Topic 8 · page 13 of 29

Thermal stability: size and charge of the cation

A small, highly charged cation has a high charge density and polarises the carbonate or nitrate ion, so the compound decomposes more easily. Stability to heat rises down each group, and Group 1 is more stable than Group 2.

13 min read8.16

Read first: Ionisation energy and reactivity down Groups 1 and 2

Thermal decomposition 8.16

Carbonates and nitrates of Groups 1 and 2 can be decomposed by heating. How easily they decompose depends on the metal: some need only a Bunsen flame and others will not decompose at all. A compound that decomposes at a higher temperature is more thermally stable.

The pattern, which is explained on this page and shown experimentally on the next one:

Charge density and polarisation 8.16

The explanation depends on two properties of the cation: its charge and its size.

Ion Charge Ionic radius / nm Charge ÷ radius
Li⁺ +1 0.076 13.2
Na⁺ +1 0.102 9.8
K⁺ +1 0.138 7.2
Mg²⁺ +2 0.072 27.8
Ca²⁺ +2 0.100 20.0
Sr²⁺ +2 0.113 17.7
Ba²⁺ +2 0.136 14.7
A bar chart of charge divided by ionic radius, in units of elementary charge per nanometre. Group 1: lithium 13.2, sodium 9.8, potassium 7.2. Group 2: magnesium 27.8, calcium 20.0, strontium 17.7, barium 14.7. Bars fall down each group, and Group 2 bars are taller than Group 1 bars.
Charge density falls down each group because the radius increases while the charge stays the same. Group 2 ions have twice the charge, so higher charge density.

A cation with a high charge density attracts the electrons of a nearby anion strongly. It polarises the anion: it distorts its electron cloud, pulling it towards the cation.

Two panels. Left: a small cation with a high charge, labelled 2 plus, next to a carbonate ion. The electron cloud of the carbonate is pulled strongly towards the cation, drawn as a stretched, lopsided outline, and the C to O bond nearest the cation is marked as weakened. Right: a large cation with a lower charge, labelled 1 plus, next to a carbonate ion. The carbonate cloud is only slightly distorted.
A high charge density cation polarises the anion strongly, which weakens a bond in the anion. A low charge density cation hardly distorts it.

When the anion is strongly polarised, electrons are drawn from the anion towards the cation. A bond within the anion (a C–O bond in carbonate, an N–O bond in nitrate) is weakened. The anion breaks up more easily on heating, so the compound decomposes at a lower temperature.

  1. Down a group the ionic radius increases and the charge stays the same, so charge density falls. Polarisation falls. The compound becomes more thermally stable: magnesium carbonate decomposes more easily than barium carbonate.
  2. Group 1 against Group 2. A Group 2 ion has twice the charge and a smaller radius than the Group 1 ion next to it, so a much higher charge density. Group 2 carbonates and nitrates decompose at lower temperatures than their Group 1 equivalents. Sodium carbonate does not decompose in a Bunsen flame, magnesium carbonate does.
  3. Lithium is the exception in Group 1. Li⁺ is the smallest Group 1 ion, so its charge density (13.2) is well above that of Na⁺ (9.8) and close to that of the larger Group 2 ions (Ba²⁺ 14.7). It polarises carbonate and nitrate enough to behave like a Group 2 ion, so lithium carbonate and nitrate decompose more easily than the other Group 1 compounds, and give products like Group 2 compounds.

Worked exampleAn exam answer

Explain why barium carbonate is more thermally stable than magnesium carbonate.

  1. Mg²⁺ and Ba²⁺ have the same charge (2+), but Mg²⁺ is smaller, so it has a higher charge density.
  2. Mg²⁺ polarises the carbonate ion more: it distorts the electron cloud and weakens a C–O bond.
  3. So less energy is needed to decompose MgCO₃. MgCO₃ decomposes at a lower temperature, and BaCO₃ is more stable.

Check yourself

  1. Calculate the charge density, as charge divided by radius, of Na⁺ (0.102 nm) and Mg²⁺ (0.072 nm).
  2. Explain why magnesium carbonate decomposes at a lower temperature than calcium carbonate.
  3. Explain why sodium carbonate is more stable to heat than magnesium carbonate.
  4. State the trend in thermal stability of the Group 2 nitrates down the group, and give the reason.
  5. Lithium carbonate decomposes in a Bunsen flame but sodium carbonate does not. Explain.
  6. A student says "potassium carbonate is more stable than sodium carbonate because K⁺ has more charge". Correct this.

Answers cover these until you have tried the questions

  1. Na⁺: 1 ÷ 0.102 = 9.8. Mg²⁺: 2 ÷ 0.072 = 27.8.
  2. Both cations have a 2+ charge, but Mg²⁺ is smaller than Ca²⁺, so it has a higher charge density and polarises the carbonate ion more. A C–O bond is weakened more, so MgCO₃ decomposes more easily.
  3. Mg²⁺ has twice the charge of Na⁺ and is smaller, so its charge density (27.8) is much higher than that of Na⁺ (9.8). It polarises the carbonate ion more, so MgCO₃ decomposes at a lower temperature.
  4. The nitrates become more thermally stable down the group. The cation gets larger, the charge density falls, polarisation of the nitrate ion decreases, so the N–O bond is weakened less.
  5. Li⁺ is smaller than Na⁺, so it has a higher charge density and polarises the carbonate ion more, weakening it enough to decompose in a Bunsen flame. Na⁺ has a lower charge density and does not.
  6. Sodium and potassium ions have the same charge (1+). K⁺ is larger, so it has a lower charge density and polarises the carbonate ion less. That makes potassium carbonate more stable.