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Every claim a chemist makes about bonding rests, ultimately, on an experimental observation. We say that solid sodium chloride is ionic because we can melt it and watch electricity flow through the liquid; we say that ice contains discrete H₂O molecules held together by hydrogen bonds because X-ray diffraction shows us the geometric pattern of oxygen atoms; we say a compound is an alcohol because its infrared spectrum carries the broad O-H stretch characteristic of that functional group. This lesson catalogues the experimental toolkit that turns models into evidence-based statements. Electrolysis (Davy, Faraday) established that ions are real, separable entities. X-ray crystallography (the Braggs, father and son, 1913; Nobel 1915 — names only) revealed lattice geometry and let chemists measure ionic and covalent radii directly. Mass spectrometry pins down molecular formulae and isotope patterns. Infrared spectroscopy identifies functional groups via characteristic bond-stretching frequencies. NMR — the heavyweight structural-connectivity technique — is signposted here and developed in full in §3.3.15.
Spec mapping (AQA 7405): This lesson maps to §3.1.3 (bonding evidence and structure determination at the conceptual level). It cross-references §3.3.15 (NMR and infrared spectroscopy, developed in detail in course 10), §3.3.16 (chromatography, also course 10), §3.1.1 (mass spectrometry, originally introduced in course 1 lesson 1), and §3.2 (inorganic chemistry — Group 2 and Group 7 reactions provide chemical evidence for ionic behaviour). Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: AO1 covers recall of which technique answers which structural question (mass spec → Mᵣ and isotope pattern; IR → functional groups; X-ray → lattice geometry; NMR → connectivity). AO2 tests interpretation of a simplified spectrum or diffraction dataset given in a question stem. AO3 — the synoptic top-band skill — asks students to combine data from two or three techniques to deduce a structure, or to reason about which technique is appropriate to answer a given experimental question.
It helps to picture the experimental techniques as rungs on a ladder of increasing fineness. At the bottom rung sit bulk physical-property measurements: melting point, electrical conductivity (in the solid and the molten/aqueous states), solubility behaviour, hardness, volatility. These were covered in lesson 6. They are cheap, fast, and informative at the gross level: a substance that has a very high melting point, conducts only when molten, and dissolves in water to give a conducting solution is almost certainly ionic. A substance that melts below room temperature and does not conduct in any state is almost certainly a small covalent molecule. The bulk measurements quickly sort substances into the four big categories — ionic, simple molecular, giant covalent, metallic — but they do not tell you where the atoms are or what the bonds look like in detail.
The higher rungs of the ladder are the spectroscopic and diffraction techniques, which probe matter at the atomic and bond scale. X-ray crystallography reveals the geometric arrangement of atoms in a crystal — the lattice parameters, the bond angles, the ionic and covalent radii. Mass spectrometry pins down molecular formulae and signposts halogen content via isotope ratios. Infrared spectroscopy identifies functional groups via their bond-stretching frequencies. Nuclear magnetic resonance unpicks the connectivity of atoms inside a molecule. Each technique answers a question that bulk-property measurements cannot. Together they constitute the modern toolkit of structural chemistry, and the A-Level student is expected to know which question each one answers.
A short orienting list:
When you write Na⁺Cl⁻ for sodium chloride, you are making a strong physical claim: that the solid contains separate, oppositely-charged particles. How do we know? The most direct evidence comes from electrolysis. In 1807 Humphry Davy passed an electric current through molten potassium hydroxide and isolated potassium metal at the negative electrode (cathode). He did the same with sodium hydroxide and isolated sodium. The very existence of these metals as isolable elements was demonstrated only when an external electric field could pull positively-charged ions apart from negatively-charged ones in the molten state.
A generation later, Michael Faraday systematised the quantitative side. His laws of electrolysis (1834) state, in modern language, that the mass of element deposited at an electrode is directly proportional to the quantity of electric charge passed, and that the mass of element deposited per unit charge depends on the molar mass and the ion charge. In symbols, the moles of electrons required to deposit n moles of a singly-charged ion is exactly n; for a doubly-charged ion it is 2n. The Faraday constant F = 96 485 C mol⁻¹ is the charge on one mole of electrons. This quantitative fit between electric charge and chemical change provided overwhelming evidence that ions are discrete carriers of integer-multiple electric charge.
A simple demonstration sometimes shown in the laboratory makes the point visible. A piece of moist filter paper is laid on a glass slide and a d.c. voltage applied to its two ends via crocodile clips. A drop of copper(II) sulfate solution placed in the middle slowly migrates as a blue stain toward the cathode, while a drop of potassium permanganate placed alongside migrates as a purple stain toward the anode. The two coloured ions, [Cu(H₂O)₆]²⁺ and MnO₄⁻, move in opposite directions because they carry opposite charges. The directionality and stoichiometry of these observations would be inexplicable on a purely covalent picture of NaCl or CuSO₄; they are routine on the ionic model.
For A-Level purposes the key inference is qualitative: electrolysis evidence demonstrates that ions are real entities, not merely a bookkeeping notation. The Group 1 metals, the halogens (Cl₂ at the anode of brine electrolysis), and the transition metal cations were all first isolated or characterised via electrolytic methods. Linking this back to lesson 0 (the ionic-bond model) closes the loop: the ionic model is not an arbitrary postulate, it is the structural inference forced on us by electrochemical evidence.
The single most powerful structural technique in modern chemistry is X-ray diffraction. When a beam of monochromatic X-rays (wavelength typically 0.5–2 Å, comparable to interatomic spacings) is directed at a crystalline sample, the regular three-dimensional array of atoms scatters the X-rays in specific directions, producing a diffraction pattern. The Bragg condition, formulated by William Henry Bragg and his son William Lawrence Bragg in 1913 (Nobel Prize 1915), gives the constraint:
nλ = 2d sin θ
where λ is the X-ray wavelength, d is the spacing between planes of atoms in the crystal, θ is the angle of incidence (and reflection), and n is an integer (the order of diffraction). State this equation; you are not asked to derive it at A-Level. The crucial point is that the diffraction angles θ allow chemists to compute the interatomic spacings d directly. From a complete set of diffraction angles and intensities, the full three-dimensional arrangement of atoms in the unit cell can be reconstructed.
X-ray crystallography is the technique that gave chemistry its first direct view of atomic positions. From it, Linus Pauling and others assigned values to ionic radii (e.g. Na⁺ ≈ 102 pm, Cl⁻ ≈ 181 pm) and to covalent radii (e.g. C ≈ 77 pm, O ≈ 73 pm). The same technique revealed that diamond is a giant covalent lattice of tetrahedrally-bonded carbon atoms with C-C distance 154 pm, that graphite consists of parallel hexagonal layers with C-C distance 142 pm (and a much longer inter-layer separation of 335 pm — the empirical signature of weak London forces between layers), and that NaCl is a face-centred cubic array with Na⁺ and Cl⁻ alternating along each crystallographic axis.
Famously, X-ray fibre diffraction patterns recorded by Rosalind Franklin in 1952 (and Maurice Wilkins) provided the data from which James Watson and Francis Crick deduced the double-helix structure of DNA in 1953 (names only). The same technique, in modern single-crystal form, is the workhorse for protein structural biology and for routine identification of new organic and inorganic compounds in research laboratories.
Limitations matter at A-Level. First, X-ray crystallography requires a single crystal (or a finely crystalline powder, in the case of powder diffraction) — amorphous solids, liquids, and gases cannot be studied this way. Second, hydrogen atoms are notoriously hard to locate because X-rays are scattered by electrons, and hydrogen has only one electron — so hydrogen positions are often determined by complementary neutron diffraction or by inference from chemistry. Third, growing a high-quality single crystal can take weeks or months, and the technique is therefore not "quick" in any reasonable sense; mistaking it for a rapid analytical method is a common A-Level error.
The distinguishing power of X-ray crystallography is particularly important when comparing giant covalent and giant ionic structures. Both are high-melting and hard, and bulk properties alone may not unambiguously separate them. X-ray diffraction reveals two qualitatively different features. In a giant ionic lattice (e.g. NaCl), the electron-density map shows two clearly separated regions of high electron density centred on the cation and anion positions — there is no electron density bridging them, because the bonding is electrostatic and not shared. In a giant covalent lattice (e.g. diamond), the electron-density map shows continuous electron density bridging neighbouring atoms — the shared-electron bonds are visible directly as electron density between atomic centres.
Mass spectrometry was introduced in course 1 lesson 1 as the technique used to measure relative atomic masses and isotope abundances. At this stage of the syllabus it returns in its molecular role: mass spectrometry is used to determine the molecular formula of an unknown compound, and to provide connectivity clues via the fragmentation pattern.
A simplified picture of the experiment: the sample is vaporised and ionised (in electron-impact mode, by bombardment with high-energy electrons; in electrospray mode, by spraying a dilute solution through a charged needle). The resulting positive ions are accelerated through a potential, deflected by a magnetic field, and detected at positions determined by the mass-to-charge ratio m/z. For most A-Level purposes, the charge z = +1, so the position on the spectrum is read as the mass m of the ion in atomic mass units.
The highest m/z value in the spectrum (with finite intensity) is the molecular ion peak, M⁺•, and its mass equals the relative molecular mass Mᵣ of the compound. This is the single most useful piece of mass-spec data: it pins down a molecular formula candidate. For example, a peak at m/z = 46 is consistent with C₂H₆O (M = 46.0), CH₂O₂ (M = 46.0 — formic acid), or NO₂ (M = 46.0).
The isotope pattern sharpens the picture for halogen-containing compounds. Chlorine has two stable isotopes, ³⁵Cl and ³⁷Cl, with natural abundances approximately 75% and 25%. A compound containing one chlorine atom therefore shows a molecular-ion peak at m/z = M and a satellite peak at m/z = M + 2 with intensity ratio approximately 3:1. Bromine has two isotopes, ⁷⁹Br and ⁸¹Br, with natural abundances close to 50% each. A compound containing one bromine atom therefore shows molecular-ion peaks at m/z = M and m/z = M + 2 in approximately 1:1 ratio. The presence of an M+2 satellite with the characteristic 3:1 or 1:1 ratio is a diagnostic signature of chlorine or bromine respectively. (Iodine has only one stable isotope so gives no such pattern; fluorine likewise.)
The fragmentation pattern — peaks at m/z lower than the molecular ion — reveals connectivity. When the molecular ion is formed it carries excess internal energy and may break apart at the weaker bonds. Each fragment that retains the positive charge appears in the spectrum. For example, the molecular ion of ethanol C₂H₅OH (Mᵣ 46) loses a methyl radical (CH₃•, mass 15) to give a fragment at m/z = 31 (CH₂OH⁺), and loses water (H₂O, mass 18) to give a fragment at m/z = 28. The mass differences between observed peaks (46 − 31 = 15; 46 − 28 = 18) point directly to the chemical groups being lost.
A common discrimination problem: distinguish ethanol C₂H₅OH (Mᵣ 46) from methoxymethane CH₃OCH₃ (also Mᵣ 46). Mass spectrometry can do this because the fragmentation patterns differ. Ethanol loses 1 H to give a strong peak at m/z = 45 (CH₃CHOH⁺) and loses CH₃ to give a peak at m/z = 31. Methoxymethane fragments preferentially by C-O cleavage, giving a dominant peak at m/z = 15 (CH₃⁺) and at m/z = 29 (CHO⁺) but no strong peak at 45. The mass-spec patterns are characteristic, and the technique therefore separates structural isomers with identical molecular formulae.
A modern variant worth mentioning briefly: electrospray ionisation is a "soft" technique that does not fragment the molecule. It produces predominantly the protonated molecular ion [M+H]⁺ at m/z = Mᵣ + 1. This is the standard ionisation method for proteins and other large biological molecules, which would be shattered by electron impact. A-Level may ask only that you recognise [M+H]⁺ as the M+1 peak in an electrospray spectrum.
Covalent bonds vibrate — they stretch and bend — at frequencies characteristic of the bond type and the masses of the atoms involved. Infrared spectroscopy exploits this: a beam of infrared radiation passed through a sample is absorbed at frequencies that match the natural vibrational frequencies of the bonds, leaving dips ("absorption peaks") in the transmitted spectrum. The peak positions are reported on an inverse wavelength scale, the wavenumber ν̃ in cm⁻¹, with the spectrum traditionally plotted from 4000 cm⁻¹ on the left to 400 cm⁻¹ on the right.
A bond can only absorb infrared radiation if its vibration causes a change in dipole moment of the molecule. Polar bonds (C=O, O-H, N-H, C-Cl) absorb strongly. Non-polar bonds in symmetric environments (the N≡N stretch in N₂; the symmetric O=C=O stretch in CO₂) absorb very weakly or not at all. This is the selection rule for IR activity. Homonuclear diatomic molecules (H₂, N₂, O₂, Cl₂) are completely IR-silent because they have no dipole moment and the stretching motion does not create one. This connects directly back to lesson 3 on electronegativity and polarity.
The IR spectrum is conventionally split into two regions:
The A-Level student is expected to memorise a small number of diagnostic stretches:
| Bond | Functional group | Wavenumber range (cm⁻¹) | Character |
|---|---|---|---|
| O-H | Alcohol | 3200–3550 | Broad, strong (hydrogen-bonded) |
| O-H | Carboxylic acid | 2500–3300 | Very broad, often extending into C-H region |
| N-H | Amine, amide | 3300–3500 | Medium, may be split into two peaks for primary amines |
| C-H | Alkyl, alkenyl, aryl | ~3000 | Sharp, just below or just above 3000 depending on hybridisation |
| C≡N | Nitrile | 2200–2260 | Sharp, medium |
| C=O | Ketone, aldehyde | 1680–1720 | Sharp, strong |
| C=O | Carboxylic acid | 1700–1725 | Sharp, strong (often slightly higher than ketone) |
| C=O | Ester | 1735–1750 | Sharp, strong (typically the highest of the three) |
| C-O | Alcohol, ether, ester | 1000–1300 | Strong, often several peaks |
Worked example: distinguishing ethanol from propanone (acetone) by IR. Ethanol C₂H₅OH carries the broad O-H absorption at ~3300 cm⁻¹ (very characteristic of hydrogen-bonded alcohols) and a strong C-O stretch in the 1000–1100 cm⁻¹ region; there is no C=O peak. Propanone CH₃COCH₃ carries no O-H absorption at 3300 cm⁻¹ (no -OH group); the dominant peak is the C=O stretch at approximately 1715 cm⁻¹. A glance at the two spectra is enough to assign each.
Another worked example: distinguishing a ketone from a carboxylic acid (both contain C=O). A ketone shows the C=O peak at ~1715 cm⁻¹ and no O-H absorption. A carboxylic acid shows the C=O peak slightly higher (~1710 cm⁻¹) and a very broad O-H absorption spanning approximately 2500–3300 cm⁻¹. The very broad acid O-H is unmistakable once recognised.
Nuclear magnetic resonance is the structural-connectivity technique par excellence. In a strong magnetic field, the magnetic moments of nuclei with non-zero spin (¹H, ¹³C, ¹⁹F, ³¹P, others) precess at frequencies that depend on the local electronic environment of the nucleus. By irradiating the sample at radio frequencies and detecting the resonance, an NMR spectrum is recorded that reveals the number of chemically distinct environments for the chosen nucleus, the relative number of nuclei in each environment, and (through spin-spin coupling) which environments are adjacent to which.
For A-Level the technique is treated qualitatively in §3.1.3 and quantitatively in §3.3.15 (course 10). Three orienting ideas are worth stating now:
For now, the only thing the student needs to know is: NMR exists, it is the technique that resolves connectivity, and full detail follows in §3.3.15.
A typical A-Level (and undergraduate) structural-deduction problem provides several lines of spectroscopic evidence and asks the student to integrate them. The standard workflow is:
Worked example. An unknown organic compound has:
Deduction. Mᵣ = 60 is consistent with C₃H₈O. The IR identifies an alcohol (O-H broad at 3300, C-O at 1050, no C=O). The mass-spec loss of 15 (-CH₃) suggests a CH₃ group on the molecule. The NMR shows only two environments with a 1:3 ratio (using a simplifying assumption that the question lumps several environments). The compound is consistent with propan-2-ol (CH₃)₂CHOH — which contains two methyl groups equivalent by symmetry, a CH next to OH, and an OH; with judicious approximation a 1:3 integration ratio between one set of H atoms (CH + OH, accounting for 2H) and the methyl protons (6H) gives 2:6 = 1:3. (At full A-Level the assignment would be checked more carefully against integration; at this introductory stage the exercise is to combine M, IR and NMR consistently.)
A simpler worked example: M = 46, IR has broad O-H at 3300, C-O at 1050, no C=O; NMR has peaks in 1:3 ratio (or, more carefully, a triplet : quartet : broad-singlet pattern with integration 3:2:1). Compound = ethanol C₂H₅OH.
Each technique has its own sample-preparation requirements and characteristic artefacts:
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