Enlace quimico

1,332 views

Published on

0 Comments
2 Likes
Statistics
Notes
  • Be the first to comment

No Downloads
Views
Total views
1,332
On SlideShare
0
From Embeds
0
Number of Embeds
4
Actions
Shares
0
Downloads
0
Comments
0
Likes
2
Embeds 0
No embeds

No notes for slide

Enlace quimico

  1. 1. Chapter 8 Basic Concepts of Chemical Bonding Prentice Hall © 2003 Chapter 8 CHEMISTRY The Central Science 9th Edition David P. White
  2. 2. Prentice Hall © 2003 Chapter 8 <ul><li>Chemical bond : attractive force holding two or more atoms together. </li></ul><ul><li>Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. </li></ul><ul><li>Ionic bond results from the transfer of electrons from a metal to a nonmetal. </li></ul><ul><li>Metallic bond : attractive force holding pure metals together. </li></ul>Chemical Bonds, Lewis Symbols, and the Octet Rule
  3. 3. Prentice Hall © 2003 Chapter 8 <ul><li>Lewis Symbols </li></ul><ul><li>As a pictorial understanding of where the electrons are in an atom, we represent the electrons as dots around the symbol for the element. </li></ul><ul><li>The number of electrons available for bonding are indicated by unpaired dots. </li></ul><ul><li>These symbols are called Lewis symbols. </li></ul><ul><li>We generally place the electrons one four sides of a square around the element symbol. </li></ul>Chemical Bonds, Lewis Symbols, and the Octet Rule
  4. 4. Prentice Hall © 2003 Chapter 8 Lewis Symbols Chemical Bonds, Lewis Symbols, and the Octet Rule
  5. 5. Prentice Hall © 2003 Chapter 8 <ul><li>The Octet Rule </li></ul><ul><li>All noble gases except He has an s 2 p 6 configuration. </li></ul><ul><li>Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). </li></ul><ul><li>Caution : there are many exceptions to the octet rule. </li></ul>Chemical Bonds, Lewis Symbols, and the Octet Rule
  6. 6. Prentice Hall © 2003 Chapter 8 Consider the reaction between sodium and chlorine: Na( s ) + ½Cl 2 ( g )  NaCl( s )  H º f = -410.9 kJ Ionic Bonding
  7. 7. Prentice Hall © 2003 Chapter 8 <ul><li>The reaction is violently exothermic. </li></ul><ul><li>We infer that the NaCl is more stable than its constituent elements. Why? </li></ul><ul><li>Na has lost an electron to become Na + and chlorine has gained the electron to become Cl  . Note: Na + has an Ne electron configuration and Cl  has an Ar configuration. </li></ul><ul><li>That is, both Na + and Cl  have an octet of electrons surrounding the central ion. </li></ul>Ionic Bonding
  8. 8. Prentice Hall © 2003 Chapter 8 <ul><li>NaCl forms a very regular structure in which each Na + ion is surrounded by 6 Cl  ions. </li></ul><ul><li>Similarly, each Cl  ion is surrounded by six Na + ions. </li></ul><ul><li>There is a regular arrangement of Na + and Cl  in 3D. </li></ul><ul><li>Note that the ions are packed as closely as possible. </li></ul><ul><li>Note that it is not easy to find a molecular formula to describe the ionic lattice. </li></ul>Ionic Bonding
  9. 9. Prentice Hall © 2003 Chapter 8 Ionic Bonding
  10. 10. Prentice Hall © 2003 Chapter 8 Ionic Bonding <ul><li>Energetics of Ionic Bond Formation </li></ul><ul><li>The formation of Na + ( g ) and Cl  ( g ) from Na( g ) and Cl( g ) is endothermic. </li></ul><ul><li>Why is the formation of Na( s ) exothermic? </li></ul><ul><li>The reaction NaCl( s )  Na + ( g ) + Cl  ( g ) is endothermic (  H = + 788 kJ/mol). </li></ul><ul><li>The formation of a crystal lattice from the ions in the gas phase is exothermic: </li></ul><ul><li>Na + ( g ) + Cl  ( g )  NaCl( s )  H =  788 kJ/mol </li></ul>
  11. 11. Prentice Hall © 2003 Chapter 8 Ionic Bonding <ul><li>Energetics of Ionic Bond Formation </li></ul><ul><li>Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions. </li></ul><ul><li>Lattice energy depends on the charges on the ions and the sizes of the ions: </li></ul><ul><li> is a constant (8.99 x 10 9 J·m/C 2 ), Q 1 and Q 2 are the charges on the ions, and d is the distance between ions. </li></ul>
  12. 12. Prentice Hall © 2003 Chapter 8 Ionic Bonding <ul><li>Energetics of Ionic Bond Formation </li></ul><ul><li>Lattice energy increases as </li></ul><ul><ul><li>The charges on the ions increase </li></ul></ul><ul><ul><li>The distance between the ions decreases. </li></ul></ul>
  13. 14. Prentice Hall © 2003 Chapter 8 Ionic Bonding <ul><li>Electron Configurations of Ions of the Representative Elements </li></ul><ul><li>These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital. </li></ul><ul><li>Electron configurations can predict stable ion formation: </li></ul><ul><ul><li>Mg: [Ne]3 s 2 </li></ul></ul><ul><ul><li>Mg + : [Ne]3 s 1 not stable </li></ul></ul><ul><ul><li>Mg 2+ : [Ne] stable </li></ul></ul><ul><ul><li>Cl: [Ne]3 s 2 3 p 5 </li></ul></ul><ul><ul><li>Cl  : [Ne]3 s 2 3 p 6 = [Ar] stable </li></ul></ul>
  14. 15. Prentice Hall © 2003 Chapter 8 Ionic Bonding <ul><li>Transition Metal Ions </li></ul><ul><li>Lattice energies compensate for the loss of up to three electrons. </li></ul><ul><li>In general, electrons are removed from orbitals in order of decreasing n (i.e. electrons are removed from 4 s before the 3 d ). </li></ul><ul><li>Polyatomic Ions </li></ul><ul><li>Polyatomic ions are formed when there is an overall charge on a compound containing covalent bonds. E.g. SO 4 2  , NO 3  . </li></ul>
  15. 16. Prentice Hall © 2003 Chapter 8 Covalent Bonding <ul><li>When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. </li></ul><ul><li>When similar atoms bond, they share pairs of electrons to each obtain an octet. </li></ul><ul><li>Each pair of shared electrons constitutes one chemical bond. </li></ul><ul><li>Example: H + H  H 2 has electrons on a line connecting the two H nuclei. </li></ul>
  16. 17. Prentice Hall © 2003 Chapter 8 Covalent Bonding
  17. 18. Prentice Hall © 2003 Chapter 8 Covalent Bonding <ul><li>Lewis Structures </li></ul><ul><li>Covalent bonds can be represented by the Lewis symbols of the elements: </li></ul><ul><li>In Lewis structures, each pair of electrons in a bond is represented by a single line: </li></ul>
  18. 19. Prentice Hall © 2003 Chapter 8 Covalent Bonding <ul><li>Multiple Bonds </li></ul><ul><li>It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): </li></ul><ul><ul><li>One shared pair of electrons = single bond (e.g. H 2 ); </li></ul></ul><ul><ul><li>Two shared pairs of electrons = double bond (e.g. O 2 ); </li></ul></ul><ul><ul><li>Three shared pairs of electrons = triple bond (e.g. N 2 ). </li></ul></ul><ul><li>Generally, bond distances decrease as we move from single through double to triple bonds. </li></ul>
  19. 20. Prentice Hall © 2003 Chapter 8
  20. 21. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity <ul><li>In a covalent bond, electrons are shared. </li></ul><ul><li>Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. </li></ul><ul><li>There are some covalent bonds in which the electrons are located closer to one atom than the other. </li></ul><ul><li>Unequal sharing of electrons results in polar bonds. </li></ul>
  21. 22. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity <ul><li>Electronegativity </li></ul><ul><li>Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. </li></ul><ul><li>Pauling set electronegativities on a scale from 0.7 ( Cs ) to 4.0 ( F ). </li></ul><ul><li>Electronegativity increases </li></ul><ul><ul><li>across a period and </li></ul></ul><ul><ul><li>down a group. </li></ul></ul>
  22. 23. Bond Polarity and Electronegativity Electronegativity
  23. 24. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity <ul><li>Electronegativity and Bond Polarity </li></ul><ul><li>Difference in electronegativity is a gauge of bond polarity: </li></ul><ul><ul><li>electronegativity differences around 0 result in non-polar covalent bonds (equal or almost equal sharing of electrons); </li></ul></ul><ul><ul><li>electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons); </li></ul></ul><ul><ul><li>electronegativity differences around 3 result in ionic bonds (transfer of electrons). </li></ul></ul>
  24. 25. Prentice Hall © 2003 Chapter 8
  25. 26. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity <ul><li>Electronegativity and Bond Polarity </li></ul><ul><li>There is no sharp distinction between bonding types. </li></ul><ul><li>The positive end (or pole) in a polar bond is represented  + and the negative pole  -. </li></ul>
  26. 27. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity <ul><li>Dipole Moments </li></ul><ul><li>Consider HF: </li></ul><ul><ul><li>The difference in electronegativity leads to a polar bond. </li></ul></ul><ul><ul><li>There is more electron density on F than on H. </li></ul></ul><ul><ul><li>Since there are two different “ends” of the molecule, we call HF a di pole. </li></ul></ul><ul><li>Dipole moment ,  , is the magnitude of the dipole: </li></ul><ul><li>where Q is the magnitude of the charges. </li></ul><ul><li>Dipole moments are measured in debyes , D. </li></ul>
  27. 28. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity <ul><li>Bond Types and Nomenclature </li></ul><ul><li>In general, the least electronegative element is named first. </li></ul><ul><li>The name of the more electronegative element ends in – ide . </li></ul><ul><li>Ionic compounds are named according to their ions, including the charge on the cation if it is variable. </li></ul><ul><li>Molecular compounds are named with prefixes. </li></ul>
  28. 29. Prentice Hall © 2003 Chapter 8
  29. 30. Prentice Hall © 2003 Chapter 8 Bond Polarity and Electronegativity Bond Types and Nomenclature Ionic Molecular MgH 2 Magnesium hydride H 2 S Hydrogen sulfide FeF 2 Iron(II) fluoride OF 2 Oxygen difluoride Mn 2 O 3 Manganese(III) oxide Cl 2 O 3 Dichlorine trioxide
  30. 31. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Add the valence electrons. </li></ul><ul><li>Write symbols for the atoms and show which atoms are connected to which. </li></ul><ul><li>Complete the octet for the central atom the complete the octets of the other atoms. </li></ul><ul><li>Place leftover electrons on the central atom. </li></ul><ul><li>If there are not enough electrons to give the central atom an octet, try multiple bonds. </li></ul>
  31. 32. Prentice Hall © 2003 Chapter 8
  32. 33. Prentice Hall © 2003 Chapter 8
  33. 34. Prentice Hall © 2003 Chapter 8
  34. 35. Prentice Hall © 2003 Chapter 8
  35. 36. Prentice Hall © 2003 Chapter 8
  36. 37. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Formal Charge </li></ul><ul><li>It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. </li></ul><ul><li>To determine which structure is most reasonable, we use formal charge. </li></ul><ul><li>Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity. </li></ul>
  37. 38. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Formal Charge </li></ul><ul><li>To calculate formal charge: </li></ul><ul><ul><li>All nonbonding electrons are assigned to the atom on which they are found. </li></ul></ul><ul><ul><li>Half the bonding electrons are assigned to each atom in a bond. </li></ul></ul><ul><li>Formal charge is: </li></ul><ul><li>valence electrons - number of bonds - lone pair electrons </li></ul>
  38. 39. Prentice Hall © 2003 Chapter 8
  39. 40. Prentice Hall © 2003 Chapter 8
  40. 41. Prentice Hall © 2003 Chapter 8
  41. 42. Prentice Hall © 2003 Chapter 8
  42. 43. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Formal Charge </li></ul><ul><li>Consider: </li></ul><ul><li>For C : </li></ul><ul><ul><li>There are 4 valence electrons (from periodic table). </li></ul></ul><ul><ul><li>In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. </li></ul></ul><ul><ul><li>Formal charge: 4 - 5 = -1. </li></ul></ul>
  43. 44. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Formal Charge </li></ul><ul><li>Consider: </li></ul><ul><li>For N : </li></ul><ul><ul><li>There are 5 valence electrons. </li></ul></ul><ul><ul><li>In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. </li></ul></ul><ul><ul><li>Formal charge = 5 - 5 = 0. </li></ul></ul><ul><li>We write: </li></ul>
  44. 45. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Formal Charge </li></ul><ul><li>The most stable structure has: </li></ul><ul><ul><li>the lowest formal charge on each atom, </li></ul></ul><ul><ul><li>the most negative formal charge on the most electronegative atoms. </li></ul></ul><ul><li>Resonance Structures </li></ul><ul><li>Some molecules are not well described by Lewis Structures. </li></ul><ul><li>Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms </li></ul>
  45. 46. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Resonance Structures </li></ul><ul><li>Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single ( longer ) and one double bond ( shorter ). </li></ul>
  46. 47. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures Resonance Structures
  47. 48. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Resonance Structures </li></ul><ul><li>Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities. </li></ul>
  48. 49. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Resonance Structures </li></ul><ul><li>Example: in ozone the extreme possibilities have one double and one single bond. The resonance structure has two identical bonds of intermediate character. </li></ul><ul><li>Common examples: O 3 , NO 3 - , SO 4 2- , NO 2 , and benzene. </li></ul>
  49. 50. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Resonance in Benzene </li></ul><ul><li>Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom. </li></ul><ul><li>There are alternating double and single bonds between the C atoms. </li></ul><ul><li>Experimentally, the C-C bonds in benzene are all the same length. </li></ul><ul><li>Experimentally, benzene is planar. </li></ul>
  50. 51. Prentice Hall © 2003 Chapter 8 Drawing Lewis Structures <ul><li>Resonance in Benzene </li></ul><ul><li>We write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring: </li></ul><ul><li>Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor). </li></ul>
  51. 52. Prentice Hall © 2003 Chapter 8 Exceptions to the Octet Rule <ul><li>There are three classes of exceptions to the octet rule: </li></ul><ul><ul><li>Molecules with an odd number of electrons; </li></ul></ul><ul><ul><li>Molecules in which one atom has less than an octet; </li></ul></ul><ul><ul><li>Molecules in which one atom has more than an octet. </li></ul></ul><ul><li>Odd Number of Electrons </li></ul><ul><li>Few examples. Generally molecules such as ClO 2 , NO, and NO 2 have an odd number of electrons. </li></ul>
  52. 53. Prentice Hall © 2003 Chapter 8 Exceptions to the Octet Rule <ul><li>Less than an Octet </li></ul><ul><li>Relatively rare. </li></ul><ul><li>Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. </li></ul><ul><li>Most typical example is BF 3 . </li></ul><ul><li>Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds. </li></ul>
  53. 54. Prentice Hall © 2003 Chapter 8
  54. 55. Prentice Hall © 2003 Chapter 8 Exceptions to the Octet Rule <ul><li>More than an Octet </li></ul><ul><li>This is the largest class of exceptions. </li></ul><ul><li>Atoms from the 3 rd period onwards can accommodate more than an octet. </li></ul><ul><li>Beyond the third period, the d -orbitals are low enough in energy to participate in bonding and accept the extra electron density. </li></ul>
  55. 56. Prentice Hall © 2003 Chapter 8
  56. 57. Prentice Hall © 2003 Chapter 8
  57. 58. Prentice Hall © 2003 Chapter 8
  58. 59. Prentice Hall © 2003 Chapter 8
  59. 60. Prentice Hall © 2003 Chapter 8 Strengths of Covalent Bonds <ul><li>The energy required to break a covalent bond is called the bond dissociation enthalpy , D . That is, for the Cl 2 molecule, D (Cl-Cl) is given by  H for the reaction: </li></ul><ul><li>Cl 2 ( g )  2Cl( g ). </li></ul><ul><li>When more than one bond is broken: </li></ul><ul><li>CH 4 ( g )  C( g ) + 4H( g )  H = 1660 kJ </li></ul><ul><li>the bond enthalpy is a fraction of  H for the atomization reaction: </li></ul><ul><li>D (C-H) = ¼  H = ¼(1660 kJ) = 415 kJ. </li></ul><ul><li>Bond enthalpies can either be positive or negative. </li></ul>
  60. 62. Prentice Hall © 2003 Chapter 8 Strengths of Covalent Bonds <ul><li>Bond Enthalpies and the Enthalpies of Reactions </li></ul><ul><li>We can use bond enthalpies to calculate the enthalpy for a chemical reaction. </li></ul><ul><li>We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed. </li></ul><ul><li>The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken less the sum of bond enthalpies for bonds formed. </li></ul>
  61. 63. Prentice Hall © 2003 Chapter 8 Strengths of Covalent Bonds <ul><li>Bond Enthalpies and the Enthalpies of Reactions </li></ul><ul><li>Mathematically, if  H rxn is the enthalpy for a reaction, then </li></ul><ul><li>We illustrate the concept with the reaction between methane, CH 4 , and chlorine: </li></ul><ul><li>CH 4 ( g ) + Cl 2 ( g )  CH 3 Cl( g ) + HCl( g )  H rxn = ? </li></ul>
  62. 64. Strengths of Covalent Bonds
  63. 65. Prentice Hall © 2003 Chapter 8 Strengths of Covalent Bonds <ul><li>Bond Enthalpies and the Enthalpies of Reactions </li></ul><ul><li>In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed. </li></ul><ul><li>The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken. </li></ul><ul><li>The above result is consistent with Hess’s law. </li></ul>
  64. 66. Prentice Hall © 2003 Chapter 8 Strengths of Covalent Bonds <ul><li>Bond Enthalpy and Bond Length </li></ul><ul><li>We know that multiple bonds are shorter than single bonds. </li></ul><ul><li>We can show that multiple bonds are stronger than single bonds. </li></ul><ul><li>As the number of bonds between atoms increases, the atoms are held closer and more tightly together. </li></ul>
  65. 67. Strengths of Covalent Bonds
  66. 68. End of Chapter 8: Basic Concepts of Chemical Bonding Prentice Hall © 2003 Chapter 8

×