1. Chemical Bonding I: Basic Concepts Chapter 9

2. 9.1 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. 1A 1____ 2A __ns 2 ____ 3ns 2 np 1 4A 4ns 2 np 2 5A 5______ 6A 6ns 2 np 4 ___ 7ns 2 np 5 Group# of valence e - e - configuration

3. 9.1

4. 9.2 Li + F Li + F - The Ionic Bond 1s 2 2s 1 1s 2 2s 2 2p 5 1s 2 1s 2 2s 2 2p 6 [He][Ne] Li Li + + e - e - + FF - F - Li + + Li + F -

5. 9.3 Lattice energy (E) _______ as Q ______ and/or as r ______. cmpd lattice energy MgF 2 MgO LiF LiCl 2957 3938 1036 853 Q= ??? Electrostatic (Lattice) Energy E = k Q+Q-Q+Q- r Q + is the charge on the _______ Q - is the charge on the ________ r is the distance between the ______ Lattice energy (E) is the energy required to completely _______ one mole of a solid ______ compound into ________ ions.

6. 9.3 Born-Haber Cycle for Determining Lattice Energy  H overall =  H 1 +  H 2 +  H 3 +  H 4 +  H 5 oooooo

7. 9.3

8. A ___________is a chemical bond in which two or more electrons are ________ by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 _____ pairs ____ pairs ______ pairs _______covalent bond ______covalent bond 9.4

9. 8e - H H O ++ O HH O HHor 2e - Lewis structure of water _______bond – _________ share two pairs of electrons single covalent bonds O C O or O C O 8e - double bonds _______ bond – two atoms share________ of electrons N N 8e - N N triple bond or 9.4

10. Bond Type Bond Length (pm) C-CC-C 154 CCCC 133 CCCC 120 C-NC-N 143 CNCN 138 CNCN 116 Lengths of Covalent Bonds Bond Lengths _____ bond < ________ Bond < _____ Bond 9.4

12. H F F H ____________________ or _________ is a covalent bond with _____________ electron ______ around one of the two atoms electron _____ region electron _____ region e - ???e - ???? ++ -- 9.5

13. ______________ is the ability of an atom to ______ toward itself the ______ in a chemical bond. Electron Affinity _______, Cl is ______ Electronegativity - _______, F is ________ X (g) + e - X - (g) 9.5

16. Covalent _____ ________ partial transfer of e - _____ transfer e - Increasing difference in electronegativity Classification of bonds by difference in electronegativity DifferenceBond Type 0-.5Covalent  2 Ionic 0 < and <1.7 Polar Covalent 9.5

17. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0.7Cl – 3.03.0 – 0.7 = 2.3______ H – 2.1S – 2.52.5 – 2.1 = 0.4_________ N – 3.0 3.0 – 3.0 = 0________ 9.5

18. 1.Draw __________ of compound showing what atoms are bonded to each other. Put _______electronegative element in the center. 2.Count _____________of valence e -. Add 1 for each ________ charge. Subtract 1 for each positive charge. 3.Complete an _____ for all atoms except hydrogen 4.If structure contains too many electrons, form ____________bonds on ______ atom as needed. Writing Lewis Structures 9.6

19. Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 1 – __ is less electronegative than___, put___ in center FNF F Step 2 – Count _______electrons N - ______ and F ______ ???? Step 3 – Draw _____bonds between N and F atoms and complete _____ on N and F atoms. Step 4 - Check, are # of e - in structure equal to number of valence e - ? ___ single bonds +____ lone pairs _____= ? 9.6

20. Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is ______ electronegative than O, put ____ in center OCO O Step 2 – Count valence electrons -_____________________ ????? = 24 valence electrons Step 3 – Draw ______ bonds between C and O atoms and complete ______ on C and O atoms. Step 4 - Check, are # of e - in structure ____ to number of valence e - ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons 9.6 Step 5 - Too many electrons, form _____bond and re-check # of e - 2 ____ bonds (2x2) = 4 1 _____bond = 4 8 ______ pairs (8x2) = 16 Total = 24

21. 9.7 Two possible skeletal structures of formaldehyde (______) HCOH H CO H An atom’s __________ is the difference between the number of _______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 valence electrons in the free atom - total number of nonbonding electrons - The ____________charges of the atoms in a molecule or ion must _______ the ______ on the molecule or ion.

22. HCOH C – 4 e - O – 6 e - 2H – 2x1 e - 12 e - 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 formal charge on C = 4 -2 -2 -½ x 6 = -1 formal charge on O = 6 -2 -2 -½ x 6 = +1 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 - +1 9.7

23. C – 4 e - O – 6 e - 2H – 2x1 e - 12 e - 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 H CO H ______charge on C = 4 -0 -0 -_______ = 0 _____charge on O = 6 -_______ = 0 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 - 00 9.7

24. Formal Charge and Lewis Structures 9.7 1.For ______ molecules, a Lewis structure in which there are_______________ is _______to one in which formal charges are present. 2.Lewis structures with ______ formal charges are less_______ than those with _____ formal charges. 3.Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which ________ formal charges are placed on the more ____________atoms. Which is the most likely Lewis structure for CH 2 O? HCOH +1 H CO H 00

25. A ___________________is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only _______ Lewis structure (_____ formal charge has been determined!). More possible structures gives the overall structure more validity. 9.8 The true structure is an___________ of all the possible structures.

26. Ozone OOO + - OOO + - 9.8 What are the resonance structures of the carbonate (CO 3 2 -) ion?

27. Violations of the Octet Rule Usually occurs with B and elements of higher periods and most nonmetals. Common exceptions are: Be, B, P, S, Xe, Cl, Br, and As. How do you know if it’s an EXPANDED octet? –More than ______ bonds –????????? doesn’t work out with just 8 BF 3 SF 4 ____: 4 ____: 6 _____: 8 OR 10 ____: 8, 10, OR 12

28. Exceptions to the Octet Rule The Incomplete Octet HHBe Be – 2e - 4e - BeH 2 BF 3 B – 3e - - 24e - FBF F 3 single bonds (3x2) = 6 Total = 24 9.9

29. Exceptions to the Octet Rule _____________Molecules N – 5e - O – 6e - 11e - NO N O The __________ (central atom with principal quantum number n > 2) SF 6 S – 6e - 6F – 42e - S F F F F F F 6 single bonds (6x2) = 12 Total = 48 9.9

30. The _______ change required to break a particular bond in one mole of ________ molecules is the ___________ H 2 (g) H (g) +  H 0 = 436.4 kJ Cl 2 (g) Cl (g) +  H 0 = 242.7 kJ HCl (g) H (g) +Cl (g)  H 0 = 431.9 kJ O 2 (g) O (g) +  H 0 = 498.7 kJ OO N 2 (g) N (g) +  H 0 = 941.4 kJ N N Bond Energy Bond Energies _____ bond <______ bond < ______ bond 9.10

31. Bond Energies (BE) and Enthalpy changes in reactions  H 0 = _______________________________ Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products. 9.10 endothermic exothermic

32. 9.10 H 2 (g) + Cl 2 (g) 2HCl (g)2H 2 (g) + O 2 (g) 2H 2 O (g)

33. Use bond energies to calculate the enthalpy change for: H 2 (g) + F 2 (g) 2HF (g) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) HH1436.4 FF 1156.9 Type of bonds formed Number of bonds formed Bond energy (kJ/mol) Energy change (kJ) HF2568.21136.4 9.10