Year 12 Chemistry Module 6 — Acid/Base Reactions ⏱ ~45 min Lesson 19 of 19 IQ3

Acid/Base Analysis Techniques — Industrial & Digital

A food laboratory analyses hundreds of vinegar samples per day — digital pH probes, automated titration instruments, and back titration all running simultaneously. Understanding which method to use when, and why the glass electrode works at all, is what distinguishes a trained analytical chemist from someone who just follows a protocol.

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Think First

A food safety inspector visits three different laboratories, each analysing the same batch of commercial vinegar for acidity.

Laboratory A uses a glass electrode pH probe calibrated with two buffer solutions and reads pH 2.87 directly.
Laboratory B performs a manual titration with phenolphthalein and 0.1000 mol/L NaOH, calculating % acetic acid from the titre.
Laboratory C performs a conductometric titration, reading the EP from the minimum of the conductance curve.

All three methods give the same answer to within ±0.5% for the acetic acid concentration. Before you read on, write down:

Question 1: What are the different advantages each laboratory is exploiting?

Question 2: When would you choose a direct pH reading over a titration?

Question 3: Which method is most appropriate for a continuous production line monitoring scenario?

📐

Key Relationships — This Lesson

E = E° + (RT/F) × ln[H⁺]

Nernst equation (simplified for pH probe)

E ≈ E° − 0.0592 × pH (at 25°C)

E in volts; coefficient = 0.0592 V/pH unit

Slope = (V₂ − V₁) / (pH₂ − pH₁) ≈ −0.0592 V/pH

Two-point linear calibration of a pH meter

% acetic acid = (mass(CH₃COOH) / (volume × density)) × 100%

Direct titration calculation for vinegar

📖 Know

How a glass electrode pH meter works physically and electrically.
The two-point calibration process using buffer solutions.
Specific acid/base analysis techniques used in the food, pharmaceutical, and environmental industries.

💡 Understand

Why calibration is mandatory to correct for electrode aging and temperature.
The precision differences between a pH probe reading and a volumetric titration.
Why different industries select different analytical methods for quality control.

✅ Can Do

Compare the advantages and limitations of digital probes vs chemical indicators.
Calculate the percentage of acetic acid in vinegar from titration data.
Calculate approximate concentration from a direct pH reading using a reverse Ka calculation.

📚 Core Content

Key Terms — scan these before reading

Brønsted-Lowry acidA proton (H⁺) donor in an acid-base reaction.

Brønsted-Lowry baseA proton (H⁺) acceptor in an acid-base reaction.

Conjugate acid-base pairTwo species differing by one H⁺ that interconvert.

pHThe negative logarithm of hydronium ion concentration.

BufferA solution resisting pH change upon addition of small amounts of acid or base.

TitrationA technique to determine concentration by reaction with a standard solution.

A glass electrode pH meter is not a chemical indicator in disguise — it is an electrochemical device that converts the concentration of H⁺ ions in solution into a measurable electrical voltage, and the physical principle behind this conversion connects pH measurement directly to thermodynamics.

A glass electrode consists of a thin glass membrane with a special composition (high in SiO₂, Na₂O, and CaO) that is selectively permeable to H⁺ ions. The membrane has hydrated gel layers on both surfaces. Inside the electrode is a reference solution of known [H⁺] (typically 0.1 mol/L HCl with a fixed, known pH).

When the electrode is placed in a test solution, H⁺ ions from the test solution exchange into the outer gel layer of the glass membrane, while H⁺ ions from the inner reference solution occupy the inner gel layer. Because the H⁺ concentrations on the two sides of the membrane are different, a potential difference (voltage) develops across the membrane.

This voltage (E) is proportional to the difference in [H⁺] between the test solution and the reference solution, described by the Nernst equation: at 25°C, E ≈ E° − 0.0592 × pH. A higher [H⁺] in the test solution (lower pH) produces a larger voltage difference. The pH meter's electronics measure this voltage and convert it to a pH reading using the calibration relationship.

A second electrode (the reference electrode — typically an Ag/AgCl or calomel electrode) provides a stable, constant potential against which the glass electrode potential is measured. The combination of glass electrode + reference electrode constitutes the pH probe.

Role

Selective H⁺ exchange

Known fixed [H⁺]

Stable external potential

Voltage → pH conversion

Nernst slope correction

Physical process

H⁺ ions partition into the gel layer; potential difference develops

Provides stable inner reference for voltage measurement

Maintains constant reference; allows voltage difference to be measured

Uses Nernst relationship and calibration slope to convert voltage to pH

Temperature changes Nernst slope; modern probes compensate automatically

Must DoIn HSC extended response questions on how a pH meter works, the minimum complete answer includes: (1) the glass membrane is selectively permeable to H⁺ ions; (2) a potential difference (voltage) develops across the membrane proportional to the difference in [H⁺]; (3) the voltage is converted to pH by the meter electronics using the Nernst relationship; (4) calibration with buffer solutions establishes the slope and intercept. All four elements are required for full marks.

Common ErrorStudents say "the pH meter measures the concentration of H⁺ directly." The probe does not measure concentration directly — it measures a voltage (electrical potential difference). The conversion from voltage to pH is done by the meter's electronics. Direct concentration measurement is not how the glass electrode works.

InsightThe Nernst slope of −0.0592 V/pH unit at 25°C means that a change of 1 pH unit corresponds to exactly 59.2 mV of potential difference. Modern pH meters measure voltage to ±0.1 mV precision, corresponding to ±0.002 pH units — far more precise than any indicator or universal indicator paper.

Calibration — Why It Is Mandatory and How It Is Done

Correcting for electrode aging and temperature drift

A pH probe that has not been calibrated that day is not reliable — because the glass membrane properties, the reference electrode potential, and the Nernst slope all drift with time, temperature, and use, and calibration with buffer solutions corrects for all of these drifts simultaneously.

Calibration is the process of establishing the relationship between the voltage measured by the probe and the true pH of the solution, using solutions of precisely known pH (buffer solutions). Without calibration, the relationship is unknown because:

The glass membrane ages: the hydrated gel layer properties change over time, shifting the electrode potential.
Temperature affects the Nernst slope: at 25°C the slope is −0.0592 V/pH, but at 37°C it is −0.0614 V/pH.
The reference electrode potential can drift slightly depending on conditions.

Two-point calibration procedure: (1) rinse the electrode with distilled water and blot dry; (2) immerse in Buffer Solution 1 (e.g. pH 4.00); (3) wait for the reading to stabilise; (4) set the meter reading to pH 4.00 (this sets the intercept); (5) rinse and blot; (6) immerse in Buffer Solution 2 (e.g. pH 7.00); (7) wait for stability; (8) set the meter reading to the buffer's pH (this sets the slope). The two calibration points define a linear relationship between voltage and pH.

What is adjusted

Intercept of voltage-pH line

Slope of voltage-pH line

Theoretical Nernst slope

Removes previous solution

Physical significance

Sets absolute position of calibration line (offset correction)

Sets sensitivity: mV per pH unit (Nernst slope correction)

Corrects for temperature-dependent slope change

Prevents carry-over contamination between calibration points

Must DoIn any practical investigation using a pH probe, describe the calibration step explicitly: "The pH probe was calibrated with two buffer solutions of known pH (pH 4.00 and pH 7.00) before any measurements were taken. Calibration establishes the slope and intercept of the voltage-pH relationship." A practical report that omits the calibration description loses marks.

Common ErrorStudents say "calibration sets the pH meter to read zero at pH 7." This is wrong — calibration sets both the intercept (using one buffer) and the slope (using a second buffer). Setting the meter to read zero at pH 7 (one-point calibration) only adjusts the intercept — it does not correct for slope errors. Two-point calibration is always required.

Comparing Methods — When to Use Each Technique

Matching the analytical tool to the sample and purpose

No single analytical technique is best for every situation — the choice between a pH probe, a direct indicator titration, a back titration, and a conductometric titration depends on the properties of the sample, the required precision, the available equipment, and whether the measurement needs to be continuous or single-point.

Method 1 — Direct pH probe reading: measures current [H⁺] in the solution; gives pH directly in seconds; continuous monitoring possible; suitable for any solution (coloured, turbid, or clear). Does not give concentration directly — must use Ka and Henderson-Hasselbalch or compare to a standard if concentration is needed. Best for: rapid screening, continuous monitoring.
Method 2 — Indicator titration (direct): gives moles of acid or base reacted directly; suitable when the analyte is soluble and fast-reacting; requires a suitable indicator (matching EP pH); gives concentration to ±0.1% with care. Not suitable for: coloured solutions, turbid solutions, weak acid + weak base, very dilute solutions. Best for: vinegar % acidity, drug purity.
Method 3 — Back titration: essential for insoluble or slow-reacting analytes; gives mass and percentage of active ingredient; more steps and more sources of error than direct titration. Best for: CaCO₃ in antacid/eggshell/limestone.
Method 4 — Conductometric titration: no indicator required; works for coloured or turbid solutions, very dilute solutions, weak acid + weak base; equivalence point determined objectively from graph rather than subjective colour change. Requires conductance meter. Best for: weak acid + weak base, coloured solutions, automated industrial applications.

pH probe (direct)

Gives directly: pH; concentration needs Ka

Best for: Rapid screening; continuous monitoring; any solution type

Not suitable for: Precise concentration without Ka data

Indicator titration

Gives directly: Concentration (directly)

Best for: Clear, soluble, fast-reacting analytes; sharp EP

Not suitable for: Coloured/turbid solutions; weak/weak; very dilute

Back titration

Gives directly: Mass and % of analyte

Best for: Insoluble, slow, or gaseous analytes

Not suitable for: Rapidly reacting soluble analytes

Conductometric

Gives directly: EP volume without indicator

Best for: Coloured solutions; weak/weak; dilute; automated

Not suitable for: Requires specialised equipment; sensitive to temperature

Must DoIn any HSC question asking you to select or justify an analytical method, state: (1) which method you choose; (2) why it is appropriate for this specific sample (linking to a property of the sample — solubility, colour, strength, concentration); (3) why at least one alternative method would be less suitable. All three elements are required for full marks.

Common ErrorStudents select the pH probe as "always the most accurate method." A calibrated pH probe is highly precise for pH measurements — but it does not directly give concentration. To determine % acidity of vinegar, a titration is needed — not a pH reading alone. A pH probe gives pH; a titration gives concentration; a back titration gives mass percentage of insoluble analyte. These are different measurements for different purposes.

Industrial Applications — Acid/Base Analysis Across Sectors

How the real world monitors acidity and purity

Every industry that produces or uses acidic or basic substances — from food and beverage to pharmaceuticals to environmental monitoring — has a specific, regulated acid-base analysis technique tailored to the properties of its analyte and the precision required by its quality specifications.

Food and beverage industry: Wine acidity is measured by titrating total acidity (primarily tartaric acid) with NaOH — the result is expressed as g/L tartaric acid equivalents. Dairy acidity in yoghurt and cheese is measured by titrating lactic acid with NaOH (Dornic method). Citric acid in fruit juice and ascorbic acid (vitamin C) are all analysed by NaOH titration with phenolphthalein.
Pharmaceutical industry: Aspirin tablets are quality controlled by dissolving a known mass in ethanol-water, titrating with NaOH, and comparing to the stated acetylsalicylic acid content. The purity test verifies that no hydrolysis has occurred (aspirin → acetic acid + salicylic acid) — hydrolysis would change the Ka and the titration result.
Environmental monitoring: Water quality pH is measured continuously by glass electrode probes at treatment plants and discharge points. Total alkalinity of natural water (primarily HCO₃⁻) is determined by titration with H₂SO₄. Acid rain characterisation: SO₂ absorbed into water forms H₂SO₄; pH < 5.6 indicates acidic deposition. Industrial effluent must be neutralised to pH 6.5–8.5 before discharge — continuous pH monitoring with automated dosing of neutralising agents.

Industry	Analyte	Analysis method	Key specification
Wine	Tartaric/malic/acetic acid	NaOH titration; phenolphthalein	Total acidity 5–8 g/L tartaric acid equivalents
Dairy	Lactic acid	NaOH titration (Dornic method)	Yoghurt: 70–140°D; cheese: varies
Pharmaceuticals	Acetylsalicylic acid (aspirin)	NaOH titration; pH probe for purity	≥99.0% purity by mass
Environmental	pH, total alkalinity (HCO₃⁻)	Glass electrode; H₂SO₄ titration	Drinking water: pH 6.5–8.5; alkalinity >50 mg/L
Wastewater	pH, dissolved acids/bases	Continuous pH probe + automated dosing	Discharge pH 6.5–8.5 (legal requirement)

Must DoFor any industrial acid-base analysis question, your answer must include: (1) the name of the specific acid or base being analysed; (2) the analytical method used and why it is appropriate; (3) a specific quantitative standard or regulatory requirement that the analysis must meet.

Common ErrorStudents describe acid-base analysis using only litmus paper or universal indicator, ignoring the digital and titration methods specified by the NESA syllabus. Always describe the glass electrode pH probe as the standard instrument, with litmus and universal indicator as outdated qualitative tools suitable only for rough pH estimation.

Household Substance Analysis — Vinegar and Antacids

The HSC prescribed investigations in full detail

The HSC prescribed investigation requires the chemical analysis of a common household substance — and the two most commonly specified substances (vinegar and antacid tablets) bring together every calculation and technique from L14–L19 in a single integrated practical.

Household substance 1 — Vinegar (% acetic acid by direct titration): Vinegar is a dilute aqueous solution of acetic acid (CH₃COOH), typically labelled 4–8% acidity.

Procedure: (1) rinse and fill burette with 0.5000 mol/L NaOH (standardised); (2) pipette exactly 10.00 mL of vinegar into a conical flask; (3) dilute with approximately 20 mL of distilled water (reduces intensity of yellow vinegar colour); (4) add 3 drops of phenolphthalein; (5) titrate with NaOH until first permanent faint pink, 30 seconds; (6) record concordant titres.
Calculation: n(NaOH) = c × V; n(CH₃COOH) = n(NaOH) (1:1); mass(CH₃COOH) = n × 60.06; % = (mass/volume of vinegar × density) × 100%.
Digital probe alternative: measure pH of undiluted vinegar with a calibrated pH probe → use Ka = 1.8 × 10⁻⁵ and ICE table in reverse: [H⁺] = 10⁻ᵖᴴ; c(CH₃COOH) ≈ [H⁺]²/Ka. This gives approximate concentration without a full titration — useful for screening but less precise than titration.

Household substance 2 — Antacid tablet (back titration for CaCO₃): Full procedure from L18 applies — crush tablet, add excess standard HCl, allow complete reaction, drive off CO₂, back-titrate excess with standard NaOH, calculate via four-step method.

Analysis	Sample	Method	Key calculation step	Indicator or probe
Vinegar % acidity	10.00 mL vinegar	Direct NaOH titration	n(CH₃COOH) = n(NaOH); % = mass/volume × 100%	Phenolphthalein
Vinegar pH	Undiluted vinegar	Glass electrode, calibrated	pH reading; c estimated from Ka	pH probe (calibrated)
Antacid % CaCO₃	Crushed tablet	Back titration: excess HCl + NaOH back-titration	n(CaCO₃) = n(HCl)_reacted/2	Phenolphthalein (back-titration step)
Antacid base test	Dissolved antacid	pH probe	pH > 7 confirms basic; cannot quantify CaCO₃	pH probe

Must DoIn a practical report comparing the digital probe method and the titration method for vinegar analysis, the key points are: (1) both methods give the acetic acid concentration; (2) titration is more precise (±0.1% vs ±2–5% for pH probe calculation); (3) pH probe is faster and gives continuous readings but requires an ICE table reverse calculation with Ka (introducing Ka uncertainty); (4) for regulatory purposes (food labelling compliance), titration is the accepted standard method.

Common ErrorStudents say "both methods give the same precision because they both give a number." Precision is not determined by getting a numerical answer — it is determined by the uncertainty of each step in the measurement. A pH probe calibrated to ±0.01 pH units gives [H⁺] with ~2.3% uncertainty; this propagates through the Ka-based reverse calculation to give c(acid) with ~5–10% uncertainty at best. A titration with concordant titres to ±0.05 mL gives concentration with ~0.1–0.3% uncertainty.

Worked Example 1 — Vinegar analysis by two methods

A student analyses white vinegar (density = 1.005 g/mL) using two methods. Method A: they pipette 10.00 mL of vinegar and titrate with 0.5000 mol/L NaOH, obtaining concordant titres of 18.55, 18.60, 18.60 mL. Method B: they measure pH of undiluted vinegar with a calibrated probe and read pH 2.40. Ka(CH₃COOH) = 1.8 × 10⁻⁵.
(a) Calculate the % acetic acid by mass from Method A.
(b) Calculate the approximate % acetic acid from Method B using the reverse Ka calculation.
(c) Compare the two results and explain which is more reliable for regulatory compliance.

(a) Method A (Titration):
Average titre = (18.55 + 18.60 + 18.60)/3 = 18.58 mL = 0.01858 L.
n(NaOH) = 0.5000 × 0.01858 = 9.29 × 10⁻³ mol.
n(CH₃COOH) = n(NaOH) = 9.29 × 10⁻³ mol (1:1 ratio).
mass(CH₃COOH) = 9.29 × 10⁻³ × 60.06 = 0.558 g.
mass(vinegar) = 10.00 mL × 1.005 g/mL = 10.05 g.
% CH₃COOH = (0.558 / 10.05) × 100% = 5.55%.

(b) Method B (pH Probe):
[H⁺] = 10⁻²·⁴⁰ = 3.98 × 10⁻³ mol/L.
Using Ka = [H⁺]² / (c − [H⁺]), assuming c ≈ [H⁺]² / Ka + [H⁺]:
c = (3.98 × 10⁻³)² / (1.8 × 10⁻⁵) + 3.98 × 10⁻³ = 0.880 + 0.00398 = 0.884 mol/L.
mass(CH₃COOH) in 10.00 mL = 0.884 × 0.01000 × 60.06 = 0.531 g.
% = (0.531 / 10.05) × 100% = 5.28%.

(c) Comparison:
Method A (titration) gives a more reliable result for regulatory compliance. Concordant titres to ±0.05 mL give concentration uncertainty of ~0.3%. Method B relies on a pH reading propagated through the Ka reverse calculation, introducing Ka uncertainty (~5–10% overall). Food labelling regulations specify titratable acidity as the accepted method of measurement.

Answer: (a) 5.55%. (b) 5.28%. (c) Titration (Method A) is more reliable due to lower uncertainty (~0.3% vs ~5–10%) and direct stoichiometric measurement.

Worked Example 2 — Intermediate: Comparing analytical methods for a pharmaceutical tablet

A pharmaceutical chemist analyses aspirin tablets (acetylsalicylic acid, Ka = 3.0 × 10⁻⁴, M = 180.2 g/mol, labelled 300 mg per tablet). Three analytical methods are available: (A) dissolve tablet in ethanol-water and titrate with 0.1000 mol/L NaOH using phenolphthalein; (B) dissolve tablet in water and measure pH with a calibrated probe; (C) perform a conductometric titration of the dissolved tablet with NaOH.
(a) Describe one specific advantage of Method A over Method B for this application.
(b) Describe one specific advantage of Method C over Method A for this application.
(c) The student finds that some aspirin has hydrolysed to acetic acid and salicylic acid during storage. Would this hydrolysis cause the Method A titration to overestimate or underestimate the aspirin content, and why?

(a) Method A advantage over B:
Method A (titration) directly measures the moles of acetylsalicylic acid by stoichiometric reaction with NaOH, giving concentration to ±0.1–0.3% uncertainty. Method B (pH probe) gives pH from which c can be estimated via the reverse Ka calculation — but this introduces ~5–10% uncertainty from Ka imprecision and temperature effects. For pharmaceutical quality control (purity ≥99.0%), Method A is more precise and appropriate for regulatory compliance.

(b) Method C advantage over A:
Method C (conductometric titration) does not require an indicator. Aspirin solutions can be slightly turbid (aspirin has limited solubility in water) — a turbid solution may obscure the phenolphthalein colour change in Method A. Conductometric titration uses electrical measurement rather than visual observation, and is unaffected by solution turbidity or colour.

(c) Effect of aspirin hydrolysis:
Hydrolysis reaction: aspirin (monoprotic) → acetic acid (monoprotic) + salicylic acid (monoprotic). One mole of aspirin produces one mole of acetic acid AND one mole of salicylic acid. If 1 mole of aspirin hydrolyses, it gives 2 moles of acid products that both react with NaOH. The titration measures total titratable acid — it cannot distinguish between intact aspirin and its hydrolysis products. A hydrolysed sample consumes MORE NaOH than the same mass of intact aspirin would. The calculated n(aspirin) = n(NaOH) overestimates the true amount of intact aspirin.

Answer: (a) Method A gives higher precision (~0.3% vs 5–10%) required for pharmaceutical purity. (b) Method C is unaffected by turbidity/colour. (c) Hydrolysis produces two monoprotic acids per aspirin molecule → titre is too large → Method A overestimates intact aspirin content.

Worked Example 3 — Hard: Extended Response Design

(8 marks) A Year 12 student is asked to determine the acetic acid content of three different commercial vinegar brands and report their results to a food safety authority. They have access to a calibrated glass electrode pH probe, a standard 0.5000 mol/L NaOH solution, a conductance meter, and phenolphthalein indicator.
(a) Recommend the most appropriate analytical method for this application and justify your choice over the two alternatives.
(b) Describe in detail how the glass electrode pH probe works, including the role of the glass membrane, the Nernst equation, and the calibration step.

(a) Method recommendation:
Recommended method: direct NaOH indicator titration with phenolphthalein. Justification over pH probe: the food safety report requires a precise, quantitative result (±0.1%). The pH probe calculation method introduces uncertainty from Ka variation (~5–10% total uncertainty) — unacceptable for regulatory reporting. Justification over conductometric titration: conductometric titration requires more equipment and more time per sample. For three brands, direct titration with visual endpoint is faster, simpler, and equally precise because commercial white vinegar is pale and can be diluted to allow phenolphthalein endpoint detection.

(b) Glass electrode mechanism:
The glass electrode consists of a thin glass membrane selectively permeable to H⁺ ions. The inner surface is in contact with a reference solution of known, fixed [H⁺]. The outer surface is immersed in the test solution. H⁺ ions exchange into the hydrated gel layers on both sides. Because [H⁺] differs on the two sides, a potential difference (voltage, E) develops across the membrane. This is described by the Nernst equation: at 25°C, E ≈ E° − 0.0592 × pH. The meter's electronics measure this voltage and convert it to a pH value. This conversion requires calibration — two buffer solutions of known pH (e.g. pH 4.00 and pH 7.00) are used to establish the slope and intercept of the linear voltage-pH relationship for this specific probe at the current temperature, correcting for electrode aging and temperature drift.

Answer: Full 8-mark response covering method justification (titration is more precise than probe, faster than conductometric) and the electrochemical mechanism of the glass electrode (membrane exchange, Nernst voltage, two-point calibration).

✅ Comprehensive Answers

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🔬 Activity 1 — Method Selection

1. Indicator Titration: High precision (~0.3%) is ideal for determining exact concentration of a clear, colourless, fast-reacting solution.

2. Direct pH Probe: Allows for continuous, automated, 24/7 electrical monitoring without consuming reagents.

3. Conductometric Titration: Weak acid + weak base titrations do not have a sharp pH jump, so indicators fail. Conductance provides an objective endpoint (minimum gradient/intersection).

4. Back Titration: CaCO₃ is insoluble and reacts slowly. It must be dissolved in excess acid first to ensure complete reaction and a clean endpoint.

⚙️ Activity 2 — Calibration Sequence

1. Rinse electrode with distilled water and blot dry. 2. Immerse in pH 4.00 buffer, wait for stability, and set meter. This adjusts the intercept (offset). 3. Rinse and blot dry. 4. Immerse in pH 7.00 buffer, wait for stability, and set meter. This adjusts the slope (sensitivity, mV per pH unit). This two-point process corrects for electrode aging and temperature drift.

❓ Multiple Choice

1. D — ΔV = −185 − (−155) = −30 mV. ΔpH = −30 / −59.2 = +0.51. pH = 4.00 + 0.51 = 4.51.

2. C — Direct titration of NaHCO₃ produces CO₂ gas, which causes bubbling and obscures the indicator. Back titration avoids this.

3. B — The dark red colour of wine masks the faint pink phenolphthalein endpoint. Conductance is unaffected by colour.

4. C — Two points are mathematically required to define a line (slope and intercept), correcting for physical changes in the electrode.

5. D — Titration is a direct stoichiometric measurement with high precision. The pH probe method relies on Ka, which introduces significant uncertainty.

📝 Short Answer Model Answers

Q6 (4 marks): The glass membrane is selectively permeable to H⁺ ions. [1] A potential difference (voltage) develops across the membrane proportional to the difference in [H⁺] between the test solution and the inner reference solution. [1] The Nernst equation (E ≈ E° − 0.0592 × pH) mathematically relates this voltage to the pH. [1] The meter's electronics measure the voltage and convert it to a pH reading using a calibrated slope and intercept. [1]

Q7 (4 marks): Concordant titres: 16.40, 16.35, 16.45 (16.90 is excluded). Average = 16.40 mL. [1]
n(NaOH) = 0.5000 × 0.01640 = 8.200 × 10⁻³ mol. [1]
n(CH₃COOH) = 8.200 × 10⁻³ mol (1:1 ratio).
mass(CH₃COOH) = 8.200 × 10⁻³ × 60.06 = 0.4925 g. [1]
mass(vinegar) = 10.00 mL × 1.005 g/mL = 10.05 g.
% CH₃COOH = (0.4925 / 10.05) × 100% = 4.90%. [1]

Q8 (3 marks): Hydrolysis of one mole of aspirin (monoprotic) produces one mole of acetic acid and one mole of salicylic acid (both monoprotic). [1] This means 1 mole of original aspirin turns into 2 moles of titratable acid. [1] Because the sample consumes twice as much NaOH per hydrolysed molecule, the calculated aspirin content will be overestimated. [1]

Acid/Base Analysis Techniques — Industrial & Digital

Download this lesson's worksheet

Key Relationships — This Lesson

📖 Know

💡 Understand

✅ Can Do

Calibration — Why It Is Mandatory and How It Is Done

Comparing Methods — When to Use Each Technique

pH probe (direct)

Indicator titration

Back titration

Conductometric

Industrial Applications — Acid/Base Analysis Across Sectors

Household Substance Analysis — Vinegar and Antacids

⚠️ Common Misconceptions — Module 6 Lesson 19

📓 Copy Into Your Books

Glass Electrode pH Meter

Calibration (Mandatory)

Method Comparison

Industrial Applications

Match the Method to the Scenario

The Calibration Process

Test Your Knowledge

Extended Questions

Revisit Your Thinking

✅ Comprehensive Answers

🔬 Activity 1 — Method Selection

⚙️ Activity 2 — Calibration Sequence

❓ Multiple Choice

📝 Short Answer Model Answers

Mark lesson as complete

Questions