Bronsted-Lowry and Lewis Acid-Base Theory
Brønsted-Lowry Theory & Conjugate Pairs
Core Definition • Acid: A proton (H⁺) donor • Base: A proton (H⁺) acceptor
Proton Transfer Visualization
Key Principles
- Conjugate pairs differ by exactly one proton (HA ⇌ H⁺ + A⁻)
- Every acid produces a conjugate base; every base produces a conjugate acid
- Inverse Strength: A strong acid yields a weak conjugate base
- Equilibrium position is determined by the relative strength of species
Example: HA + H₂O ⇌ H₃O⁺ + A⁻ 1. Acid (HA) donates H⁺ to Water. 2. Water acts as a Base, accepting the H⁺. 3. Products are the Conjugate Base (A⁻) and Conjugate Acid (H₃O⁺).
Strong Acids & Complete Dissociation
Relative Strength Scale
Key Characteristics
- Strong acids completely dissociate in aqueous solution (Kₐ >> 1)
- Assume 100% ionization; equilibrium lies far to the right
- Conjugate bases of strong acids are extremely weak (negligible basicity)
Common Strong Acids HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄ Examples of Conjugate Bases Cl⁻, Br⁻, I⁻, NO₃⁻, ClO₄⁻, HSO₄⁻ (do not accept protons significantly)
Weak Acids and Partial Dissociation
Core Definition • Weak acids only partially dissociate • Ka is small but measurable (Ka < 1) • Equilibrium lies to the left (reactants)
Equilibrium & Examples
HA + H₂O ⇌ H₃O⁺ + A⁻
Common Weak Acids Acetic AcidCH₃COOH Formic AcidHCOOH PhenolC₆H₅OH
Conjugate Relationship Inverse Strength:Weak acids have relatively strong conjugate bases. • Acetate (CH₃COO⁻) and Formate (HCOO⁻) are weak bases. • Significant basicity compared to conjugates of strong acids.
Ka Expression and Equilibrium Constant
Equilibrium Shift & Acid Strength
Larger Ka ➔ Favors Products
Smaller Ka ➔ Favors Reactants
Equilibrium balance directly correlates with acid dissociation
The Ka Expression
- Ka = [H₃O⁺][A⁻] / [HA] at equilibrium
- Water concentration is omitted
- Quantifies proton donation tendency
Acid Strength & Magnitude Spans many orders of magnitude. Strong Acids: Ka > 1 Weak Acids: Ka < 1 Units: M (monoprotic)
Henderson-Hasselbalch Equation and pH Calculation
pH = pKa + log([A⁻]/[HA]) Relates pH to the acid dissociation constant (pKa) and the concentration ratio of conjugate base to acid.
Concentration Ratios & pH
- [A⁻] = [HA]pH = pKa (Half-equivalence point)
- [A⁻] > [HA]pH > pKa (More basic solution)
- [A⁻] < [HA]pH < pKa (More acidic solution)
Buffer Properties & Example Buffer Capacity: Maximum when pH = pKa; resists pH change most effectively. Acetic Acid Example (pKa = 4.74) If [CH₃COO⁻]/[CH₃COOH] = 10: pH = 4.74 + log(10) = 5.74
pKa Scale and Acid Strength Ranking
pKa = -log(Ka) Inverse logarithmic relationship: lower pKa signifies a stronger acid.
Comparative pKa Values Formic acid: 3.74 Acetic acid: 4.74 Phenol: 10.0
General Rule
- Strong Acids: pKa < 0
- Weak Acids: pKa > 0
The Half-Equivalence Point At pH = pKa, the concentrations of acid [HA] and conjugate base [A-] are equimolar ([HA] = [A-]).
Conjugate Base Stability: Electronegativity & Induction
Stronger Acid = More Stable Conjugate Base Stability is dictated by the atom's ability to accommodate negative charge.
Key Factors
- Electronegativity:Atoms bonded to the acidic proton pull density towards themselves.
- Electron-Withdrawing:High EN groups (F, Cl, O, N) stabilize charge by dispersion.
- Induction:C and H donate density, while halogens/electronegative atoms withdraw it.
Comparative Example Trifluoroacetic Acid (CF₃COOH) pKa = 0.2 Acetic Acid (CH₃COOH) pKa = 4.74 The 3 fluorine atoms in CF₃COOH exert a massive inductive effect, greatly increasing acid strength.
Subtitle
Subtitle
Polarizability & Size Effects on Stability
Larger Conjugate Bases = Better Stability Polarizability increases down a group, reducing charge density across a larger electron cloud.
The Haloacid Series HF < HCl < HBr < HI Fluoride (F⁻): Smallest, least polarizable ion. Iodide (I⁻): Largest, most polarizable ion.
Key Trends
- Charge Dispersion: Larger anions distribute charge over a greater volume, reducing density.
- Periodic Trend: Polarizability increases down a group in the periodic table.
- Acid Strength: Increases with conjugate base size in the haloacid series.
Subtitle
Conjugate Base Stability: Resonance and Hybridization
Stability is enhanced by charge delocalization and s-character.
1. Resonance Effects Delocalization:Spreading negative charge over multiple atoms stabilizes the base. Carboxylic Acids vs. Alcohols:RCOOH are stronger acids than ROH because RCOO⁻ has two resonance forms.
2. Hybridization Effects S-Character:Higher s-character (sp > sp² > sp³) brings electrons closer to the nucleus. Comparative Example:Acetylene (pKa 25, sp) is much stronger than Ethane (pKa > 50, sp³).
Subtitle
Subtitle
Equilibrium Position & Proton Transfer
Proton transfer proceeds from stronger acid to stronger base. Equilibrium always favors the formation of the weaker acid and weaker base.
Directional Rules
- Compare pKa values: transfers from lower pKa to higher pKa species.
- If pKa(HA) < pKa(HB): HA is the stronger acid; equilibrium favors B⁻.
- Equilibrium Constant (K):K = 10^(pKa(HB) - pKa(HA))
Practical Example CH₃COOH + NH₃ ⇌ CH₃COO⁻ + NH₄⁺ pKa Analysis: • CH₃COOH: pKa = 4.74 • NH₄⁺: pKa = 9.25 Outcome: Favors Products Since 4.74 < 9.25, the reaction proceeds forward.
Polyprotic Acids and Sequential Dissociation
Polyprotic acids donate multiple protons sequentially; each dissociation has its own pKa value.
Key Principles Dissociation Trend: pKa1 < pKa2 < pKa3 Why it happens: Removing a proton from a neutral molecule is significantly easier than from a negatively charged species.
Example: Phosphoric Acid (H₃PO₄) pKa Values:
- pKa1 = 2.12 (H₃PO₄)
- pKa2 = 7.21 (H₂PO₄⁻)
- pKa3 = 12.67 (HPO₄²⁻)
Dominant Species: H₃PO₄, H₂PO₄⁻, HPO₄²⁻, PO₄³⁻ dominate at varying pH levels.
Summary: Predicting Acid Strength & Reaction Outcomes
1. Acid Strength Foundations Depends on conjugate base stability via electronegativity, resonance, and hybridization.
2. Equilibrium Prediction Lower pKa acids transfer protons to higher pKa bases; favors weaker species.
3. Buffer Dynamics Henderson-Hasselbalch equation relates pH, pKa, and [Base]/[Acid] ratios.
4. Strong vs. Weak Acids Strong acids (pKa < 0) dissociate fully. Weak acids (pKa > 0) reach equilibrium.
5. Polyprotic Dissociation Protons are donated sequentially with decreasing Ka (increasing pKa) values.
- Example: H₃PO₄ (Phosphoric Acid)
- pKa1 < pKa2 < pKa3
Subtitle
Lewis Acid-Base Theory Fundamentals
Lewis theory focuses on electron pair transfer, expanding the definition of acid-base reactions beyond proton exchange.
Lewis Acid An electron pair acceptor that seeks electrons from another species.
Lewis Base An electron pair donor that provides electrons to another species.
The Result Formation of a coordinate covalent bond.
Theoretical Scope Explains reactions without hydrogen not accounted for by Brønsted-Lowry theory.
Nucleophiles and Electrophiles
Electrophiles (E⁺) A Lewis acid that attracts electron-rich sites, seeking electrons.
- Positively charged (cations)
- Contain empty orbitals
- Electropositive carbon atom in a polar covalent bond (C-X, C=O)
Nucleophiles (Nu⁻) A Lewis base that attacks electron-poor sites, seeking positive charge.
- Typically negatively charged (anions)
- Contain lone pairs of electrons
Chemical Context The terms nucleophile and electrophile are used to describe reactivity specifically within organic chemistry mechanistic pathways.
Examples of Lewis Acid-Base Reactions
Adduct Formation Boron trifluoride (BF₃) accepts an electron pair from ammonia (NH₃) to form a BF₃-NH₃ complex.
Industrial Catalysis Aluminum chloride (AlCl₃) acts as a Lewis acid in Friedel-Crafts reactions with aromatic compounds.
Organic Electrophiles Carbocations (R₃C⁺) are strong electrophiles that accept electron pairs from nucleophiles like alkenes.
Cyanide Nucleophilicity Cyanide ion (CN⁻) is a strong nucleophile that donates electrons to electrophilic carbon atoms.
Metal Coordination Metal ions like Ag⁺ and Zn²⁺ function as Lewis acids, coordinating with electron-rich ligands.
Comparing Chemical Frameworks
Brønsted-Lowry Specific cases of Lewis interactions involving protons (H⁺). All Brønsted-Lowry acids are Lewis acids, but the reverse is not always true.
Lewis Theory The broadest framework. It encompasses all electron pair transfers, providing a universal view of chemical interactions.
Nu- / E+ Nucleophile-electrophile terminology emphasizes reactivity and mechanistic pathways, primarily within organic reactions.
Core Interconnectivity
- Every Brønsted-Lowry acid acts as a Lewis acid by accepting an electron pair.
- Lewis theory is essential for explaining reactions that do not involve hydrogen.
Acid-Base - university
Kim Conlon
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Transcript
Bronsted-Lowry and Lewis Acid-Base Theory
Brønsted-Lowry Theory & Conjugate Pairs
Core Definition • Acid: A proton (H⁺) donor • Base: A proton (H⁺) acceptor
Proton Transfer Visualization
Key Principles
Example: HA + H₂O ⇌ H₃O⁺ + A⁻ 1. Acid (HA) donates H⁺ to Water. 2. Water acts as a Base, accepting the H⁺. 3. Products are the Conjugate Base (A⁻) and Conjugate Acid (H₃O⁺).
Strong Acids & Complete Dissociation
Relative Strength Scale
Key Characteristics
- Strong acids completely dissociate in aqueous solution (Kₐ >> 1)
- Assume 100% ionization; equilibrium lies far to the right
- Conjugate bases of strong acids are extremely weak (negligible basicity)
Common Strong Acids HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄ Examples of Conjugate Bases Cl⁻, Br⁻, I⁻, NO₃⁻, ClO₄⁻, HSO₄⁻ (do not accept protons significantly)Weak Acids and Partial Dissociation
Core Definition • Weak acids only partially dissociate • Ka is small but measurable (Ka < 1) • Equilibrium lies to the left (reactants)
Equilibrium & Examples
HA + H₂O ⇌ H₃O⁺ + A⁻
Common Weak Acids Acetic AcidCH₃COOH Formic AcidHCOOH PhenolC₆H₅OH
Conjugate Relationship Inverse Strength:Weak acids have relatively strong conjugate bases. • Acetate (CH₃COO⁻) and Formate (HCOO⁻) are weak bases. • Significant basicity compared to conjugates of strong acids.
Ka Expression and Equilibrium Constant
Equilibrium Shift & Acid Strength
Larger Ka ➔ Favors Products
Smaller Ka ➔ Favors Reactants
Equilibrium balance directly correlates with acid dissociation
The Ka Expression
Acid Strength & Magnitude Spans many orders of magnitude. Strong Acids: Ka > 1 Weak Acids: Ka < 1 Units: M (monoprotic)
Henderson-Hasselbalch Equation and pH Calculation
pH = pKa + log([A⁻]/[HA]) Relates pH to the acid dissociation constant (pKa) and the concentration ratio of conjugate base to acid.
Concentration Ratios & pH
Buffer Properties & Example Buffer Capacity: Maximum when pH = pKa; resists pH change most effectively. Acetic Acid Example (pKa = 4.74) If [CH₃COO⁻]/[CH₃COOH] = 10: pH = 4.74 + log(10) = 5.74
pKa Scale and Acid Strength Ranking
pKa = -log(Ka) Inverse logarithmic relationship: lower pKa signifies a stronger acid.
Comparative pKa Values Formic acid: 3.74 Acetic acid: 4.74 Phenol: 10.0
General Rule
The Half-Equivalence Point At pH = pKa, the concentrations of acid [HA] and conjugate base [A-] are equimolar ([HA] = [A-]).
Conjugate Base Stability: Electronegativity & Induction
Stronger Acid = More Stable Conjugate Base Stability is dictated by the atom's ability to accommodate negative charge.
Key Factors
Comparative Example Trifluoroacetic Acid (CF₃COOH) pKa = 0.2 Acetic Acid (CH₃COOH) pKa = 4.74 The 3 fluorine atoms in CF₃COOH exert a massive inductive effect, greatly increasing acid strength.
Subtitle
Subtitle
Polarizability & Size Effects on Stability
Larger Conjugate Bases = Better Stability Polarizability increases down a group, reducing charge density across a larger electron cloud.
The Haloacid Series HF < HCl < HBr < HI Fluoride (F⁻): Smallest, least polarizable ion. Iodide (I⁻): Largest, most polarizable ion.
Key Trends
Subtitle
Conjugate Base Stability: Resonance and Hybridization
Stability is enhanced by charge delocalization and s-character.
1. Resonance Effects Delocalization:Spreading negative charge over multiple atoms stabilizes the base. Carboxylic Acids vs. Alcohols:RCOOH are stronger acids than ROH because RCOO⁻ has two resonance forms.
2. Hybridization Effects S-Character:Higher s-character (sp > sp² > sp³) brings electrons closer to the nucleus. Comparative Example:Acetylene (pKa 25, sp) is much stronger than Ethane (pKa > 50, sp³).
Subtitle
Subtitle
Equilibrium Position & Proton Transfer
Proton transfer proceeds from stronger acid to stronger base. Equilibrium always favors the formation of the weaker acid and weaker base.
Directional Rules
Practical Example CH₃COOH + NH₃ ⇌ CH₃COO⁻ + NH₄⁺ pKa Analysis: • CH₃COOH: pKa = 4.74 • NH₄⁺: pKa = 9.25 Outcome: Favors Products Since 4.74 < 9.25, the reaction proceeds forward.
Polyprotic Acids and Sequential Dissociation
Polyprotic acids donate multiple protons sequentially; each dissociation has its own pKa value.
Key Principles Dissociation Trend: pKa1 < pKa2 < pKa3 Why it happens: Removing a proton from a neutral molecule is significantly easier than from a negatively charged species.
Example: Phosphoric Acid (H₃PO₄) pKa Values:
- pKa1 = 2.12 (H₃PO₄)
- pKa2 = 7.21 (H₂PO₄⁻)
- pKa3 = 12.67 (HPO₄²⁻)
Dominant Species: H₃PO₄, H₂PO₄⁻, HPO₄²⁻, PO₄³⁻ dominate at varying pH levels.Summary: Predicting Acid Strength & Reaction Outcomes
1. Acid Strength Foundations Depends on conjugate base stability via electronegativity, resonance, and hybridization.
2. Equilibrium Prediction Lower pKa acids transfer protons to higher pKa bases; favors weaker species.
3. Buffer Dynamics Henderson-Hasselbalch equation relates pH, pKa, and [Base]/[Acid] ratios.
4. Strong vs. Weak Acids Strong acids (pKa < 0) dissociate fully. Weak acids (pKa > 0) reach equilibrium.
5. Polyprotic Dissociation Protons are donated sequentially with decreasing Ka (increasing pKa) values.
Subtitle
Lewis Acid-Base Theory Fundamentals
Lewis theory focuses on electron pair transfer, expanding the definition of acid-base reactions beyond proton exchange.
Lewis Acid An electron pair acceptor that seeks electrons from another species.
Lewis Base An electron pair donor that provides electrons to another species.
The Result Formation of a coordinate covalent bond.
Theoretical Scope Explains reactions without hydrogen not accounted for by Brønsted-Lowry theory.
Nucleophiles and Electrophiles
Electrophiles (E⁺) A Lewis acid that attracts electron-rich sites, seeking electrons.
Nucleophiles (Nu⁻) A Lewis base that attacks electron-poor sites, seeking positive charge.
Chemical Context The terms nucleophile and electrophile are used to describe reactivity specifically within organic chemistry mechanistic pathways.
Examples of Lewis Acid-Base Reactions
Adduct Formation Boron trifluoride (BF₃) accepts an electron pair from ammonia (NH₃) to form a BF₃-NH₃ complex.
Industrial Catalysis Aluminum chloride (AlCl₃) acts as a Lewis acid in Friedel-Crafts reactions with aromatic compounds.
Organic Electrophiles Carbocations (R₃C⁺) are strong electrophiles that accept electron pairs from nucleophiles like alkenes.
Cyanide Nucleophilicity Cyanide ion (CN⁻) is a strong nucleophile that donates electrons to electrophilic carbon atoms.
Metal Coordination Metal ions like Ag⁺ and Zn²⁺ function as Lewis acids, coordinating with electron-rich ligands.
Comparing Chemical Frameworks
Brønsted-Lowry Specific cases of Lewis interactions involving protons (H⁺). All Brønsted-Lowry acids are Lewis acids, but the reverse is not always true.
Lewis Theory The broadest framework. It encompasses all electron pair transfers, providing a universal view of chemical interactions.
Nu- / E+ Nucleophile-electrophile terminology emphasizes reactivity and mechanistic pathways, primarily within organic reactions.
Core Interconnectivity