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Bonding

Chapter 3: Bonding

Updated Apr 10, 2026
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1. (3.1) Atomic orbitals are best described as:
A. Orbitals are mathematical solutions (wavefunctions) of the Schrödinger equation whose squared magnitude gives the probability density.
2. (3.1) A 2p orbital is shaped like:
C. Each of the three p orbitals (pxp_{x}, pyp_{y}, pzp_{z}) is a two-lobed dumbbell oriented along one axis.
3. (3.2) Molecular orbital theory treats bonds as:
B. Constructive interference gives bonding MOs; destructive interference gives antibonding MOs.
4. (3.2) The MO bond order is calculated as:
D. Bond order = 0 implies no bond. O₂ bond order is 2, matching its observed double bond (and paramagnetism, via the MO picture).
5. (3.3) Orbital hybridization is:
A. Hybridization is a mathematical convenience that matches predicted geometry (from VSEPR) to orbital overlap.
6. (3.3) An sp² carbon has:
C. The unhybridized p forms the π bond of a double bond. Examples: alkene carbons, carbonyl carbon.
7. (3.4) Sigma (σ) bonds are:
B. σ bonds have cylindrical symmetry about the bond axis and permit free rotation in single bonds.
8. (3.4) Every single bond in organic molecules is:
D. Multiple bonds have one σ plus additional π (double = σ + π; triple = σ + 2π).
9. (3.5) Pi (π) bonds are:
A. π bonds restrict rotation around a double bond, which is why cis/trans isomerism exists.
10. (3.5) A double bond consists of:
C. A triple bond = one σ + two π. The σ bond supplies most of the bond energy.
11. (3.6) Bond length:
B. C-C ≈ 1.54 Å; C=C ≈ 1.34 Å; C≡C ≈ 1.20 Å. Triple bonds are shortest and strongest.
12. (3.6) Among C-C bond types, bond strength increases in the order:
D. Triple > double > single in bond energy. A triple bond averages ~230 kcal/mol; a single bond ~83 kcal/mol.
13. (3.7) Conjugation refers to:
A. 1,3-butadiene is the simplest conjugated system. Lone pairs adjacent to a π bond can also conjugate.
14. (3.7) Conjugated systems are generally:
C. The extra stability is called resonance or delocalization energy.
15. (3.8) Resonance structures:
B. Only π electrons and lone pairs move; σ framework stays the same. More stable contributors contribute more to the hybrid.
16. (3.8) The most important resonance contributor typically:
D. Full octets and low charge separation are energetically favored. If charges are unavoidable, negative charge prefers oxygen over carbon.
17. (3.9) An aromatic compound:
A. Pyridine, furan, and tropylium cation are aromatic with heteroatoms or non-benzene rings.
18. (3.9) Benzene's resonance (aromatic) stabilization is approximately:
C. Measured by comparing heats of hydrogenation.
19. (3.10) VSEPR theory predicts molecular geometry from:
B. Lone pairs are "larger" than bonding pairs, compressing bond angles.
20. (3.10) An sp³ central atom with 4 bonds and 0 lone pairs has:
D. CH₄ is the classic example. With 3 bonds + 1 lone pair, geometry becomes pyramidal (NH₃).
21. (3.11) A molecule is polar when:
A. Both bond polarity and molecular shape matter. CCl₄ (tetrahedral) is nonpolar because its four C-Cl dipoles cancel.
22. (3.11) CO₂ is:
C. Individual C=O bonds are polar, but the linear shape cancels their dipoles.
23. (3.12) The strongest common IMF in organic molecules is typically:
B. H-bonds account for water's elevated boiling point, DNA base pairing, and much of protein structure.
24. (3.12) London dispersion forces scale with:
D. This is why straight-chain alkanes boil at successively higher temperatures as chain length grows.

Everything in organic chemistry comes down to electrons. Where they live, how they move, which atoms share them, and how easily they leave. Reactions happen because electrons move. Molecules have shape because electrons fill specific regions of space. Polarity, boiling points, solubility, and reactivity all trace back to a single question: where are the electrons?

This chapter is the playbook for answering that question. You will learn how atomic orbitals combine to form molecular orbitals, how hybridization explains the tetrahedral methane of your first chemistry class and the trigonal planar benzene of your biochem exam, and why some molecules are planar, conjugated, and astonishingly stable (aromatic) while others of the same electron count distort themselves out of planarity to avoid being destabilized (antiaromatic).

Once you see bonding through the lens of orbitals, resonance, and hybridization, organic reactions stop looking like magic. Nucleophiles attack electrophiles because one has an accessible electron pair and the other has an empty orbital waiting to accept it. Acids are acidic because their conjugate bases stabilize the extra electrons. It all fits together.

The Central Analogy

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