Chapter 5 — Key Takeaways

What you should leave Chapter 5 with

1. Alkanes are the simplest hydrocarbons: all $sp^3$ carbons, all single bonds, essentially unreactive. They are the baseline against which other molecules' reactivity is measured.

2. Conformations are 3D arrangements accessible by single-bond rotation (no bonds broken). Different conformations have different energies; molecules interconvert rapidly at room temperature.

3. Newman projection is the standard tool for visualizing conformations along a single bond. The front carbon is shown as a dot; the back carbon as a circle.

4. Staggered vs eclipsed (ethane): staggered is more stable than eclipsed by 3 kcal/mol, due to torsional strain in the eclipsed form (electron-electron repulsion + lost hyperconjugation).

5. Anti vs gauche (butane): anti (180° methyl-methyl) is more stable than gauche (60° methyl-methyl) by ~0.9 kcal/mol, due to steric strain in gauche.

6. Cyclohexane in the chair conformation is essentially strain-free. The chair has all bond angles ~109.5° (no angle strain) and all adjacent C-H bonds staggered (no torsional strain).

7. Axial vs equatorial: each carbon of chair cyclohexane has one axial bond (perpendicular to ring plane) and one equatorial (roughly in plane). Axial substituents have 1,3-diaxial interactions (steric clashes with axial H's at C3 and C5).

8. Substituents prefer equatorial position. The energy cost of axial vs equatorial is the A-value (e.g., -CH₃: 1.7 kcal/mol; -tBu: 4.9 kcal/mol). At equilibrium, the chair with the largest substituents equatorial dominates.

9. Ring flip: a chair can interconvert to its alternate chair by a ring flip (~10.5 kcal/mol barrier; fast at room temperature). Each ring flip swaps every axial substituent to equatorial and vice versa.

10. Other rings:

    • Cyclopropane (3-membered): ~27 kcal/mol strain (angle + torsional).
    • Cyclobutane (4): ~26 kcal/mol strain.
    • Cyclopentane (5): ~6 kcal/mol strain.
    • Cyclohexane (6): essentially strain-free.
    • Cycloheptane and larger: small transannular strain.

11. Thermodynamics ($\Delta G = \Delta H - T\Delta S$) determines whether a reaction can happen:

    • $\Delta G < 0$: thermodynamically favorable.
    • $K_{eq} = e^{-\Delta G/RT}$.
    • Rule of thumb: every 1.4 kcal/mol of $\Delta G$ changes $K_{eq}$ by a factor of 10.

12. Bond dissociation energies (BDEs) let you estimate $\Delta H$ of a reaction: $$\Delta H_{rxn} \approx \Sigma \text{BDE(bonds broken)} - \Sigma \text{BDE(bonds formed)}$$ Gets within 5-10 kcal/mol for most reactions; useful for predicting whether a reaction is enthalpically favorable.

13. Heats of combustion of cycloalkanes reveal ring strain. Cyclohexane is the reference (no strain); cyclopropane has ~9 kcal/mol higher combustion enthalpy per CH₂ unit (corresponding to ~27 kcal/mol total ring strain).

14. Kinetics (Arrhenius equation, $k = A e^{-E_a/RT}$) determines how fast a reaction runs:

    • $E_a$ in the 15-25 kcal/mol range is the sweet spot for laboratory reactions at ambient conditions.
    • Every 10 °C increase ~doubles the rate (typical $E_a$).

15. Transition state (TS) is the highest-energy configuration on the reaction path; not a real molecule (lasts ~10⁻¹³ s); drawn with partial bonds.

16. Hammond postulate: the TS resembles whichever of reactants or products is closer to it in energy.

    • Exothermic ($\Delta G < 0$): early TS (resembles reactants).
    • Endothermic ($\Delta G > 0$): late TS (resembles products).
    • Highly exothermic reactions have TSs almost identical to reactants.

17. Thermodynamic vs kinetic control:

    • Thermodynamic: reaction reversible; products equilibrate; major product is most stable.
    • Kinetic: reaction irreversible; major product reflects lowest-$E_a$ pathway.
    • Low T → kinetic; high T → thermodynamic.

18. Catalysis lowers $E_a$ without changing $\Delta G$. A thermodynamically favorable but kinetically blocked reaction becomes viable when a catalyst lowers the barrier.

19. Independence of $\Delta G$ and $E_a$: a reaction can be very favorable ($\Delta G$ very negative) but very slow ($E_a$ high). Example: hydrogen + oxygen at room temperature ($\Delta G \approx -50$ kcal/mol but no reaction without spark).

20. Mastery of Chapter 5 gives you the energetic framework for every subsequent reaction. Every mechanism in the book will be analyzed using:

    • Bond changes → $\Delta H$ estimates.
    • Conformational preferences → which TS is favored.
    • A-values and steric arguments → product distributions.
    • Hammond postulate → predicting TS structure.
    • Thermodynamic vs kinetic → choosing conditions for desired product.

Cross-references

• Chapter 2 — Bonding (foundation; sp³ hybridization, bond angles).
• Chapter 4 — Functional groups (alkane nomenclature).
• Chapter 6 — Spectroscopy (IR identifies C-H stretches at ~2900 cm⁻¹).
• Chapter 10 — SN2 (conformational analysis of TS).
• Chapter 12 — E2 elimination (anti-periplanar geometry; chair-flip arguments).
• Chapter 15 — Alkene addition (Hammond postulate predicts Markovnikov regiochemistry).
• Chapter 19 — Diels-Alder (kinetic vs thermodynamic adducts).
• Chapter 27 — Enolates (kinetic vs thermodynamic enolate).
• Chapter 32 — Carbohydrates (glucose pyranose chair, all-equatorial preference).
• Appendix A — Bond dissociation energy table.

Study tip

For each new reaction you encounter, work through the 5-question energetic checklist:

1. What bonds break? What bonds form? Sum BDEs to estimate $\Delta H$.
2. Is $\Delta S$ favorable or unfavorable? Bond formation typically decreases entropy; ring opening increases entropy.
3. Is $\Delta G < 0$ at the temperature of interest? Combine $\Delta H$ and $\Delta S$.
4. What is the $E_a$? Is it accessible at the chosen temperature? Use the rule of thumb (15-25 kcal/mol = lab conditions).
5. What does the TS look like (Hammond postulate)? Use this to predict regio/stereoselectivity.

If you can do this for any reaction, you've internalized Chapter 5.

The habit to leave with: every time you see a new reaction, ask both whether it is thermodynamically favorable and whether the activation energy is low enough at the temperature you are running the reaction. Every useful reaction satisfies both conditions. Failing to separate the two is the single most common error in reasoning about reactivity.

Chapter 6 — the last chapter of Part I — is IR and mass spectrometry. Tools for identifying molecules by their vibrations and mass fragments.