Unit 10 · Physical Chemistry

Indicators & Titration Curves

Acid-base indicators, choosing the right indicator, 4 types of titration curves, equivalence point pH, buffer regions, polyprotic acids.

10.1

Acid-Base Indicators

IndicatorA weak acid (HIn) or weak base whose conjugate base (In⁻) has a different colour. Colour change occurs when pH ≈ pKa(HIn). pH range of colour change ≈ pKa ± 1.
IndicatorpKa (HIn)Acid colourAlkaline colourpH range
Methyl orange3.5RedYellow3.1–4.4
Methyl red5.1RedYellow4.4–6.2
Litmus6.5RedBlue5.0–8.0
Bromothymol blue7.1YellowBlue6.0–7.6
Phenolphthalein9.2ColourlessPink/red8.2–10.0
Alizarin yellow11YellowRed10.1–12.0

Choosing the Right Indicator

The indicator must change colour within the vertical portion of the titration curve (the steep region near the equivalence point). If the pH jump is large (e.g. 4 to 10), many indicators work. If small (e.g. weak acid/weak base), few indicators are suitable.

10.2

Titration Curves

Strong Acid + Strong Base (e.g. HCl + NaOH)

Before equivalence: pH = −log[H⁺] (excess acid) Equivalence point: pH = 7 (salt NaCl is neutral) After equivalence: pH = 14 + log[OH⁻] (excess base) pH jump at equivalence: very steep (~pH 4 to 10) Suitable indicators: methyl orange, methyl red, phenolphthalein, bromothymol blue

Weak Acid + Strong Base (e.g. CH₃COOH + NaOH)

Initial pH: higher than strong acid (less dissociation) Half-equivalence point: pH = pKa of the acid (buffer region) Equivalence point: pH > 7 (basic! — CH₃COO⁻ hydrolyses: CH₃COO⁻ + H₂O ⇋ CH₃COOH + OH⁻) After equivalence: excess NaOH pH jump at equivalence: smaller than strong/strong; on alkaline side Suitable indicator: phenolphthalein (8.2–10) — NOT methyl orange (changes too early) Buffer region: pH = pKa + log([A⁻]/[HA]) (Henderson-Hasselbalch)

Strong Acid + Weak Base (e.g. HCl + NH₃)

Equivalence point: pH < 7 (acidic! — NH₄⁺ hydrolyses) pH jump: smaller, on acidic side Suitable indicator: methyl orange or methyl red (NOT phenolphthalein)

Weak Acid + Weak Base (e.g. CH₃COOH + NH₃)

No significant pH jump at equivalence — the titration curve is nearly flat The equivalence point pH ≈ (pKa + pKb + pKw)/2 (approximately neutral, slightly acidic or basic depending on relative Ka/Kb) NO suitable simple indicator — potentiometric titration (pH electrode) must be used
10.3

Polyprotic Acids

Phosphoric Acid H₃PO₄

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Three ionisation steps: H₃PO₄ ⇋ H⁺ + H₂PO₄⁻ Ka₁ = 7.5 × 10⁻³ pKa₁ = 2.12 H₂PO₄⁻ ⇋ H⁺ + HPO₄²⁻ Ka₂ = 6.2 × 10⁻⁸ pKa₂ = 7.21 HPO₄²⁻ ⇋ H⁺ + PO₄³⁻ Ka₃ = 4.8 × 10⁻¹³ pKa₃ = 12.32 Three equivalence points visible on titration curve with NaOH. Buffer regions at half-equivalence: pH = pKa₁, pKa₂, pKa₃ respectively. Used to make pH buffer standard solutions (pKa₂ ≈ 7 — physiological pH).
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Exercises

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Quiz — 25 Questions

Unit 10: Indicators & Titration Curves

25 Qs
Q1

Phenolphthalein is colourless below pH 8.2 and pink above pH 10. It should NOT be used for:

HCl + weak base (NH₃): equivalence point is acidic (pH~5). Phenolphthalein changes at pH 8.2–10, which is well past the equivalence point. The indicator would never reach its end point at the equivalence point — it would need excess NaOH added to reach pH >8.2. Use methyl orange instead.
Q2

At the half-equivalence point of a weak acid titration, pH =

At half-equivalence: [HA] = [A⁻]. Henderson-Hasselbalch: pH = pKa + log(1) = pKa. This directly gives the pKa of the weak acid. A practical method for measuring pKa by titration.
Q3

The equivalence point of a strong acid + weak base titration is:

At equivalence, all weak base (NH₃) is converted to conjugate acid (NH₄⁺). NH₄⁺ hydrolyses: NH₄⁺ + H₂O ⇋ NH₃ + H⁺O⁺ → acidic solution. pH < 7. Use methyl orange (changes in acid range, pH 3.1–4.4).
Q4

For a weak acid/weak base titration:

Weak acid/weak base: NO large pH jump at equivalence. Gradual, flat curve. No indicator has a sharp enough colour change to give a precise end point. Potentiometric method with a pH electrode is required.
Q5

Which indicator is best for strong acid + strong base?

Strong/strong: large pH jump from ~4 to ~10. Many indicators work because their colour change range falls within this jump. Methyl orange (3.1–4.4), phenolphthalein (8.2–10), bromothymol blue (6–7.6) all give sharp end points.
Q6

Methyl orange turns red below pH 3.1 and yellow above pH 4.4. The pKa of methyl orange is approximately:

pKa ≈ midpoint of pH range = (3.1+4.4)/2 ≈ 3.5–3.8. An indicator changes colour when [HIn] = [In⁻], i.e. pH = pKa. The transition range is pKa ± 1.
Q7

The buffer region in a weak acid/strong base titration occurs:

Buffer region: mixture of weak acid (HA) and its conjugate base (A⁻ from neutralisation) present. Resists pH change. pH = pKa + log([A⁻]/[HA]). Maximum buffering at [A⁻] = [HA], i.e. pH = pKa (half-equivalence point).
Q8

A titration curve shows three equivalence points. The acid being titrated is most likely:

Three equivalence points → three ionisable protons → triprotic acid. H₃PO₄ (phosphoric acid) has pKa₁=2.12, pKa₂=7.21, pKa₃=12.32 and shows three steps in its NaOH titration curve.
Q9

The steep (vertical) portion of a titration curve corresponds to:

The steep/vertical portion is near the equivalence point. The pH changes dramatically with tiny additions of titrant. This is where the indicator should change colour. Large vertical region (strong acid/base) → many indicators work. Small vertical region (weak acid/weak base) → difficult to find a suitable indicator.
Q10

Why does the weak acid/strong base titration curve start at a higher pH than strong acid/strong base?

Partial dissociation: HA ⇋ H⁺ + A⁻ with Ka << 1. [H⁺] << C (initial concentration). pH > −log(C). Strong acid: fully dissociated, pH = −log(C). At same concentration, weak acid has higher initial pH.
Q11

Select the correct statement about indicators:

Indicators should be used in small quantities (a few drops) so they don't significantly affect the pH of the solution being titrated. They are themselves weak acids/bases; if too much is added, they consume significant titrant. Best practice: 2–3 drops of indicator.
Q12

In a diprotic acid titration with NaOH, the FIRST equivalence point corresponds to:

First equivalence: H₂A + NaOH → NaHA + H₂O. One mole NaOH per mole acid. Second equivalence: NaHA + NaOH → Na₂A + H₂O (second proton neutralised). Two steps visible on curve with diprotic acids like H₂SO₃, H₂C₂O₄, carbonic acid.
Q13

Bromothymol blue (pH range 6–7.6) is ideal for:

Bromothymol blue's range (6–7.6) is ideal for near-neutral pH systems (biological fluids, aquariums, pools). For titrations, it works for strong acid/strong base (pH jump 4–10 includes 6–7.6), but is impractical for weak acid/strong base (equivalence at pH~9, outside its range).
Q14

The salt formed at equivalence in HCl + NH₃ titration is:

NH₄Cl: NH₄⁺ is the conjugate acid of weak base NH₃. NH₄⁺ + H₂O ⇋ NH₃ + H₃O⁺. Hydrolysis gives H⁺ → acidic solution at equivalence. This is why you use methyl orange (acid range) as indicator.
Q15

If phenolphthalein (range 8.2–10) is used for HCl + NaOH titration:

The steep region of HCl/NaOH titration spans pH ~4 to ~10. Phenolphthalein changes at pH 8.2–10, well within this range. It gives a sharp end point (colourless to pink). All common indicators work for strong/strong titrations because the steep region is so large.
Q16

The pH of 0.1 mol/L CH₃COOH (pKa=4.74) before titration begins is:

pH = ½(pKa − log C) = ½(4.74 − log 0.1) = ½(4.74 + 1) = ½(5.74) = 2.87. Or: [H⁺] = √(Ka·C) = √(10⁻⁴·⁷¹ × 0.1) = √(1.82×10⁻⁶) = 1.35×10⁻³ → pH = 2.87.
Q17

What happens to the pH at the very end of a weak acid/strong base titration (after equivalence)?

After equivalence: excess NaOH dominates. pH controlled by [OH⁻]excess. pH = 14 + log[OH⁻]. As more NaOH is added, pH increases slowly (logarithmic — levels off). Curve is similar to strong/strong titration in this region.
Q18

A student adds too many drops of indicator. This could:

Too much indicator = significant amount of weak acid (HIn) present = consumes more NaOH to change colour. Gives a volume of titrant that is too high (positive error). Good practice: 2–3 drops maximum.
Q19

A buffer solution formed during a weak acid + strong base titration:

A buffer (mixture of weak acid HA and its conjugate base A⁻): resists pH change because: adding H⁺ → reacts with A⁻ (A⁻ + H⁺ → HA); adding OH⁻ → reacts with HA (HA + OH⁻ → A⁻ + H₂O). Effective when pH ≈ pKa ± 1. Best buffer at pH = pKa.
Q20

What is the most important consideration when selecting an indicator for a titration?

The indicator’s transition range must fall in the vertical part of the curve. If the transition range overlaps with the steep section, a sharp end point (small half-drop change) is achieved. If the range is outside the steep section, the colour change occurs too early (before equivalence) or too late (after equivalence) — both give errors.
Q21

The Henderson-Hasselbalch equation pH = pKa + log([A⁻]/[HA]) applies:

Henderson-Hasselbalch applies whenever both weak acid and its conjugate base are present in non-negligible quantities — throughout the buffer region of the titration. Valid when Ka/C << 1 and when neither [HA] nor [A⁻] is negligible.
Q22

Alizarin yellow (pH range 10.1–12.0) would be used for:

Alizarin yellow is only useful when the equivalence point occurs at pH > 10 — very weak acids (e.g. phenol pKa=10) titrated with NaOH. The jump on the curve must include pH 10–12 for this indicator to work. Not useful for common titrations where equivalence is at pH 7–9.
Q23

The pH at the equivalence point of CH₃COOH (pKa=4.74) + NaOH is:

At equivalence: all CH₃COO⁻. Kb(CH₃COO⁻) = Kw/Ka = 10⁻¹⁴/10⁻⁴·⁷¹ = 10⁻⁹·²⁶. For 0.1 mol/L: [OH⁻] = √(Kb·C) = √(10⁻⁹·²⁶×0.05) ≈ 1.17×10⁻⁵ → pOH = 4.93 → pH = 9.07 ≈ 8.9–9. Basic at equivalence.
Q24

For a triprotic acid (H₃A), the second equivalence point corresponds to:

Second equivalence: NaH₂A + NaOH → Na₂HA + H₂O. The second of the three protons is removed. In H₃PO₄/NaOH: second equivalence at pH≈pKa₂≈9.6 (midpoint of pKa₂ and pKa₃).
Q25

A titration is performed twice: once using methyl orange and once using phenolphthalein. Both give the same volume of titrant. The titration is most likely:

If both methyl orange (3.1–4.4) and phenolphthalein (8.2–10) give the same end point, the pH jump must include both ranges simultaneously — i.e. a very large jump from <3.1 to >10. This is characteristic of strong acid + strong base where the pH jump spans from ~2 to ~12.

Unit 10 Quiz — Indicators & Titration (25 Questions)

Select one answer each
Q1

An acid-base indicator changes colour because:

HIn ⇌ H⁺ + In⁻. Low pH: HIn form (one colour); high pH: In⁻ form (different colour). Colour change occurs near pKIn.
Q2

The pH range of colour change for an indicator is approximately:

Indicator changes colour when [In⁻]/[HIn] changes from 1/10 to 10/1 — this spans pH = pKIn ± 1.
Q3

Phenolphthalein changes from colourless to pink in the pH range:

Phenolphthalein: colourless in acid/neutral, pink in alkali. Colour change pH 8.2–10.0. Good for strong base titrations.
Q4

Methyl orange is suitable for titrating:

Methyl orange changes colour pH 3.1–4.4. The equivalence point of strong acid/weak base titration has pH < 7 — within range.
Q5

At the equivalence point of a strong acid–strong base titration, pH is:

At 25°C, strong acid + strong base → water + neutral salt → pH = 7. At other T, Kw differs so neutral pH ≠ 7.
Q6

The equivalence point of a weak acid–strong base titration has pH:

Conjugate base A⁻ hydrolyses water → OH⁻ produced → alkaline equivalence point (pH > 7).
Q7

A titration curve for strong acid vs strong base shows a near-vertical portion at the equivalence point because:

Near equivalence, almost all acid is consumed. A tiny excess of base raises [OH⁻] dramatically → steep pH rise.
Q8

For a weak acid–strong base titration, the buffer region occurs:

As NaOH is added, HA → A⁻. When [HA] = [A⁻], pH = pKa. The flat region of the curve is the buffer region.
Q9

A diprotic acid (H₂A) titrated with NaOH shows:

H₂A + OH⁻ → HA⁻ (first equivalence), then HA⁻ + OH⁻ → A²⁻ (second equivalence). Two steps, two inflection points.
Q10

Back titration is used when:

Add known excess of reagent → allow complete reaction → back-titrate excess to find how much was consumed.
Q11

The indicator used for a strong acid–strong base titration can be:

The steep vertical section of the strong acid/base curve spans pH ~3–11 — both MO (3.1–4.4) and PP (8.2–10) fall within it.
Q12

Why is phenolphthalein NOT suitable for strong acid–weak base titrations?

Strong acid/weak base equivalence point is acidic (pH < 7). The steep region ends well below pH 8 — PP misses it.
Q13

A standard solution is prepared accurately by:

Primary standard: pure, stable, high molar mass, accurately weighable. Dissolved to exact volume — concentration known precisely.
Q14

Standardisation of NaOH solution uses potassium hydrogen phthalate (KHP) because:

KHP (KHC₈H₄O₄) is stable, non-hygroscopic, high molar mass (204 g/mol) — ideal primary standard for NaOH.
Q15

The titre in a titration is:

Titre = volume (cm³ or mL) of titrant required to reach the endpoint, read from the burette.
Q16

In a complexometric titration using EDTA, EDTA:

EDTA (ethylenediaminetetraacetic acid) has 6 donor atoms — forms stable 1:1 complexes with most metal cations.
Q17

Hardness of water is caused by Ca²⁺ and Mg²⁺ ions. It is measured by titration with:

At pH 10 (NH₃/NH₄⁺ buffer), EDTA titrates Ca²⁺ + Mg²⁺. EBT indicator: wine-red (metal-EBT complex) → blue at endpoint.
Q18

Redox titrations using KMnO₄ are carried out in:

In acid: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O. In neutral/alkaline, MnO₄⁻ → MnO₂ (brown precipitate) — not suitable.
Q19

The endpoint is shown in KMnO₄ titrations because:

KMnO₄ is deep purple; Mn²⁺ is almost colourless. The first excess MnO₄⁻ gives a persistent pale pink colour.
Q20

An iodometric titration uses thiosulfate to titrate:

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻. Starch forms intense blue-black complex with I₂ — endpoint: blue → colourless.
Q21

Argentometric titration (Mohr method) uses AgNO₃ to determine:

AgNO₃ + Cl⁻ → AgCl (white precipitate). After all Cl⁻ is consumed, excess Ag⁺ reacts with CrO₄²⁻ → red Ag₂CrO₄.
Q22

The term 'endpoint' in a titration refers to:

Endpoint = indicator colour change. Equivalence point = stoichiometric point. Small difference = titration error.
Q23

In a precipitation titration, the Fajans method uses:

Fluorescein in solution: yellow-green. Adsorbed on AgCl surface after equivalence: pink. Colour change indicates endpoint.
Q24

Potentiometric titration determines the endpoint by:

A pH electrode or ion-selective electrode measures potential throughout. The steep inflection in the plot = equivalence point.
Q25

Why is a magnetic stirrer/constant swirling used during titration?

Mixing ensures the added titrant reacts completely and indicator equilibrium is reached — gives a more accurate endpoint.
📝 Go to Unit Test →
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Unit Test — 50 marks

Section A

30 marks
Q1 [5]

Describe the theory of acid-base indicators. Explain why pKa of the indicator should match the equivalence point pH, and derive the pH range of colour change. [5]

An indicator HIn is a weak acid where HIn and In⁻ have different colours. Equilibrium: HIn ⇋ H⁺ + In⁻, Ka(HIn). When [H⁺] = Ka(HIn): [HIn] = [In⁻] — midpoint of colour change, pH = pKa(HIn). Human eye detects colour change when [In⁻]/[HIn] > 10 or < 1/10. pH range: pKa ± log(10) = pKa ± 1. The indicator's pKa should match the pH at the equivalence point so that the colour change occurs when exactly stoichiometric amounts have reacted. If pKa(indicator) is far from equivalence pH, the end point does not coincide with equivalence point → titration error.
Q2 [5]

Draw and describe the shape of the titration curve for 25 mL of 0.1 mol/L ethanoic acid (pKa=4.74) titrated with 0.1 mol/L NaOH. Mark the initial pH, half-equivalence point, equivalence point, and suggest a suitable indicator. [5]

Initial pH: CH₃COOH 0.1 mol/L: [H⁺]=√(Ka×C)=√(1.82×10⁻⁵)=1.35×10⁻³ → pH=2.87. As NaOH added: pH rises through buffer region. Half-equivalence (12.5 mL NaOH added): pH = pKa = 4.74. Buffer region: pH changes slowly (flat region). Equivalence point (25 mL NaOH): all CH₃COOH → CH₃COO⁻ (0.05 mol/L). pH≈9.07 (basic). Steep section: 24–26 mL NaOH, pH rises from ~7 to ~11. After equivalence: excess NaOH dominates. Suitable indicator: phenolphthalein (8.2–10) — colour change from colourless to pink at pH~9, within the steep section. NOT methyl orange (changes at pH 3–4, during buffer region before equivalence).
Q3 [5]

Compare titration curves for: (a) 0.1 mol/L HCl + 0.1 mol/L NaOH; (b) 0.1 mol/L NH₃ + 0.1 mol/L HCl. Include shape, equivalence point pH, pH jump magnitude, and indicator choice. [5]

(a) HCl + NaOH (strong/strong): Starts pH~1, gradual rise, sharp S-jump pH 4–10 at equivalence, equivalence pH=7 (neutral NaCl salt). Large vertical section spanning pH 4–10. Indicators: any (methyl orange, phenolphthalein, BTB all work). (b) NH₃ + HCl (weak base/strong acid): Starts pH~11 (0.1 mol/L NH₃, pKb=4.74 → pH=11.13). As HCl added: pH falls through buffer region (NH₃/NH₄⁺). Half-equivalence: pH = 14 − pKb = 14 − 4.74 = 9.26 (pKa of NH₄⁺). Equivalence: NH₄Cl formed; NH₄⁺ hydrolyses → pH<7 (~5). Steep section smaller, on acidic side (pH~6 to pH~3). Indicator: methyl orange or methyl red (changes in acid range). NOT phenolphthalein (already colourless at pH~5–6, won’t show endpoint).
Q4 [5]

For a phosphoric acid (H₃PO₄) titration with NaOH: (a) write the equation for each step; (b) state the pH at each half-equivalence point; (c) state the pH at the first equivalence point. [5]

(a) Step 1: H₃PO₄ + NaOH → NaH₂PO₄ + H₂O (pKa₁=2.12). Step 2: NaH₂PO₄ + NaOH → Na₂HPO₄ + H₂O (pKa₂=7.21). Step 3: Na₂HPO₄ + NaOH → Na₃PO₄ + H₂O (pKa₃=12.32).
(b) Half-equivalence points: pH = pKa₁ = 2.12; pH = pKa₂ = 7.21; pH = pKa₃ = 12.32.
(c) First equivalence pH ≈ (pKa₁ + pKa₂)/2 = (2.12 + 7.21)/2 = 4.67 (for diprotic salt NaH₂PO₄ in aqueous solution).
Q5 [5]

A sample of aspirin (acetylsalicylic acid, M=180 g/mol) is dissolved in ethanol and titrated with 0.100 mol/L NaOH. It takes 23.45 mL to reach the equivalence point (pink phenolphthalein). Calculate the mass of aspirin in the sample. [5]

Aspirin + NaOH → sodium aspirin + H₂O (1:1 ratio). Moles NaOH = 0.100 × 23.45/1000 = 2.345×10⁻³ mol. Moles aspirin = 2.345×10⁻³ mol (1:1). Mass = 2.345×10⁻³ × 180 = 0.4221 g aspirin. Note: phenolphthalein is appropriate because aspirin (weak acid, pKa≈3.5) + strong base → equivalence point is basic (pH~8–9), within phenolphthalein range.
Q6 [5]

Explain why: (a) HCl/NaOH has a large pH jump at equivalence; (b) CH₃COOH/NaOH has a smaller jump on the alkaline side; (c) CH₃COOH/NH₃ has virtually no jump. [5]

(a) HCl/NaOH: both strong. Near equivalence, tiny excess acid gives very low pH (e.g. half-drop excess HCl = 0.02 mL × 0.1 = 0.002 mmol excess in ~50 mL → [H⁺]=4×10⁻⁵ mol/L → pH=4.4). Symmetric on alkaline side. Huge swing because no buffer to resist.
(b) CH₃COOH/NaOH: before equivalence, buffer (HA/A⁻) resists pH change → gradual rise. After equivalence, excess NaOH dominates as with strong/strong. pH jump shifted to alkaline side (pH~7 to ~11), smaller in range because the alkaline jump is normal but the acid side is buffered.
(c) CH₃COOH/NH₃: BOTH weak. The buffer action operates on BOTH sides of the equivalence point (weak acid before → buffer; conjugate acid/base pair near equivalence → buffer; weak base after → buffer). The pH changes very gradually throughout — essentially no steep section. The curve is almost a straight line, making endpoint detection impossible with simple indicators.

Section B

20 marks
Q7 [10]

(a) What is meant by a ‘primary standard’ and why is Na₂CO₃ preferred to NaOH for standardising HCl? [3] (b) A student standardises HCl with 0.1000 mol/L Na₂CO₃. Titration data: 25.00 mL Na₂CO₃ requires 24.75 mL HCl to reach the first equivalence point using bromocresol green. Calculate [HCl]. [3] (c) Why would using phenolphthalein instead of bromocresol green for the Na₂CO₃/HCl titration give a different volume of HCl? [4]

(a) Primary standard: pure, stable solid with known exact molar mass; reacts cleanly with no side reactions; NaOH absorbs CO₂ from air (forms Na₂CO₃) and water, so its concentration changes over time — not a primary standard. Na₂CO₃ is anhydrous, stable, available with high purity, M=106 precisely known — ideal primary standard.
(b) Na₂CO₃ + HCl → NaHCO₃ + NaCl (first equivalence, 1:1 ratio per CO₃²⁻ to HCO₃⁻). Moles Na₂CO₃ = 0.1000 × 25.00/1000 = 2.500×10⁻³ mol. Moles HCl = 2.500×10⁻³ mol (1:1 here). [HCl] = 2.500×10⁻³ / 24.75×10⁻³ = 0.1010 mol/L.
(c) Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂ is the SECOND equivalence (complete neutralisation). Bromocresol green detects first equivalence (CO₃²⁻→HCO₃⁻, pH~8.3). Phenolphthalein also detects this first equivalence (changes at pH 8.2–10). HOWEVER: if the student is targeting the SECOND equivalence (pH~4, total reaction), methyl orange would be needed (gives double the HCl volume). Using bromocresol green vs methyl orange: half as much HCl at first vs second equivalence. If phenolphthalein is used, it gives the first equivalence (same as bromocresol green, same volume). But the equivalence point of the second step (pH~4) is below phenolphthalein range — phenolphthalein would have already turned colourless before equivalence in the second step.
Q8 [10]

Critically evaluate the use of acid-base titrations in pharmacy (purity testing of drugs), food science (acidity of fruit juice), and environmental monitoring (alkalinity of river water). For each: describe the titration used, the indicator chosen, and one potential source of error. [10]

Pharmacy (aspirin purity): aspirin (acetylsalicylic acid, pKa≈3.5) dissolved in ethanol, titrated with standardised NaOH solution. Indicator: phenolphthalein (equivalence point basic, pH~9). Error: aspirin slowly hydrolyses to salicylic acid + acetic acid in solution — if sample stands too long before titration, extra acid consumes extra NaOH → high result for acidity → low apparent purity. Mitigated by: fresh solution, cold ethanol solvent, rapid titration.
Food science (fruit juice acidity): juice titrated with 0.1 mol/L NaOH. Indicator: phenolphthalein (mixture of citric, malic, ascorbic acids — all weak → equivalence at basic pH). Report as % citric acid equivalent. Error: CO₂ dissolved in juice acts as weak acid — titrates as carbonic acid → overestimates total acidity. Mitigated by: boiling juice briefly to remove dissolved CO₂ before titration.
Environmental monitoring (river alkalinity): water sample titrated with standardised H₂SO₄ or HCl. Alkalinity due to HCO₃⁻, CO₃²⁻, OH⁻. First equivalence (phenolphthalein endpoint): CO₃²⁻ → HCO₃⁻. Second equivalence (methyl orange/bromocresol green): HCO₃⁻ → CO₂. Report as mg/L CaCO₃ equivalent. Error: dissolved CO₂ from air enters sample during titration (open vessel), forming H₂CO₃ which is titrated as if it were alkalinity — underestimates true alkalinity. Mitigated by: sealed sample bottle, slow careful titration, minimal agitation.
← Unit 9: pHUnit 11: Solubility Product →

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