Chapter 4: Acids and Bases – Organic Chemistry Study Notes

Acids and Bases

Arrhenius, Brønsted-Lowry, and Lewis Definitions

Organic chemistry relies on several definitions of acids and bases, each with unique criteria and applications. Understanding these definitions is essential for analyzing chemical reactivity and mechanisms.

• Arrhenius Definition:

  • Acid: Increases the concentration of H+ ions (protons) in aqueous solution.

  • Base: Increases the concentration of OH- ions in aqueous solution.

  • Limitations: Only applies to aqueous solutions; does not account for hydronium ion formation ($\ce{H3O^+}$) or bases without hydroxide ions.

  • Example Equations:

    • $\ce{HCl + H2O \rightarrow H3O^+ + Cl^-}$

    • $\ce{NaOH + H2O \rightarrow Na^+ + OH^- + H2O}$

• Brønsted-Lowry Definition:

  • Acid: Donates H+ ions (protons).

  • Base: Accepts H+ ions.

  • Example Equations:

    • $\ce{HCl + H2O \rightarrow H3O^+ + Cl^-}$

    • $\ce{Na^+ + OH^- + H2O \rightarrow Na^+ + H2O + OH^-}$

• Lewis Definition:

  • Acid: Accepts an electron pair.

  • Base: Donates an electron pair (usually has a lone pair).

  • Example: $\ce{AlCl3}$ (Lewis acid) reacts with $\ce{(CH3)2O}$ (Lewis base) to form a new bond.

Arrow Pushing and Mechanisms

Arrow pushing is a technique used to illustrate the movement of electrons during acid-base reactions. Curved arrows show the flow of electron pairs, which is fundamental for understanding reaction mechanisms.

• Key Point: The arrow starts at the electron source (lone pair or bond) and points to the electron sink (atom or bond formation).

• Example: Proton transfer between water and ammonia, or between alcohol and amine.

Acid Strength and pKa

pKa and Acid Strength

The acid dissociation constant (pKa) is a quantitative measure of acid strength. It is defined as:

• Smaller pKa: Stronger acid (more likely to lose a proton).

• Stronger acid: Weaker conjugate base.

Example Table:

Acid	Name	pKa	Conjugate Base	Strength
$\ce{HCl}$	Hydrochloric acid	-7	$\ce{Cl^-}$	Strongest
$\ce{CH3COOH}$	Acetic acid	4.76	$\ce{CH3COO^-}$	Moderate
$\ce{CH3CH2OH}$	Ethanol	16	$\ce{CH3CH2O^-}$	Weak

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation relates pKa and pH, allowing prediction of the ratio of acid and base forms in solution:

• $\mathrm{p}K_a = \mathrm{pH} + \log\left(\frac{[HA]}{[A^-]}\right)$

• pH < pKa: Compound exists primarily in acidic form (protonated).

• pH > pKa: Compound exists primarily in basic form (deprotonated).

• Example: For a molecule with pKa values 2.5, 5.5, and 11, the structure will change as pH varies.

Laboratory Application

Acid-base properties are used in organic labs for separation techniques. For example, separating anisole and benzoic acid by exploiting their differing solubilities in acid/base conditions.

Factors Affecting Acid Strength

Electronegativity

Increasing electronegativity of the atom bonded to hydrogen increases acid strength. The more electronegative the atom, the easier it is for the hydrogen to leave as a proton.

• Example: $\ce{HF}$ is a stronger acid than $\ce{CH4}$ due to fluorine's high electronegativity.

Size

As atomic size increases (down a group), acid strength increases because the negative charge on the conjugate base is more stable when spread over a larger volume.

• Example: $\ce{HI}$ is a stronger acid than $\ce{HCl}$.

Hybridization

Greater 's' character in the atom's hybridization leads to stronger acids. This is due to shorter, stronger bonds that make proton loss easier.

Compound	Conjugate Base	Hybridization	% s Character	pKa
$\ce{CH3CH3}$	$\ce{CH3CH2^-}$	sp3	25%	~50
$\ce{NH3}$	$\ce{NH2^-}$	sp3	25%	~38
$\ce{HC\equiv CH}$	$\ce{HC\equiv C^-}$	sp	50%	~25
$\ce{ROH}$	$\ce{RO^-}$	sp3	25%	~16

Inductive Electron Withdrawal

Electron-withdrawing groups (substituents) increase acid strength by stabilizing the negative charge on the conjugate base. The effect is stronger when the group is more electronegative or closer to the acidic hydrogen.

• Example: Chloroacetic acid is stronger than acetic acid due to the electron-withdrawing chlorine atom.

Electron Delocalization (Resonance)

Delocalization of electrons on the conjugate base (resonance) increases acid strength by stabilizing the negative charge.

• Example: Acetate ion ($\ce{CH3COO^-}$) is stabilized by resonance, making acetic acid a stronger acid than ethanol.

Structural Effect	Representative Acids	Conjugate Bases	Description
Electronegativity	$\ce{H2S}$, $\ce{HCl}$	$\ce{HS^-}$, $\ce{Cl^-}$	Chloride is more electronegative; negative charge is held more tightly.
Size	$\ce{H2S}$, $\ce{HBr}$, $\ce{HI}$	$\ce{HS^-}$, $\ce{Br^-}$, $\ce{I^-}$	Larger anions spread charge over greater volume.
Resonance	$\ce{CH3COOH}$, $\ce{C6H5COOH}$	$\ce{CH3COO^-}$, $\ce{C6H5COO^-}$	Resonance stabilizes the conjugate base.
Induction	$\ce{CH3COOH}$, $\ce{CCl3COOH}$	$\ce{CH3COO^-}$, $\ce{CCl3COO^-}$	Electron-withdrawing substituents stabilize the base.
Hybridization	$\ce{C6H5OH}$, $\ce{C6H5NH2}$	$\ce{C6H5O^-}$, $\ce{C6H5NH^-}$	More s character leads to stronger acids.

Common Organic Acids and Bases

Carboxylic Acids

Carboxylic acids are the most common organic acids, characterized by the –COOH group. Their pKa is typically 3–5. They can act as acids or bases, and their strength depends on the R group, which influences inductive withdrawal and resonance stabilization.

• Example: Acetic acid ($\ce{CH3COOH}$), chloroacetic acid ($\ce{ClCH2COOH}$).

Alcohols

Alcohols are weaker acids than carboxylic acids due to less inductive withdrawal and localized electrons. The –OH group is key, with pKa values around 15–16. Alcohols can act as acids or bases.

• Example: Ethanol ($\ce{CH3CH2OH}$).

Amines

Amines are the most common organic bases and rarely act as acids. They are weaker than alcohols due to the lower electronegativity of nitrogen compared to oxygen. pKa values are typically 30–40.

• Example: Methylamine ($\ce{CH3NH2}$).

Protonated Compounds

Protonated compounds have an additional proton and are stronger acids than their non-protonated counterparts. Their pKa is usually less than 1.

• Example: Protonated acetic acid ($\ce{CH3COOH2^+}$), protonated methyl alcohol ($\ce{CH3OH2^+}$).

Conjugate Acids and Bases

Definitions and Identification

In acid-base reactions, removing a proton yields a conjugate base, while adding a proton yields a conjugate acid. The species with the extra hydrogen is always the acid.

• Example: $\ce{HCO_3^-}$ (bicarbonate) loses a proton to become $\ce{CO_3^{2-}}$ (conjugate base); gains a proton to become $\ce{H_2CO_3}$ (conjugate acid).

Acid	Conjugate Base
$\ce{H_3O^+}$	$\ce{H_2O}$
$\ce{NH_4^+}$	$\ce{NH_3}$
$\ce{CH_3COOH}$	$\ce{CH_3COO^-}$
Base	Conjugate Acid
$\ce{CH_3O^-}$	$\ce{CH_3OH}$
$\ce{NH_2^-}$	$\ce{NH_3}$

Practice Examples

• Identify: What is the conjugate base of $\ce{HSO_4^-}$? Answer: $\ce{SO_4^{2-}}$

• Identify: What is the conjugate acid of $\ce{HCO_3^-}$? Answer: $\ce{H_2CO_3}$

Summary Table: Factors Affecting Acid Strength

Factor	Effect	Example
Electronegativity	Higher electronegativity increases acid strength	$\ce{HF}$ > $\ce{CH4}$
Size	Larger atom stabilizes conjugate base	$\ce{HI}$ > $\ce{HCl}$
Hybridization	More s character = stronger acid	$\ce{HC\equiv CH}$ > $\ce{CH2=CH2}$ > $\ce{CH3CH3}$
Inductive Effect	Electron-withdrawing groups stabilize base	Chloroacetic acid > acetic acid
Resonance	Delocalization stabilizes base	Acetate ion

Additional info: These notes expand on the original slides by providing definitions, equations, and context for each concept, as well as reconstructed tables for clarity and completeness.