Lecture 12 - Lewis Structures II

Tuesday, February 20, 2024

1:30 PM

"The nature of the chemical bond is the problem at the heart of all chemistry." - ﷟HYPERLINK "https://www.nobelprize.org/prizes/chemistry/1954/ceremony-speech/"Linus Pauling

Class notes for Lecture 1-11: https://bricejurban.github.io/CHEM111/
Assignments this week:
﷟HYPERLINK "https://boisestatecanvas.instructure.com/courses/28699/assignments/998317"HW 6 - Periodic Trends and Ionic Bonding (Fri 2/23)
﷟HYPERLINK "https://boisestatecanvas.instructure.com/courses/28699/assignments/998318"HW 7 - Lewis Structures (Mon 2/26)
Reading for today's lecture: Chapter 7.5-7.10 
Reading for next lecture: Chapter 8
Assignments next week:
﷟HYPERLINK "https://boisestatecanvas.instructure.com/courses/28699/assignments/945547"Midterm 2 (Tuesday (2/27) in-class)
﷟HYPERLINK "https://boisestatecanvas.instructure.com/courses/28699/assignments/993257"Emerging Tech in Chemistry (Friday (3/1) w/ replies by Sunday (3/3)
Office Hours: 
Friday 11-1 CIC 
﷟HYPERLINK "https://calendly.com/bricejurban/office-hours"By appointment

Today (2/20)
Lewis Structures
Formal Charges
Not enough electrons
Too many electrons
Resonance Structures
End of content Midterm 2
Thursday (2/22) 
Valence Shell Electron Pair Repulsion Theory of Molecular Geometry (VSEPR)
Drawing Covalent Molecular Structures (Lewis Structures)
Algorithm for Drawing Lewis Structures

Arrange the symbols of the atoms that are bonded together in the molecule
next to one another.
Compute the total number of valence electrons in the molecule by adding
the number of valence electrons for all the atoms in the molecule. If the
species is an ion rather than a molecule, then you must take the charge of
the ion into account by adding electrons if it is a negative ion or subtracting
electrons if it is a positive ion
Represent a two-electron covalent bond by placing a line between the
atoms that are assumed to be bonded to each other
Arrange the remaining valence electrons as lone pairs about each atom so
that the octet rule is satisfied for each one
Check the formal charge on each atom to help with competing structures 
If there is not enough electrons, you may need to use double bonds, triple bonds or rings or the central atom(s) has a deficient octet or you have a radical
If there are too many electrons, you may need to put electron pairs on the central atom and/or form double bonds to reduce formal charges (this will not be assessed in CHEM 101)

 Tips: 
The central atom is usually the least electronegative unless hydrogen which is always terminal.
Some compounds will not have a central atom, but several.
Be, B, and Al have deficient octets
Elements in the 3rd row (S, P . . . )can expand their octet to have 10, 12, or more e-.
If the molecule is an ion, bracket the structure and put a charge in the corner
Molecules with an odd number of electrons will form a free radical, a reactive species.
If there are too many electrons, you may need to put electron pairs on the central atom and/or form double bonds to reduce formal charges (this will not be assessed in CHEM 101)
Draw any valid resonance structures if requested
Assigning Formal Charges help with determining the best structure (Step 5)
 Sometimes more than one possible structure may be possible. In that case it is necessary to assign a formal charge to the atoms in the structure to help aid us in choosing the correct one. We assume each pair of shared electrons are shared equally and assign one of the electrons to each atom. Lone electron pairs are assigned to the atom they are located on. Use this equation:

Untitled picture.png Machine generated alternative text:
formal charge 
on an atom In 
a Lewis formula 
total number of 
— valence electrons — 
in the free atom 
total number 
of lone-pair 
electrons 
total number 
of bonds 


Which structure is preferable for hydroxylamine?

NH3O








No formal charges (best structure)
 NH2OH







Which structure is preferable for hydrogen peroxide?
H2OO



HOOH
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When there are not enough electrons (Step 6)
Some elements are electron deficient are do not require an octet when forming a covalent bond. 
This is restricted to H, Li, Mg, Be, B, and Al.
What is more likely is that you need to use double or triple bonds, especially if there is C or O present.
If the molecule is larger, you can also form ring structures to use up 2 electrons.
Lastly, if you have an odd number of electrons radical (unpaired electron) compounds are also possible.
H2 






Li2 




BeCl2








MgH2





BH3







AlBr3







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CO2











N2
HCN
C2H6 (ethane)






C2H4 (ethene)
C2H2 (ethyne)
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CH3• 
(methyl radical)





•NO 
(nitric oxide)


















When there are too many electrons (Step 7)
Some elements allow for an expanded octet when forming covalent bonds. 
This is restricted to elements in the third shell or higher including the elements: P, S, Cl, As, Se, Te, Br, I, and even Xe.
Never will an element such as C, N, O, or F have an expanded octet


 SF6
XeF2
IF5
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 SF6
 
 
 
 
 
 

 
 
XeF2
IF5



When resonance structures are needed (Step 8)
Resonance refers to areas in the molecule where the electrons in double/triple bonds are delocalized across more than two atoms.
Especially for polyatomic ions and organic molecules, multiple structures can be needed to fully capture the molecule's complexity.
A double headed arrow is used to indicate a resonance structure ↔.
Molecules do not transition between resonance structures but exist as them all simultaneously.
They are the closest approximation we can get to representing the molecule.
HCO3– (hydrogen carbonate)
 
 
 
 
 
 

 

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C2O42– (oxalate) 



 




















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C6H6 (benzene)











 (C6H5)3C• (trityl radical)
 First organic radical discovered. Moses Gomberg in 1900 at University of Michigan (my alma mater)






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6 other resonance structures are possible with the other two rings systems.
This delocalization is what gives rise to the stability of the radical.




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