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Bonding and Chemical Interactions

Chapter 3: Bonding and Chemical Interactions

5 min read Updated Mar 26, 2026
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🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

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1. (3.1) Ionic bonds form when:
C. Ionic bonding occurs between species with large ΔEN (typically > 1.7). The resulting lattice energy stabilizes the crystal.
2. (3.1) Ionic compounds typically:
D. Strong electrostatic forces require lots of energy to break. In solid form the ions are locked; they can move (and carry current) only as melt or solution.
3. (3.2) Pure covalent (nonpolar covalent) bonds form when:
B. H₂, O₂, N₂ are perfectly nonpolar (ΔEN = 0). C-H bonds are essentially nonpolar (ΔEN ≈ 0.4).
4. (3.2) Compared with ionic compounds, covalent compounds generally:
A. Only intermolecular forces hold molecular crystals together, so less energy is needed to melt them. No free ions means no conduction.
5. (3.3) A polar covalent bond results when:
C. The more electronegative atom carries a partial negative charge (δ-), the other a δ+. Typical polar covalent range: ΔEN ≈ 0.5-1.7.
6. (3.3) The molecular dipole moment is:
D. CO₂ has polar C=O bonds but zero net dipole because the two bond vectors cancel. H₂O does not cancel, so it is polar.
7. (3.4) A Lewis structure shows:
B. Total valence electrons, octets on each non-H atom, and a formal charge check are the recipe.
8. (3.4) In a Lewis structure of CO₂, the central carbon has:
A. Two double bonds give each atom an octet with zero formal charge, yielding the well-known linear O=C=O structure.
9. (3.5) An example of an incomplete octet is:
C. Boron is happy with six electrons in BF₃. It is a strong Lewis acid because it readily accepts a lone pair to complete the octet.
10. (3.5) Expanded octets are possible for:
D. 3d orbitals provide the extra space. Second-period elements (C, N, O, F) are strictly capped at 8 electrons.
11. (3.6) VSEPR theory predicts molecular geometry by:
B. Lone pairs repel more than bonding pairs, so they push other groups closer together (e.g., H-O-H bond angle ≈ 104.5°, less than ideal tetrahedral 109.5°).
12. (3.6) A central atom with four bonding pairs and zero lone pairs has:
A. Four identical groups point to vertices of a tetrahedron with 109.5° angles. CH₄, CCl₄, and NH₄⁺ fit this template.
13. (3.7) Molecular geometry differs from electron geometry when:
C. Lone pairs occupy space but do not add a vertex to the shape. H₂O has tetrahedral electron geometry but bent molecular geometry.
14. (3.7) NH₃ has which molecular geometry?
D. Four electron domains → tetrahedral electron geometry; one lone pair leaves a trigonal pyramid. Bond angle is about 107°.
15. (3.8) sp³ hybridization involves:
B. sp³ gives the tetrahedral 109.5° geometry of CH₄. sp² (three hybrids) gives trigonal planar; sp (two hybrids) gives linear.
16. (3.8) An sp-hybridized carbon (e.g., in C≡C) has:
A. The two unhybridized p orbitals form the two π bonds of the triple bond. Result: linear geometry plus two π bonds.
17. (3.9) Sigma (σ) bonds are:
C. σ bonds are always the first bond between two atoms. They have cylindrical symmetry and allow free rotation.
18. (3.9) Pi (π) bonds:
D. Rotation about a double bond would require breaking the π overlap. This is why E/Z (cis/trans) isomers are distinct compounds.
19. (3.10) Bond order is defined as:
B. Single = 1, double = 2, triple = 3. Resonance can give fractional bond orders (e.g., 1.5 in benzene).
20. (3.10) As bond order increases:
A. More shared electrons pull the nuclei closer and take more energy to break. Triple > double > single in bond energy.
21. (3.11) Among common non-ionic intermolecular forces, the strongest is:
C. H-bonds explain the high boiling points of H₂O, NH₃, and HF and are crucial in DNA and protein structure.
22. (3.11) London dispersion forces are:
D. Larger electron clouds are more polarizable, giving stronger instantaneous dipole attractions. This is why iodine (I₂) is solid while fluorine (F₂) is gaseous at room temperature.
23. (3.12) Ion-dipole interactions:
B. Water's oxygen (δ-) orients toward cations; hydrogens (δ+) orient toward anions. This "hydration shell" helps dissolve salts.
24. (3.12) The high solubility of ionic compounds in water reflects:
A. ΔGdissolutionG_{\text{dissolution}} = ΔHlatticeH_{\text{lattice}} - ΔHhydrationH_{\text{hydration}} + entropy. Solubility is governed by whether hydration releases enough energy to break the lattice.

Think about the last time you added salt to a pot of boiling water. The crystals dissolved almost instantly, the water bubbled a little differently, and the pasta cooked just a bit faster. That one kitchen moment involves ionic bonding (the crystal lattice of NaCl), ion-dipole interactions (water molecules ripping those ions apart), colligative properties (the boiling point rising), and the polar nature of water itself. All of that from a pinch of salt.

Bonding is the reason matter exists as something more interesting than a cloud of isolated atoms. It explains why diamond is the hardest natural material while graphite crumbles under your pencil - even though both are pure carbon. It explains why oil and water refuse to mix, why DNA holds its double-helix shape, and why your cell membranes can be both flexible and waterproof at the same time.

For the MCAT, bonding is not just one chapter - it is the foundation that makes every other chemistry and biochemistry topic make sense. If you understand why atoms share, steal, or pool their electrons, you can predict molecular shape, polarity, solubility, reactivity, and physical properties without memorizing a single table.

Bonds Are a Spectrum, Not Categories

Before we dive in, throw away the idea that bonds come in neat little boxes. Bonding is a spectrum. On one end, you have purely ionic bonds where electrons are fully transferred. On the other end, you have nonpolar covalent bonds where electrons are shared equally. Everything in between - and that includes most of the molecules you will see on the MCAT - falls somewhere along this continuum. The key variable is electronegativity difference between the bonded atoms.

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