pH, Buffers & Henderson Equation Why pH matters (NEET-level view) pH is a compact way to express acidity. It guides everything from stomach acid to drug formulations to blood chemistry. For NEET, you must be fast with: strong vs weak acid/base pH, buffer pH using Henderson–Hasselbalch, buffer capacity ideas, indicator choice, and reading titration curves. The pH scale (0–14) with everyday substances. Use this mental map to estimate acidity/basicity quickly. pH Negative logarithm (base 10) of hydrogen ion concentration: lower pH → more acidic; higher pH → more basic. Negative logarithm (base 10) of hydroxide ion concentration. pH + pOH = 14 at 25 C . pOH Definition of pH Definition of pOH Valid at 25 C because water auto-ionization product has this value at this temperature. pH–pOH relation ( 25 C ) Ionic product of water ( 25 C ) Temperature-dependent; equals 1.0 × 10⁻¹⁴ at 25 C . Hence pH + pOH = 14 at 25 C . Pure water has [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M at 25 C ⇒ pH = 7 (neutral). If temperature changes, Kw changes, so neutral pH is not always 7. remember Substance pH pH of common substances (approximate, 25 C ) Lemon juice 2.4 Vinegar (acetic/ethanoic acid solution) Coffee Milk 6.6 Pure water Blood (plasma) 7.4 Baking soda (NaHCO3) Ammonia solution 11 Sodium hydroxide (strong base) 14 Quick calculations: strong acids and bases For a strong monoprotic acid (e.g., HCl) or strong monobasic base (e.g., NaOH), assume complete dissociation. The concentration of H⁺ (or OH⁻) equals the analytical concentration C (in mol L⁻¹). Strong monoprotic acid of concentration C: [H⁺] = C ⇒ pH = −log C. Strong monobasic base of concentration C: [OH⁻] = C ⇒ pOH = −log C ⇒ pH = 14 − pOH (at 25 C ). For polyprotic/polybasic, the first step may dominate; detailed treatment is beyond this concept’s boundary. Examples: 1) 0.010 M HCl: pH = 2.00. 2) 0.040 M NaOH: pOH = 1.40 ⇒ pH = 12.60 ( 25 C ). Weak acids and weak bases: smart approximations Weak acids and bases partially ionize. At low concentrations with weak ionization ( ≪ 1), we can use the square-root approximation. Think of it like a gentle spring: it stretches a little (ionizes a little) and then resists more change. Valid when is small so (C − x) C. Weak acid approximation Weak acid pH in log form Similarly for a weak base with base dissociation constant Kb and concentration C: [OH⁻] = K b C , pOH = 1 2 ( pK b - C), and pH = 14 − pOH (at 25 C ). Example (weak acid): Ethanoic acid (acetic acid; IUPAC: ethanoic acid; SMILES: CC(=O)O), Ka = 1.8 × 10⁻⁵. For C = 0.10 M, pH = 0.5(4.74 − log 0.10) = 0.5(4.74 − (−1))? No — careful: log 0.10 = −1.00, so pH = 0.5(4.74 − (−1.00)) = 0.5(5.74) = 2.87. Buffers: what they are and how they work A buffer solution resists pH change when a small amount of acid or base is added. Two common types: • Acidic buffer: a weak acid (HA) + its salt with a strong base (A⁻) — e.g., ethanoic acid (acetic acid; CC(=O)O) + sodium acetate (CC(=O)[O-].[Na+]). • Basic buffer: a weak base (B) + its salt with a strong acid (BH⁺) — e.g., ammonia (azane; N) + ammonium chloride ([NH4+].[Cl-]). Solution that resists large pH changes on addition of small amounts of acid/base. Buffer Acidic buffer Made from a weak acid and its conjugate base (salt with strong base). Made from a weak base and its conjugate acid (salt with strong acid). Basic buffer How an acidic buffer works: added H⁺ is consumed by A⁻; added OH⁻ is consumed by HA — keeping pH nearly steady. Mechanism idea: the conjugate pair acts like two shock absorbers. If you add H⁺, the conjugate base A⁻ quickly ties it up to make HA. If you add OH⁻, the weak acid HA neutralizes it to form A⁻ and water. Because the ratio [A⁻]/[HA] changes only a little, pH changes only slightly. Acidic buffer (Henderson–Hasselbalch) Basic buffer (Henderson–Hasselbalch) Use Henderson–Hasselbalch only when both components (acid and its salt, or base and its salt) are present in appreciable amounts. You can use molarities or mole ratios directly (if total volume stays the same, ratio of moles = ratio of concentrations). neet-alert Acidic buffer HA + A⁻ (e.g., CH3COOH + CH3COO⁻) Around pK a ± 1 H2CO3/HCO3⁻ (blood plasma) Basic buffer B + BH⁺ (e.g., NH3 + NH4⁺) Around 14 − pK b ± 1 H2PO4⁻/HPO4²⁻ (intracellular) Type Typical pair Effective pH range Biological example Acidic vs basic buffers — quick comparison Worked example (acidic buffer): 0.20 M ethanoic acid + 0.30 M sodium acetate, pK a = 4.74. pH = 4.74 + log(0.30/0.20) = 4.74 + log(1.5) = 4.74 + 0.176 = 4.92. Effect of adding small strong acid: If 0.010 mol HCl is added to 1.0 L of the above buffer, A⁻ decreases by 0.010 mol and HA increases by 0.010 mol. New ratio = (0.30 − 0.010)/(0.20 + 0.010) = 0.29/0.21. New pH = 4.74 + log(0.29/0.21) ≈ 4.74 + 0.140 = 4.88. Only a slight drop. Buffer capacity Amount of strong acid/base a buffer can neutralize with minimal pH change; maximum when [acid] ≈ [salt] (i.e., pH ≈ pK a for acidic buffer). Effective buffer range: about pH = pK a ± 1 (acidic buffers). Best capacity when [A⁻] = [HA] so pH = pK a . remember Biological buffers: blood and cells The bicarbonate buffer (H2CO3/HCO3⁻) maintains blood pH tightly near 7.35–7.45. Blood plasma uses the H2CO3/HCO3⁻ buffer ( pK a ≈ 6.35) to keep pH at 7.35–7.45. Inside cells, the HPO4²⁻/H2PO4⁻ pair ( pK a ≈ 7.20) is important. Even small deviations can affect enzymes and oxygen transport (acidosis pH < 7.35; alkalosis pH > 7.45). In metabolic acidosis, added H⁺ is buffered by HCO3⁻ forming H2CO3, which equilibrates with CO2 + H2O. Lungs help by exhaling CO2, assisting pH recovery. clinical Indicators and choosing the right one Common acid–base indicators Indicator pH range Acidic color Basic color Methyl orange 3.1 – 4.4 Red Yellow Litmus ≈ 4.5 – 8.3 Red Blue Phenolphthalein 8.3 – 10.0 Colorless Pink 2026-05-26T17:04:34.542Z Indicator color changes across their transition ranges: methyl orange, litmus, and phenolphthalein. gpt-image-2 Indicator color-change diagram: three horizontal bars for methyl orange (pH 3.1–4.4 red→orange→yellow), litmus (pH 4.5–8.3 red→purple→blue), phenolphthalein (pH 8.3–10 colorless→faint pink→pink). Clean vector style, pH axis 0–14, arrows in red, no text inside image. Titration curves: where is the equivalence point? The equivalence point pH depends on acid/base strengths: • Strong acid vs strong base: equivalence near pH 7. • Weak acid vs strong base: equivalence above 7 (often ~8–9). • Strong acid vs weak base: equivalence below 7 (often ~5–6). • Weak vs weak: no sharp jump — avoid simple indicators. Titration types and suitable indicators Acid + Base Equivalence pH Suitable indicator(s) Strong acid + Strong base ≈ 7 Methyl orange or Phenolphthalein Weak acid + Strong base > 7 (≈ 8–9) Phenolphthalein Strong acid + Weak base < 7 (≈ 5–6) Methyl orange Weak acid + Weak base No sharp change Avoid — use pH meter Set of 4 titration curves on white background: (1) HCl vs NaOH (S-shaped crossing pH 7), (2) CH3COOH vs NaOH (buffer region, pH at half-equivalence = pKa, equivalence >7), (3) HCl vs NH3 (equivalence <7), (4) weak–weak shallow curve. Mark suitable indicator ranges as colored bands. Vector, clean axes, labels. gpt-image-2 Four titration curves: SA–SB, WA–SB, SA–WB, and WA–WB with equivalence regions and indicator windows. 2026-05-26T17:04:35.030Z Acid + base → salt + water underlies titration and buffer action. The steep pH change near equivalence in strong–strong titrations comes from rapid neutralization. Industrial and clinical relevance Blood-buffer maintenance: bicarbonate system guards 7.35–7.45 (acidosis < 7.35; alkalosis > 7.45). Pharmaceuticals: buffer formulations stabilize drug pH for stability and absorption. Quality control: acid–base titrations determine purity and concentration. pH-sensitive drug release: coatings dissolve at set pH (e.g., enteric coatings). Wastewater: pH control prevents corrosion/precipitation; e.g., high pH (> 9) may require treatment (e.g., alum-based processes) before discharge. Buffer recipe: Acidic = HA + salt of HA (A⁻). Basic = B + salt of BH⁺. Buffers only resist significant pH changes within their buffer capacity; small but real pH shifts still occur. Buffers completely prevent any change in pH. A buffer solution is always neutral with pH = 7. Buffer pH depends on the acid/base pair and their ratio; it can be acidic, neutral, or basic. Very strong acid solutions can have pH < 0 (e.g., ~10 M HCl has pH around −1). pH cannot be less than 0. Use phenolphthalein only when equivalence pH lies in 8.3–10 (e.g., weak acid–strong base). It is wrong for strong acid–weak base where equivalence pH < 7. Phenolphthalein works for all titrations. Key terms at a glance −log10[H⁺]; acidity scale at 25 C neutral = 7. pH −log10[OH⁻]; at 25 C , pH + pOH = 14. pOH Relates buffer pH to pKa (or pKb) and ratio of conjugate pair. Henderson–Hasselbalch equation Amount of strong acid/base a buffer can absorb with minimal pH change; max near [acid] ≈ [salt]. Buffer capacity Weak acid + its conjugate base (salt). Acidic buffer Weak base + its conjugate acid (salt). Basic buffer A weak acid/base dye with different colors in two forms and a characteristic transition range. Acid–base indicator Plot of pH vs volume of titrant added, showing buffer regions and equivalence point. Titration curve Stoichiometric completion of reaction (acid moles = base moles for monoprotic). Equivalence point Where half the acid is neutralized; for weak acid–strong base, pH = pKa. Half-equivalence point Bicarbonate system H2CO3/HCO3⁻ maintaining pH 7.35–7.45. Blood buffer Before using Henderson, convert masses/volumes to moles and update moles after any added strong acid/base. Then take the ratio. Do not forget dilution only if volumes change — but ratios of moles remain valid if both components are diluted equally. neet-alert High-yield checkpoints Strong acid/base: pH from −log C (monoprotic/monobasic). Weak acid: [H⁺] = K a C ; pH = 0.5( pK a − log C). Weak base analogous for pOH. Henderson: pH = pK a + log([salt]/[acid]); pOH = pK b + log([salt]/[base]). Buffer capacity max when [salt] ≈ [acid]; effective range ≈ pK a ± 1. Choose indicator to match the equivalence pH window.