compounds are called Salts or Crystals NOT MOLECULES
  • 34. • Always formed between metals and non- metals [METALS ]+ [NON-METALS ] - Lost e- Gained e-
  • 35. • Electronegativity difference > 2.0 – Look up e-neg of the atoms in the bond and subtract NaCl CaCl2 • Compounds with polyatomic ions NaNO3
  • 36. Ionic Bond, A Sea of Electrons
  • 37. The formations of ionic bond governed by the following factors: 1.Ionization energy: • Formation of ionic bond metal atom loses electron to form cation • Energy required for this equal to ionization energy • Alkali metals have lowest ionization energy, thus have more tendency to form cation 2.Electron gain enthalpy: • Electron released in the formation of cation are to be accepted by the other atom taking part in the ionic bond formation • Electron accepting tendencies depend on upon the electron gain enthalpy • Defined as energy released when isolated gaseous atom takes up an electron to form anion. • Greater the negative enthalpy, easier the formation of anion
  • 38. 3. Lattice energy: • Combination of oppositely charged ions to form ionic crystal, with release of energy is referred as lattice energy • Higher value of lattice energy, greater will be the stability of compound • Magnitude of lattice energy gives idea about the strength of interionic forces • Size of ions: • In case of similar ions inter-nuclear distance is lesser due to which inter-ionic attraction is greater and hence the magnitude of lattice energy will be larger • Charge on the ions: • Ions have higher charge exerts stronger forces of attraction and hence larger amount of energy is released. Thus value of lattice energy is higher
  • 39. • Ionic compound exist in solid state • The network of ions have a definite geometric pattern which depends on the size and charge of ions • Posses high melting and boiling points due to strong electrostatic force of attraction between the ions • Good conductor of electricity in molten or dissolved state • Does not conduct electricity in solid state as ions are not free to move • Are soluble in polar solvent like water as solvent interacts with the ions of ionic solid •The chemical reactions between ionic compounds in aqueous solution involves the combination between their ions, such reactions are called ionic reactions.
  • 40. This thermochemical cycle was devised by Born and Haber in 1919. It relates the lattice energy of a crystalline substance to other thermochemical data. The Born-Haber cycle is the application of Hess's law to the enthalpy of formation of an ionic solid at 298 K.
  • 41. Formation of crystalline sodium chloride form sodium metal and chlorine gas can be described by the reaction. Na(s) + ½ Cl2 (g) → NaCl (crystal) ∆r H = ∆f H° = − 411 kJ mol−1 (energy evolved) This overall reaction can be considered to proceed in a stepwise manner as follows The signs of the energy involved in each step follow the rule that energy evolved is negative and energy absorbed is positive. These steps are summarized in Fig. 6.3. From the Born-Haber cycle the value for any one of the steps can be calculated if data for all the other steps are known.
  • 42. • Pairs of e- are shared between non- metal atoms • electronegativity difference < 2.0 • forms polyatomic ions
  • 43. • Between nonmetallic elements of similar electronegativity. • Formed by sharing electron pairs • Stable non-ionizing particles, they are not conductors at any state • Examples; O2, CO2, C2H6, H2O, SiC
  • 44. • Bonds in all polyatomic ions and diatomic are all covalent bonds • Two types of covalent bond: • Non-Polar Covalent Bond • Polar Covalent Bond
  • 45. when electrons are shared equally H2 or Cl2
  • 46. Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom Oxygen Molecule (OOxygen Molecule (O2)2)
  • 47. when electrons are shared but shared unequally H2O
  • 48. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.
  • 49. •Covalent bond can be represented by line drawn between symbols of element whose atoms are involved in sharing of electron •Single Bond : Sharing single electron pairs between two atoms •Double Bond : Sharing two electron pairs between two atoms •Triple Bond : Sharing three electron pairs between two atoms
  • 50. Hydrogen (H2) has a single bond between atoms. Oxygen (O2) has a double bond between atoms, indicated by two lines (=). Nitrogen (N2) has a triple bond between atoms, indicated by three lines ( ). Each bond represents an electron pair.≡
  • 51. • Compounds formed exist as discrete molecules •Weak intermolecular force due to small molecular size •Mainly exist in liquid or gaseous state •Sugar, urea, starch etc. exist in solid state •Low melting and Boiling points due to weak attractive forces •Poor conductor of electricity in fused or dissolved state •Less soluble in water •Gives molecular reaction
  • 52. VSEPR (Valence Shell Electron Pair Repulsion) Theory is a model for understanding and predicting the shape of molecules. This theory is based on the principle that electrons will spread out as far as possible from each other as a result of their mutual repulsion.
  • 53. We refer to the electron pairs as electron domains. A double or triple bond counts as one electron domain. This molecule has four electron domains.
  • 54. These are the electron- domain geometries for two through six electron domains around a central atom.
  • 55. The molecular geometry is defined by the positions of only the atoms in the molecules, not the nonbonding pairs (lone pairs).
  • 56. Within each electron domain, there may be several molecular geometries.
  • 57. In this domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.
  • 58. There are two molecular geometries: • Trigonal planar, if all the electron domains are bonding • Bent, if one of the domains is a nonbonding pair
  • 59. There are three molecular geometries: • Tetrahedral, if all are bonding pairs • Trigonal pyramidal if one is a nonbonding pair • Bent if there are two nonbonding pairs
  • 60. There are four distinct molecular geometries: • Trigonal bipyramidal • Seesaw • T-shaped • Linear
  • 61. In the trigonal bipyramidal electron domain geometry, the electron domains have two distinct positions: • Axial • Equatorial
  • 62. Nonbonding electron pairs occupy the equatorial, rather than axial, positions in this geometry.
  • 63. All positions are equivalent in the octahedral domain. There are three molecular geometries: • Octahedral • Square pyramidal • Square planar
  • 64. Predicting the Molecular Shape of a Molecule: Examples What are the molecular shapes of CO2 and NF3?
  • 65. Step 1: Draw the structural formula (using the outside atoms to determine the bonds). Oxygen needs two bonds Fluorine needs one bond
  • 66. Step 2: Add in lone pairs to the central atom (first determine how many valence electrons it has). Carbon has four valence electrons. Each bond includes one of these valence electrons, so there are none left over. Nitrogen has five valence electrons. Each bond includes one of these valence electrons, so there are two left over (one lone pair).
  • 67. Step 3: Count the total number of electron domains, and the number of lone pairs. Use this to determine the shape. 2 electron domains, 0 lone pairs 4 electron domains, 1 lone pair
  • 68. Nonbonding electron pairs (lone pairs) have a stronger repulsion than bonding electrons. This causes the actual bond angles to be slightly less than otherwise predicted. CH4 NH3 H2O
  • 69. In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.
  • 70. This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.
  • 71. To tell if bonds are polar or nonpolar, we simply look at the difference in electronegativity for the two atoms. To tell if a molecule is polar or nonpolar, we need to know its shape.
  • 72. Nonpolar molecules have a symmetrical distribution of charge (even if the bonds themselves are polar). Polar molecules have an asymmetrical distribution of charge (even if the bonds themselves are nonpolar). H2O CO2
  • 73. • Cannot explain shape of very much polar compound • Is also unable to explain the shape of very much polar in which delocalized p electrons are very much • VSEPR theory does not explain the shapes of molecules having inert electron pair
  • 74. Electronegativity is a measure of an atoms ability to attract electrons from an atom to which it is bonded
  • 75. Mulliken defined electronegativity as: • ½ (electron affinity + ionization energy) • This approach yields very high electronegativity values for He and Ne, even though they do not form compounds
  • 76. Electronegative atoms bonded to a less electronegative central atom tend to draw electron density away from the central atom, thus lowering repulsion. PFPF33 PClPCl33 PBrPBr33 97.897.8oo 100.3100.3oo 101101oo
  • 77. • The dipole moment (µ) of a molecule is the product of the magnitude of the charge (ε) and the distance (d) that separates the centers of positive and negative charge µ= ε∗d • A unit of dipole moment is the debye (D) • One debye (D) is equal to 1 x 10–10 esu Α
  • 78. • A polar covalent bond has a bond dipole; a separation of positive and negative charge centers in an individual bond • Bond dipoles have both a magnitude and a direction (they are vector quantities) • Ordinarily, a polar molecule must have polar bonds, BUT … polar bonds are not sufficient • A molecule may have polar bonds and be a nonpolar molecule – IF the bond dipoles cancel.
  • 79. • CO2 has polar bonds, but is a linear molecule; the bond dipoles cancel and it has no net dipole moment (µ = 0 D). • The water molecule has polar bonds also, but is an angular molecule. • The bond dipoles do not cancel (µ = 1.84 D), so water is a polar molecule. Net dipole No net dipole
  • 80. The first quantum mechanical model to explain the nature and stability of a covalent bond was formulated by Heitler and London 1927. This theory was then modified by Pauling and Slater in 1931. This theory is commonly known as the Valence-bond theory. The main postulates of the valence bond theory are: i. A covalent bond is formed due to the overlap of the outermost half-filled orbitals of the combining atoms. The strength of the bond is determined by the extent of overlap. ii. The two half-filled orbitals involved in the covalent bond formation should contain electrons with opposite spins. The two electrons then move under the influence of both the nuclei. iii. The completely-filled orbitals (orbitals containing two paired electrons) do not take part in the bond formation.
  • 81. iv. An s-orbital does not show any preference for direction. The non- spherical orbitals such as, p- and d-orbitals tend to form bonds in the direction of the maximum overlap, i.e., along the orbital axis. v. Between the two orbitals of the same energy, the orbital which is non-spherical (e.g., p- and d- orbitals forms stronger bonds than the orbital which is spherically symmetrical, e.g., s-orbital. vi. The valence of an element is equal to the number of half-filled orbitals present in it. In the valence bond model, the stability of a molecule is explained in terms of the following types of interactions. a. electron - nuclei attractive interactions, i.e., the electrons of one atom are attracted by the nucleus of the other atom also. b. electron - electron repulsive interactions, i.e., electrons of one atom are repelled by the electrons of the other atom. c. nucleus - nucleus repulsive interactions, i.e., nucleus of one atom is repelled by the nucleus of the other atom.
  • 82. The attractive and the repulsive interactions oppose each other. When the attractive interactions are stronger than the repulsive interactions, certain amount of energy is released. Due to the lowering of energy the molecule becomes stable. Various interactions which act between the two atoms are shown in Fig.
  • 83. Valence Bond Description of Hydrogen Molecule
  • 84. TYPES OF OVERLAPPING Various types of atomic orbital overlap leading to the formation of covalent bond are: 1. s − s overlap . In this type of overlap, half-filled s -orbitals of the two combining atoms overlap each other. This is shown in Fig.
  • 85. 2. s − p overlap . Here a half-filled s -orbital of one atom overlaps with one of the p -orbitals having only one electron in it. This is shown in Fig. 3. p−p overlap along the orbital axis. This is called head on, end-on or end- to-end linear overlap. Here, the overlap of the two half-filled p-orbitals takes place along the line joining the two nuclei. This is shown in Fig.
  • 86. 4. p−p sideways overlap. This is also called lateral overlap. In this types of overlap, two p-orbitals overlap each other along a line perpendicular to the internuclear axis, i.e., the two overlapping p- orbitals are parallel to each other. This is shown in Fig. TYPES OF COVALENT BONDS: SIGMA (σ) AND PI (π) BONDS The overlapping of orbitals is possible in two ways. (i) along their orbital axis so that the electron density along the axis is maximum. (ii) along a direction perpendicular to the bond axis due to sideways overlap of the orbitals. Depending upon the manner in which the two atomic orbitals overlap with each other, two types of bonds are formed. These are called, sigma (σ) bond, and pi (π) bond.
  • 87. A covalent bond formed due to the overlap of orbitals of the two atoms along the line joining the two nuclei (orbital axis) is called sigma ( σ ) bond. For example, the bond formed due to s-s and s-p, and p-p overlap along the orbital axis are sigma bonds, (by convention Z-axis is taken as inter-nuclear axis. A covalent bond formed between the two atoms due to the sideways overlap of their p -orbitals is called a pi ( π ) bond
  • 88. Sigma (σ) bond Pi (π) bond 1. It is formed due to axial overlap of the twoorbitals. The overlap may be of s-s, s-p, p- p orbitals. 1. This bond is formed by the lateral (sideways) overlap of two p-orbitals. 2. There can be only one sigma bond between atoms. 2. There can be more than one π-bond between the two atoms. 3. The electron density is maximum and cylindrically symmetrical about the bond axis. 3. The electron density is high along a direcion tion at right angle to the bond axis. 4. The bonding is relatively strong. 4. The bonding due to a π-bond is weak. 5. Free rotation of atoms about sigma (σ) bond is possible. 5. Free rotation about a π bond is not possible. 6. It can be formed independently, i.e., there can be a sigma (σ) bond without having a π bond in any molecule. 6. The π bond is formed only after σ bond has been formed.
  • 89. The concept of hybridization is used to explain the nature of bonds, and shape of the polyatomic molecules. For an isolated atom hybridization has no meaning. According to the concept of hybridization, certain atomic orbitals of nearly the same energy undergo mixing to produce equal number of new orbitals. The new orbitals so obtained are called hybrid orbitals. The process of mixing of the atomic orbitals to form new hybrid orbitals is called hybridization. All hybrid orbitals of a particular kind have equal energy, identical shapes and are symmetrically oriented in space. The types of atomic orbitals involved in hybridization, and the nature of hybridization depends upon the requirements of the reaction. For example, carbon in methane (CH4 ) shows sp3 hybridization. In ethene (ethylene, C2 H4 ), it exhibits sp2 hybridization. In ethyne (acetylene, C2 H2 ) it shows sp hybridization.
  • 90. The hybrid orbitals are designated according to the type and the number of atomic orbitals merging together. For example, Mixing orbitals Hybrid orbital Hybridisation one s and three p- four sp3 orbitals sp3 hybridisation one s and two p- three sp2 orbitals sp2 hybridisation one s and one p- two sp orbitals sp hybridisation Conditions Necessary For Hybridisation Atomic orbitals undergo hybridisation only if the following conditions are satisfied. (i) Atomic orbitals of the same atom participate in hybridisation. Electrons present in these atomic orbitals do not participate in the hybridisation, and occupy the hybrid orbitals as usual. (ii) The atomic orbitals participating in hybridisation should have nearly equal energy. (ii) Characteristics of Hybrid Orbitals The characteristics of hybrid orbitals are: i. The number of hybrid orbitals formed is equal to the number of the atomic orbitals participating in hybridisation.
  • 91. ii. All hybrid orbitals are equivalent in shape, and energy, but different from the participating atomic orbitals. iii. A hybrid orbital which takes part in the bond formation must contain only one electron in it. iv. A hybrid orbital, like atomic orbitals, cannot have more than two electrons. The two electrons should have their spins paired. v. Due to the electronic repulsions between the hybrid orbitals, they tend to remain at the maximum distance from each other Types of Hybridisation Depending upon the nature of the orbitals involved in hybridisation, different types of hybridisation become possible. The type of hybridisation shown by an atom depends upon the requirements of the reaction. • sp3 hybridisation: In any atom, corresponding to energy levels (or shells) for which n ≥ 2, there is one s orbital and three p orbitals. For example, for n = 2, we have one 2s and three 2p orbitals; for n = 3, we have one 3s, and three 3p orbitals. These four orbitals undergo mixing to provide four new hybrid orbitals. s + (px + py + pz ) → sp3 one s-orbital three p-orbitals four hybrid orbitals
  • 92. Hybridisation sp3 , geometry tetrahedral, bond angle 109.5’
  • 93. • sp2 hybridisation: In certain reactions, one s and two p (say px and py ) orbitals of an atom undergo mixing to produce three equivalent sp2 hybridised orbitals. The three sp2 hybrid orbitals are oriented in a plane along the three corners of an equilateral triangle, i.e., they are inclined to each other at an angle of 120°. The third p-orbital (say pz here) remains unchanged. Each hybrid orbital has 33.3% s-character and 66.7% p-character. Formation of sp2 -hybrid orbitals from one s and two p-orbitals is shown in Fig. 6.43. s + (px + py ) → sp2 one s-orbital two p-orbitals three hybrid orbitals
  • 94. Boron trifluoride has a plane trigonal shape in which all three bonds are identical. Hybridisation sp2 , geometry trigonal planar, bond angle 120’
  • 95. • sp hybridisation. In this type of hybridisation, one s and one p (say pz ) orbitals belonging to the same main energy level hybridise to give two sp hybrid orbitals. These sp hybrid orbitals are oriented at an angle of 180° to each other. Each hybrid orbital has 50% s- and 50% p- character. The other two p-orbitals (say 2px and 2py ) remain unhybridised and are oriented at right angles to each other and to the internuclear axis. Hybridisation sp, geometry Linear, bond angle 180’
  • 96. Formation of ethene molecule.
  • 97. Formation of ethyne (acetylene) molecule.
  • 98. •When some compounds cannot be represented by a single definite structure rather more than one structure •Thus, the various structure written for a compound to explain the known properties of the compound are called as resonating or contributing or canonical structure •This phenomenon is called resonance •The real structure is a resonance hybrid of all the resonating structure •The actual energy of the resonance hybrid and the most stable one of the resonating structures is called resonance energy
  • 99. Conditions of Resonance : 1.Different contributing structures should have the same position of the constituent, though may have different electronic arrangement 2.The number of unpaired electrons should be the same in all resonating structures 3.The contributing structure should have nearly the same energy 4.The bond length and bond angles should closer to the real structure
  • 100. Characteristics of Resonance : 1.Resonating structure are imaginary and do not have real existence 2.Resonance hybrid is more stable i.e. its energy is least among different resonating forms 3.The difference between the energy of resonance hybrid and most stable resonating form is called resonance energy. 4.Larger the value of resonance energy greater the stability of hybrid resonance 5.Bond length in hybrid structures are intermediate of the bond lengths in various resonating forms
  • 101. This is the Lewis structure we would draw for ozone, O3. - +
  • 102. • But this is at odds with the true, observed structure of ozone, in which… – …both O—O bonds are the same length. – …both outer oxygens have a charge of −1/2.
  • 103. • One Lewis structure cannot accurately depict a molecule such as ozone. • We use multiple structures, resonance structures, to describe the molecule.
  • 104. Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.
  • 105. • In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.
  • 106. • The organic compound benzene, C6H6, has two resonance structures. • It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.
  • 107. • A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. • In a simple covalent bond, each atom supplies one electron to the bond - but that doesn't have to be the case. • A co-ordinate bond is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.
  • 108. Example: NH4 +
  • 109. Properties of Coordinate bond : 1.Are generally soluble in water and organic solvents 2.Boiling and melting points of these compounds are less than electrovalent compounds but are higher than covalent compounds 3.Compounds ionize in aqueous solution giving simple and complex ions 4.These bonds are also directional and stereoisomerism is also found 5.Molecules possess definite shape and definite bond angles, thus have definite geometry
  • 110. A B ψA ψB ψAB = N(cA ψA + cBψB) ψ2 AB = (cA2 ψA2 + 2cAcB ψA ψB + cB2 ψB 2 ) Overlap integral The wave function for the molecular orbitals can be approximated by taking linearThe wave function for the molecular orbitals can be approximated by taking linear combinations of atomic orbitals.combinations of atomic orbitals. Probability density c – extent to which each AO contributes to the MO
  • 111. cA = cB = 1 +. +. . .+ bondingψg Amplitudes of wave functions added ψg = N [ψA + ψB] Constructive interferenceConstructive interference
  • 112. ψ2 AB = (cA2 ψA2 + 2cAcB ψA ψB + cB2 ψB 2 ) electron density on original atoms,electron density on original atoms, density between atomsdensity between atoms
  • 113. The accumulation of electron density between the nuclei put the electron in a position where it interacts strongly with both nuclei. The energy of the molecule is lower Nuclei are shielded from each other
  • 114. Amplitudes of wave functions subtracted. Destructive interferenceDestructive interference Nodal plane perpendicular to the H-H bondNodal plane perpendicular to the H-H bond axis (en density = 0)axis (en density = 0) Energy of the en in this orbital is higher.Energy of the en in this orbital is higher. +. -. .. node antibonding ψu = N [ψA - ψB] cA = +1, cB = -1 ψu + - ΨA-ΨB
  • 115. The electron is excluded from internuclear regionThe electron is excluded from internuclear region  destabilizingdestabilizing AntibondingAntibonding
  • 116. When 2 atomicWhen 2 atomic orbitalsorbitals combine there are 2combine there are 2 resultantresultant orbitalsorbitals.. low energy bonding orbitallow energy bonding orbital high energyhigh energy antibondingantibonding orbitalorbital 1sb 1sa σ1s σ* E 1s MolecularMolecular orbitalsorbitals EgEg. s. s orbitalsorbitals
  • 117. Molecular potential energy curve shows the variation of the molecular energy with internuclear separation.
  • 118. Looking at the Energy Profile • Bonding orbital • called 1s orbital • s electron • The energy of 1s orbital decreases as R decreases • However at small separation, repulsion becomes large • There is a minimum in potential energy curve
  • 119. 11.4 eV 109 nm H2H2 Location of Bonding orbital 4.5 eV LCAO of n A.O ⇒ n M.O.
  • 120. The overlap integral ∫= τψψ dS BA * The extent to which two atomic orbitals on different atom overlaps : theThe extent to which two atomic orbitals on different atom overlaps : the overlap integraloverlap integral
  • 121. S > 0 Bonding S < 0 anti S = 0 nonbonding Bond strength depends on theBond strength depends on the degree of overlapdegree of overlap
  • 122. Homonuclear Diatomics • MOs may be classified according to: (i) Their symmetry around the molecular axis. (ii) Their bonding and antibonding character. ∀σ1s< σ1s*< σ2s< σ2s*< σ2p< πy(2p) = πz(2p) <πy*(2p) =πz*(2p)<σ2p*.
  • 123. dx2 -dy2 and dxy Cl4Re ReCl4 2-
  • 124. First period diatomic moleculesFirst period diatomic molecules σ1s2 H Energy HH2 1s 1s σg σu* Bond order = ½ (bonding electrons – antibonding electrons) Bond order: 1
  • 125. σ1s2 , σ* 1s2He Energy HeHe2 1s 1s σg σu* Molecular Orbital theory is powerful because it allows us to predict whether molecules should exist or not and it gives us a clear picture of the of the electronic structure of any hypothetical molecule that we can imagine. Diatomic molecules: The bonding in He2 Bond order: 0
  • 126. Second period diatomic moleculesSecond period diatomic molecules σ1s2 , σ* 1s2 , σ2s2 Bond order: 1 Li Energy LiLi2 1s 1s 1σg 1σu* 2s 2s 2σg 2σu*
  • 127. σ1s2 , σ* 1s2 , σ2s2 , σ* 2s2 Bond order: 0 Be Energy BeBe2 1s 1s 1σg 1σu* 2s 2s 2σg 2σu* Diatomic molecules: Homonuclear Molecules of the Second Period
  • 128. SimplifiedSimplified
  • 129. SimplifiedSimplified
  • 130. Molecular orbital theory was put forward by R.S. Mulliken to explain the nature of bonding in the molecules of covalent compounds. Mulliken was awarded Nobel Prize for Chemistry in 1966. Major postulates of the theory are: (i) The wavefunction of an electron in a molecule is called molecular orbital (MO). The molecular orbital surrounds all the nuclei in the molecule, i.e., MO’s are polycentric. (ii) The atomic orbitals (AO’s) of nearly equal energy, and appropriate symmetry combine to give equal number of MO’s. The MO’s are constructed by the linear combination of the atomic orbitals (LCAO method). (iii) MO of lower energy is called bonding molecular orbital ( ψb ), while that of higher energy as antibonding molecular orbital (ψ a ),
  • 131. (v) The electrons of the constituent atoms of a molecule are distributed over all the available MO’s in accordance with the Aufbau principle, the Pauli's exclusion principle and Hund’s rule. (vi) Like atomic orbitals (AO’s), the molecular orbitals can also be arranged according to their energies. The internuclear axis is taken to be in the z-direction. For the molecule or molecular ions formed from Li, Be, B, C, and N, the energies of 2s and 2p orbitals are quite close to each other. Because of the repulsion between the electrons that occupy 2s and 2p orbitals, the energy of the σ2p molecular orbital gets raised. Relative to π 2p orbitals.
  • 132. Splitting patterns for the second row Diatomic If we combine the splitting schemes for the 2s and 2p orbitals, we can predict bond order in all of the diatomic molecules and ions composed of elements in the first complete row of the periodic table. Remember that only the valence orbitals of the atoms need be considered.
  • 133. One minor complication that you should be aware of is that the relative energies of the σ and π bonding molecular orbitals are reversed in some of the second-row diatomics.
  • 134. The presence of one or more unpaired electrons accounts for the paramagnetic nature of the molecule. The electronic configuration in which all the electrons are paired indicate the diamagnetic nature of the species. The strength of a chemical bond is described in terms of a parameter called bond order. As per definition, the bond order is expressed as, Bond order = (No. of electrons in BMO-No. of electrons in ABMO)/2=(Nb-Na)/2 where, N b is the total number of electrons in bonding MOs. N ais the total number of electrons in antibonding MOs. (a) When, N b > N a : Bond order > 0 (+ ve). Then, a stable bond formation is indicated. (b) When, N b =N a : Bond order =0. Then, the bond is unstable. In fact. such a bond is not formed. Conditions For the Formation of MOs From the Atomic Orbitals Formation of MOs by the combination of atomic orbitals takes place only if the following conditions are satisfied: (i) The combining atomic orbitals should have nearly equal energies. Only the atomic orbitals of nearly the same energy combine to form MOs. For example, 1s atomic orbitals of two atoms can combine to form one bonding (σ 1s ) and one antibonding (σ * 1s ) orbitals. The 1s atomic orbital of one atom cannot combine with 2s or 2p atomic orbital of the other atom.
  • 135. (ii) The combining atomic orbitals should have the same symmetry. The atomic orbitals are oriented in space. Only those atomic orbitals can combine to form molecular orbitals which have the same symmetry about the molecular axis. For example, a pxorbital of an atom can combine with a p xorbital of another atom. A p xorbital cannot combine with a pzorbital. (iii) The combining atomic orbitals should overlap effectively. MOs are formed only if the combining atomic orbitals overlap to a reasonable extent. In-phase and out-of-phase wave combinations “Matter waves” corresponding to the two separate hydrogen 1s orbitals interact; both in-phase and out-of-phase combinations are possible, and both occur. One of the resultants is the bonding orbital that we just considered. The other, corresponding to out-of-phase combination of the two orbitals, gives rise to a molecular orbital that has its greatest electron probability in what is clearly the antibonding region of space. This second orbital is therefore called an antibonding orbital.
  • 136. Dicarbon Dioxygen
  • 137. If a hydrogen atom is bonded to a highly electronegative element such as fluorine, oxygen, nitrogen, then the shared pair of electrons lies more towards the electronegative element. This leads to a polarity in the bond in such a way that a slight positive charge gets developed on H-atom, viz., H+δ : O−δ H+δ : F−δ H+δ : N−δ This positive charge on hydrogen can exert electrostatic attraction on the negatively charged electronegative atom of the same or the other molecule forming a bridge-like structure such as Xδ− − Hδ+ × × × × × × Yδ− − Hδ+ where X and Y are the atoms of strongly electronegative elements. The bond between the hydrogen atom of one molecule and a more electronegative atom of the same or another molecule is called hydrogen bond.
  • 138. A hydrogen bond is shown by a dotted line (.....). Hydrogen bond is not a covalent bond as the 1s orbital of hydrogen is already completed, and the 2s level is high up in its energy. Conditions Necessary For the Formation of Hydrogen Bond Hydrogen bond is formed only when the following conditions are satisfied. (i) Only the molecules in which hydrogen atoms is linked to an atom of highly electro-negative element, are capable of forming hydrogen bonds. (ii) The atom of the highly electronegative element should be small. These conditions are met by fluorine, oxygen and nitrogen atoms. As a results, all compounds containing hydrogen atom linked to an atom of either N, O, or F exhibit hydrogen bonding. Some Typical Compounds Showing Hydrogen Bonding Hydrogen fluoride (HF).
  • 139. Water (H2 O). Ice (H2 O(s)). Each Oxygen is "linked" in by a combination of a covalent bond and a hydrogen bond to 4 other Oxygens.
  • 140. Notice that each Oxygen can be linked to Hydrogen in one of two ways. Or Types of Hydrogen Bonding There are two types of hydrogen bonding, viz., (a) Intermolecular hydrogen bonding (b) Intermolecular hydrogen bonding When the hydrogen bonding is between the H-atom of one molecule and an atom of the electronegative element of another molecule, it is termed as intermolecular hydrogen bonding. For example, hydrogen bonding in water, ammonia etc., is intermolecular hydrogen bonding.
  • 141. The intramolecular hydrogen bonding is between the hydrogen of one functional group, and the electronegative atom of the adjacent functional group in the same molecule. For example, the molecule of o-nitrophenol, shows intramolecular hydrogen bonding. The p- nitrophenol shows intermolecular hydrogen bonding.