(L6) chemical bonding

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(L6) chemical bonding

  1. 1. 1 CHAPTER 7  Chemical Bonding
  2. 2. 2 Introduction  Attractive forces that hold atoms together in compounds are called chemical bonds.  The electrons involved in bonding are usually those in the outermost (valence) shell.
  3. 3. 3 Introduction  Chemical bonds are classified into two types: o Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another. o Covalent bonding results from sharing one or more electron pairs between two atoms.
  4. 4. 4 Lewis Dot Formulas of Atoms  Lewis dot formulas or Lewis dot representations are a convenient bookkeeping method for tracking valence electrons.  Valence electrons are those electrons that are transferred or involved in chemical bonding. • They are chemically important.
  5. 5. 5 Lewis Dot Formulas of Atoms Li Be B C N O F Ne H . He Li Be B C N O F Ne .. HeH . Li Be B C N O F Ne .. HeH . . Li Be B C N O F Ne .. .. HeH . . Li Be B C N O F Ne .. .. .. HeH . . . Li Be B C N O F Ne .. .. .. .. HeH . . . . .Li Be B C N O F Ne .... .. .. .. HeH . . . . . . . . Li Be B C N O F Ne .... .. .. .. HeH . . . . . .. . . . . .. .Li Be B C N O F Ne .... .. .. .. HeH . . . . . .. .. .. . . . . . . . . . .Li Be B C N O F Ne .... .. .. .. HeH . . . . . .. .. .. .. . . . . . . . . . . . . . .. .
  6. 6. 6 Lewis Dot Formulas of Atoms  Elements that are in the same periodic group have the same Lewis dot structures. Li & Na . . N & P .. .. . . . . . . F & Cl.. . .. . . . . .. .. .
  7. 7. 7 Ionic Bonding Formation of Ionic Compounds  An ion is an atom or a group of atoms possessing a net electrical charge.  Ions come in two basic types: 1. positive (+) ions or cations • These atoms have lost 1 or more electrons. 2. negative (-) ions or anions • These atoms have gained 1 or more electrons.
  8. 8. 8 Covalent Bonding  Covalent bonds are formed when atoms share electrons.  If the atoms share 2 electrons a single covalent bond is formed.  If the atoms share 4 electrons a double covalent bond is formed.  If the atoms share 6 electrons a triple covalent bond is formed.  The attraction between the electrons is electrostatic in nature • The atoms have a lower potential energy when bound.
  9. 9. 9 Formation of Covalent Bonds  This figure shows the potential energy of an H2 molecule as a function of the distance between the two H atoms.
  10. 10. 10 Formation of Covalent Bonds  Representation of the formation of an H2 molecule from H atoms.
  11. 11. 11 Formation of Covalent Bonds  We can use Lewis dot formulas to show covalent bond formation. 1. H molecule formation representation. +H. H. H H.. or H2 H Cl H Cl+. .. . . .. .. .. .. .. . or HCl 2. HCl molecule formation
  12. 12. 12 Lewis Formulas for Molecules and Polyatomic Ions  First, we explore Lewis dot formulas of homonuclear diatomic molecules.  Two atoms of the same element. 1. Hydrogen molecule, H2. H HorH H. . F F .. . . . . . . .. .. .. F F . . . . .. .. .. .. or N N········ ·· N N·· ··or 2. Fluorine, F2. 3. Nitrogen, N2.
  13. 13. 13 Lewis Formulas for Molecules and Polyatomic Ions  Next, look at heteronuclear diatomic molecules.  Two atoms of different elements. • Hydrogen halides are good examples. 1. hydrogen fluoride, HF or ··H F ·· ·· H F. . ·· ·· ·· or ··H Cl ·· ·· H Cl. . ·· ·· ·· or ··H Br ·· ·· H Br . . ·· ·· ·· 2. hydrogen chloride, HCl 3. hydrogen bromide, HBr
  14. 14. 14 Lewis Formulas for Molecules and Polyatomic Ions  Now we will look at a series of slightly more complicated heteronuclear molecules.  Water, H2O H H O ·· ·· ·· ··
  15. 15. 15 Lewis Formulas for Molecules and Polyatomic Ions  Ammonia molecule , NH3 H H N ·· ·· ·· ·· H
  16. 16. 16 Lewis Formulas for Molecules and Polyatomic Ions  Lewis formulas can also be drawn for molecular ions.  One example is the ammonium ion , NH4 +. H H N ·· ·· ·· ·· H H + •Notice that the atoms other than H in these molecules have eight electrons around them.
  17. 17. 17 Polar and Nonpolar Covalent Bonds  Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds.  Nonpolar covalent bonds have a symmetrical charge distribution.  To be nonpolar the two atoms involved in the bond must be the same element to share equally.
  18. 18. 18 Polar and Nonpolar Covalent Bonds  Some examples of nonpolar covalent bonds.  H2 H HorH H. . N N········ ·· N N·· ··or N2
  19. 19. 19 Polar and Nonpolar Covalent Bonds  Covalent bonds in which the electrons are not shared equally are designated as polar covalent bonds  Polar covalent bonds have an asymmetrical charge distribution  To be a polar covalent bond the two atoms involved in the bond must have different electronegativities.
  20. 20. 20 Polar and Nonpolar Covalent Bonds  Some examples of polar covalent bonds.  HF bondpolarvery1.9Difference 4.02.1ativitiesElectroneg FH 1.9 
  21. 21. 21 Polar and Nonpolar Covalent Bonds  Polar molecules can be attracted by magnetic and electric fields.
  22. 22. Bond Polarity 22  Recall that a polar bond has an asymmetric distribution of electrons  X-X is nonpolar  X-Y is polar  Polarity of a bond increases with increasing difference in electronegativity between the two atoms  Bond is a dipole  One end is (δ+), while the other is (δ-)
  23. 23. 23
  24. 24. Electronegativity Scales  Pauling Scale  Relative scale used to measure electronegativity  Atoms with vastly different electronegativity  ionic cmpds.  Atoms with somewhat different electronegativity  polar covalent cmpd.  Nonpolar cmpds are electronically symmetrical 24
  25. 25. 25 Writing Lewis Formulas: The Octet Rule  The octet rule states that representative elements usually attain stable noble gas electron configurations in most of their compounds.  Lewis dot formulas are based on the octet rule.  We need to distinguish between bonding (or shared) electrons and nonbonding (or unshared or lone pairs) of electrons.
  26. 26. 26 Writing Lewis Formulas: The Octet Rule  N - A = S rule  Simple mathematical relationship to help us write Lewis dot formulas.  N = number of electrons needed to achieve a noble gas configuration.  N usually has a value of 8 for representative elements.  N has a value of 2 for H atoms.  A = number of electrons available in valence shells of the atoms.  A is equal to the periodic group number for each element.  A is equal to 8 for the noble gases.  S = number of electrons shared in bonds.  A-S = number of electrons in unshared, lone, pairs.
  27. 27. 27 Writing Lewis Formulas: The Octet Rule  For ions we must adjust the number of electrons available, A.  Add one e- to A for each negative charge.  Subtract one e- from A for each positive charge.  The central atom in a molecule or polyatomic ion is determined by:  The atom that requires the largest number of electrons to complete its octet goes in the center.  For two atoms in the same periodic group, the less electronegative element goes in the center.
  28. 28. 28 Writing Lewis Formulas: The Octet Rule  Example 7-2: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.  N = 2 (H) + 8 (C) + 8 (N) = 18  A = 1 (H) + 4 (C) + 5 (N) = 10  S = 8  A-S = 2  This molecule has 8 electrons in shared pairs and 2 electrons in lone pairs. H C N·· ·· ···· H C N··or··
  29. 29. 29 Writing Lewis Formulas: The Octet Rule  Example 7-3: Write Lewis dot and dash formulas for the sulfite ion, SO3 2-. N = 8 (S) + 3 x 8 (O) = 32 A = 6 (S) + 3 x 6 (O) + 2 (- charge) = 26 S = 6 A-S = 20  Thus this polyatomic ion has 6 electrons in shared pairs and 20 electrons in lone pairs.  Which atom is the central atom in this ion? You do it!
  30. 30. 30 Writing Lewis Formulas: The Octet Rule  What kind of covalent bonds, single, double, or triple, must this ion have so that the six shared electrons are used to attach the three O atoms to the S atom? O S O O ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· 2- O S O O ·· ·· ·· ···· ·· ·· ·· ···· 2- or
  31. 31. 31 Resonance  Example 7-4: Write Lewis dot and dash formulas for sulfur trioxide, SO3. You do it! N = 8 (S) + 3 x 8 (O) = 32 A = 6 (S) + 3 x 6 (O) = 24 S = 8 A-S = 16 orO S O O ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· O S O O ·· ···· ·· ·· ·· ····
  32. 32. 32 Resonance  There are three possible structures for SO3.  The double bond can be placed in one of three places. O S O O ·· ···· ·· ·· ·· ···· OS O O ·· ···· ·· ·· ·· ···· ·· O S O O ·· ···· ·· ·· ·· ·· oWhen two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule’s structure. oDouble-headed arrows are used to indicate resonance formulas.
  33. 33. 33 Resonance  Resonance is a flawed method of representing molecules.  There are no single or double bonds in SO3 . • In fact, all of the bonds in SO3 are equivalent.  The best Lewis formula of SO3 that can be drawn is: SO O O
  34. 34. 34 Formal Charge • Calculation of a formal charge on a molecule is a mechanism for determining correct Lewis structures • The formal charge is the hypothetical charge on an atom in a molecule or polyatomic ion. • The best Lewis structures will have formal charges on the atoms that are zero or nearly zero.
  35. 35. 35 Formal Charge Rules for Assigning Formal Charge 1. Formal Charge = grp number – (number of bonds + number of unshared e-) 2. An atom that has the same number of bonds as its periodic group number has a formal charge of 0. 3. a. The formal charges of all atoms must sum to 0 in molecules. b. The formal charges must sum to the ion’s charge for a polyatomic ion.
  36. 36. 36 Formal Charge Cl 7 – (2+4) = +1 N 5 – (3+2) = 0 O 6 – (1+6) = -1 Cl 7 – (1+6) = 0 N 5 – (3+2) = 0 O 6 – (2+4) = 0 Cl N O Cl N O Consider nitrosyl chloride, NOCl
  37. 37. 37 Writing Lewis Formulas: Limitations of the Octet Rule  There are some molecules that violate the octet rule.  For these molecules the N - A = S rule does not apply: 1. The covalent compounds of Be. 2. The covalent compounds of the IIIA Group. 3. Species which contain an odd number of electrons. 4. Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents. 5. Compounds of the d- and f-transition metals.
  38. 38. 38 Writing Lewis Formulas: Limitations of the Octet Rule  In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations.  The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5).
  39. 39. 39 Writing Lewis Formulas: Limitations of the Octet Rule  Example 7-5: Write dot and dash formulas for BBr3.  This is an example of exception #2. You do it! B ·· . Br·· ·· ·· . BBr Br Br ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· Br B Br Br ·· ·· ·· ·· ·· ·· ·· ·· ·· or
  40. 40. 40 Writing Lewis Formulas: Limitations of the Octet Rule  Example 7-6: Write dot and dash formulas for AsF5 . You do it! As ·· . . . F·· ·· ·· . ·· As F F F F F ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· or ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· ·· AsF F F F F ·· ·· ·· ·· ·· ··
  41. 41. Lewis Structures 41
  42. 42. Lewis Structures 42
  43. 43. 43 The Continuous Range of Bonding Types  Covalent and ionic bonding represent two extremes. 1. In pure covalent bonds electrons are equally shared by the atoms. 2. In pure ionic bonds electrons are completely lost or gained by one of the atoms.  Most compounds fall somewhere between these two extremes.
  44. 44. 44 Continuous Range of Bonding Types  All bonds have some ionic and some covalent character.  For example, HI is about 17% ionic  The greater the electronegativity differences the more polar the bond.
  45. 45. -End- 45

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