The Mistake Bank
Chapter 4: Chemical Bonding
Bonds break, bonds form. Make sure your logic holds together.
Shape vs. Geometry
VSEPR TheoryScenario: What is the shape of the water molecule (\(H_2O\))?
Student sees \(sp^3\) hybridization and 4 electron pairs.
Answer: "Tetrahedral".
(Confusing geometry with shape!)
Distinguish Shape from Geometry!
Geometry (Electron Geometry): Includes lone pairs. For \(H_2O\), it IS Tetrahedral.
Shape (Molecular Geometry): Ignores lone pairs, looks only at atoms.
$$ \text{Shape of } H_2O = \mathbf{Bent \text{ or } V-shaped} $$
The Dipole Moment Trap
PolarityScenario: Which has a higher dipole moment: \(NH_3\) or \(NF_3\)?
Student thinks: "Fluorine is the most electronegative element, so the polarity must be higher."
$$ \mu_{NF_3} > \mu_{NH_3} $$
(Incorrect Vector Addition!)
Look at the Vectors!
In \(NH_3\), the orbital dipole (lone pair) and bond dipoles (N-H) point in the same direction (additive).
In \(NF_3\), the lone pair points up, but N-F bonds point down (subtractive).
$$ \mu_{NH_3} > \mu_{NF_3} $$
Counting Pi Bonds
Sigma & PiScenario: How many \(\sigma\) and \(\pi\) bonds are in a triple bond ($C \equiv C$)?
Student assumes a triple bond means 3 \(\pi\) bonds.
Or forgets to count the sigma bond.
The First Bond is Always Sigma!
Any multiple bond starts with one strong head-on overlap (\(\sigma\)). The rest are lateral overlaps (\(\pi\)).
- Single Bond: 1 \(\sigma\)
- Double Bond: 1 \(\sigma\), 1 \(\pi\)
- Triple Bond: 1 \(\sigma\), 2 \(\pi\)
Oxygen's Magnetism
Molecular Orbital TheoryScenario: Explain the magnetic nature of \(O_2\).
Student draws the Lewis structure:
$$ :\: \ddot{O} = \ddot{O} \: : $$
All electrons look paired. Answer: "Diamagnetic".
(Valence Bond Theory fails here!)
Use M.O.T!
Molecular Orbital Theory reveals two unpaired electrons in the antibonding orbitals (\(\pi^* 2p_x, \pi^* 2p_y\)).
Therefore, \(O_2\) is Paramagnetic.
PCl5 Bond Lengths
HybridizationScenario: Are all 5 P-Cl bonds in \(PCl_5\) equivalent?
Student assumes symmetry implies equality.
Answer: "Yes, all are same length."
Axial vs Equatorial!
In Trigonal Bipyramidal geometry:
- 3 Equatorial bonds: \(120^\circ\) apart.
- 2 Axial bonds: \(90^\circ\) to the equatorial plane.
The Axial bonds suffer more repulsion from the equatorial pairs, making them longer and weaker.
Formal Charge vs. Reality
Lewis StructuresScenario: Calculate the formal charge on Oxygen in Ozone (\(O_3\)). Does it mean the atom is actually an ion?
Student calculates a +1 charge on the central Oxygen and assumes it behaves like a cation.
Or forgets the formula entirely.
Formal Charge \(\neq\) Real Charge
Formal charge is just a bookkeeping tool to find the most stable Lewis structure (lowest energy).
$$ FC = V - L - \frac{1}{2}S $$
V=Valence, L=Lone pair e-, S=Shared e-
Confess Your Sins!
"A bond is only as strong as its weakest link... usually your memory of exceptions."
Did one of these catch you? Or do you have a different horror story from your last exam?
Scroll down to the comments section below and tell us:
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