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Basic Concepts of Chemical Bonding

Basic Concepts of Chemical Bonding. Chapter 8. Three Types of Chemical Bonds. Ionic bond Transfer of electrons Between metal and nonmetal ions Metallic bond Bonding electrons relatively free to move Covalent bond Sharing of electrons Between nonmetal atoms. F. H. F. H.

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Basic Concepts of Chemical Bonding

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  1. Basic Concepts of Chemical Bonding Chapter 8

  2. Three Types of Chemical Bonds • Ionic bond • Transfer of electrons • Between metal and nonmetal ions • Metallic bond • Bonding electrons relatively free to move • Covalent bond • Sharing of electrons • Between nonmetal atoms

  3. F H F H Polar covalent bond ≡ a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e- poor e- rich d+ d- Electronegativity: the ability of an atom in a molecule to attract electrons to itself

  4. Fig 8.6 Electronegativities of the Elements

  5. Figure 8.7

  6. Polar Covalent Bonds Table 8.3 The greater the difference in electronegativity, the more polar is the bond. Fig 8.9 Least polar Most polar

  7. Writing Lewis Structures (p 314) • Sum up all valence electrons. Add 1 for each negative charge. Subtract 1 for each positive charge. • Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. • Complete octets of atoms connected to central atom. • Place remaining electrons on central atom. • If not enough electrons to give central atom an octet, try multiple bonds.

  8. Write the Lewis structure of nitrogen trifluoride (NF3). F N F F Step 1 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 2 – N is less electronegative than F, put N in center Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

  9. Write the Lewis structure of the carbonate ion (CO32-). O C O 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 O Total = 24 • Step 1 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) • -2 charge – 2e- 4 + (3 x 6) + 2 = 24 valence electrons Step 2 – C is less electronegative than O, put C in center Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e-

  10. formal charge on an atom in a Lewis structure total number of valence electrons in the free atom total number of nonbonding electrons ( total number of bonding electrons ) 1 - - = 2 Formal Charges Formal charge - the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

  11. Formal Charges • The best Lewis structure… • …is the one with the fewest charges • …puts a negative charge on the most electronegative atom.

  12. Draw Lewis structure for ozone, O3 or

  13. But this is at odds with the true, observed structure of ozone, in which… • …both O−O bonds are the same length:

  14. Resonance • Just as green is a synthesis of blue and yellow… • …ozone is a synthesis of these two resonance structures.

  15. What are the resonance structures of the carbonate (CO32-) ion? O O O O O O O O O C C C O O O O O O Resonance structure - one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. e.g., ozone 2- 2- 2-

  16. Resonance 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.

  17. Be – 2e- 2H – (2)1e- H Be H 4e- B – 3e- 3 single bonds (3x2) = 6 3F – (3)7e- F B F 24e- Total = 24 9 lone pairs (9x2) = 18 F Exceptions to the Octet Rule Too few electrons BeH2 BF3

  18. N – 5e- S – 6e- N O 6F – 42e- O – 6e- 11e- 48e- F 6 single bonds (6x2) = 12 F F Total = 48 S 18 lone pairs (18x2) = 36 F F F Exceptions to the Octet Rule Odd number of electrons NO Too many electrons(central atom with principal quantum number n > 2) SF6

  19. Average Bond Enthalpies Bond Enthalpies Single bond < Double Bond < Triple Bond

  20. Table 8.4 Average bond Enthalpies (kJ/mol) • These are average bond enthalpies, not absolute bond enthalpies • The C−H bonds in methane, CH4, will be a bit different than the C−H bond in chloroform, CHCl3

  21. Fig 8.14 Estimating Enthalpies of Reaction Hrxn = (bond enthalpies of bonds broken) - (bond enthalpies of bonds formed)

  22. Fig 8.14 Estimating Enthalpies of Reaction CH4 (g) + Cl2 (g)  CH3Cl (g) + HCl (g) In this example: • one C-H bond and one Cl-Cl bond are broken • one C-Cl and one H-Cl bond are formed

  23. Fig 8.14 Estimating Enthalpies of Reaction CH4 (g) + Cl2 (g)  CH3Cl (g) + HCl (g) So, H = [D(C−H) + D(Cl−Cl)] − [D(C−Cl) + D(H−Cl)] = [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)] = (655 kJ) - (759 kJ) = -104 kJ

  24. Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond

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