Lewis Structures and Molecular Shapes

Announcements

No lab this week due to the end-of-block quiz.
All demonstrators will be available in specific rooms for the first hour of the lab session to answer questions and help with the quiz.
Room details are posted on Canvas under announcements.
Attendance is not mandatory; students can visit any demonstrator.
These sessions are a trial, and continuation depends on student interest.
The demonstrator is essentially a private tutor available during labs.

The planned demonstration for today's lecture failed.
An ammonia fountain experiment will be demonstrated on Thursday to illustrate amines.
Feedback on the course will be collected at the start of each lecture.

The main concept of this lecture: Lewis structures are useful in organic chemistry because they indicate the shape of molecules.
Tetrahedral carbons ($\text{sp}^3$ carbons)
Trigonal planar carbons ($\text{sp}^2$ carbons)
Linear carbons ($\text{sp}$ carbons)

Rob Chapman is a senior lecturer in chemistry.
Background: Engineering degree followed by a Ph.D. in chemistry.
Research interests: Designing polymers to mimic the therapeutic action of proteins, particularly anticancer drugs like trail.

Encouragement to think deeply about science and its implications.
Historical context: Scientists were once considered natural philosophers.
The Ph.D. degree is a vestige of this tradition, representing a broad understanding of the world.
Importance of asking deeper question like Why is the world made of carbon?
- Carbon's electronic structure (four valence electrons) allows it to form many different shapes and bond with multiple elements.
- Carbon sits well within the covalent part of the periodic table.

Electronic structure for hydrogen: 1s1
Electronic structure for carbon: 1s2 2s2 2p2
The periodic table mirrors the electronic structure of elements.
Number of electrons in the p-block: 6
Number of electrons in the s-block: 2
Number of electrons in the d-block: 10
Lewis diagrams show a dot for each electron in the outer shell.
- Carbon: 4 electrons
- Nitrogen: 5 electrons
- Oxygen: 6 electrons
- Fluoride: 7 electrons
Atoms form covalent molecules by sharing electrons to achieve a complete octet.
- Hydrogen shares 1 electron.
- Carbon shares 4 electrons.
- Nitrogen shares 3 electrons.
- Oxygen shares 2 electrons.
Each line in a covalent molecule represents two shared electrons.
- Every hydrogen forms one bond.
- Every carbon forms four bonds.
- Every nitrogen forms three bonds.
- Every oxygen forms two bonds.

Example: Ethanol ($\text{CH}3\text{CH}2\text{OH}$)
- Stick diagram for ethanol.
- Lewis structure includes lone pairs on the oxygen atom.
  - Oxygen has six valence electrons initially.
Example: Ethanal ($\text{CH}_3\text{CHO}$)
- Lewis structure includes a double bond to the oxygen and lone pairs on the oxygen atom.
Example: Sulfate ($\text{SO}_4^{2-}$)
- Total number of electrons: (6 \times 4) + (6 \times 1) + 2 = 32
Formal Charge Calculation:
- Formula: Formal Charge = Valence Electrons - (Lone Pair Electrons + 1/2 Bonded Electrons)
- For oxygen: 6 - 7 = -1
- For sulfur: 6 - 4 = +2
Adjusting the Lewis Structure:
- Move electrons to minimize formal charges.
- Correct Lewis structure for sulfate includes double bonds to two of the oxygen atoms.
Formal Charges in Corrected Structure:
- Formal charge on oxygen with single bond: -1
- Formal charge on oxygen with double bond: 0
- Formal charge on sulfur: 0
- Overall charge: -2

Ammonia ($\text{NH}_3$):
- Lewis structure: Nitrogen bonded to three hydrogens with one lone pair, totaling eight electrons.
- Shape: Trigonal pyramid because there are four groups of electrons, including the lone pair.
Ammonium ($\text{NH}_4^+$):
- Lewis structure: Nitrogen bonded to four hydrogens, totaling eight electrons.
- Formal charge on nitrogen: +1
- Shape: Tetrahedral, as all four hydrogens get as far away from each other as possible.
Amide ($\text{NH}_2^-$):
- Lewis structure: Nitrogen bonded to two hydrogens with two lone pairs, totaling eight electrons.
- Formal charge on nitrogen: -1
- Shape: Bent, similar to water.
$\text{CH}3\text{CH}2\text{O}^-$:
- Total valence electrons: (4 \times 2) + (1 \times 5) + 6 + 1 = 20
- Six electrons remaining, three lone pairs on oxygen

Methane ($\text{CH}_4$): Tetrahedral due to four electron groups around carbon.
Ethane ($\text{C}2\text{H}4$): Planar because each carbon has three electron groups (no lone pairs).
Ethyne ($\text{C}2\text{H}2$): Linear due to two electron groups around each carbon.

Challenge: Explaining tetrahedral methane ($\text{CH}_4$) with carbon's electronic structure ($\text{1s}^2 \text{2s}^2 \text{2p}^2$).
VSEPR theory: Atomic orbitals (one 2s and three 2p) hybridize to form four $\text{sp}^3$ hybrid orbitals of equal energy.
Hybridization: Mixing of atomic orbitals to create new hybrid orbitals that explain molecular shapes.
For ethylene: One p orbital is left out of hybridization; one s orbital and two p orbitals mix to create three $\text{sp}^2$ hybrid orbitals (trigonal planar).
Double bond: Composed of a sigma ($\sigma$) and a pi ($\pi$) bond.
- Sigma bond: Formed by the hybrid orbitals.
- Pi bond: Formed by overlap of leftover p orbitals.
For linear alkyne: One s and one p orbital mixing to make a hybrid orbital.