Chemical Bonding: Chapter 7. Chapter Outline. Chapter Outline

Transcription

1 Chemical Bonding: Chapter 7 Chapter utline Lewis Dot Formulas of Atoms Ionic Bonding Formation of Ionic Compounds Covalent Bonding Formation of Covalent Bonds Lewis Formulas for Molecules and Polyatomic Ions Writing Lewis Formulas: The ctet Rule Chapter utline Resonance Writing Lewis Formulas: Limitations of the ctet Rule Polar and onpolar Covalent Bonds Dipole Moments The Continuous Range of Bonding Types 1

2 Formal Charge An accounting device, not real charge on the atoms - - valence e s valence e s in free atom in structure Sum of FC s = zero for a molecule, or charge on an ion Minimize FC s to get best structure Resonance Write Lewis dot and dash formulas for sulfur trioxide, S 3 = 8 (S) + 3 x 8 () = 32 A = 6 (S) + 3 x 6 () = 24 S = 8 A-S = 16 S or S Resonance There are three possible structures for S 3 The double bond can be placed in one of three places S S S When 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 Double-headed arrows are used to indicate resonance formulas 2

3 Bond Lengths: An Experimental Test Å (Å) = = 1150 Å (both) Resonance Resonance is a flawed method of representing molecules There are no single or double bonds in S 3 In fact, all of the bonds in S 3 are equivalent The best Lewis formula of S 3 that can be drawn is: S Lewis Structures: Violation of the ctet Rule There are some molecules that violate the octet rule For these molecules the - 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 3

4 Lewis Structures: Violation of the ctet 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) Lewis Structures: Violation of the ctet Rule Write dot and dash formulas for BBr 3 This is an example of exception #2 Br B Br B Br or Br Br B Br Br Lewis Structures: Violation of the ctet Rule Write dot and dash formulas for AsF 5 As F F F F As or F F As F F F F F 4

5 Polar and on-polar Covalent Bonds Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds onpolar 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 H H or HH or Polar and on-polar 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 Polar and on-polar Covalent Bonds Consider HF H F Electronegativities 21 LM Difference = 19 very polar bond Shown below is an electron density map of HF Blue areas indicate low electron density Red areas indicate high electron density 5

6 Polar and on-polar Covalent Bonds Compare HF with HI Electronegativities Difference = 04 Shown below is an electron density map of HI otice that the charge separation is not as big as for HF H slightly polar bond I Dipole Moments Molecules whose centers of positive and negative charge do not coincide, have an asymmetric charge distribution, and are polar These molecules have a dipole moment The dipole moment has the symbol µ µ is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q Dipole Moments Molecules that have a small separation of charge have a small µ Molecules that have a large separation of charge have a large µ For example, HF and HI: + - δ H - Fδ 191 Debye units + - δ H - I δ 038 Debye units 6

7 Dipole Moments There are some nonpolar molecules that have polar bonds There are two conditions that must be true for a molecule to be polar 1) There must be at least one polar bond present or one lone pair of electrons 2) The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another Continuous Range of Bonding Types Covalent and ionic bonding represent two extremes In pure covalent bonds electrons are equally shared by the atoms In pure ionic bonds electrons are completely lost or gained by one of the atoms Most compounds fall somewhere between these two extremes 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 7

8 ext Class: Exam 2 Finish work on WL homework Chapter 7 Finish Reading Chapter 7 D PRACTICE EXAM! 8