1. BONDING

2. Bonding Generalities Unlike Charges Attract Unlike Charges Attract Electrons will Be in Pairs Electrons will Be in Pairs Only Valence Electrons are Involved Only Valence Electrons are Involved

3. Two Main Types of Bonds 1. Ionic Bonds- Electrostatic Attraction Between a + and – Ion That Holds Together Ionic Compounds 2. Covalent Bond – Sharing of an Valence Electron Pair Between two Nonmetal Atoms to Form a Molecule

4. In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds. FK

5. FK + _

6. FK + _ The compound potassium fluoride consists of potassium (K + ) ions and fluoride (F - ) ions

7. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule).

8. Octet Rule When Nonmetal Atoms Share Valence Electrons to Form a Covalent Bond- When Nonmetal Atoms Share Valence Electrons to Form a Covalent Bond- –It Will Be to Have the Same Number of Valence Electrons as the Closest Noble Gas  H  2 electrons  Everything Else  8 Electrons

9. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair.

10. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule). But rather than losing or gaining electrons, atoms now share an electron pair. The shared electron pair is called a bonding pair

11. Cl 2 Chlorine forms a covalent bond with itself

12. Cl How will two chlorine atoms react?

13. Cl Each chlorine atom wants to gain one electron to achieve an octet

14. Cl Neither atom will give up an electron – chlorine is highly electronegative. What’s the solution – what can they do to achieve an octet?

15. Lewis Dot Symbols

16. A covalent bond is a chemical bond in which two or more electrons are shared by two nonmetal atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond 9.4

17. 1.Draw skeletal structure of compound showing what atoms are bonded to each other. 2.Count total number of valence e -. Add 1 for each negative charge. Subtract 1 for each positive charge. 3.Connect surrounding atoms to central atom with single bonds. Add remaining electrons (2 at a time) such that surrounding atoms follow octet rule (H -2 electrons). Stop adding electrons once the number exceeds value calculated in step 2. Writing Lewis Structures for Molecular Compounds 9.6

18. Writing Lewis Structures (cont) 4. Make sure that every atom satisfies octet rule and the total number of valence electrons in Lewis Structure is correct. 5. If a surrounding atom does not have an octet; move lone pair to bonding position from an adjacent atom to form a double or triple bond.

19. Double bond – two atoms share two pairs of electrons O C O or O C O 8e - double bonds Triple bond – two atoms share three pairs of electrons N N 8e - N N triple bond or

20. Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 1 – N is less electronegative than F, put N in center FNF F Step 2 – Count valence electrons N - 5 and F - 7 5 + (3 x 7) = 26 valence electrons 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 9.6

21. Exception to Generality of Bonds and Lone Pairs for Atom 1. Resonance Structures 2. Polyatomic Ions 3. Exceptions to the Octet Rule Incomplete Octet Incomplete Octet Expanded Octet Expanded Octet Odd # Electrons Odd # Electrons

22. OCO O -- OCO O - - OCO O - - 9.8 What are the resonance structures of the carbonate (CO 3 2 -) ion?

23. Coordinate Covalent Bonds In a coordinate covalent bonds, both electrons in the bond come from one of the atoms the bond is between. In a coordinate covalent bonds, both electrons in the bond come from one of the atoms the bond is between. Usually, coordinate covalent bond is between Usually, coordinate covalent bond is between –An atom with incomplete octet and –An atom with a lone pair

24. Exceptions to the Octet Rule (on Central Atom) The Expanded Octet (central atom with principal quantum number n≥3 SF 6 S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 9.9

25. Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electronegativity - relative, F is highest 9.5

27. Lewis Dot Structure for NO (cont) O is more electronegative than N O is more electronegative than N The atom with a lower electronegativity has a lower attraction for the electrons in the bond and will be the atom with 7 rather than 8 electrons The atom with a lower electronegativity has a lower attraction for the electrons in the bond and will be the atom with 7 rather than 8 electrons N O Free radical – unpaired electron

28. Bond Type Bond Length (pm) C-CC-CC-CC-C154 CCCCCCCC133 CCCCCCCC120 C-NC-NC-NC-N143 CNCNCNCN138 CNCNCNCN116 Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond 9.4

29. In an IONIC bond, electrons are lost or gained, resulting in the formation of IONS in ionic compounds. FK

30. FK + _

31. FK + _ The compound potassium fluoride consists of potassium (K + ) ions and fluoride (F - ) ions

32. In covalent bonding, atoms still want to achieve a noble gas configuration (the octet rule).

33. Resonance Structures 1. RS are Imaginary; the Real Structure is Intermediate All the RS. 2. Only Electrons Move Between RS; Atoms Positions Never Change 3. At Least One Atom in RS Will Have a Non-zero Formal Charge and Won’t Follow Generality About # Bonds and Lone Pairs.

34. 9.7 An atom’s formal charge is 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. formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons () total number of valence electrons in the free atom - total number of nonbonding electrons - The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

35. Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is less electronegative than O, put C in center OCO O Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e - 4 + (3 x 6) + 2 = 24 valence electrons 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 9.6 Step 5 - Too many electrons, form double bond and re-check # of e - 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24

36. Check Lewis Dot Structures 1. Correct Number of Valence Electrons in Structure 2. Every Atom Follows the Octet Rule 3. Every Atom Follows Generality About # Bonds and Lone Pairs UNLESS it Follows One of the Exceptions

37. Exceptions to the Octet Rule (on Central Atom) The Incomplete Octet HHBe Be – 2e - 2H – 2x1e - 4e - BeH 2 BH 3 B – 3e - 3H – 3x1e - 6e - 9.9 Incomplete Octet ! 4 not 8 e- Incomplete Octet! 6 not 8 e-

38. Exceptions to the Octet Rule Odd-Electron Molecules N – 5e - O – 6e - 11e - NO * When there is an odd # of valence electrons; One atom will have 7 rather than 8 electrons * Use the electronegativity of atoms to determine which atom has 7 rather than 8 electrons