1. 1 8.7 Covalent Bonding Molecules arise from localized attractive forces between atoms, which we call covalent bondsMolecules arise from localized attractive forces between atoms, which we call covalent bonds Atoms in the molecule are connected strongly, but molecules are not strongly attracted to each other.Atoms in the molecule are connected strongly, but molecules are not strongly attracted to each other. Molecular compounds are usually gases or liquids unless the molecules are very largeMolecular compounds are usually gases or liquids unless the molecules are very large

2. 2 8.7 Covalent Bonds Solids are usually soft low melting points low boiling points low heats of fusion low heats of vaporizationSolids are usually soft low melting points low boiling points low heats of fusion low heats of vaporization Properties arise because molecules are not strongly attracted to other moleculesProperties arise because molecules are not strongly attracted to other molecules Usually found with nonmetalsUsually found with nonmetals

3. 3 Single Covalent Bonds Sharing of 1 pair of electronsSharing of 1 pair of electrons Each atom has one half-filled valence orbital that overlap one anotherEach atom has one half-filled valence orbital that overlap one another H. +. H  H:HH. +. H  H:H Single bond represented as H:H or H–H Called a Lewis formula or electron-dot formulaSingle bond represented as H:H or H–H Called a Lewis formula or electron-dot formula

4. 4 8.10 Lewis Symbols Lewis Symbols: The number of valence electrons available for bonding are indicated by unpaired dots.Lewis Symbols: The number of valence electrons available for bonding are indicated by unpaired dots.

5. 5 Lewis Symbols We generally place the electrons on four sides of a square around the element symbol.We generally place the electrons on four sides of a square around the element symbol.

6. 6 Single bonds between like atoms HalogensHalogens

7. 7 Single bonds between unlike atoms HF Figure 7.10HF Figure 7.10

8. 8 Some atoms can form bonds with more than one atom CCl 4 Figure 7.10 How many valence electrons are supplied by each atom?CCl 4 Figure 7.10 How many valence electrons are supplied by each atom?

9. 9 Lewis Model of Covalent Bonding (Localized electron model) 1. 1.A covalent bond is formed by the sharing of at least two valence electrons between two atoms. 2. 2.At normal temperatures and pressures, for non- metals, the maximum number of valence electrons per atom is eight (with the exception of two for H and He). Each shared electron is counted twice (one per atom). 3.A stable molecule is formed when sharing of valence electrons lowers the chemical energy.

10. 10 Procedure for Drawing Lewis Structures 1.Arrange the atoms (symbols) around the central atom. 2.Count up total valence electrons 3.Draw a line to represent a single bond (2 shared electrons) between the central atom and each surrounding atom. This is a 2-electron bond between the cores. 4.Place remaining electrons on the outside atoms, then the central atom 5.Shift electrons, as necessary, to make multiple bonds and satisfy the octet rule

11. 11 Whiteboard Work #1: Draw the Lewis Structures for the following molecules:Draw the Lewis Structures for the following molecules: BF 3BF 3 NH 3NH 3 NO 2NO 2 CO 2CO 2

12. 12 8.11 Exceptions to the Octet “Rule” Incomplete OctetsIncomplete Octets Usually only seen for Be, B, Al Leads to stable, yet reactive molecules! Coordinate Covalent BondingCoordinate Covalent Bonding Molecules with too few electron pairs can bond with molecules with unshared electron pairs to form a new shared-electron-pair bondMolecules with too few electron pairs can bond with molecules with unshared electron pairs to form a new shared-electron-pair bond BF 3 + NH 3  F 3 BNH 3BF 3 + NH 3  F 3 BNH 3

13. 13 Exceptions to the Octet Rule Why is BH 4 - more stable than BH 3 ?Why is BH 4 - more stable than BH 3 ? Why is BF 4 - more stable than BF 3 ?Why is BF 4 - more stable than BF 3 ? Why does aluminum chloride exist in the gaseous state as Cl 2 AlCl 2 AlCl 2 (that is, Al 2 Cl 6 ) instead of AlCl 3 ?Why does aluminum chloride exist in the gaseous state as Cl 2 AlCl 2 AlCl 2 (that is, Al 2 Cl 6 ) instead of AlCl 3 ?

14. 14 8.11 Exceptions to the Octet Rule Odd-Electron MoleculesOdd-Electron Molecules Also stable, yet reactive!Also stable, yet reactive! Examine the Lewis formula for NO 2Examine the Lewis formula for NO 2 Why does NO 2 combine with itself to form N 2 O 4 ?Why does NO 2 combine with itself to form N 2 O 4 ?

15. 15 CO 2

16. 16 Multiple Bonds Can share more than one pair of electrons to form double or triple bondsCan share more than one pair of electrons to form double or triple bonds 2 electron pairs 3 electron pairs

17. 17 Whiteboard Work #2: SO 2Draw the Lewis structure for SO 2

18. 18 8.12 Resonance Structures Lewis formulas don’t always accurately represent bonds. Sometimes it takes two formulas to adequately represent the bonds.Lewis formulas don’t always accurately represent bonds. Sometimes it takes two formulas to adequately represent the bonds.

19. 19 Resonance Structures In the sulfur dioxide example, we could move the electron pair from either oxygen to form the double bond. The two structures we drew were equivalent. These are examples of equivalent resonance structures. The Lewis-dot representation suggests that the molecule exists as either one structure or the other, and that it sort of “flips” back and forth. This is NOT the case.

20. 20 Rather the electron pair of the double bond is sort of “smeared out” over both bond positions. This “smearing out” is referred to as delocalization, and can be represented using a dashed line. Each S–O bond is neither a single or a double bond. The three electron pairs distribute themselves among the 2 bonding positions and is basically a 3/2 = 1.5 bond.

21. 21 Resonance How many different valid Lewis formulas can you write for the following molecules or ions? How do they differ? SO 2 SO 3 H 2 SO 4 CO 3 2 - NO 3 - HNO 3 NCS -How many different valid Lewis formulas can you write for the following molecules or ions? How do they differ? SO 2 SO 3 H 2 SO 4 CO 3 2 - NO 3 - HNO 3 NCS - The different resonance forms represent delocalized bonding.The different resonance forms represent delocalized bonding.

22. 22 Write Lewis formulas for the following molecules or ions: SO 2SO 2 SO 3SO 3 H 2 SO 4H 2 SO 4 CO 3 2 -CO 3 2 - NO 3 - HNO 3 NCS -

23. 23 Formal Charge Can be used to decide between alternate Lewis structuresCan be used to decide between alternate Lewis structures There is considerable controversy as to whether the concept of formal charge dictating electron distribution is in fact correct.There is considerable controversy as to whether the concept of formal charge dictating electron distribution is in fact correct. You can read more about it in your text.You can read more about it in your text. Is helpful for evaluating Lewis structuresIs helpful for evaluating Lewis structures

24. 24 Whiteboard Work #3 Draw the Lewis structure for I 3 1– Draw the Lewis structure for SbF 5

25. 25 Exception to the Octet Rule Some bonding situations result in the central atom having > 4 electron pairs ( > 8 electrons) around it:Some bonding situations result in the central atom having > 4 electron pairs ( > 8 electrons) around it: S can form SH 2 with S having 8 valence electronsS can form SH 2 with S having 8 valence electrons S can form SCl 4 with S having 10 valence electronsS can form SCl 4 with S having 10 valence electrons S can form SCl 6 with S having 12 valence electronsS can form SCl 6 with S having 12 valence electrons How is this possible??? (Think QM…)How is this possible??? (Think QM…)

26. 26 Expanded Octets Non-Metal atoms in Period 3 or higher can have more than 8 electrons around them (i.e. in their valence shell or top shelf) Why is this? Recall period 3 atoms (and period 4, 5, 6… atoms) have d-orbitals that can hold those extra electrons!

27. 27 lWhen do we find expanded octets? What is the origin of the octet rule? Exceptions to the Octet Rule Expanded Valence ShellsExpanded Valence Shells What do you do if there are too many electrons to be accommodated by octets?What do you do if there are too many electrons to be accommodated by octets? Write Lewis formulas for the following: SF 4 SF 6 IF 4 + XeF 4 XeF 2 PF 5 BrF 3 BrF 5Write Lewis formulas for the following: SF 4 SF 6 IF 4 + XeF 4 XeF 2 PF 5 BrF 3 BrF 5

28. 28 Modern View of Lewis Bonding Model 1.A stable molecule is formed when sharing of electrons lowers the chemical energy. 2.For non-metals, the number of valence electrons equals the column number. The maximum number of valence electrons per core is 8 (2 for H and He). 3.Usually a covalent bond is formed by sharing 2 valence electrons between 2 cores. 4.At low temperatures, a non-reactive substance usually has the maximum number of electrons per core (octet rule). There are some exceptions.

29. 29 Website:Website: http://www.kentchemistry.com/links/bondin g/lewisdotstruct.htm

30. 30 Write Lewis formulas for the following molecules or ions: NH 3NH 3 NH 4 +NH 4 + CCl 2 F 2CCl 2 F 2 SOCl 2SOCl 2 SO 2SO 2 CO 2CO 2 CO 3 2 - SO 3 2 - H 2 SO 4 HCN CN - NCS -

31. 31 Lewis Structures

32. 32 Ch. 8 Group Quiz Write Lewis formulas for the following molecules or ions:Write Lewis formulas for the following molecules or ions: 1.PCl 3 2.CO 3.H 2 O 2 (HOOH) 4.SO 4 2 - Put your names and class date in the heading!!!

33. 33 Origins of the “Octet Rule” Octet rule: we know that s 2 p 6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs). (there are many exceptions)Octet rule: we know that s 2 p 6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs). (there are many exceptions)

34. 34 Octet Rule What ion or compound is formed from the following to approximate a noble gas electronic configuration? What is the configuration?What ion or compound is formed from the following to approximate a noble gas electronic configuration? What is the configuration? Na1s 2 2s 2 2p 6 3s 1Na1s 2 2s 2 2p 6 3s 1 Na + 1s 2 2s 2 2p 6Na + 1s 2 2s 2 2p 6 H 1s 1H 1s 1 H + or H - 1s 0 or 1s 2H + or H - 1s 0 or 1s 2

35. 35 Cl1s 2 2s 2 2p 6 3s 2 3p 5Cl1s 2 2s 2 2p 6 3s 2 3p 5 Cl - 1s 2 2s 2 2p 6 3s 2 3p 6Cl - 1s 2 2s 2 2p 6 3s 2 3p 6 O1s 2 2s 2 2p 4O1s 2 2s 2 2p 4 O 2 - 1s 2 2s 2 2p 6O 2 - 1s 2 2s 2 2p 6 H + O1s 1 + 1s 2 2s 2 2p 4H + O1s 1 + 1s 2 2s 2 2p 4 H 2 O1s 2 2s 2 2p 6 for O, 1s 2 for HH 2 O1s 2 2s 2 2p 6 for O, 1s 2 for H Na + O1s 2 2s 2 2p 6 3s 1 + 1s 2 2s 2 2p 4Na + O1s 2 2s 2 2p 6 3s 1 + 1s 2 2s 2 2p 4 Na 2 O1s 2 2s 2 2p 6 for Na and ONa 2 O1s 2 2s 2 2p 6 for Na and O Octet Rule

36. 36 C + H 1s 2 2s 2 2p 2 + 1s 1C + H 1s 2 2s 2 2p 2 + 1s 1 CH 4 1s 2 2s 2 2p 6 for C, 1s 2 for HCH 4 1s 2 2s 2 2p 6 for C, 1s 2 for H C + Cl1s 2 2s 2 2p 2 + 1s 2 2s 2 2p 6 3s 2 3p 5C + Cl1s 2 2s 2 2p 2 + 1s 2 2s 2 2p 6 3s 2 3p 5 CCl 4 1s 2 2s 2 2p 6 for C, 1s 2 2s 2 2p 6 3s 2 3p 6 for ClCCl 4 1s 2 2s 2 2p 6 for C, 1s 2 2s 2 2p 6 3s 2 3p 6 for Cl C + O1s 2 2s 2 2p 2 + 1s 2 2s 2 2p 4C + O1s 2 2s 2 2p 2 + 1s 2 2s 2 2p 4 CO 2 1s 2 2s 2 2p 6 for C and OCO 2 1s 2 2s 2 2p 6 for C and O Octet Rule

37. 37 Write Lewis formulas for the following molecules or ions: 1.NO 1.NO 2 2.N 2 O 4 2.BH 3 3.BH 4 - 3.AlCl 3 4.SF 4 4.SF 6 5.IF 4 + 5.XeF 4 6.XeF 2 6.PF 5 BrF 5