Investigation involving quantitative analysis
Internal 4 credits Titration Moles Uncertainty MethodA practical internal: you carry out a quantitative analysis (usually a titration) and process the data with proper mole calculations and uncertainty. Merit/Excellence comes from precise technique and honest evaluation of reliability.
1 Reading the notation
Titration and gravimetric work is mostly units. Getting these straight is most of the accuracy.Symbols and units in quantitative analysis
9 termsCover the right column and check you can say what each one means.
- n
- Number of moles (mol). n = c × V for a solution, n = m ÷ M for a solid.
- c
- Concentration in mol L⁻¹ (moles per litre), sometimes written M. If a volume is given in mL you MUST convert to litres before using n = cV.
- V
- Volume. In n = cV it must be in litres, not millilitres.
- m and M
- Lower-case m = mass in grams. Capital M = molar mass in g mol⁻¹. n = m ÷ M. Case matters: m and M are different quantities.
- mol L⁻¹ / g mol⁻¹
- The ⁻¹ means “per”. mol L⁻¹ = moles per litre; g mol⁻¹ = grams per mole.
- ± (plus-or-minus)
- The absolute uncertainty on a measurement. 25.00 ± 0.05 mL means the true value lies between 24.95 and 25.05.
- Concordant
- Titres agreeing within 0.10 mL of each other. Only concordant titres are averaged.
- Aliquot / titre
- Aliquot = the measured volume pipetted into the flask. Titre = the volume delivered from the burette.
- 3 s.f. (significant figures)
- How many digits carry real information. Match the least precise measurement; quote uncertainty to one significant figure.
The unit trap
n = cV needs LITRES. A burette reads millilitres. Dividing by 1000 at the right moment is the single most common place marks are lost in this standard.2 How this connects to your other standards
This internal is where the theory becomes something you physically do, and it borrows from almost every other standard.How this connects
What this internal draws on:
Aqueous equilibria (91392): the theory behind your titration Your titration curve, equivalence point and indicator choice are all 91392 content. Titrating a weak acid with a strong base gives a BASIC equivalence point (pH > 7), which is why phenolphthalein is correct and methyl orange is not. Do the theory and you stop guessing indicators. Spectroscopic data (91388): the other analysis internal Two halves of analytical chemistry: this one measures HOW MUCH (concentration), 91388 identifies WHAT (structure). Both require careful evidence, uncertainty and evaluation. If your unknown needs identifying first, colorimetry and spectroscopy overlap directly. Redox (91393): if you do a redox titration MnO₄⁻ (self-indicating, first permanent pink) or iodine/thiosulfate titrations need balanced half-equations to get the mole ratio right. Get the half-equation wrong and every calculation after it is wrong. Organic compounds (91391): what you are often titrating Ethanoic acid, citric acid and aspirin are common analytes. Knowing which proton is acidic (–COOH) tells you the mole ratio per molecule. Vital when a diprotic or triprotic acid appears.3 Core mole calculations
Watch: Core mole calculations Opens a YouTube search: pick a video that clicks for youThe relationships you’ll use constantly
Moles from concentration n=cV V in litres. c in mol L⁻¹. Moles from mass n=\frac{m}{M} M = molar mass (g mol⁻¹). Titration link \frac{n(\text{A})}{n(\text{B})}=\text{ratio from the balanced equation} Use the equation’s coefficients to convert between the two reactants. Worked example A standard titration calculation 25.0 mL of NaOH is neutralised by 22.4 mL of 0.100 mol L⁻¹ HCl. Find [NaOH]. n(HCl) = c×V = 0.100 × 0.0224 = 2.24×10⁻³ mol. HCl + NaOH → NaCl + H₂O, ratio 1:1, so n(NaOH) = 2.24×10⁻³ mol. c(NaOH) = n/V = 2.24×10⁻³ / 0.0250 = 0.0896 mol L⁻¹. ✓ [NaOH] = 0.0896 mol L⁻¹ (≈ 0.090 mol L⁻¹).4 Titration technique (for Merit/Excellence)
Good practice that earns marks
- Rinse burette with the titrant, pipette with the analyte (avoids dilution errors).
- Do a rough titration first, then concordant titres (within 0.1 mL) and average only those.
- Read the burette to the bottom of the meniscus at eye level (0.05 mL precision).
- Swirl constantly; add dropwise near the endpoint; the endpoint is the first permanent colour change.
- Use an appropriate indicator for the acid/base strength (e.g. phenolphthalein for strong base–weak acid).
5 Uncertainty & evaluation
Combining uncertainties (for × and ÷) % uncertainty of result = Σ (% uncertainties of each measurement) Add the relative (%) uncertainties of every measured quantity. Absolute from percent \text{absolute}=\frac{\%}{100}\times\text{value} Report the final answer as value ± absolute uncertainty.Evaluating reliability
- Identify the largest source of uncertainty (often the smallest volume measured) and how to reduce it.
- Distinguish random errors (scatter, reduce by repeats) from systematic errors (bias, e.g. uncalibrated glassware).
- Comment on concordance of titres as evidence of precision.
Colorimetric / gravimetric options
The internal may instead use colorimetry (absorbance ∝ concentration, calibration curve, Beer–Lambert) or gravimetric analysis (precipitate, filter, dry to constant mass). The mole reasoning is the same. Measure a quantity, convert to moles, use the ratio.Practice
Recall it first, then use it — testing yourself builds the memory, where rereading the notes above only feels easier.
Recall and application shuffled together — you have to work out what each item wants before you can answer it, which is the part the exam actually tests. Both kinds feed the same spacing schedule, so anything you miss comes back sooner.
Say the answer out loud before flipping — that attempt is where the learning happens, even when you get it wrong.
Exam-style only. Instant feedback, shuffled each attempt, best score saved on this device.