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Analytical chemistry can feel like a long list of unrelated techniques: titrations, chromatography, half a dozen kinds of spectroscopy. It becomes much easier once you see the pattern: every technique answers what is it?, how much is there?, or both, and every technique relies on a property that differs between substances. This guide organises the subject into six units, gives the core ideas and checkpoints for each, and links to the full articles.
How to use this guide
- Read the core ideas for a unit.
- Try the checkpoint questions without notes.
- For anything you can’t answer confidently, read the linked article and try again.
- Use the technique-selection table near the end to practise choosing methods.
Unit 1: Foundations and vocabulary
Core ideas
- Qualitative analysis identifies substances; quantitative analysis measures amounts. See qualitative vs quantitative analysis.
- Accuracy is closeness to the true value; precision is closeness of repeated results to each other.
- Random errors scatter results; systematic errors shift them all in one direction. See experimental errors.
- The analyte is the substance being measured; the matrix is everything else in the sample.
- A representative sample is essential. See sampling.
Checkpoint
- A balance that reads 0.02 g too high on every weighing causes which type of error?
- Three titres of 24.10, 24.15 and 24.10 cm³ are precise. Does that prove they’re accurate?
Unit 2: Volumetric and gravimetric analysis
Core ideas
- A titration measures the volume of a standard solution that reacts exactly with the analyte. See volumetric analysis.
- The calculation always runs: moles of titrant → mole ratio → moles of analyte → concentration or mass. See titration calculations.
- Four main types: acid–base, redox, complexometric (EDTA) and precipitation. See EDTA titrations and precipitation titrations.
- A back titration adds a known excess and titrates what’s left.
- End points can be found instrumentally: potentiometric and conductometric titrations.
- Gravimetric analysis finds an amount by weighing a pure product of known formula.
- Glassware choice matters: see measuring volume accurately and volumetric glassware tolerances.
Checkpoint 3. 25.00 cm³ of NaOH needs 20.00 cm³ of 0.100 mol/dm³ HCl. Find [NaOH]. 4. Why can’t sodium hydroxide be used as a primary standard? 5. In a back titration, what quantity do you subtract from what?
Unit 3: Chromatography
Core ideas
- All chromatography separates by distribution between a stationary phase and a mobile phase. See chromatography.
- Rf = distance moved by spot ÷ distance moved by solvent front. See Rf value calculations.
- Paper and thin-layer chromatography are quick and qualitative.
- Column chromatography separates larger amounts for purification.
- Gas chromatography suits volatile compounds; HPLC suits non-volatile and heat-sensitive ones.
- Components are identified by retention time and quantified by peak area.
Checkpoint 6. A spot moves 3.6 cm while the solvent moves 8.0 cm. Calculate Rf. 7. Why would you choose HPLC rather than GC for a protein?
Unit 4: Spectroscopy
Core ideas
- Different regions of the electromagnetic spectrum cause different changes in molecules. See the spectroscopy overview.
- UV-visible and colorimetry measure concentration using the Beer–Lambert law: A = εcl.
- Infrared spectroscopy identifies functional groups (O–H broad around 3200–3550 cm⁻¹; C=O strong around 1700 cm⁻¹).
- NMR shows hydrogen and carbon environments, their ratios and their neighbours (n + 1 rule).
- Raman spectroscopy complements IR.
- Atomic methods measure elements: AAS, ICP-MS, XRF and flame tests.
- Calibration curves convert instrument signals into concentrations.
Checkpoint 8. A solution has absorbance 0.450 in a 1.00 cm cell; ε = 1500 dm³ mol⁻¹ cm⁻¹. Find c. 9. An IR spectrum shows a strong band at 1715 cm⁻¹ and no broad band above 3000 cm⁻¹. Which group is likely present, and which is ruled out?
Unit 5: Mass spectrometry
Core ideas
- A mass spectrometer ionises, separates by m/z, and detects. See mass spectrometry.
- For elements, peak heights give isotope abundances and relative atomic mass. See mass spectrometry and isotopes.
- For molecules, the molecular ion gives relative molecular mass; fragments give structural clues.
- M+2 peaks reveal chlorine (3 : 1) or bromine (1 : 1).
Checkpoint 10. An element has isotopes of mass 63 (69.2%) and 65 (30.8%). Calculate its relative atomic mass.
Unit 6: Putting it together
Core ideas
- Real problems combine techniques. See how chemists identify an unknown compound and elemental analysis.
- Inorganic ions are identified by a logical sequence of tests. See identifying ions.
- Applications: forensics, water testing, food analysis, drug testing and art authentication.
Checkpoint 11. A compound has M⁺ at 88 and empirical formula C₂H₄O. What’s its molecular formula? 12. Which technique would you use to show that a white solid contains sulfate ions?
Technique-selection table
| Question | Suitable technique(s) |
|---|---|
| How much acid is in this vinegar? | acid–base titration |
| Which dyes are in this ink? | paper or thin-layer chromatography |
| How much caffeine is in this drink? | HPLC with UV detection |
| Is there ethanol in this blood sample, and how much? | headspace gas chromatography |
| Which functional groups does this compound have? | infrared spectroscopy |
| How are the atoms connected? | NMR |
| What is the molecular mass? | mass spectrometry |
| How much lead is in this water? | ICP-MS or AAS |
| What is the concentration of this coloured solution? | colorimetry / UV-visible with a calibration curve |
| Is this solid pure? | melting point, TLC |
| How hard is this water? | EDTA titration |
| What’s the empirical formula of this new compound? | CHN elemental analysis |
For practice at choosing and justifying, cover the right-hand column and answer from memory.
Checkpoint answers
- Systematic error.
- No. Precise results can all be wrong by the same amount if there’s a systematic error.
- n(HCl) = 0.100 × 0.02000 = 2.00 × 10⁻³ mol; [NaOH] = 2.00 × 10⁻³ ÷ 0.02500 = 0.0800 mol/dm³.
- It absorbs water and carbon dioxide from the air, so its mass doesn’t correspond to a known amount of NaOH.
- Moles of reagent left over (from the titration) are subtracted from moles of reagent originally added.
- Rf = 3.6 ÷ 8.0 = 0.45.
- Proteins are large, non-volatile and heat-sensitive; GC requires the sample to be vaporised.
- c = A ÷ (εl) = 0.450 ÷ 1500 = 3.00 × 10⁻⁴ mol/dm³.
- A C=O group (ketone or aldehyde likely); an O–H group, and so a carboxylic acid or alcohol, is ruled out.
- (63 × 69.2 + 65 × 30.8) ÷ 100 = 63.6 (copper).
- 88 ÷ 44 = 2, so C₄H₈O₂.
- Dissolve, add dilute HCl then barium chloride solution; a white precipitate shows sulfate.
A two-week revision plan
| Day | Focus |
|---|---|
| 1 | Unit 1 vocabulary; errors and uncertainty |
| 2–3 | Titration calculations, including back titrations |
| 4 | Redox, EDTA and precipitation titrations |
| 5 | Gravimetric analysis and elemental analysis |
| 6 | Chromatography and Rf values |
| 7 | Review days 1–6; redo missed checkpoints |
| 8 | UV-visible, colorimetry and Beer–Lambert |
| 9 | IR spectroscopy |
| 10 | NMR |
| 11 | Mass spectrometry |
| 12 | Combined spectroscopy problems |
| 13 | Qualitative ion tests |
| 14 | Mixed past-paper practice and the technique-selection table |
Short, spaced sessions with active recall work better than long rereading sessions.
Key takeaways
- Every technique answers “what?”, “how much?” or both, using a property that differs between substances.
- Master the titration calculation pattern; it appears in every volumetric method.
- Learn Rf, Beer–Lambert and relative atomic mass calculations until they’re automatic.
- Know which technique suits which question, and why.
- Test yourself regularly; reading alone isn’t revision.
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