bonding in chemistry-govt. model science college jabalpur [ Manorama Singh ]


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bonding in chemistry-govt. model science college jabalpur [ Manorama Singh ]

  1. 1. Bonding and MolecularStructure: Fundamental Concepts Chapter 9 Chapter 9 1
  2. 2. Valence Electrons- The electrons involved in bonding are called valence electrons.- Valence electrons are found in the incomplete, outermost orbital shell of an atom.- We can represent the electrons as dots around the symbol for the element.- These pictorial representations are called Lewis Structures or Lewis Dot Structures. Chapter 9 2
  3. 3. Lewis Symbols and the Octet Rule Chapter 9 3
  4. 4. Chemical Bond Formation- There are three types of chemical bonds Ionic Bond - electrostatic attraction between ions of opposite charge (NaCl). Covalent Bond - sharing of electrons between two atoms (Cl2). Metallic Bond - sharing of electrons between several atoms (Ag). Chapter 9 4
  5. 5. Ionic BondingConsider the reaction between sodium and chlorine: Na(s) + ½Cl2(g) → NaCl(s) Chapter 9 5
  6. 6. Ionic Bonding Na(s) + ½Cl2(g) → NaCl(s) ∆H°f = -410.9 kJ- This reaction is very exothermic- Sodium loses an electron to become Na+- Chlorine gains an electron to become Cl−- Na+ has an Ne electron configuration and Cl− has an Ar configuration Chapter 9 6
  7. 7. Ionic BondingEnergetics of Ionic Bond FormationLattice Energy (∆Hlattice) – The energy required tocompletely separate one mole of a solid ionic compoundinto its gaseous ions.Lattice energy depends on -the charge on the ions -the size of the ionsCoulomb’s equation: Q1Q2 E=k d Q1, Q2 = charge on ions k = 8.99 x 109 J-m/c2 d = distance between ions Chapter 9 7
  8. 8. Covalent BondingWhen similar atoms bond, they share pairs of electrons to each obtain an octet.Example Cl + Cl Cl Cl a pair of electrons connect the two nuclei. Chapter 9 8
  9. 9. Covalent BondingMultiple Bonds- It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds) - One shared pair of electrons - single bond (H2) - Two shared pairs of electrons - double bond (O2) - Three shared pairs of electrons - triple bond (N2). H H O O N N- Generally, bond distances decrease as we move from single through double to triple bonds. Chapter 9 9
  10. 10. Lewis Symbols and the Octet RuleOctet rule – Atoms tend to gain, lose or share electronsuntil they are surrounded by eight valence electrons. Chapter 9 10
  11. 11. Drawing Lewis Structures1) Draw a skeleton structure of the molecule or ion showing the arrangement of the atoms and the connect each atom to another with a single bond.2) Determine the total number of valence elections in the molecule or ion.3) Deduct 2 electrons for each single bond used in step 1.4) Distribute the rest of the electrons so that each atom (except H) has 8 electrons. - If you are “short” electrons, form multiple bonds - If you have “extra” electrons, one of the heavy atoms may be able to hold more that eight electrons. Chapter 9 11
  12. 12. Drawing Lewis StructuresPCl3 Cl P Cl Cl Element Number Electrons Total P 1 5 5 Cl 3 7 21 Total Electrons 26 Chapter 9 12
  13. 13. Drawing Lewis StructuresPCl3 Cl P Cl Cl Element Number Electrons Total P 1 5 5 Cl 3 7 21 Total Electrons 26 Electrons used 6 Electrons remaining 20 Cl P Cl Cl Chapter 9 13
  14. 14. Drawing Lewis StructuresIsoelectronic Species Molecules or ions having the same number of valence electrons and the same Lewis structure. N O + N N C O Chapter 9 14
  15. 15. Drawing Lewis StructuresResonance Structures- Some molecules are not well described by Lewis Structures.Example: Ozone O O 1 O O3 1 O O3 Chapter 9 15
  16. 16. Drawing Lewis StructuresResonance Structures- Experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single and one double bond. O O O Chapter 9 16
  17. 17. Drawing Lewis StructuresResonance Structures- Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities.- Each Lewis structure is call a Resonance FormResonance Form – Two or more Lewis structures having the same arrangements of atoms but a different arrangement of electrons Chapter 9 17
  18. 18. Drawing Lewis StructuresResonance Structures- In ozone the resonance forms have one double and one single bond. O O O O O O- The actual structure of O3 is a combination (or average) of the individual forms called a resonance hybrid. Chapter 9 18
  19. 19. Exceptions to the Octet RuleThere are three classes of exceptions to the octet rule: - Molecules with an odd number of electrons - Molecules in which one atom has less than an octet - Molecules in which one atom has more than an octet Chapter 9 19
  20. 20. Exceptions to the Octet RuleOdd Number of Electrons- there are few molecules which fit this categoryExamples. ClO2, NO, and NO2 O N O Chapter 9 20
  21. 21. Exceptions to the Octet RuleLess than an Octet- This refers to the central molecule- Typical for compounds of Groups 1A, 2A, and 3A.Examples: LiH, BeH2, BF3 F Li F H Be H B F F Chapter 9 21
  22. 22. Exceptions to the Octet RuleMore than an Octet- This starts for atoms in the 3rd period onwards.- This is due to vacant d orbitals which can hold the “extra” electrons.- Another factor is the size of the central atom, as they get bigger, it gets easier to place additional atoms around the central atom. Chapter 9 22
  23. 23. Molecular ShapesLewis structures give atomic connectivity (which atoms are connected to which). Chapter 9 23
  24. 24. Molecular ShapesMolecular Shapes are determined by: Bond Distance – Distance between the nuclei of two bonded atoms along a straight line. Bond Angle – The angle between any two bonds containing a common atom. Chapter 9 24
  25. 25. Molecular Shapes Chapter 9 25
  26. 26. Molecular ShapesValence Shell Electron Pair Repulsion Theory (VSEPR)- VSEPR theory is based on the idea that electrostatic repulsion of the electrons are reduced to a minimum when the various regions of high electron density assume positions as far apart as possible. Chapter 9 26
  27. 27. Molecular ShapesPredicting Molecular Geometries - draw the Lewis structure - count the total number of bonding regions and lone pairs around the central atom - arrange the bonding regions and lone pairs in one of the standard geometries to minimize e− -e− repulsion - multiple bonds count as one bonding region Chapter 9 27
  28. 28. Molecular ShapesPredicting Molecular GeometriesCommon Configuration for saturated molecules. Regions of Electron-Pair Bond Angle Density Geometry 2 (AX2) Linear 180o 3 (AX3) Trigonal Planar 120o 4 (AX4) Tetrahedral 109.5o 5 (AX5) Trigonal 90o / 120o Bipyramidal 6 (AX6) Octahedral 90o Chapter 9 28
  29. 29. Molecular Shapes- The “region of electron density” refers to: - Lone pairs - Covalent bonds (single, double, triple)- Remember, you can’t “see” lone-pairs but they do take-up space. Chapter 9 29
  30. 30. Molecular ShapesPredicting Molecular Geometries Chapter 9 30
  31. 31. Molecular ShapesPredicting Molecular Geometries Chapter 9 31
  32. 32. Molecular ShapesMolecules with Expanded Valence Shells Chapter 9 32
  33. 33. Molecular Shapes To minimize e− −e− repulsion, lone pairs are always placed in equatorial positions. Chapter 9 33
  34. 34. Molecular ShapesMolecules with Expanded Valence Shells Chapter 9 34
  35. 35. Charge DistributionFormal ChargeUsed to predict the correct Lewis Structure. 1) Half of the electrons in a bond are assigned to each atom in a bond. 2) Both electrons of an unshared pair of electrons are assigned to the atoms to which the unshared pair belong. 3) The formal charge of an atom is equal to the valence electrons minus the number of electrons assigned to each atom.Formal Charge = (group number) – (assigned electrons) 4) The sum of the formal charges will equal the charge on the molecule or polyatomic ion. Chapter 9 35
  36. 36. Charge DistributionUsing Formal Charge1) A Lewis structure in which all formal charges in a molecule are equal to zero is preferable to one in which some formal charges are not zero.2) If a Lewis structure has non-zero formal charges, the one with the fewest nonzero formal charges is preferred.3) A Lewis structure with one large formal charge is preferable to one with several small formal charges.4) A Lewis structure with adjacent formal charges should have opposite signs.5) When choosing between several Lewis structures, the structure with negative formal charges on the more electronegative atom is preferable. Chapter 9 36
  37. 37. Bond Polarity and Electronegativity- Electrons in a covalent bond may not be shared evenly.Electronegativity – The ability of an atom in a molecule to attract electrons to itself.- The periodic trend for electronegativity is up and to the right across the periodic table. Chapter 9 37
  38. 38. Bond Polarity and ElectronegativityElectronegativity Chapter 9 38
  39. 39. Bond Polarity and ElectronegativityElectronegativity and Bond Polarity- A chemical bond between elements with large differences in eletronegativity will shift the electrons to the atom with the higher electronegativity.- The positive end (or pole) in a polar bond is represented δ+ and the negative pole δ-.- This is called a polar covalent bond.- If the electronegativity difference is small, the bond is nonpolar; if it is large, it is a polar bond. Chapter 9 39
  40. 40. Polarity of Molecules- To determine if a molecule is polar, you need to know two things: - polarity of the bonds in a molecule - how the bonds are arranged- A molecule is considered polar if its center of negative and positive charge do not coincide. δ+ δ− H F H F- Polar molecules have a dipole (a vector quantity)- If these dipoles act equally and in opposition to each other, the dipoles cancel-out and the molecule is considered nonpolar. Chapter 9 40
  41. 41. Polarity of MoleculesDipole Moments of Polyatomic MoleculesExample:CO2, each C-O dipole is canceled because themolecule is linear.H2O, the H-O dipoles do not cancel because themolecule is bent. Chapter 9 41
  42. 42. Polarity of MoleculesDipole Moments of Polyatomic MoleculesTwo simple rules to help determine molecular polarity (most of the time)1. If there are lone pairs on the central atom – the molecule is polar.2. If there is more than one type of bond on the central atom – the molecule is polar. Chapter 9 42
  43. 43. Polarity of MoleculesDipole Moments of Polyatomic Molecules Chapter 9 43
  44. 44. Strengths of Covalent BondsBond Enthalpy (Energy) - The energy required to break a covalent bond of a gaseous substance. Cl2(g) → 2Cl(g) ∆H = DCl-Cl- When more than one bond is broken the bond enthalpy is a fraction of ∆H for the atomization reaction : CH4(g) → C(g) + 4H(g) ∆H = 1660 kJ DC-H = ¼∆H = ¼(1660 kJ) = 415 kJ.- Bond enthalpies can either be positive or negative. Chapter 9 44
  45. 45. Strengths of Covalent Bonds Chapter 9 45
  46. 46. Strengths of Covalent BondsBond Enthalpies and the Enthalpies of Reaction- Bond enthalpies can be used to calculate ∆Hrxn. ∆Hrxn = ∑D(bonds broken) − ∑D(bonds formed). Chapter 9 46
  47. 47. Bond Enthalpies and the Enthalpies of Reaction H H H H C C + H O O H H O C C O H H H H H Bonds Broken Bonds Formed C=C C-C 614 kJ/mol 348 kJ/mol O-O C-O 146 kJ/mol 358 kJ/mol∆Hrxn = [1mol(614kJ/mol)+1mol(146kJ/mol)]-[2mol(358kJ/mol)+1mol(348kJ/mol)] = -304 kJ Chapter 9 47
  48. 48. Bond Enthalpies and the Enthalpies of Reaction Cl N Cl 2 N N + 3 Cl Cl Cl Bonds Broken Bonds Formed N-Cl N=N 200 kJ/mol 941 kJ/mol Cl-Cl 242 kJ/mol ∆Hrxn = [2(3mol(200kJ/mol))]-[1mol(941kJ/mol)+3mol(242kJ/mol)] = -467 kJ Chapter 9 48
  49. 49. Homework 2, 6, 18, 22, 26, 32, 46, 57 Chapter 9 49