8.6.A.1  Acid-Base Strength and Conjugate Stability Based on Molecular Structure:

1. Acid-Base Theory & Conjugates:

i. Brønsted–Lowry Acid-Base Theory:

Definition:

* Acid: A compound that donates a proton (H⁺).
* Base: A compound that accepts a proton (H⁺).

This theory is an extension of the Arrhenius definition in that it does not demand that the compounds be in aqueous solution.

Example Reaction:

Reaction:

$$\text{HCl (acid)}+\text{H₂O (base)}\to \text{Cl⁻ (conjugate base)}+\text{H₃O⁺ (conjugate acid)}$$

ii. Conjugate Acid-Base Pairs:

In each Brønsted–Lowry acid-base reaction:

* The acid is transformed into its conjugate base after releasing a proton.
* The base is transformed into its conjugate acid after gaining a proton.

Example Pair:

* NH₃ / NH₄⁺ → NH₃ is a base; NH₄⁺ is its conjugate acid.

iii. Relationship Between Strength and Conjugate Stability:

a.  Strong Acids → Weak Conjugate Bases

* A strong acid fully dissociates in solution.
* Its conjugate base is very weak and stable, with minimal tendency to re-accept a proton.

Example:

* HCl (strong acid) → Cl⁻ (weak conjugate base)

b. Weak Acids → Stronger Conjugate Bases:

* A weak acid only partially dissociates.
* Its conjugate base is stronger and less stable, with more tendency to accept a proton.

Example:

* CH₃COOH (weak acid) ⇌ CH₃COO⁻ (stronger conjugate base)

iv. Key Takeaways:

* The more powerful the acid, the less capable is its conjugate base.
* The more powerful the base, the less capable is its conjugate acid.
* Stability of conjugates (usually because of resonance, electronegativity, or size) is an important consideration:

* More stable conjugates = weaker conjugate species.
* Negative charge delocalization = increased conjugate base stability = more powerful parent acid.

2. Molecular Structure & Acidity:

a. Identifying Acidic Protons:

# Why is a Proton Acidic?

An acidic proton is one that may be donated as H⁺, with the resulting stable conjugate base.

b. How to Identify:

* Check for H atoms bonded to:

* Strongly electronegative atoms (N, O, halogens).
* Carbon atoms next to electron-withdrawing groups (EWGs).
* Atoms part of resonance systems.

c. Typical Acidic Groups:

d. Factors Stabilizing the Conjugate Base:

The more stable the conjugate base, the stronger the acid. Three broad factors:

A. Electronegativity:

* More electronegative atoms can stabilize negative charge on the conjugate base better.
* Conjugate base (A⁻) will be more stable if the negative charge resides on a highly electronegative atom.

Example:

* H–F > H–O > H–N in acidity

* F⁻ is more stable due to increased electronegativity.

B. Inductive Effects (Electron Withdrawal Through Sigma Bonds):

* Electronegative atoms or groups withdraw electron density away, distributing the negative charge.
* More EWGs around = more stable conjugate base = stronger acid.

Example:

* CH₃COOH (acetic acid) vs. CF₃COOH (trifluoroacetic acid)

* CF₃ group stabilizes the conjugate base by inductive withdrawal → stronger acid.

 C. Resonance Delocalization\:

* Resonance distributes the negative charge over several atoms.
* This delocalization significantly stabilizes the conjugate base.

Example:

* Acetic acid: CH₃COO⁻ conjugate base has resonance between both oxygen atoms → stable → stronger acid.

* Phenol: Phenoxide ion (C₆H₅O⁻) resonance-stabilized over the aromatic ring.

e. Summary Table:

3. Functional Groups & Common Examples:

i. Acidic Functional Groups:

These groups donate protons (H⁺) and give rise to conjugate bases.

ii. Basic Functional Groups:

These groups accept protons (H⁺) and give rise to conjugate acids.

iii. Summary of Relative Strengths:

a.  Acids (Strong → Weak):
HCl, HNO₃ > H₂SO₄ > TsOH > Carboxylic acids > Phenols > Alcohols > Ammonium ions > Alkynes

b. Bases (Strong → Weak):
Alkoxides > Hydroxide > Amines ≈ Ammonia > Carboxylates > Aromatic amines

4. Periodic Trends & Electronegativity:

i. Electronegativity and Acid/Base Strength:

Electronegativity (EN) refers to the atom’s capacity to pull electrons in a bond.

A. Effect on Acidity:

* The more electronegative the atom that holds the negative charge in the conjugate base, the more stable.
* More stable conjugate base → stronger acid.

Example (across a period):

CH₄ < NH₃ < H₂O < HF

* H–F is the strongest acid since F is the most electronegative.

B. Influence on Basicity:

* Bases with less electronegative atoms are more ready to share electrons (donate a lone pair).
* Thus, they are stronger bases.

Example:

$$\text{CH₃⁻}>\text{NH₂⁻}>\text{OH⁻}>\text{F⁻}$$

* CH₃⁻ is stronger as a base since C is less electronegative than N, O, or F.

ii. Periodic Trends That Affect Acidity and Basicity:

 Across a Period (left → right):

Example (acids):

* H–C < H–N < H–O < H–F → increases in acidity

iii. Down a Group (top → bottom):

Example (Group 17 acids):

$$\text{HF}<\text{HCl}<\text{HBr}<\text{HI}$$

* Although F is most electronegative, HI is the strongest acid because H–I bond is weaker → easier to donate H⁺.

iii. Bond Strength vs Electronegativity:

* For binary acids (HX):

* Down a group: Acid strength increases because H–X bond becomes weaker.
* Across a period: Acid strength increases because of increasing electronegativity of X.

4. Summary Table: