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Chapter 8: Ionic and Covalent Bonding

RVCC Fall 2009 CHEM 103 – General Chemistry I. Chapter 8: Ionic and Covalent Bonding. Chemistry: The Molecular Science, 3 rd Ed. by Moore, Stanitski, and Jurs. Bonding – What holds atoms together?.

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Chapter 8: Ionic and Covalent Bonding

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  1. RVCC Fall 2009 CHEM 103 – General Chemistry I Chapter 8:Ionic and Covalent Bonding Chemistry: The Molecular Science, 3rd Ed. by Moore, Stanitski, and Jurs

  2. Bonding – What holds atoms together? Octet rule: To form bonds, atoms gain, lose, or share e- to achieve a valence shell of 8 (or isoelectronic with a noble gas). • Ionic bond – an electrostatic attraction between a cation and an anion that forms when electrons transfer from one atom to another. • Covalent (Molecular) bond – the net attractive force that results from the sharing of electrons between atoms.

  3. : . : Cl : Ionic Bonds An ionic bond is formed by the transfer of electrons from one atom (metal with low EA) to another (nonmetal with high EA). The resultant ions are held together by electrostatic attraction. Na . [Ne]3s1 [Ne]3s23p5 Na+ Cl- [Ne] [Ne]3s23p6 = [Ar] Each atom has satisfied the octet rule.

  4. : F . : : : F Mg . . . : : : : : : : : : : Ionic Bonds MgF2 - - F Mg 2+ F

  5. Ionic Compounds - Properties • crystalline • high melting point • high boiling point • soluble in water • electrolytes Crystal Lattice • hard • brittle

  6. Attraction Stable bond Repulsion Covalent Bonding - G.N. Lewis (1916) Some atoms share e- to form bonds. When two nonmetals bond, they often share electrons since they have similar attractions (EA) for them. This sharing of valence electrons is called the covalent bond. • Number of bonds = Number shared e- pairs.

  7. Covalent Bonds H2 . . : + H H H H

  8. Covalent Bonds

  9. Covalent Bonds - Epotential “well”

  10. Single Covalent Bonds Lewis structures: show ALL valence electrons dot = 1 e- line = 1 pair of e- single bond- one shared pair of e- H − H H H

  11. H H – C – H H .. .. .. .. .. .. .. F – N – F F .. .. .. .. .. H – O – H .. .. .. .. H – F H – F Single Covalent Bonds # of e- shared to form an octet (8-A group#) # of e- shared Group # of to form an octet Example valence e- (8 - A group#) 4A 4 4 C in CH4 5A 5 3 N in NF3 6A 6 2 O in H2O 7A 7 1 F in HF ..

  12. : : : : H Cl Lewis Structures Lewis electron-dot formulas or Lewis structures. bonding pair lone pairs : : : bonding pair H Cl : lone pair • An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared).

  13. H H : : : : C C : : H H Multiple Bonds In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared. • It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms. C has octet. H OK with 2. or

  14. ::: H H C C or : : Multiple Bonds Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms. Elements that form multiple bonds: C, O , N, S

  15. The Procedure • Using the molecular formula, count the total number of valence electrons available (bonding + lone pairs). • Valence electrons for each atom corresponds to group # • Adjust for charge (add electron for each minus, delete electron for each plus) 2(1) + 6 3(1) + 6 - 1

  16. The Procedure • Make a skeleton by connecting the atoms with single bonds only. When connecting atoms, remember… • Put the least electronegative atom in the center. (Usually the first listed in the chemical formula.) • Hydrogen is ALWAYS a terminal atom. More electronegative atoms are terminal (F, O…) • Make the structure symmetric. 4 pairs 4 pairs 2 pairs left 1 pair left

  17. The Procedure • Put the left over electrons as lone pairs, preferably on the more electronegative atoms Is the octet rule satisfied? • If YES, then you’re done…

  18. The Procedure • If you are electron-deficient(not enough electrons to complete an octet), then some atoms must share more than two electrons. “If you have a lone pair, make those two atoms share.” • Ex. C2H4

  19. The Procedure • If you have excess electrons, at least one atom must have an expanded valence • Must be element from third period or lower • Usually the central atom • e.g. SF4

  20. General Rule • The total number of valence electrons on an atom (from bonds & lone pairs) cannot exceed that atom’s maximum valence. • First period: 2 electrons (s) • Second period: 8 electrons (s,p) • Third period & below: prefer to have 8, but can expand when necessary (s,p,d)

  21. Writing Lewis Dot Formulas 20 e- total or 10 pairs SCl2 8 left 0 left : : : : : Cl S Cl : : :

  22. O : : : : : : C : : : Cl Cl Writing Lewis Dot Formulas 24 e- total 12 pairs COCl2 9 left 0 left Note that the carbon has only 6 electrons.

  23. 9 e- left O : : : : : : C : : : Cl Cl Writing Lewis Dot Formulas COCl2 12 pairs 0 e- left To fulfill the octet rule… “If you have a lone pair, make those two atoms share!”

  24. : : : : : : Writing Lewis Dot Formulas COCl2 24 e- total 18 e- left : : O 0 e- left C Cl Cl Note that the octet rule is now obeyed.

  25. Writing Lewis Dot Formulas Practice N2 SF4 O2 ClO3- HCN ClO2-1 PO4-3 NO3-1

  26. Question no no no no no • First evaluate the total valence electrons: • 24 e- • 26 e-, wrong • 24 e-, looks OK • 24 e-, one F has too many • 24 e-, N not enough • 24 e-, but least electronegative has to be in the center • 24 e-, no bond between two N.

  27. Exceptions to the Octet Rule Although many molecules obey the octet rule, there are exceptions where the central atom has less or more than eight electrons. Incomplete octet – B, H Boron has 3 valence electrons BF3 .. .. :F – B – F: .. .. :F: ..

  28. : : : F : F : : F : : : : : F : : F : : Exceptions to the Octet Rule If a nonmetal is in the third period or greater it can accommodate as many as twelve electrons as the central atom. PF5 P

  29. : : : : F : F : : F : F : : : : Exceptions to the Octet Rule In sulfur tetrafluoride, SF4, the sulfur atom must accommodate two extra lone pairs for a total of 5 electron pair (10 electrons) : S

  30. Formal Charge and Lewis Structures In certain instances, more than one feasible Lewis structure can be illustrated for a molecule. For example, : : H C N H N C or The concept of “formal charge” can help us decide which structure is correct.

  31. Formal Charge and Lewis Structures formal charge = valence e- before bonding– valence e- after bonding = valence e- - [1/2 bonding e- + lone pair e-] : : H C N H N C or H: 1-½(2) = 0 H: 1-½(2) = 0 C: 4 - ½(8) = 0 C: 4 – (½(6)+2) = -1 N: 5 – (½(8) + 2) = 0 N: 5 – (½(8)) = +1

  32. formal charges : : H C N H N C or 0 0 0 0 +1 -1 Formal Charge and Lewis Structures • Smaller formal charges are more favorable • More electronegative (or higher EA) atom should have negative formal charges • Like charges should not be on adjacent atoms • Net formal charge should be the overall charge on the molecule/ion.

  33. Practice Determine the most stable structure for dinitrogen oxide. (All structures have 16 valence electrons.) N=N=O N-N≡O N ≡ N - O -1 +1 0 -2 +1 +1 0 +1 -1 formal charge= valence e- - [1/2 bonding e- + lone pair e-]

  34. Practice - Formal Charge • Which structure is correct? : O N Cl O N Cl or 0 0 0 -1 0 +1

  35. : : : : : : O O : : : : O O O O : : Delocalized Bonding: Resonance The structure of ozone, O3, can be represented by two different Lewis electron-dot formulas. The bond lengths for the above structures are: O – O 132 pm O = O 112 pm However, experiments show that both bonds are identical.

  36. O O O Delocalized Bonding: Resonance According to theory, one pair of bonding electrons is spread (delocalized) over the region of all three atoms. • In fact, the actual bond length is 127.8 pm (in between 132 and 112pm). • The actual molecule is a hybrid or composite structure and not different structures that change back and forth… although, we often represent it that way.

  37. O H O N O O H O N O O H O N O Delocalized Bonding: Resonance Lewis resonance structures, have the same atoms in the same positions. Only an electron pair position is different.

  38. Resonance Structures Which pair does NOT represent resonance structures? O=S O O S=O S=S O S O=S

  39. -2 O O C O -2 O O C O Resonance Structures Draw resonance structure(s) for the following: -2 O O C O

  40. “All covalent bonds are created equal but some are more equal than others.” (We assumed equal sharing when we calculated formal charge.)

  41. Electronegativity… …is a measure of the ability of an atom in a molecule to draw bonding electrons to itself when bonded. Periodic Trend - Electronegativity increases across a period decreases down a group

  42. Electronegativity Notice, there are NO values for EN for the noble gases.

  43. Types of Bonds 0.9 3.0 Na+ Cl- 2.5 2.1 C - H 1.6 1.6 Zn Zn • Ionic: • ΔEN >1.8 • electron transfer • Covalent: • ΔEN <1.8 • electron sharing • Metallic: • electron-sea model or band theory

  44. Covalent Bonds (EN<1.8) • Non-polarcovalent- ΔEN = 0 – 0.5 • Examples: H-H, Cl-Cl, C-H bonds • Polarcovalent - ΔEN = 0.5 – 1.8 • Examples: H-O, C-Cl, C-O bonds

  45. Bond Polarity In HCl we have a partial negative charge on the chlorine (denoted d-)and a partial positive charge on the hydrogen (denoted d+) The bond is polar covalent. d+ d- : H :Cl: :

  46. Bond Polarity Arrange the following bonds from the most to the least polar: HH, HCl, HF, HI, HBr Compare the electronegativity of Cl, F, I and Br: Least Most Electronegative Electronegative I Br Cl F Determine polarity: HH HI HBr HCl HF non polar most polar

  47. Practice Which of the following bonds in each pair are more polar? C-S or C-O Cl-Cl or O=O N-H or C-H

  48. Bond Length Bond length (or bond distance) is the distance between the nuclei of two bonded atoms (the sum of atomic radii).

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