Summarized Org Chem 3

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Last updated 12:22 PM on 9/1/26

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18 Terms

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Substitution
Addition
Elimination
Rearrangement

What are the 4 general types of organic reactions?

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$<ul><li>Homolysis ($A:B \rightarrow A\cdot + B\cdot$): Bond cleavage where each atom takes one electron, forming radicals.</li><li>Heterolysis ($A:B \rightarrow A^+ + :B^-$): Bond cleavage where one atom takes both electrons, forming ions. Almost always occurs at polar bonds.</li></ul>$

Homolysis ( $A:B \rightarrow A\cdot + B\cdot$ ): Bond cleavage where each atom takes one electron, forming radicals.
Heterolysis ( $A:B \rightarrow A^+ + :B^-$ ): Bond cleavage where one atom takes both electrons, forming ions. Almost always occurs at polar bonds.

What is the difference between Homolytic and Heterolytic bond cleavage?

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Electrophiles ("electron-loving"): Electron-deficient species that seek electrons to obtain a stable valence shell (e.g., carbocations).
Nucleophiles ("nucleus-loving"): Electron-rich species that seek a proton or positively charged center (e.g., carbanions).

Define Electrophiles and Nucleophiles.

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Carbocation: 6 valence electrons, positive charge ($C^+$).
Carbanion: 8 valence electrons, negative charge ($C:^-$)

What are the valence electron counts and charges for Carbocations and Carbanions?

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They show the movement/flow of electrons. An arrow starts at a site of higher electron density (lone pair or covalent bond) and points toward an electron-deficient site

What do curved arrows represent in reaction mechanisms?

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Brønsted-Lowry: Acid = proton ($H^+$) donor; Base = proton acceptor.
Lewis: Acid = electron-pair acceptor; Base = electron-pair donor.

Compare the Brønsted-Lowry vs. Lewis definitions of Acids and Bases.

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$<ul><li>$pK_a = -\log(K_a)$</li><li>Larger $K_a$ = Stronger acid</li><li>Smaller $pK_a$ = Stronger acid (Larger $pK_a$ means weaker acid).</li></ul>$

$pK_a = -\log(K_a)$
Larger $K_a$ = Stronger acid
Smaller $pK_a$ = Stronger acid (Larger $pK_a$ means weaker acid).

How are $K_a$ and $pK_a$ related to acid strength?

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The stronger the acid, the weaker its conjugate base. Conversely, a weak acid has a strong conjugate base.

What is the relationship between the strength of an acid and its conjugate base?

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Acid-base reactions always favor the formation of the weaker acid and weaker base (equilibrium shifts away from the stronger acid/base). The weaker pair is always on the same side.

How do you predict the outcome/equilibrium of an acid-base reaction?

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<ul><li><p><span><strong>Across a Row (Left to Right):</strong> Acidity increases due to increasing <strong>electronegativity</strong> (polarizes bond and stabilizes </span><span style="line-height: 1.15;">$A^-$</span><span>).</span></p></li><li><p><span><strong>Down a Column (Top to Bottom):</strong> Acidity increases due to decreasing <strong>bond strength</strong> to hydrogen.</span></p></li></ul><p></p>

Across a Row (Left to Right): Acidity increases due to increasing electronegativity (polarizes bond and stabilizes $A^-$).
Down a Column (Top to Bottom): Acidity increases due to decreasing bond strength to hydrogen.

How does periodic table position affect acidity in rows vs. columns?

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$Hydrogens on orbitals with more $s$-character are more acidic ($sp > sp^2 > sp^3$) because $s$-orbitals hold electrons closer to the nucleus, stabilizing the resulting anion.<ul><li>Relative acidity: $HC\equiv CH > H_2C=CH_2 > H_3C-CH_3$</li></ul>$

Hydrogens on orbitals with more $s$-character are more acidic ($sp > sp^2 > sp^3$) because $s$-orbitals hold electrons closer to the nucleus, stabilizing the resulting anion.

Relative acidity: HC\equiv CH > H_2C=CH_2 > H_3C-CH_3

How does orbital hybridization affect hydrocarbon acidity?

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<p><span>The transmission of charge through space/bonds. Electron-Withdrawing Groups (EWGs, e.g., </span><span style="line-height: 1.15;">$F, Cl$</span><span>) pull electron density away, highly polarizing the acidic </span><span style="line-height: 1.15;">$H$</span><span> and stabilizing the conjugate base, thereby <strong>increasing acidity</strong>.</span></p>

The transmission of charge through space/bonds. Electron-Withdrawing Groups (EWGs, e.g., $F, Cl$) pull electron density away, highly polarizing the acidic $H$ and stabilizing the conjugate base, thereby increasing acidity.

What is the inductive effect on acidity?

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<ul><li><p><span><strong>Resonance Stabilization:</strong> The carboxylate anion is resonance-stabilized across two equivalent oxygen atoms.</span></p></li><li><p><span><strong>Inductive Effect:</strong> The highly polarized carbonyl group (</span><span style="line-height: 1.15;">$C=O$</span><span>) draws electron density away from the hydroxyl group.</span></p></li></ul><p></p>

Resonance Stabilization: The carboxylate anion is resonance-stabilized across two equivalent oxygen atoms.
Inductive Effect: The highly polarized carbonyl group ($C=O$) draws electron density away from the hydroxyl group.

Why are carboxylic acids significantly more acidic than alcohols?

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<p><span>Any base stronger than hydroxide (</span><span style="line-height: 1.15;">$OH^-$</span><span>) converts to </span><span style="line-height: 1.15;">$OH^-$</span><span> in water. To use stronger bases (like </span><span style="line-height: 1.15;">$NH_2^-$</span><span> or alkyl lithiums), <strong>nonaqueous solvents</strong> like liquid </span><span style="line-height: 1.15;">$NH_3$</span><span> or hexane must be used.</span></p>

Any base stronger than hydroxide ($OH^-$) converts to $OH^-$ in water. To use stronger bases (like $NH_2^-$ or alkyl lithiums), nonaqueous solvents like liquid $NH_3$ or hexane must be used.

What is the leveling effect of water, and how are stronger bases used?

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$<ul><li>$\Delta G^\circ = -RT \ln(K_{eq})$</li><li>$\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$</li><li>If $\Delta G^\circ < 0$, products are favored ($K_{eq} > 1$).</li></ul>$

$\Delta G^\circ = -RT \ln(K_{eq})$
$\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$
If \Delta G^\circ < 0, products are favored ($K_{eq} > 1$).

What are the mathematical relationships for standard free energy ( $\Delta G^\circ$ )?

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$<ul><li>Organic molecules with lone pairs on heteroatoms ($O, N$) can accept protons (e.g., alcohols, ethers, ketones).</li><li>$\pi$-electrons in alkenes can act as bases by reacting with strong acids to form carbocations.</li></ul>$

Organic molecules with lone pairs on heteroatoms ($O, N$) can accept protons (e.g., alcohols, ethers, ketones).
$\pi$ -electrons in alkenes can act as bases by reacting with strong acids to form carbocations.

How can organic compounds without negative charges act as bases?

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Step 1: Brønsted acid-base reaction (alcohol oxygen gets protonated by hydronium).
Step 2: Reverse Lewis acid-base reaction / Heterolytic cleavage (alboxonium ion loses water to form a carbocation).
Step 3: Lewis acid-base reaction ($Cl^-$ nucleophile attacks carbocation electrophile)

What are the 3 reaction steps in the conversion of tert-butyl alcohol to tert-butyl chloride using $HCl$?

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Deprotonate the terminal alkyne using a strong base (e.g., $NH_2^-$ in liquid $NH_3$). Titrate the resulting alkynide anion with heavy water ($D_2O$ or $T_2O$).

How are Deuterium ($^2H$) and Tritium ($^3H$) isotopic labels added to alkynes?