1. Chapter 8
Chemical Bonding

2. Overview Lewis Symbols & Octet Ionic Bonding Covalent Bonding
Electron Configurations and Ions
Ion Sizes
Covalent Bonding
Multiple Bonds
Bond Polarity
Electronegativity

3. Exceptions to Lewis Rules
Lewis Structures
Formal Charge
Resonance Structures
Exceptions to Lewis Rules
Electron Deficient
Expanded Valence
Radicals
Covalent Bond Strengths
Bond Enthalpies
Bond Length
Oxidation Numbers

4. Lewis Symbols and Octet Rule
symbols showing valence electrons for an atom or ion
for active metals & representative elements
number of valence electron = group number
for representative elements
stability requires 8 electrons -- an octet
like the nobel gases

5. Lewis Dot Symbols: C Na N F O 1s22s2p63s1 1s22s22p2 1s22s22p3C Na N
1s22s2p63s1 1s22s22p2 1s22s22p3 F O
1s22s22p s22s22p5

6. Ionic Bonding Ionic Bond
strong attractive, electrostatic force between cations and anions
Ions form to achieve noble gas configuration – an octet
ion symbols shown as Lewis symbols
generally forms between metals and non-metals
overall, the formation of ionic bonds releases energy, exothermic

7. ® Cl Cl Na Na Atoms Ions Na(s) + ½Cl2(g) ® NaCl(s) DHfo = - 411 kJ/molCl + Cl - Na Na +
both have octets
electropositive metal
nonmetal with high electron affinity
Na(s) + ½Cl2(g) ® NaCl(s) DHfo = kJ/mol

8. ® Cl Cl Mg Mg Cl Atoms Ions _ 2+ _ electropositive metalMg Cl 2+ Mg _ Cl
electropositive metal
all have octets
nonmetals with high electron affinity

9. Energetics of Ionic Bond Formation
formation of ionic compounds is exothermic
electrostatic attraction E = k Q1Q2 lattice energy d
Q1 charge of cation, Q2 charge of anion, d distance between cation and anion
steps for formation
change Na from solid to gas -- endothermic
dissociate Cl2(g) -- endothermic
remove electron from Na -- endothermic
add electron to Cl -- exothermic
combine ions, electrostatic attraction -- very exothermic
sum of all above is exothermic

10. Practice Ex. 8.1:
Which would you expect to have the greatest lattice energy:
AgCl CuO CrN
AgCl Q1 = +1 Q2 = -1
CuO Q1 = +2 Q2 = -2
CrN Q1 = +3 Q2 = -3
CrN because E = k Q1Q2 is largest d

11. Problem:
If NaCl has a high lattice E and is very stable, would you expect NaCl2 to be even more stable?
NaCl2 does not form -- even though the lattice energy would be higher, the formation of Na2+ would be so costly, so endothermic, that it would overwhelm the exothermic lattice energy

12. Formation of Ions Representative Elements Transition Metals
last electron entered in an electron configuration is the first electron lost in ion formation
11Na 1s2 2s2 2p6 3s Na+ 1s2 2s2 2p6 3s0
Transition Metals
the s electrons are always lost first, before any d electrons are lost
26Fe [Ar] 4s2 3d Fe2+ [Ar] 4s0 3d Fe3+ [Ar] 4s0 3d5

13. Sizes of Ions Cations Þ smaller than atoms from which they are derived
removal of one or more electrons will
increase Zeff ‘felt’ by remaining electrons
causes contraction of cation
Anions Þ larger than atoms from which they are derived
addition of one or more electrons will
decrease Zeff ‘felt’ by outer electrons
causes some repulsion between electrons creating some expansion of size

14. Covalent Bonding A pair of electrons shared between two atoms
Generally occurs between two non-metals so that atoms can attain an octet
Releases energy upon formation -- exothermic
Two atoms can share:
one pair of electrons , single bond
two pair of electrons, double bond
three pair of electrons, triple bond

15. H H H
One Shared Pair of Electrons = Single Bond

16. O O O
Two Shared Pair of Electrons = Double Bond

17. N N N N Three Shared Pair of Electrons = Triple Bond · · · · · · · · ·

18. The more electrons, the shorter & stronger the bond
Bond strength
single < double < triple
Bond length
single > double > triple
N - N N = N NN Å Å Å 163 kJ kJ kJ

19. H O H Cl Two types of covalent bonds non-polar · polar ·
electrons are shared equally
atoms sharing electrons have equal attraction for them H O
polar
electrons are not shared equally
atoms sharing electrons have different attractions for them H Cl
Cl has a greater attraction for the electrons

20. Bond Polarity & Electronegativity
ability of an atom in a bond to attract the shared electrons
related to electron affinity & ionization energy but not the same as
decreases down a group (except for transition elements)
increases across a row
EN of atom are relative to one another
range from 0.7 to 4.0

21. Bond Polarity
electrons shared between two atoms with different EN are shared unequally
unequal sharing creats a separation of charge -- polar bond
greater the DEN of the atoms, the greater the polarity of the bond
EN EN 3.0
H Cl d+
d -
H Cl
partial (+) charge
partial (-) charge
shared electrons spend more time around Cl

22. Drawing Lewis Structures
Rules:
write Lewis symbols for each atom in formula
count total no. of electrons
count total no. of unpaired electrons divide by 2 = no. of covalent bonds
arrange atoms (more electropositive in the center)
place correct no. of covalent bonds
place remaining electrons around atoms so that all atoms have octets

23. N H NH3 N H 8 total electrons 6 upe- ¸ 2 = 3 cov. bonds
NH3
8 total electrons 6 upe- ¸ = 3 cov. bonds N H
hydrogens are always terminal

24. C O CO2 = C = O 16 total electrons 8 upe- ¸ 2 = 4 cov. bonds
CO2
16 total electrons 8 upe- ¸ = 4 cov. bonds
= C = O
more electropositive atom

25. C O
CO32-
24 total electrons 8 upe- ¸ = 4 cov. bonds C O

26. Formal Charge
Bookkeeping method to keep track of electrons in Lewis structures
Rules
all unshared electrons assigned to the atom on which they reside
half of all shared electrons are assigned to each atom in bond
no. valence e- on free atom
- no. valence e- on bound atom
Formal Charge

27. NH3 N H H H
N 5 e e H N H H
H 1 e e

28. CO2 C O O
= C = O
C 4 e e
O 6 e e

29. CO32- C O O O O C O O O 6 e- 6 e- 0 C 4 e- 4 e- 0 O 6 e- 7 e- -1 · · ·
O 6 e e
C 4 e e -2 O C O O -1
O 6 e e

30. O = C = - C º Formal charge must equal any charge on the structure
prediction of stability
most stable structures have lowest sum of absolute values of formal charges O
= C =
- C º -1 +1
= 0
= 2
preferred structure