Explain decomposition through cation polarisation, compare Group 1 and 2 product patterns, and connect flame colours to discrete electron transitions.
Small, highly charged cations destabilise large oxyanions
Thermal decomposition is the breakdown of a substance into simpler products on heating. A compound that requires more severe heating to decompose is more thermally stable. This is different from having a high melting temperature: melting alone does not change chemical identity.
A small cation with a high charge has high charge density and strongly distorts the electron cloud of a nearby carbonate or nitrate ion. This polarisation weakens bonding within the oxyanion and makes decomposition easier. Down Groups 1 or 2, cation size increases while charge stays the same; polarisation decreases and thermal stability increases.
For comparable cation sizes, a 2+ ion polarises an anion more strongly than a 1+ ion. Lithium is an important exception within Group 1 because Li⁺ is particularly small: its nitrate and carbonate decompose in patterns more like Group 2 compounds. The explanation concerns charge density and distortion, not electrons being removed from the metal during heating.
Carbonate decomposition releases carbon dioxide
Group 2 carbonates decompose to the oxide and carbon dioxide. Magnesium carbonate decomposes more readily than calcium carbonate, with the trend towards greater stability continuing down the group. The residual oxide is a new compound, not molten carbonate.
Lithium carbonate also decomposes to its oxide and carbon dioxide. Other common Group 1 carbonates are stable under ordinary laboratory heating; avoid stating that no possible temperature or condition could ever decompose them.
Carbon dioxide turns limewater cloudy because insoluble calcium carbonate forms. The gas test confirms a product of decomposition; seeing a gas alone would not uniquely identify CO₂.
Nitrate products depend on the group and lithium exception
Brown NO₂ gas is evidence for the Group 2/lithium decomposition pattern. Oxygen can support combustion, but the gases from nitrate decomposition may include toxic nitrogen dioxide, so gas testing must use a prescribed supervised method in a fume cupboard. Do not smell gases or put your face near the apparatus.
The sodium product NaNO₂ is sodium nitrate(III), commonly called sodium nitrite: nitrogen is +3 rather than +5. For Group 2 nitrates the nitrogen goes to +4 in NO₂, while some oxygen changes from −2 to 0 in O₂. These equations are redox as well as thermal decomposition.
Worked prediction: strontium nitrate follows 2Sr(NO₃)₂ → 2SrO + 4NO₂ + O₂. Four nitrogen atoms and twelve oxygen atoms occur on each side. Replacing Sr with Na without changing the product type would incorrectly ignore the different Group 1 pattern.
| Starting nitrate | Products | Balanced pattern |
|---|---|---|
| Group 2 M(NO₃)₂ | Metal oxide, nitrogen dioxide, oxygen | 2M(NO₃)₂ → 2MO + 4NO₂ + O₂ |
| Lithium nitrate | Lithium oxide, nitrogen dioxide, oxygen | 4LiNO₃ → 2Li₂O + 4NO₂ + O₂ |
| Other common Group 1 nitrates | Nitrate(III) / nitrite and oxygen | 2NaNO₃ → 2NaNO₂ + O₂ |
A comparison needs controlled heating and a meaningful endpoint
For a carbonate comparison, place equal mole amounts of comparable dry powders into identical hard-glass tubes. Heat using the same controlled position and heat input, and pass the evolved gas into the same volume and concentration of limewater. Record time to a defined persistent cloudiness or, preferably, gas production at measured temperature. Repeat and compare the spread.
Keep particle size, tube dimensions and the sample’s position in the heated region similar. Equal masses would not be equal amounts for different carbonates and could give different total gas volumes. Leaks, moisture and different heat transfer can change the observed time; a shorter time alone is partly a rate/heat-transfer measurement rather than a perfect measure of thermodynamic stability.
An electrically controlled heating block or furnace with temperature measurement gives a stronger comparison than an uncontrolled flame, especially when decomposition temperatures differ widely. Use small amounts, eye protection and heat-resistant apparatus; never seal a heated gas-producing system. Remove the delivery tube from limewater before heating stops to reduce suck-back as gas cools.
For nitrates, use a fume cupboard and prescribed collection/exhaust arrangement because NO₂ is toxic. Record the onset of gas evolution and, where appropriate, mass loss after cooling to a stable mass. Heating again after cooling helps check completion; handling or weighing a hot tube introduces burn risk and unreliable mass readings.
Characteristic colours identify some metal ions
Flame testing concerns a compound containing the ion, not combustion of the pure metal. In particular, magnesium salts give no characteristic visible flame-test colour, while magnesium metal burns with intense white light. The 2023 AS Q2(b) report explicitly identifies this distinction.
Francium and radium are not routine classroom flame-test samples; their radioactivity and scarcity prevent treating them as ordinary practical unknowns. Use supplied information for unusual or unfamiliar species rather than inventing a memorised colour.
| Ion in the tested salt | Typical colour | Caution |
|---|---|---|
| Li⁺ | Crimson / scarlet red | Distinguish with a clean wire or spectrum |
| Na⁺ | Intense yellow | Trace contamination can mask other colours |
| K⁺ | Lilac | Can be masked by sodium |
| Rb⁺ | Red-violet | Less common classroom sample |
| Cs⁺ | Blue-violet | Less common classroom sample |
| Be²⁺, Mg²⁺ | No characteristic visible flame colour | Not the same as burning the metal |
| Ca²⁺ | Brick-red / orange-red | Compare with a known sample |
| Sr²⁺ | Crimson red | Distinguish from lithium using further evidence |
| Ba²⁺ | Apple-green | Barium compounds require hazard controls |
Explain excitation and build contamination controls into the method
Clean a nichrome or platinum loop using the prescribed hydrochloric acid and heat it in a non-luminous flame until it gives no persistent colour. Test a small amount of sample and compare its colour with known references. Use a clean loop or a separate clean applicator for each sample; concentrated acid requires appropriate eye protection and careful handling under the laboratory procedure.
Energy from the flame excites electrons to higher allowed energy levels. As electrons return to lower levels, photons with particular energies and wavelengths are emitted. Different elements have different energy spacings, so they give characteristic line spectra and overall colours. The visible light is emitted on the downward transition, not while the electron absorbs energy to rise.
A hand-held spectroscope distinguishes line positions more reliably than a subjective description of a mixed colour. Repeat a doubtful result, test a blank and compare a known sample under the same conditions. A yellow trace in every sample may indicate sodium contamination; it does not demonstrate every unknown is a pure sodium salt.
Quick checks
Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.
Q1. Why is BaCO₃ more thermally stable than MgCO₃?Show answer
Ba²⁺ is larger at the same charge, so has lower charge density and polarises CO₃²⁻ less. Bonding within carbonate is less weakened, so stronger heating is needed for decomposition.
Q2. Write the balanced thermal decomposition equations for Ca(NO₃)₂ and KNO₃.Show answer
2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂. For potassium, 2KNO₃ → 2KNO₂ + O₂. Calcium follows the Group 2 oxide pattern; potassium forms nitrite/nitrate(III).
Q3. Why does using equal masses of two carbonates weaken a fair comparison of gas evolution?Show answer
Different molar masses mean different mole amounts and different possible CO₂ totals. Use equal mole amounts, similar particle size and controlled heating, then compare a defined endpoint or gas volume.
Q4. A salt gives no characteristic flame colour. Is this inconsistent with magnesium being present?Show answer
No. Mg²⁺ salts have no characteristic visible flame colour. The brilliant white light from burning magnesium metal is a different experiment and cannot be used as the salt’s flame-test colour.
Q5. A learner says electrons emit coloured light as they rise to a higher level. Correct the sequence and explain why the light has particular wavelengths.Show answer
Electrons absorb energy to rise. They emit photons when falling to lower allowed levels. Discrete energy differences give specific photon energies and therefore specific wavelengths.
Sources
Sources and examiner guidance (reviewed 9 October 2026)
- Pearson Edexcel 9CH0 specification, Issue 3 — Topic 4, printed pp. 13–14, checked against 8CH0 Topic 4, printed pp. 11–12. Reviewed 9 October 2026.
- Chemrevise — Edexcel Inorganic Chemistry and the Periodic Table — All nine pages reviewed as a secondary coverage check; explanations and questions here are original.
- Pearson 8CH0/01 June 2023 mark scheme — Q2(b–c), Q3(b)(ii), Q6 and Q9(b); PDF pp. 8–9, 13, 20–21, 28–30. Guidance is specific to these questions.
- Pearson 8CH0/01 June 2023 examiner report — Q2–3, Q6 and Q9; PDF pp. 4–8. Reviewed with question context on 9 October 2026.
- Pearson 8CH0/01 June 2023 question paper — Context for ion tests, displacement, redox and Group 2 reasoning. Original exercises below do not copy these questions.
Finesse Tuition is not endorsed by AQA or Chemrevise. All explanations and examples here are our own.