Alkenes & Alkynes – Comprehensive Study Notes

Alkenes

• Also called olefins.

• Hydrocarbons containing at least one carbon–carbon double bond ($C=C$).

• General formula: $C<em>nH</em>{2n}$ (acyclic, one double bond).

• Abundant in nature.

  • Ethylene: natural plant hormone that induces fruit ripening; large-scale industrial feedstock.

  • $\alpha$-Pinene: major component of turpentine (obtained from pine-tree resin; used medicinally for joint, muscle, nerve pain, toothache).

  • $\beta$-Carotene: orange pigment in fruits/vegetables; provitamin A.

Alkynes

• Hydrocarbons containing at least one carbon–carbon triple bond ($C\equiv C$).

• General formula for simple, acyclic mono-ynes with no other functional groups: $C<em>nH</em>{2n-2}$.

• Traditional name "acetylenes"—but acetylene refers specifically to $C<em>2H</em>2$ (ethyne).

• Unique structural features derive from $sp$ hybridization:

  • Linear geometry ($180^{\circ}$).

  • Non-polar $\sigma$ bond strength + high $\pi$-bond energy → overall strong bond but accessible $\pi$ electrons.

  • Terminal alkynes exhibit notable acidity (easy deprotonation).

Physical Properties of Alkenes & Alkynes

• Closely resemble those of corresponding alkanes with the same carbon skeleton.

• Non-polar; intermolecular attraction limited to weak London dispersion (van der Waals) forces.

• Low densities (liquid members float on water; <1.0\,\text{g mL}^{-1}).

• Soluble in non-polar solvents (kerosene, hexane, $CHCl<em>3$, $CCl</em>4$); insoluble in water.

• Low melting and boiling points relative to isomeric compounds that possess heteroatoms or stronger intermolecular forces.

Electronic Structure of Alkenes

• Each doubly bonded carbon is $sp^2$ hybridized:

  • Three $sp^2$ orbitals lie in a plane, $120^{\circ}$ apart (trigonal planar).

  • One unhybridized $p$ orbital stands perpendicular to the $sp^2$ plane.

• Bond composition:

  • $\sigma$ bond = head-on overlap of $sp^2$ orbitals.

  • $\pi$ bond = sideways overlap of $p$ orbitals (above and below the plane).

• Consequence: restricted rotation around $C=C$.

  • $\pi$ bond must break to rotate; rotational barrier $\approx$ $350\,\text{kJ mol}^{-1}$ ($84\,\text{kcal mol}^{-1}$) vs $\sim12\,\text{kJ mol}^{-1}$ for single bonds.

Geometric (Cis–Trans) Isomerism

• Because rotation is blocked, substituted alkenes may exist as distinct spatial isomers:

  • cis isomer – two specified groups on the same side of the double bond.

  • trans isomer – groups on opposite sides.

• Example: but-2-ene shows separable cis and trans forms; they do not interconvert spontaneously at ambient conditions.

• Stability trend: \text{trans} > \text{cis} (cis suffers steric strain between same-side substituents).

• Requirement: each $C^{*}$ of $C=C$ must be bonded to two different groups.

Practice Example

• Draw cis/trans 5-chloropent-2-ene. (Student should ensure each alkene carbon bears two different groups.)

Beyond Cis/Trans: E,Z (Cahn–Ingold–Prelog) Notation

• Needed for tri- or tetrasubstituted alkenes where "cis/trans" is ambiguous.

• Steps (Sequence Rules):

  1. Rank atoms directly attached to each alkene carbon by atomic number (higher $Z$ = higher priority).

  2. If tie, move outward along each chain until first point of difference.

  3. Treat multiple bonds as equivalent "phantom" single-bonded atoms (e.g.
     $\text{C}=\text{O}$ equals $\text{C!–O, C!–O}$).

• Designations:

  • E ("entgegen", opposite): high-priority groups on opposite sides.

  • Z ("zusammen", together): high-priority groups on the same side.

Worked Example

Assign $E/Z$ to $\mathrm{CH<em>3CH=C(CH</em>3)CH<em>2OH}$. Priorities: C(OH) > C(CH3)2 on one carbon; CH3 > H on the other → high groups on same side ⇒ $Z$.

Four Fundamental Types of Organic Reactions

1. Addition – two reactants combine into one product; no atoms left over.

   • Industrial: alkene $+$ $HCl$ → alkyl chloride.

   • Biochemical: fumarate $+$ $H_2O$ → malate (citric acid cycle).

2. Elimination – reverse of addition; one reactant gives two products + small molecule ($H_2O$, $HCl$).

   • Acid-catalyzed dehydration of alcohol → alkene + water.

3. Substitution – two reactants exchange parts.

   • Alkane $+$ $Cl_2$ (UV) → alkyl chloride $+$ $HCl$.

4. Rearrangement – single reactant reorganizes bonds to isomeric product.

   • Acid-catalyzed isomerization of but-1-ene → trans-but-2-ene.

Reaction Mechanisms

• Provide step-by-step account of bond making/breaking, order of events, and relative rates.

Bond Cleavage

• Homolytic (symmetrical): one electron to each fragment → radicals.

• Heterolytic (unsymmetrical): both electrons to one fragment → ions.

Bond Formation

• Symmetrical: each reactant donates one electron (radical coupling).

• Unsymmetrical: both electrons donated by one species (polar reaction).

Two Mechanistic Classes

1. Radical reactions – odd-electron species participate; depicted with half-headed "fishhook" arrows.

2. Polar (ionic) reactions – even-electron species; full curved arrows.

Nucleophiles & Electrophiles

• Nucleophile (Nu¯ / Nu:) – electron-rich, "nucleus-loving"; donates lone pair.

  • Charged (strong): $NaOCH<em>3$, $LiCH</em>3$, $NaOH$, $KCN$, $NaNH<em>2$, $NaI$, $NaN</em>3$.

  • Neutral: $H<em>2O$, $ROH$, $H</em>2S$, $RSH$.

• Electrophile (E⁺ / E) – electron-poor, "electron-loving"; accepts electron pair.

  • Acids ($H^+$), alkyl halides, carbonyl C in aldehydes/ketones, $NO_2^+$ (nitronium).

Practice Determination

• $NO_2^+$ → electrophile (positive).

• $CH_3O^-$ → nucleophile (negative).

• $CH_3OH$ → amphoteric; can act as nucleophile (lone pairs) or electrophile (polar $C–O$, $O–H$).

Mechanistic Case Study: Electrophilic Addition of HCl to Ethylene

• Ethylene ($H<em>2C=CH</em>2$) behaves as nucleophile; $\pi$ electrons above/below plane are accessible and weaker than $\sigma$ bonds.

• $HCl$ functions as strong electrophile/proton donor.

• Two-step polar mechanism:

  1. $\pi$ electrons attack $H^+$, break $H–Cl$; generate carbocation ($CH<em>3CH</em>2^+$) + $Cl^-$.

  2. $Cl^-$ rapidly attacks carbocation → $CH<em>3CH</em>2Cl$ (chloroethane).

• Same pattern applies to other alkenes (e.g.
  cyclohexene → chlorocyclohexane).

Transition States, Intermediates & Energy Diagrams

• Reaction progress visualized with energy diagram: vertical axis = energy, horizontal = reaction coordinate.

• Transition state (‡): highest-energy point along path; cannot be isolated.

• Activation energy $E_{act}$: energy gap between reactants and transition state; determines rate.

  • Typical organic $E_{act}$: $40–125\,\text{kJ mol}^{-1}$ ($10–30\,\text{kcal mol}^{-1}$).

  • E_{act} < 80\,\text{kJ mol}^{-1} → reaction proceeds near room T; higher values often need heating.

• Intermediate: local energy minimum between two transition states (e.g.
  carbocation), longer-lived than transition state but often short-lived overall.

• Overall reaction enthalpy $\Delta E<em>{rxn} = E</em>{products}-E_{reactants}$:

  • \Delta E_{rxn} < 0 → exothermic (energy released, thermodynamically favorable).

  • \Delta E_{rxn} > 0 → endothermic (energy absorbed).

• In multistep processes, rate-determining step = slowest step (highest $E_{act}$).

Catalysis

• Catalyst provides alternate pathway with lower $E_{act}$; emerges regenerated.

  • Example: $Pd$-catalyzed hydrogenation of alkenes (vegetable oil → margarine); reaction negligible without $Pd$ even at high T, but rapid at room T with $Pd$.

• Enzymes: biological macromolecular catalysts; orchestrate reactions via numerous low-barrier steps in a finely tuned active site.

Summary of Key Principles

• Alkenes ($C=C$) & alkynes ($C\equiv C$) are unsaturated hydrocarbons, non-polar, low-density, low-bp/mp.

• $sp^2$ hybridization + $\pi$ bond prohibit rotation, giving rise to cis/trans and $E/Z$ stereoisomerism; trans (or $E$ when priorities oppose) usually more stable.

• Organic reactions classified as addition, elimination, substitution, rearrangement; understood via detailed mechanisms.

• Polar (even-electron) vs radical (odd-electron) mechanisms dictated by symmetrical vs unsymmetrical electron flow.

• Nucleophiles donate, electrophiles accept electron pairs; reaction course follows electron-rich → electron-poor direction along polar bonds.

• Mechanistic analysis employs concepts of transition states, activation energy, intermediates, reaction coordinate diagrams.

• Catalysts (including enzymes) accelerate reactions by lowering activation barriers without altering overall thermodynamics.

End of Notes