1 The NEET Cheat Sheet — Core Concepts & Exceptions
Acids and bases are two opposite kinds of chemicals — acids are sour (think lemon), bases are slippery (think soap). Scientists use three rules to decide what counts as an acid or base, each one a little broader than the one before.
Acid–Base concepts
Arrhenius
💡 The oldest, simplest rule, and it only works in water: an acid is something that lets go of a tiny positive bit (H⁺), and a base lets go of OH⁻. That's it.
Acid → furnishes H⁺ (H₃O⁺) in water · Base → furnishes OH⁻ in water.
📌 Acids: HCl, HNO₃, H₂SO₄ · Bases: NaOH, KOH, LiOH.
⚠ Limited to aqueous solutions only; cannot explain basicity of NH₃ (no OH⁻) or acidity of CO₂.
Brønsted–Lowry
💡 Picture H⁺ as a ball. The acid is the player who throws the ball; the base is the one who catches it. They always come in throw-and-catch pairs (called conjugate pairs).
Acid = proton (H⁺) donor · Base = proton acceptor. Always in conjugate pairs differing by one H⁺.
📌 HCl + H₂O → H₃O⁺ + Cl⁻ (Cl⁻ = conjugate base) · NH₃ + H₂O → NH₄⁺ + OH⁻. Water is amphoteric.
⚠ Strong acid ⇒ very weak conjugate base (Cl⁻, NO₃⁻, HSO₄⁻); strong base ⇒ weak conjugate acid. Cannot explain acids/bases without transferable H⁺.
Lewis
💡 The widest rule of all — it's about pairs of electrons (tiny shared dots), not H⁺. The acid grabs a spare pair of electrons; the base offers one. So even chemicals with no H at all can be acids.
Acid = electron-pair acceptor · Base = electron-pair donor.
📌 Lewis acids: BF₃, AlCl₃, BCl₃, H⁺, Mg²⁺, CO₂ · Lewis bases: NH₃, OH⁻, F⁻, H₂O.
⚠ Most general, but does not account for relative acid/base strength and labels even non-protonic species as acids.
How a conjugate pair works (Brønsted)
Lose a proton → become the conjugate base · gain a proton → become the conjugate acid.
Exceptions, limitations & temperature effects
- Ostwald's dilution law is valid ONLY for weak electrolytes; it fails for strong electrolytes (which are ~completely ionised, so α isn't conc-dependent this way).
- Ostwald approximation (α = √(Ka/C)) holds only when α < 0.05 (i.e. 1−α ≈ 1).
- Very dilute strong acid (e.g. 10⁻⁸ M HCl): you CANNOT ignore water's H⁺. pH ≈ 6.98 (slightly acidic), NOT 8. Include water when C ≤ 10⁻⁶ M.
- Temperature ↑ ⇒ Kw ↑ (ionisation of water is endothermic). At 298 K, Kw = 10⁻¹⁴, neutral pH = 7. At higher T, Kw > 10⁻¹⁴ ⇒ neutral water has pH < 7 (still neutral as [H⁺]=[OH⁻]).
- Neutral ≠ pH 7 in general — that equality is true only at 25 °C / 298 K.
- Salts of strong acid + strong base (NaCl, KNO₃) do NOT hydrolyse ⇒ pH = 7 (neutral).
The pH scale (with NCERT Table 6.5 values)
Lower than 7 = acidic · 7 = neutral (at 25 °C) · higher than 7 = basic. Each whole step is a ×10 change in [H⁺]. Data: NCERT Table 6.5.
2 The Formula & Approximation Bank
Every formula NEET tests, grouped by topic — with the time-saving approximation beside each. A one-line plain summary sits above each group.
pH & Water Equilibrium
💡 pH is just a short way of saying how acidic something is. Instead of writing a tiny number like 0.0000001, we write pH = 7. Smaller pH means more acidic, bigger means more basic.
| Formula | Approximation / Notes |
|---|---|
| Kw = [H⁺][OH⁻] = 10⁻¹⁴ (at 298 K) | pKw = pH + pOH = 14 at 298 K only. |
| pH = −log[H⁺] ; pOH = −log[OH⁻] | [H⁺] = 10⁻ᵖᴴ. Strong acid: [H⁺] = normality. |
| Strong base: [OH⁻] = C, pH = 14 − pOH | For C ≥ 10⁻⁶ M ignore water; below that, add water's 10⁻⁷. |
Weak Acid / Base & Ostwald's Law
💡 A weak acid only lets go of a few of its H⁺ bits. The fraction that actually splits up is called the degree of dissociation (α) — like counting how many kids in a class raised their hands.
| Formula | Approximation / Notes |
|---|---|
| Ka = Cα²/(1−α) | If α < 0.05 ⇒ 1−α ≈ 1 ⇒ Ka ≈ Cα². |
| α = √(Ka / C) (Ostwald) | Valid for weak electrolytes, α < 5%. |
| [H⁺] = √(Ka·C) ⇒ pH = ½(pKa − log C) | Monoprotic weak acid; assumes α small. |
| [OH⁻] = √(Kb·C) ⇒ pOH = ½(pKb − log C) | Weak base, α small. |
| pKa + pKb = pKw = 14 (conjugate pair, 298 K) | Use to convert between Ka and Kb. |
Salt Hydrolysis
💡 When a salt dissolves, its pieces sometimes react with water and quietly make the solution a little acidic or a little basic. That reaction with water is called hydrolysis.
| Formula | Approximation / Notes |
|---|---|
| Weak acid + strong base (e.g. CH₃COONa): pH = 7 + ½(pKa + log C) | Solution basic (pH > 7). Kh = Kw/Ka. |
| Weak base + strong acid (e.g. NH₄Cl): pH = 7 − ½(pKb + log C) | Solution acidic (pH < 7). Kh = Kw/Kb. |
| Weak acid + weak base (e.g. CH₃COONH₄): pH = 7 + ½(pKa − pKb) | Independent of concentration! Kh = Kw/(Ka·Kb). |
| Degree of hydrolysis h = √(Kh / C) | Valid when h small. |
Buffer Solutions (Henderson–Hasselbalch)
💡 A buffer is a shock-absorber for pH. Add a little acid or base and the pH barely moves. It's made of a weak acid together with its own salt working as a team.
| Formula | Approximation / Notes |
|---|---|
| Acidic buffer: pH = pKa + log([salt]/[acid]) | Best buffering when [salt]≈[acid] ⇒ pH = pKa. |
| Basic buffer: pOH = pKb + log([salt]/[base]) | pH = 14 − pOH (298 K). |
| At half-neutralisation point: pH = pKa | Half the weak acid neutralised by strong base. |
Solubility Product (Ksp)
💡 Some salts barely dissolve in water. Ksp is a number that tells you exactly how much can dissolve before the rest just stays as a solid at the bottom.
| Formula | Approximation / Notes |
|---|---|
| AB type (AgCl): Ksp = s² ⇒ s = √Ksp | s = molar solubility (mol L⁻¹). |
| AB₂ / A₂B (CaF₂, Ag₂CrO₄): Ksp = 4s³ ⇒ s = (Ksp/4)^⅓ | Watch the stoichiometric coefficients. |
| AₓBᵧ: Ksp = xˣ·yʸ·s^(x+y) | General formula. |
| Common ion (conc C): s = Ksp / Cⁿ | Solubility ↓ sharply (common ion effect). |
| Precipitation occurs when Qsp > Ksp | Qsp = Ksp ⇒ saturated; < ⇒ unsaturated. |
📈 Interactive Titration Curve
Slowly adding a base to an acid changes its pH. If you plot pH against the amount of base added, you get an S-shaped curve with a sudden jump at the “equivalence point” (where acid and base exactly cancel). Drag the sliders and watch the curve and the pH change.
Acid volume fixed at 25 mL. Strong-acid curves jump through pH 7; weak-acid curves have a buffer plateau (flat region) around pH = pKa and an equivalence point above 7.
3 The “Trap Avoidance” & Shortcut Guide
These are the exact spots where students lose easy marks. Each card shows the trap, then the quick fix.
Buffer vs Salt-Hydrolysis (the mixing trap)
After mixing, first work out millimoles (M × mL). Then read off this table:
| After mixing you are left with… | It is a | Use this |
|---|---|---|
| Weak acid in excess over strong base (salt + leftover acid) | BUFFER | pH = pKa + log([salt]/[acid]) |
| Equal moles — complete neutralisation (only salt left) | SALT HYDROLYSIS | pH = 7 + ½(pKa + log C) |
| Strong acid / base in excess | STRONG | pH from the leftover strong reagent |
⚡ Shortcut: at the half-neutralisation point (half the weak acid used up) → pH = pKa instantly.
Common-Ion Effect in Ksp problems
- Set the common-ion concentration ≈ C (ignore the tiny extra from the dissolving salt).
- For an AB salt: s = Ksp / C. For AB₂: s = √(Ksp / 4C²)-type — but in NEET just substitute C and solve.
- Result: solubility drops by orders of magnitude vs pure water.
The 10⁻⁸ M strong-acid trap
Rule: whenever the acid/base concentration is ≤ 10⁻⁶ M, you must add water’s contribution.
α-approximation & log shortcuts
- Use α = √(Ka/C) only when α < 5% (otherwise solve the full quadratic).
- Salt of a weak acid + weak base ⇒ pH is independent of concentration (pure ½(pKa − pKb)).
Memorise these log values to fly through calculations:
4 High-Yield Practice MCQs
Try each question first, then open the solution to see the quickest route — using the formulas and log shortcuts from above, not long textbook working.
10 NEET-level questions with numerically close options — 2 theory, 3 pH/Ostwald, 3 buffers/hydrolysis, 2 Ksp. Tap “Show fastest solution”.
TheoryQ1. Which of the following can act as a Lewis acid but NOT as a Brønsted acid?
- HCl
- BF₃
- CH₃COOH
- NH₄⁺
Show fastest solution
Answer: (b) BF₃BF₃ has an electron-deficient boron (empty orbital) → accepts an electron pair (Lewis acid), but it has no H⁺ to donate, so it is not a Brønsted acid. HCl, CH₃COOH and NH₄⁺ all donate protons (Brønsted acids).
TheoryQ2. The conjugate base of the dihydrogen phosphate ion (H₂PO₄⁻) is:
- H₃PO₄
- HPO₄²⁻
- PO₄³⁻
- H₂PO₄⁻
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Answer: (b) HPO₄²⁻A conjugate base is formed by removing one H⁺: H₂PO₄⁻ − H⁺ → HPO₄²⁻. (H₃PO₄ is its conjugate acid; PO₄³⁻ is two protons away.)
pH / OstwaldQ3. The pH of a 0.001 M HCl solution is:
- 2.0
- 3.0
- 3.5
- 11.0
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Answer: (b) 3.0Strong acid fully ionises: [H⁺] = 10⁻³ ⇒ pH = −log10⁻³ = 3.0. (11.0 is the pOH — a classic distractor.)
pH / OstwaldQ4. A 0.1 M weak monoprotic acid has Ka = 1.0×10⁻⁵. Its pH is approximately:
- 2.0
- 3.0
- 4.0
- 5.0
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Answer: (b) 3.0Fast method: pH = ½(pKa − log C) = ½(5 − (−1)) = ½(6) = 3.0. Or [H⁺] = √(Ka·C) = √(10⁻⁵·10⁻¹) = √10⁻⁶ = 10⁻³.
pH / OstwaldQ5. The degree of dissociation of 0.01 M acetic acid (Ka = 1.8×10⁻⁵) is about:
- 0.42 %
- 1.8 %
- 4.2 %
- 13 %
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Answer: (c) 4.2 %α = √(Ka/C) = √(1.8×10⁻⁵ / 10⁻²) = √(1.8×10⁻³) = √0.0018 ≈ 0.042 = 4.2 %. (α < 5%, so the approximation is valid.)
Buffer / HydrolysisQ6. Equal volumes of 0.2 M CH₃COOH and 0.2 M CH₃COONa are mixed (pKa = 4.74). The pH is:
- 3.74
- 4.74
- 5.74
- 7.00
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Answer: (b) 4.74This is a buffer with equal salt and acid (ratio = 1). pH = pKa + log(1) = pKa + 0 = 4.74. Mixing equal volumes keeps the ratio 1, so concentration cancels.
Buffer / HydrolysisQ7. For an acetic-acid/acetate buffer (pKa = 4.74), what is the pH when [salt]/[acid] = 10?
- 3.74
- 4.74
- 5.74
- 6.74
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Answer: (c) 5.74pH = pKa + log([salt]/[acid]) = 4.74 + log10 = 4.74 + 1 = 5.74. (Tripling distractors come from sign errors in log.)
Buffer / HydrolysisQ8. The pH of a 0.1 M sodium acetate solution (Ka of acetic acid = 1.0×10⁻⁵) is:
- 5.0
- 7.0
- 8.5
- 9.0
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Answer: (d) 9.0Salt of weak acid + strong base ⇒ pH = 7 + ½(pKa + log C) = 7 + ½(5 + (−1)) = 7 + 2 = 9.0 (basic). Choosing 5.0 forgets it's the SALT (basic), not the acid.
KspQ9. The solubility product of AgCl is 1.8×10⁻¹⁰. Its molar solubility in pure water is:
- 1.8×10⁻¹⁰
- 1.34×10⁻⁵
- 1.8×10⁻⁵
- 9.0×10⁻¹¹
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Answer: (b) 1.34×10⁻⁵AB type: Ksp = s² ⇒ s = √(1.8×10⁻¹⁰) = √1.8 ×10⁻⁵ ≈ 1.34×10⁻⁵ mol L⁻¹. (1.8×10⁻¹⁰ wrongly skips the square root.)
KspQ10. The molar solubility of AgCl (Ksp = 1.8×10⁻¹⁰) in 0.1 M NaCl is:
- 1.34×10⁻⁵
- 1.8×10⁻⁹
- 1.8×10⁻¹⁰
- 1.8×10⁻¹¹
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Answer: (b) 1.8×10⁻⁹Common-ion effect: [Cl⁻] ≈ 0.1 M (from NaCl). s = Ksp/[Cl⁻] = 1.8×10⁻¹⁰ / 0.1 = 1.8×10⁻⁹ mol L⁻¹ — far lower than in pure water (1.34×10⁻⁵), exactly as the common-ion effect predicts.