1. Chapter 1
Electronic Structure
and
Bonding
Acids and Bases

2. 1 Å = 10 picometers = 0.1 nanometers =
~ 0.1 nm
Anders Jöns Ångström
( )
1 Å = 10 picometers = 0.1 nanometers =
10-4 microns = 10-8 centimeters
Nucleus =
1/10,000
of the atom
1 nm = 10 Å
An atom vs. a nucleus
~10,000 x larger

3.    

4. Question 1 What is the electronic configuration of carbon?
A) 1s2 2s2 2px2
B) 1s2 2s2 2px1 2py12pz0
C) 1s2 2s2 2px12py12pz1
D) 1s2 1px1 1py12s2

5. Electron Configurations Noble Gases and The Rule of Eight
When two nonmetals react to form a covalent bond: They share electrons to achieve a Noble gas electron configuration.
When a nonmetal and a metal react to form an ionic compound: Valence electrons of the metal are lost and the nonmetal gains these electrons.

6. Photo Bancroft Library, University of California/LBNL Image Library
Notes from Lewis’s notebook and his “Lewis” structure.
G.N. Lewis
Photo Bancroft Library, University of California/LBNL Image Library
Footnote:
G.N. Lewis, despite his insight and contributions to chemistry, was never awarded the Nobel prize.

7. Ionic Compounds
Ionic compounds are formed when electron(s) are transferred.
Electrons go from less electronegative element to the more electronegative forming ionic bonds.

8. Covalent Compounds Share electrons. 1 pair = 1 bond.
Octet rule (“duet” for hydrogen)
Lewis structures:
Notice the charges:
In one case they balance, can you name the compound?
In the other they do not, can you name the polyatomic ion?
More about “formal” charge to come.

9. Question 2
Select the correct Lewis structure for methyl fluoride (CH3F).
A) B)
C) D)

10. Important Bond Numbers
(Neutral Atoms / Normal electron distribution)
one bond H F Cl Br I O
two bonds N
three bonds C
four bonds

11. Question 3 What is the correct Lewis structure of formaldehyde (H2CO)?
A) B)
C) D)

12. Question 4 Which of the following contains a triple bond? A) SO2
B) HCN
C) C2H4
D) NH3

13. Formal Charge
Formal charge is the charge of an atom in a Lewis structure which has a different than normal distribution of electrons. 5

14. Important Bond Numbers
(Neutral Atoms / Normal electron distribution)
one bond H F Cl Br I O
two bonds N
three bonds C
four bonds 14

15. Important Bond Numbers
(Neutral Atoms / Normal electron distribution)

16. Formal Charge
Formal charge = number of valence electrons –
(number of lone pair electrons +1/2 number of bonding electrons)
Equals the number of valence electrons (Group Number of the free atom) minus [the number of unshared valence electrons in the molecule + 1/2 the number of shared valence electrons in the molecule].
Moving/Adding/Subtracting atoms and electrons.

17. HNO3 Nitric Acid

18. Complete the following table
Complete the following table. It summarizes the formal charge on a (“central”) atom for the most important species in organic chemistry.  

19. Question 5
What is the formal charge of the carbon atom in the Lewis structure?
A) -1
B) 0
C) +1
D) +2 C

20. Question 6
What is the formal charge of the oxygen atom in the Lewis structure?
A) -1
B) 0
C) +1
D) +2 20

21. Resonance 5

22. Resonance
Resonance is a very important intellectual concept that was introduced by Linus Pauling in 1928 to explain experimental observations.
Eg. SO2
Bond order  1.5
Bond length > double bond; < single bond
TUTORIAL

23. Resonance
Two or more Lewis structures may be legitimately written for certain compounds (or ions) that have double bonds and/or free pairs of non-bonded electrons
It is a mental exercise in “pushing” or moving electrons.
Refer to Table 1.6 22

24. Rules of Resonance
Step 1: The atoms must stay in the same position. Atom connectivity is the same in all resonance structures. Only electrons move.
NON-Example: The Lewis formulas below are not resonance forms. A hydrogen atom has changed position. 2

25. Example: All structures have 18 electrons and a net charge of 0.
Rules of Resonance
Step 2: Each contributing structure must have the same total number of electrons and the same net charge.
Example: All structures have 18 electrons and a net charge of 0. 2

26. Step 3: Calculate formal charges for each atom in each structure.
Rules of Resonance
Step 3: Calculate formal charges for each atom in each structure.
Example: None of the atoms possess a formal charge in this Lewis structure. 2

27. Step 4: Calculate formal charges for the second and third structures.
Rules of Resonance
Step 4: Calculate formal charges for the second and third structures.
Example: These structures have formal charges.
NOTE: They are less favorable Lewis structures. 2

28. “Pushing” Electrons same atomic positions differ in electron positions.. C O N H : .. C O N H : + –
more stable Lewis structure
less stable Lewis structure 24

29. “Pushing” Electrons same atomic positions
differ in electron positions only .. C O N H : .. C O N H : + –
more stable Lewis structure
less stable Lewis structure 24

30. Why use Resonance Structures?
Delocalization of electrons and charges between two or more atoms helps explain energetic stability and chemical reactivity.
Electrons in a single Lewis structure are insufficient to show electron delocalization.
A composite of all resonance forms more accurately depicts electron distribution. (HYBRID) NOTE: Resonance forms are not always evenly weighted. Some forms are better than others. 25

31. Resonance Example Ozone (O3)
Lewis structure of ozone shows one double bond and one single bond O ••
• • – +
Expect: one short bond and one long bond
Reality: bonds are of equal length (128 pm) 26

32. Resonance Example Ozone (O3)
Lewis structure of ozone shows one double bond and one single bond O ••
• • – +
Resonance: O ••
• • – + O ••
• • – + 26

33. Resonance Example Ozone (O3) O + O +
Electrostatic potential map shows both end carbons are equivalent with respect to negative charge. Middle carbon is positive. O ••
• • – + O ••
• • – + 26

34. Detailed Resonance Examples26

35. Question 7 Which resonance structure contributes more to the hybrid?
A) B)

36. VSEPR Model Valence Shell Electron Pair Repulsion

37. VSEPR Model Some examples of minimizing interactions.
The molecular structure of a given atom is determined principally by minimizing electron pair (bonded &free) repulsions through maximizing separations.
Some examples of minimizing interactions.

38. Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put pairs as far apart as possible.
3. Determine positions of atoms from the way electron pairs are shared.
4. Determine the name of molecular structure from positions of the atoms.

39. Orbital
Geometry
Molecular
Geometry
Bond Angle
# of lone pairs
Chem 226
Linear Linear
Trigonal Planar Trigonal Planar
Trigonal Planar Bent
Tetrahedral Tetrahedral
Tetrahedral Trigonal Pyramidal
Tetrahedral Bent
Trigonal Bipyramidal Trigonal Bipyramidal
Trigonal Bipyramidal Seesaw
Trigonal Bipyramidal T-shape
Trigonal Bipyramidal Linear
Octahedral Octahedral
Octahedral Square Pyramidal
Octahedral Square Planar 1 2 3

40. Lewis Structures / VSEPR / Molecular Models
Computer Generated Models
Ball and stick models of ammonia, water and methane.
Worksheet 1: Bonds, Formulas, Structures & Shapes

41. Covalent Compounds
Equal sharing of electrons: nonpolar covalent bond, same electronegativity (e.g., H2)
Unequal sharing of electrons between atoms of different electronegativities: polar covalent bond (e.g., HF)

44. Question 8
Which of the following bonds is the most polar?
A) B)
C) D)

45. Bond Dipole & Dipole Moment
Dipole moments are experimentally measured.
Polar bonds have dipole moments.
dipole moment (D) = m = e x d
(e) : magnitude of the charge on the atom
(d) : distance between the two charges

46. Question 9
Which of the following bonds have the greatest dipole moment ()?
A) B)
C) D)

47. Bond Polarity
A molecule, such as HF, that has a center of positive charge and a center of negative charge is polar, and has a dipole moment. The partial charge is represented by  and the polarity with a vector arrow.

48. Question 10
In which of the compounds below is the d+ for H the greatest?
A) CH4
B) NH3
C) SiH4
D) H2O

49. Question 11
In which of the following is oxygen the positive end of the bond dipole?
A) O-F
B) O-N
C) O-S
D) O-H