The titration procedure
Titration is used to find the exact volume of one solution that reacts completely with a known volume of another solution. In acid-base titration:
- 1Fill the burette with the acid (or alkali). Record the initial burette reading to 2 decimal places (e.g. 0.00 cm³).
- 2Pipette exactly 25.0 cm³ of the alkali (or acid) into a conical flask. Add 2–3 drops of indicator (phenolphthalein or methyl orange).
- 3Add the burette solution slowly, swirling constantly. Near the endpoint, add dropwise.
- 4Stop at the endpoint — the indicator changes colour permanently. Record the final burette reading.
- 5Calculate the titre = final reading − initial reading. Repeat until two concordant results (within 0.10 cm³ of each other). Average the concordant results.
Using all titres to find the average (should use only concordant ones). Not reading the burette to 2 decimal places. Using the wrong indicator — phenolphthalein for strong acid + strong alkali; methyl orange also acceptable. Not rinsing the burette with the solution to be used (causes dilution error).
- Moles HCl = c × V = 0.100 × (22.4/1000) = 0.100 × 0.0224 = 0.00224 mol [1]
- Mole ratio HCl : NaOH = 1 : 1 (from equation). So moles NaOH = 0.00224 mol [1]
- Concentration NaOH = moles / V = 0.00224 / (25.0/1000) = 0.00224 / 0.025 [1]
- = 0.0896 mol/dm³ ≈ 0.0896 mol/dm³ [1]
- (a) Concordant results: 18.60, 18.55, 18.50 cm³ (all within 0.10 cm³ of each other). 18.90 is anomalous — exclude it. [1] Mean titre = (18.60 + 18.55 + 18.50) / 3 = 55.65 / 3 = 18.55 cm³ [1]
- (b) Moles NaOH = 0.200 × 0.0250 = 0.00500 mol [1]
- Mole ratio H₂SO₄ : NaOH = 1 : 2. Moles H₂SO₄ = 0.00500 / 2 = 0.00250 mol [1]
- Concentration H₂SO₄ = 0.00250 / (18.55/1000) = 0.00250 / 0.01855 = 0.135 mol/dm³ [1]
- Moles Ca(OH)₂ = 0.150 × 0.0400 = 0.00600 mol [1]
- Mole ratio HNO₃ : Ca(OH)₂ = 2 : 1. Moles HNO₃ = 2 × 0.00600 = 0.0120 mol [1]
- Volume HNO₃ = moles / c = 0.0120 / 0.250 = 0.0480 dm³ = 48.0 cm³ [1]
The planning framework — always use this structure
- 1State the independent variable (IV) — what you deliberately change. Be specific: "the concentration of HCl (in mol/dm³)" not "the acid."
- 2State the dependent variable (DV) — what you measure. Include HOW you measure it: "the volume of CO₂ collected in a gas syringe (cm³)" not "the amount of gas."
- 3State at least two controlled variables (CVs) — what you keep constant AND how: "the mass of CaCO₃ (5.0 g, weighed on a balance)" and "the temperature of the acid (maintained at 25°C using a water bath)."
- 4Describe the method in steps — set up apparatus, take measurements, repeat for each value of IV.
- 5State a safety precaution relevant to the specific chemicals — not just "wear goggles" (too generic). E.g. "wear safety goggles to protect eyes from acid splashes."
- 6State how to improve reliability — repeat each measurement and take an average; use a data logger; exclude anomalous results.
Always identify the specific source: "parallax error when reading the burette meniscus" or "timing error when starting the stopwatch at the moment of mixing" — vague errors earn no marks.
IV examples (specific)
"The concentration of hydrochloric acid (0.5, 1.0, 1.5, 2.0, 2.5 mol/dm³)"
"The temperature of the reaction mixture (20, 30, 40, 50, 60 °C)"
"The surface area of zinc (fine powder, small granules, large pieces)"
DV examples (specific)
"The volume of hydrogen gas collected in a gas syringe after 60 s (cm³)"
"The time for the cross to disappear through the sulfur precipitate (s)"
"The mass loss of the flask + contents every 30 s (g)"
CV examples (with control method)
"The volume of acid (25.0 cm³, measured with a pipette)"
"The mass of zinc (1.0 g, measured on a top-pan balance)"
"The concentration of acid (kept at 1.0 mol/dm³ throughout)"
Safety examples (specific)
"Wear safety goggles — HCl is corrosive and could damage eyes if it splashes"
"Work in a fume cupboard when using concentrated H₂SO₄ — fumes are toxic"
"Handle hot apparatus with tongs — risk of burns during heating"
- IV: Temperature of the sodium thiosulfate solution (e.g. 20, 30, 40, 50, 60 °C) [1]
- DV: Time taken for the cross (drawn on paper beneath the flask) to become invisible due to sulfur precipitate formation — measured with a stopwatch (s). Rate ∝ 1/time. [1]
- CVs: (i) Concentration of Na₂S₂O₃ (0.1 mol/dm³ throughout, measured with a volumetric flask). (ii) Volume of Na₂S₂O₃ (25 cm³, measured with a measuring cylinder). (iii) Concentration and volume of HCl (any stated constant value). [1 for any two with control method]
- Method: Heat 25 cm³ Na₂S₂O₃ to the required temperature in a water bath. Place the flask over a paper cross. Add 5 cm³ HCl quickly, start the stopwatch immediately. Stop timing when the cross is no longer visible from above. Record the time. Repeat at each temperature. [1]
- Repeat for reliability: Repeat each temperature measurement twice and average the times. [1]
- Safety: Wear safety goggles — HCl is corrosive; carry out in a well-ventilated area as SO₂ gas may be produced. [1]
Rules for results tables
- 1Column headers must include quantity AND unit: "Temperature / °C" or "Volume of CO₂ / cm³". Never just "Temperature" or "Volume."
- 2All readings to consistent decimal places: if burette reads to 0.05 cm³, record as 18.50, 18.55 — not 18.5 or 18.
- 3Calculated columns: show the formula used and results to appropriate significant figures — usually 3 s.f. for rate calculations.
- 4Identify anomalous results: circle them in the table and exclude from averaging/graphing.
Rate = amount of product / time (e.g. cm³/s or g/s). OR Rate = 1 / time (when measuring time for a fixed event — e.g. cross disappearance). Units of 1/time = s⁻¹.
| [HCl] (mol/dm³) | 0.5 | 1.0 | 1.5 | 2.0 | 2.5 |
|---|---|---|---|---|---|
| Volume CO₂ in 60 s (cm³) | 12 | 26 | 35 | 72 | 62 |
- (a) The anomalous result is 72 cm³ at 2.0 mol/dm³ — it is much higher than the trend and higher than the next concentration (2.5 mol/dm³). It should be circled and excluded. [1]
- (b) Rate = volume / time = volume / 60. At 1.0: 26/60 = 0.433 cm³/s. At 1.5: 35/60 = 0.583 cm³/s. At 2.5: 62/60 = 1.033 cm³/s. [1 for correct method; 1 for correct values]
- (c) As HCl concentration increases, the rate of CO₂ production increases (excluding the anomalous value at 2.0 mol/dm³). [1 — must quote data or reference trend]
Types of errors in chemistry practicals
Measurement errors (random)
Parallax error reading a meniscus (read at eye level to avoid). Reaction time error starting/stopping a stopwatch. Variable timing — use a data logger instead.
Systematic errors
Zero error on a balance (tare before use). Using contaminated glassware (rinse with the solution to be used). Incomplete reaction (wait longer or warm gently). Heat loss to surroundings in calorimetry (insulate the flask).
Validity issues
Not controlling a key variable (e.g. temperature changes during rate experiment — use a water bath). Using impure reactants — affects concentration. Not rinsing the burette (dilutes the solution).
Reliability improvements
Repeat measurements and average. Use automated equipment (data logger, colorimeter). Take more data points. Use larger volumes for greater precision. Calibrate instruments before use.
For "suggest an improvement" questions: always write as PROBLEM → SOLUTION. "The temperature of the solution varied during the experiment (problem); use a water bath set to a constant temperature to maintain 25°C throughout (solution)." Never just state the problem without the fix, or the fix without explaining what it corrects.
- Error 1: Heat loss to the surroundings through the cup and to the thermometer. The measured temperature rise is less than the true value. Improvement: Place a cardboard lid on the cup and wrap the cup in insulating material (cotton wool) to reduce heat loss to the environment. [2]
- Error 2: The thermometer may not have been stirred throughout — the reading may not reflect the maximum temperature accurately. Improvement: Stir continuously with the thermometer and record the temperature every 30 s to capture the true maximum. [2]
- The likely cause is a zero error (systematic error) — air was already present in the gas syringe before the reaction began, or the reaction had already started before she began timing. [1]
- This makes all volume readings 2 cm³ too high (a constant offset), so the total volume of CO₂ collected appears greater than it actually is. The rate calculation would also be slightly overestimated if uncorrected. The student should subtract 2 cm³ from every reading. [1]
Gas tests
| Gas | Test | Positive result |
|---|---|---|
| Hydrogen (H₂) | Hold a lit splint near the tube mouth | Squeaky pop / burns with a small explosion |
| Oxygen (O₂) | Insert a glowing splint | Glowing splint relights |
| Carbon dioxide (CO₂) | Bubble through limewater (Ca(OH)₂ solution) | Limewater turns milky/cloudy |
| Chlorine (Cl₂) | Hold damp litmus paper in the gas | Turns red then bleaches white |
| Ammonia (NH₃) | Hold damp red litmus paper in the gas | Turns blue (alkaline); pungent smell |
| Sulfur dioxide (SO₂) | Bubble through acidified potassium manganate(VII) | Purple/pink decolourises to colourless |
Cation tests — adding NaOH solution
| Ion | Precipitate colour | Behaviour with excess NaOH |
|---|---|---|
| Cu²⁺ (copper) | Blue precipitate Cu(OH)₂ | Insoluble in excess NaOH |
| Fe²⁺ (iron II) | Green precipitate Fe(OH)₂ | Insoluble; may turn brown on exposure to air |
| Fe³⁺ (iron III) | Brown/rust precipitate Fe(OH)₃ | Insoluble in excess NaOH |
| Zn²⁺ (zinc) | White precipitate Zn(OH)₂ | Dissolves in excess NaOH (amphoteric) |
| Al³⁺ (aluminium) | White precipitate Al(OH)₃ | Dissolves in excess NaOH (amphoteric) |
| Ca²⁺ (calcium) | White precipitate Ca(OH)₂ (may be slight) | Partially soluble — precipitate may dissolve slightly |
| NH₄⁺ (ammonium) | No precipitate | On warming: ammonia gas produced (pungent, turns damp red litmus blue) |
Anion tests
| Ion | Test reagent | Positive result |
|---|---|---|
| Cl⁻ (chloride) | Add dilute HNO₃ then AgNO₃(aq) | White precipitate (AgCl) |
| Br⁻ (bromide) | Add dilute HNO₃ then AgNO₃(aq) | Cream precipitate (AgBr) |
| I⁻ (iodide) | Add dilute HNO₃ then AgNO₃(aq) | Yellow precipitate (AgI) |
| SO₄²⁻ (sulfate) | Add dilute HCl then BaCl₂(aq) | White precipitate (BaSO₄), insoluble in HCl |
| CO₃²⁻ (carbonate) | Add dilute HCl; bubble gas through limewater | Effervescence; limewater turns milky (CO₂) |
| NO₃⁻ (nitrate) | Add NaOH + aluminium foil, warm gently | Ammonia gas produced (turns damp red litmus blue) |
Always acidify the solution with dilute acid before adding AgNO₃ (for halides) or BaCl₂ (for sulfates). This removes carbonate ions (CO₃²⁻) that would give false white precipitates. Use HNO₃ for halide tests (NOT HCl — would introduce Cl⁻). Use HCl for sulfate tests (NOT H₂SO₄ — would introduce SO₄²⁻).
- Test (i): white ppt with NaOH that dissolves in excess → cation is Zn²⁺ or Al³⁺ (both amphoteric). [1]
- Test (ii): cream ppt with AgNO₃ after HNO₃ → anion is Br⁻ (bromide). (White = Cl⁻; cream = Br⁻; yellow = I⁻). [1]
- Test (iii): no flame colour → not Na, K, Li, Ca, Cu. Combined with above: the solid is most likely zinc bromide (ZnBr₂) or aluminium bromide (AlBr₃). Zinc is more likely as Al³⁺ bromide is unusual. [1]