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Covalent Bonding Ionic compounds form most commonly between a metal and non-metal.

Covalent Bonding Ionic compounds form most commonly between a metal and non-metal. Now we look at compounds formed when 2 non-metals bond. We saw earlier that atoms try to gain the same number of electrons in outer shell as the closest Noble Gas.

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Covalent Bonding Ionic compounds form most commonly between a metal and non-metal.

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  1. Covalent Bonding Ionic compounds form most commonly between a metal and non-metal. Now we look at compounds formed when 2 non-metals bond...

  2. We saw earlier that atoms try to gain the same number of electrons in outer shell as the closest Noble Gas. Metals: lose electrons -> form cations (ionic bonds) Non-metals: gain electrons -> form anions (ionic bonds)

  3. Non-metals have too many electrons to lose to form cations: energy cost is too high. Solution: two non-metals each can obtain a Noble Gas set of electrons by sharing electrons!

  4. Examples Hydrogen is 1s1, so needs 1 more electron to have same electron configuration as He. Chlorine is [Ne] 3s2, 3p5. Cl needs 1 more electron to gain an electron configuration like Ar.

  5. Conclusion: If H and Cl each share a pair of electrons, they both gain what they need to have a “Noble Gas” configuration.

  6. How can we represent covalent bonding? Lewis formulas for elements give the element symbol and show how many valence level electrons each has. H• shows hydrogen has 1 valence e-.

  7. • shows carbon has 4 valence e- • C • • •• nitrogen has 5 valence e- • N • • •• oxygen has 6 valence e- • O • ••

  8. •• fluorine has 7 valence e- • • F • •• Remember: the Group Number of an element in the Periodic Table tells the number of valence electrons!

  9. Another reminder: elements in the same Group have the same electron configuration. This means... Lewis symbols for F, Cl, Br, I are same (7 dots) O, S, Se are same (6 dots) N, P are same (5 dots)

  10. Notice that the non-metals are using s- and p-type orbitals to describe the valence level electrons. Since there is but one s orbital and three p orbitals to a quantum level n, there can be at most 8 valence electrons for these elements. This is called the octet rule.

  11. Every orbital can describe but 2 electrons. So any orbital with a single electron can “contain” one more. Hence, the number of covalent bonds possible = number of single (unpaired) electrons in the element’s Lewis symbol.

  12. Thus, ___ forms __covalent bonds H 1 C 4 N 3 O 2 F 1

  13. And the same number of covalent bonds are formed by other elements of the same Group...because they have the same Lewis symbol...same electron configuration.

  14. Lewis Structures for Compounds •• • H• and • Cl combine as • •• •• • or H-Cl H••Cl • •• A shared pair of e- is shown as a dash. One single bond, three unshared pairs of e-.

  15. Ammonia is NH3. Lewis structure is... •• H••N••H • • H three single bonds, one unshared pair of e-.

  16. Water is H2O. Lewis structure is... •• H••O••H •• two single bonds, two unshared pairs of e-.

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