The Hook: Why Does Lemon Fight Stomach Acidity?
Sounds backwards, right? Lemon juice is acidic (pH ~2), yet drinking lemon water can help with acid reflux! The secret: Once metabolized, citric acid produces alkaline byproducts.
Your stomach acid (HCl), vinegar (CH₃COOH), and battery acid (H₂SO₄) are all acids—but they work differently. Some donate protons, some accept electrons. Understanding the three theories of acids and bases explains why baking soda neutralizes both stomach acid AND bee stings!
Real-world examples:
- Blood pH regulation (7.35-7.45) uses acid-base equilibrium
- Antacids neutralize HCl in your stomach
- Soap works because it’s basic (pH ~9-10)
JEE Reality: Acid-base questions appear in every JEE paper—understanding all three theories gives you flexibility to tackle any problem. Expect 2-3 direct questions plus application in buffers, salts, and organic chemistry.
The Core Concept
There are three major theories defining acids and bases, each broader than the previous:
- Arrhenius Theory (1884) - Works only in water
- Bronsted-Lowry Theory (1923) - Works in any solvent
- Lewis Theory (1923) - Broadest definition, includes non-proton transfers
Evolution: Each theory expanded the definition to include more substances.
Arrhenius Theory (The Classical View)
Definitions
Proposed by Svante Arrhenius in 1884:
Acid: Substance that produces H⁺ ions in aqueous solution
$$HCl \xrightarrow{H_2O} H^+ + Cl^-$$Base: Substance that produces OH⁻ ions in aqueous solution
$$NaOH \xrightarrow{H_2O} Na^+ + OH^-$$Neutralization: Combination of H⁺ and OH⁻ to form water
$$H^+ + OH^- \rightarrow H_2O$$Examples
| Arrhenius Acid | Ionization |
|---|---|
| HCl | $HCl \rightarrow H^+ + Cl^-$ |
| HNO₃ | $HNO_3 \rightarrow H^+ + NO_3^-$ |
| H₂SO₄ | $H_2SO_4 \rightarrow 2H^+ + SO_4^{2-}$ |
| Arrhenius Base | Ionization |
|---|---|
| NaOH | $NaOH \rightarrow Na^+ + OH^-$ |
| KOH | $KOH \rightarrow K^+ + OH^-$ |
| Ca(OH)₂ | $Ca(OH)_2 \rightarrow Ca^{2+} + 2OH^-$ |
Limitations
- Only works in aqueous solutions - Can’t explain HCl gas reacting with NH₃ gas
- Requires H⁺ or OH⁻ - Can’t explain NH₃ being a base (no OH⁻ in its formula!)
- Can’t explain neutralization without water - Like $NH_3 + HCl \rightarrow NH_4Cl$ (no water involved)
- Limited to specific ions - Doesn’t cover broader acid-base behavior
JEE Tip: If a question involves water and simple acids/bases, Arrhenius is sufficient. For anything else, use Bronsted-Lowry or Lewis.
Bronsted-Lowry Theory (The Proton Transfer View)
Definitions
Proposed independently by Johannes Bronsted and Thomas Lowry in 1923:
Acid: Proton (H⁺) donor Base: Proton (H⁺) acceptor
Key Innovation: Acid-base reactions are proton transfer reactions.
$$HA + B \rightleftharpoons A^- + BH^+$$- HA donates H⁺ → Acid
- B accepts H⁺ → Base
Important Features
- No need for water - Works in any solvent or even gas phase
- Relative definitions - Same substance can be acid or base depending on partner
- Conjugate pairs - Every acid has a conjugate base, every base has a conjugate acid
Conjugate Acid-Base Pairs
When an acid donates H⁺, it becomes its conjugate base:
$$\underbrace{HA}_{acid} \rightleftharpoons \underbrace{A^-}_{conjugate\ base} + H^+$$When a base accepts H⁺, it becomes its conjugate acid:
$$\underbrace{B}_{base} + H^+ \rightleftharpoons \underbrace{BH^+}_{conjugate\ acid}$$Complete reaction:
$$\underbrace{HA}_{acid\ 1} + \underbrace{B}_{base\ 2} \rightleftharpoons \underbrace{A^-}_{base\ 1} + \underbrace{BH^+}_{acid\ 2}$$Two conjugate pairs:
- HA / A⁻ (differ by one H⁺)
- B / BH⁺ (differ by one H⁺)
Examples of Conjugate Pairs
| Acid | Conjugate Base | Base | Conjugate Acid |
|---|---|---|---|
| HCl | Cl⁻ | NH₃ | NH₄⁺ |
| H₂SO₄ | HSO₄⁻ | H₂O | H₃O⁺ |
| CH₃COOH | CH₃COO⁻ | OH⁻ | H₂O |
| NH₄⁺ | NH₃ | CO₃²⁻ | HCO₃⁻ |
| H₃O⁺ | H₂O | SO₄²⁻ | HSO₄⁻ |
Pattern: Remove H⁺ from acid → Get conjugate base
Example Reaction Analysis
$$HCl + H_2O \rightleftharpoons H_3O^+ + Cl^-$$Identify:
- HCl → H₃O⁺: HCl donates H⁺ → HCl is acid
- H₂O → H₃O⁺: H₂O accepts H⁺ → H₂O is base
Conjugate pairs:
- HCl (acid) / Cl⁻ (conjugate base)
- H₂O (base) / H₃O⁺ (conjugate acid)
Amphiprotic (Amphoteric) Species
Amphiprotic substances can act as both acid and base depending on the partner.
Examples: H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻, HS⁻
Water as Amphiprotic
Acting as base:
$$HCl + H_2O \rightarrow H_3O^+ + Cl^-$$(H₂O accepts H⁺)
Acting as acid:
$$NH_3 + H_2O \rightarrow NH_4^+ + OH^-$$(H₂O donates H⁺)
HCO₃⁻ as Amphiprotic
Acting as base:
$$HCO_3^- + H^+ \rightarrow H_2CO_3$$(accepts H⁺)
Acting as acid:
$$HCO_3^- \rightarrow H^+ + CO_3^{2-}$$(donates H⁺)
How to identify amphiprotic species:
- Can it donate H⁺? (Has H in formula)
- Can it accept H⁺? (Has lone pairs or negative charge)
If YES to both → Amphiprotic
Common amphiprotic species for JEE:
- H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻, HPO₄²⁻, HS⁻, HSO₃⁻
NOT amphiprotic:
- Cl⁻ (can’t donate H⁺, no H)
- NH₄⁺ (can’t accept H⁺, no lone pairs after accepting one)
- SO₄²⁻ (can’t donate H⁺, no H)
Lewis Theory (The Electron Pair View)
Definitions
Proposed by Gilbert N. Lewis in 1923:
Acid: Electron pair acceptor (electrophile) Base: Electron pair donor (nucleophile)
Key Innovation: No protons needed! Focuses on electron pair transfer.
How It Works
$$\text{Base} \xrightarrow{\text{donates}} \text{electron pair} \xleftarrow{\text{accepts}} \text{Acid}$$Result: Formation of a coordinate covalent bond (dative bond)
Examples
Example 1: BF₃ + NH₃
$$BF_3 + :NH_3 \rightarrow F_3B \leftarrow NH_3$$- Lewis acid: BF₃ (accepts electron pair, electron deficient)
- Lewis base: NH₃ (donates lone pair)
Example 2: H⁺ + OH⁻
$$H^+ + :OH^- \rightarrow H-OH$$- Lewis acid: H⁺ (accepts electrons)
- Lewis base: OH⁻ (donates electrons)
Example 3: AlCl₃ + Cl⁻
$$AlCl_3 + :Cl^- \rightarrow [AlCl_4]^-$$- Lewis acid: AlCl₃ (Al has empty orbital)
- Lewis base: Cl⁻ (has lone pairs)
Identifying Lewis Acids
Lewis acids are electron deficient:
- Cations: H⁺, Ag⁺, Fe³⁺, Al³⁺
- Molecules with incomplete octet: BF₃, BCl₃, AlCl₃
- Molecules with empty d-orbitals: SiF₄, SnCl₄
- Molecules with multiple bonds to electronegative atoms: CO₂, SO₂
Identifying Lewis Bases
Lewis bases have lone pairs:
- Anions: OH⁻, Cl⁻, CN⁻, CO₃²⁻
- Molecules with lone pairs: NH₃, H₂O, amines, alcohols
- π-bond systems: Alkenes, alkynes (can donate π-electrons)
Memory Tricks & Patterns
Mnemonic for Three Theories
“Arrhenius Always Hydrates, Bronsted Bounces Protons, Lewis Loves Electrons”
- Arrhenius Always Hydrates → H⁺/OH⁻ in water
- Bronsted Bounces Protons → Proton transfer
- Lewis Loves Electrons → Electron pair transfer
Comparing the Theories
| Feature | Arrhenius | Bronsted-Lowry | Lewis |
|---|---|---|---|
| Acid definition | Produces H⁺ | Donates H⁺ | Accepts e⁻ pair |
| Base definition | Produces OH⁻ | Accepts H⁺ | Donates e⁻ pair |
| Requires water? | Yes | No | No |
| Scope | Narrowest | Medium | Broadest |
| Example acid | HCl → H⁺ | HCl → H⁺ | BF₃, H⁺, Al³⁺ |
| Example base | NaOH → OH⁻ | NH₃ | NH₃, Cl⁻ |
Pattern: Conjugate Strength
Key Rule: Stronger the acid → Weaker its conjugate base
| Strong Acid | Weak Conjugate Base | Why? |
|---|---|---|
| HCl (very strong) | Cl⁻ (very weak) | Cl⁻ has no tendency to accept H⁺ back |
| CH₃COOH (weak) | CH₃COO⁻ (fairly strong) | CH₃COO⁻ wants H⁺ back |
Inverse relationship:
- Strong acid ↔ Weak conjugate base
- Weak acid ↔ Strong conjugate base
Same for bases:
- Strong base ↔ Weak conjugate acid
- Weak base ↔ Strong conjugate acid
Interactive Demo: Visualize Titration Curves
Explore how acids and bases neutralize each other through titration.
Relative Strength of Acids and Bases
Acid Strength Order
Strong acids (complete ionization):
$$HClO_4 > HI > HBr > HCl > H_2SO_4 > HNO_3$$Weak acids (partial ionization, decreasing strength):
$$H_3O^+ > H_2SO_3 > H_3PO_4 > HF > CH_3COOH > H_2CO_3 > H_2S > NH_4^+ > HCN$$Base Strength Order
Strong bases:
$$O^{2-} > H^- > NH_2^- > OH^- > \text{Group 1 hydroxides}$$Weak bases:
$$NH_3 > CH_3COO^- > CO_3^{2-} > F^- > Cl^-$$Factors Affecting Acid Strength
Electronegativity of atom bonded to H
- More electronegative → Weaker H-X bond → Stronger acid
- Order: HF > H₂O > NH₃ > CH₄
Size of atom bonded to H (within same group)
- Larger atom → Weaker H-X bond → Stronger acid
- Order: HI > HBr > HCl > HF
Oxidation state (for oxoacids H-O-X)
- Higher oxidation state of X → Stronger acid
- Order: HClO₄ > HClO₃ > HClO₂ > HClO
Resonance stabilization of conjugate base
- More resonance → More stable A⁻ → Stronger acid HA
- CH₃COOH is stronger than CH₃OH (acetate ion has resonance)
Common Mistakes to Avoid
Wrong: “NH₃ has no OH⁻, so it can’t be a base”
Correct:
- Arrhenius: NH₃ is NOT a base (no OH⁻ in formula)
- Bronsted-Lowry: NH₃ IS a base (accepts H⁺)
- Lewis: NH₃ IS a base (donates electron pair)
Lesson: Use the appropriate theory for the context!
Common Error: Mixing up different pairs
Correct Method:
- Identify what loses H⁺ (acid) and what gains H⁺ (base)
- Acid and its conjugate base differ by exactly ONE H⁺
- Base and its conjugate acid differ by exactly ONE H⁺
Example: $H_2PO_4^- + H_2O \rightleftharpoons H_3O^+ + HPO_4^{2-}$
Pairs:
- H₂PO₄⁻ / HPO₄²⁻ (differ by 1 H⁺) ✓
- H₂O / H₃O⁺ (differ by 1 H⁺) ✓
NOT pairs:
- H₂PO₄⁻ / H₃O⁺ (don’t differ by just H⁺) ✗
Wrong: “All acids are proton donors”
Correct:
- Bronsted acids: Proton donors
- Lewis acids: Electron pair acceptors (may or may not involve protons)
Example: BF₃ is a Lewis acid but NOT a Bronsted acid (no H⁺ to donate)
Remember: Lewis theory is broader—includes all Bronsted acids plus more!
Technically:
- Amphiprotic: Can donate or accept protons (Bronsted-Lowry concept)
- Amphoteric: Can act as acid or base in Arrhenius/Lewis sense too
JEE usage: Terms are often used interchangeably, but strictly:
- All amphiprotic substances are amphoteric
- Not all amphoteric substances are amphiprotic (e.g., ZnO acts as base/acid but no H⁺ transfer)
Safe bet: Use “amphiprotic” for species that can donate/accept H⁺
Practice Problems
Level 1: Foundation (NCERT)
In the reaction: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$
Identify:
- The two acids
- The two bases
- The conjugate acid-base pairs
Solution:
Bronsted-Lowry analysis:
- NH₃ + H⁺ → NH₄⁺: NH₃ accepts H⁺ → Base
- H₂O → H⁺ + OH⁻: H₂O donates H⁺ → Acid
Answers:
- Acids: H₂O, NH₄⁺
- Bases: NH₃, OH⁻
- Conjugate pairs:
- H₂O (acid) / OH⁻ (conjugate base)
- NH₃ (base) / NH₄⁺ (conjugate acid)
Which of the following are amphiprotic?
- H₂O
- Cl⁻
- HCO₃⁻
- NH₄⁺
- HPO₄²⁻
Solution:
Amphiprotic = Can donate AND accept H⁺
H₂O - YES
- As acid: H₂O → H⁺ + OH⁻
- As base: H₂O + H⁺ → H₃O⁺
Cl⁻ - NO
- Has no H to donate
HCO₃⁻ - YES
- As acid: HCO₃⁻ → H⁺ + CO₃²⁻
- As base: HCO₃⁻ + H⁺ → H₂CO₃
NH₄⁺ - NO
- Can donate H⁺: NH₄⁺ → NH₃ + H⁺
- Cannot accept H⁺ (no lone pair available)
HPO₄²⁻ - YES
- As acid: HPO₄²⁻ → H⁺ + PO₄³⁻
- As base: HPO₄²⁻ + H⁺ → H₂PO₄⁻
Amphiprotic: H₂O, HCO₃⁻, HPO₄²⁻
Identify the Lewis acid and Lewis base:
$$BF_3 + NH_3 \rightarrow F_3B-NH_3$$Solution:
Lewis acid = Electron pair acceptor Lewis base = Electron pair donor
- BF₃: Boron has only 6 electrons (incomplete octet), can accept electrons → Lewis acid
- NH₃: Nitrogen has a lone pair, can donate electrons → Lewis base
Result: NH₃ donates its lone pair to BF₃, forming coordinate bond
Level 2: JEE Main
Which theory best explains each observation?
- HCl gas reacts with NH₃ gas to form NH₄Cl (no water involved)
- NaOH neutralizes HCl in aqueous solution
- BF₃ accepts a fluoride ion to form BF₄⁻
Solution:
HCl + NH₃ → NH₄Cl
- Bronsted-Lowry (or Lewis)
- Arrhenius fails (no water, no H⁺/OH⁻ ions)
- HCl donates H⁺ to NH₃
NaOH + HCl → NaCl + H₂O
- Arrhenius (simplest)
- H⁺ + OH⁻ → H₂O
- (Bronsted-Lowry and Lewis also work, but Arrhenius is most direct)
BF₃ + F⁻ → BF₄⁻
- Lewis only
- No proton transfer
- BF₃ accepts electron pair from F⁻
Arrange in order of increasing base strength: Cl⁻, CH₃COO⁻, F⁻, OH⁻
Solution:
Strategy: Identify conjugate acids and their strengths
| Base | Conjugate Acid | Acid Strength |
|---|---|---|
| Cl⁻ | HCl | Very strong |
| CH₃COO⁻ | CH₃COOH | Weak |
| F⁻ | HF | Weak (but stronger than CH₃COOH) |
| OH⁻ | H₂O | Very weak |
Rule: Stronger the conjugate acid → Weaker the base
Order of acids (strong → weak): HCl > HF > CH₃COOH > H₂O
Order of bases (weak → strong): Cl⁻ < F⁻ < CH₃COO⁻ < OH⁻
Answer: Cl⁻ < F⁻ < CH₃COO⁻ < OH⁻
Consider the reactions:
- $HCO_3^- + H_2O \rightarrow H_2CO_3 + OH^-$
- $HCO_3^- + H_2O \rightarrow CO_3^{2-} + H_3O^+$
What is the role of HCO₃⁻ and H₂O in each?
Solution:
Reaction 1:
- HCO₃⁻ + H⁺ → H₂CO₃: HCO₃⁻ accepts H⁺ → Base
- H₂O → H⁺ + OH⁻: H₂O donates H⁺ → Acid
Reaction 2:
- HCO₃⁻ → H⁺ + CO₃²⁻: HCO₃⁻ donates H⁺ → Acid
- H₂O + H⁺ → H₃O⁺: H₂O accepts H⁺ → Base
Conclusion:
- HCO₃⁻ is amphiprotic (acts as both acid and base)
- H₂O is amphiprotic (acts as both acid and base)
Which reaction dominates depends on solution conditions!
Level 3: JEE Advanced
Explain using Lewis theory:
$$AlCl_3 + Cl^- \rightarrow [AlCl_4]^-$$ $$2AlCl_3 \rightleftharpoons Al_2Cl_6$$Solution:
Reaction 1: AlCl₃ + Cl⁻ → [AlCl₄]⁻
Lewis acid-base:
- AlCl₃: Al has only 6 electrons in valence shell (incomplete octet) → Lewis acid
- Cl⁻: Has lone pairs → Lewis base
- Cl⁻ donates electron pair to Al, forming coordinate bond
- Result: Tetrahedral [AlCl₄]⁻
Reaction 2: 2AlCl₃ ⇌ Al₂Cl₆
Lewis acid-base (self-reaction):
- Each AlCl₃ has one Al (Lewis acid) and Cl atoms with lone pairs (Lewis base)
- Cl from one AlCl₃ donates lone pair to Al of another AlCl₃
- Forms dimer with two coordinate bonds (bridging chlorines)
Structure of Al₂Cl₆:
Cl Cl
\ /
Cl—Al Al—Cl
/ \
Cl Cl
JEE Insight: This dimerization explains why AlCl₃ vapor has lower molar mass than expected (exists as Al₂Cl₆)
Arrange in order of increasing acid strength: HClO, HClO₂, HClO₃, HClO₄
Solution:
Rule for oxoacids (HO-X-O): More oxygen atoms on X → Higher oxidation state of X → Stronger acid
Why? More O atoms → More electron withdrawal from O-H bond → Easier to lose H⁺
Oxidation states of Cl:
- HClO: Cl is +1
- HClO₂: Cl is +3
- HClO₃: Cl is +5
- HClO₄: Cl is +7
Order (increasing acid strength): HClO < HClO₂ < HClO₃ < HClO₄
Actual Ka values (approximate):
- HClO: 3×10⁻⁸ (very weak)
- HClO₂: 1×10⁻² (weak)
- HClO₃: 10³ (strong)
- HClO₄: 10⁸ (very strong)
JEE Pattern: Similar trend for HNO₂ < HNO₃, H₂SO₃ < H₂SO₄
Blood pH is maintained around 7.4 by the equilibrium:
$$CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons H^+ + HCO_3^-$$Questions:
- What happens if CO₂ increases (hyperventilation)?
- Identify the amphiprotic species.
- Why is this system effective as a buffer?
Solution:
1. Effect of increased CO₂:
More CO₂ → Equilibrium shifts right (Le Chatelier) → More H₂CO₃ forms → More H⁺ forms → pH decreases (blood becomes more acidic)
This is why hyperventilation (rapid breathing expelling CO₂) increases pH (alkalosis) Slow breathing (retaining CO₂) decreases pH (acidosis)
2. Amphiprotic species:
- H₂CO₃ can donate H⁺ (→ HCO₃⁻) or accept H⁺ (from H₂O)
- HCO₃⁻ can donate H⁺ (→ CO₃²⁻) or accept H⁺ (→ H₂CO₃)
Both H₂CO₃ and HCO₃⁻ are amphiprotic
3. Why effective buffer?
System contains:
- Weak acid: H₂CO₃ (neutralizes added base)
- Conjugate base: HCO₃⁻ (neutralizes added acid)
- Adjustable: CO₂ can be exhaled to remove acid
- Large reservoir: Body has significant HCO₃⁻ reserves
JEE Connection: This links equilibrium to biology—frequent in JEE Advanced!
Quick Revision Box
| Theory | Acid | Base | Key Feature | Limitation |
|---|---|---|---|---|
| Arrhenius | Produces H⁺ | Produces OH⁻ | Simple, historical | Only in water |
| Bronsted-Lowry | Donates H⁺ | Accepts H⁺ | Proton transfer | Needs H⁺ |
| Lewis | Accepts e⁻ pair | Donates e⁻ pair | Broadest | Complex |
Conjugate Pairs:
- Strong acid ↔ Weak conjugate base
- Weak acid ↔ Strong conjugate base
Amphiprotic: H₂O, HCO₃⁻, HSO₄⁻, H₂PO₄⁻, HPO₄²⁻, HS⁻
When to Use Which Theory
Is the reaction in water with H⁺/OH⁻?
- YES → Arrhenius (simplest)
Does the reaction involve proton (H⁺) transfer?
- YES → Bronsted-Lowry
- Works in any solvent
- Can identify conjugate pairs
Does the reaction involve electron pair donation/acceptance (no H⁺)?
- YES → Lewis (only option)
- BF₃, AlCl₃, metal ions
- Coordination chemistry
JEE Strategy: If all three apply, use the simplest (Arrhenius). If asked “according to Lewis theory,” obviously use Lewis!
Connection to Other Topics
Acid-base theories connect to:
- Ionic Equilibrium - Ka, Kb values for weak acids/bases
- pH and Buffers - Quantitative applications of acid-base equilibria
- Salts and Hydrolysis - Conjugate acid-base strength determines pH
- Coordination Compounds - Lewis acid-base forms complexes
- Organic Chemistry - Reaction mechanisms involve acid-base concepts
- Electrochemistry - pH affects electrode potentials
Teacher’s Summary
- Three theories, increasing scope - Arrhenius < Bronsted-Lowry < Lewis
- Conjugate pairs differ by one H⁺ - Acid ⇌ Conjugate base + H⁺
- Inverse strength relationship - Strong acid = Weak conjugate base
- Amphiprotic species are versatile - H₂O, HCO₃⁻, HSO₄⁻, phosphates
- Lewis is the broadest - Includes all proton transfers plus electron pair transfers
“Arrhenius gave us the basics in water, Bronsted-Lowry freed us from water, and Lewis freed us from protons. Each theory is a tool—use the right one for the job!”
JEE Strategy:
- Memorize 7 strong acids (all else are weak!)
- Practice identifying conjugate pairs (easy 2-mark questions)
- Lewis theory connects to coordination chemistry (JEE Advanced loves this)
- Amphiprotic species appear in buffer and hydrolysis problems
- Biological applications (blood pH, stomach acid) are trendy in recent JEE papers
What’s Next?
Now that you understand acid-base theories, apply them quantitatively:
- pH and Buffer Solutions - Calculate pH, design buffers using Henderson-Hasselbalch
- Solubility Product - pH affects solubility of salts
- Coordination Compounds - Lewis acid-base forms metal complexes