Flame Tests and Thermal Decomposition

In this lesson we cover the analytical identification of Group 2 cations using flame tests, as well as the thermal decomposition trends of Group 2 carbonates and nitrates, explained using polarisation theory.

1. Flame Test Colours

A flame test is a simple qualitative test used to identify specific metal cations. When metal ions are heated in a hot Bunsen burner flame, their outer electrons absorb energy and are excited to higher energy levels. As these excited electrons fall back down to their ground state (lower energy levels), they release energy in the form of light. The wavelength of the light determines the colour of the flame.

You must memorise the flame colours for Group 2 cations, along with those of Group 1 for comparison:

2. Thermal Decomposition of Group 2 Carbonates

Thermal decomposition refers to the breakdown of a compound into simpler substances by heat. All Group 2 carbonates decompose on heating to form a solid metal oxide and release carbon dioxide gas:

General Equation:

\[ \text{MCO}_3\text{(s)} \xrightarrow{\Delta} \text{MO(s)} + \text{CO}_2\text{(g)} \]

Specific Example:

\[ \text{CaCO}_3\text{(s)} \rightarrow \text{CaO(s)} + \text{CO}_2\text{(g)} \]

As you move down the group, the temperature required to decompose the carbonates increases significantly. For instance, beryllium carbonate decomposes at room temperature, magnesium carbonate decomposes at around 540 K, whereas barium carbonate requires temperatures of nearly 1600 K.

3. Thermal Decomposition of Group 2 Nitrates

All Group 2 nitrates undergo thermal decomposition to form the solid metal oxide, releasing brown nitrogen dioxide gas (\(\text{NO}_2\)) and oxygen gas (\(\text{O}_2\)):

General Equation:

\[ 2\text{M(NO}_3)_2\text{(s)} \xrightarrow{\Delta} 2\text{MO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)} \]

Specific Example:

\[ 2\text{Mg(NO}_3)_2\text{(s)} \rightarrow 2\text{MgO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)} \]

Just like the carbonates, the thermal stability of Group 2 nitrates increases down the group, requiring higher temperatures to decompose.

4. Explaining the Stability Trend: Polarisation Theory

To explain why the thermal stability of carbonates and nitrates increases down the group, we must examine the interactions between the positive metal cation (\(\text{M}^{2+}\)) and the negative molecular anion (\(\text{CO}_3^{2-}\) or \(\text{NO}_3^-\)).

The polarisation mechanism operates as follows:

1. Cation Charge Density: All Group 2 cations carry a \(2+\) charge. However, as you go down the group, the ionic radius of the cation increases. Consequently, the \(2+\) charge is spread over a larger volume, meaning the charge density decreases from \(\text{Be}^{2+}\) to \(\text{Ba}^{2+}\).
2. Polarisation of the Anion: The positively charged metal cation attracts the delocalised electron cloud of the carbonate or nitrate anion, pulling it closer. This distortion of the anion's electron cloud is called polarisation.
3. Weakening of Internal Covalent Bonds: The smaller the cation, the higher its charge density, and the more strongly it polarises the molecular anion. In a carbonate ion, this polarisation pulls electron density away from the C-O bonds and towards the cation. This weakens the C-O covalent bonds within the carbonate ion itself.
4. Decomposition Temperature: Because the C-O bonds are weakened in the presence of a highly polarising cation, less thermal energy is required to break them and split the anion into \(\text{MO}\) and \(\text{CO}_2\). Therefore, magnesium carbonate (containing the small, highly polarising \(\text{Mg}^{2+}\) ion) decomposes easily at low temperatures. Barium carbonate (containing the large, weakly polarising \(\text{Ba}^{2+}\) ion) has a stable carbonate cloud, requiring much more heat to decompose.

5. Comparing Group 1 and Group 2 Carbonates

A typical comparison question asks why Group 2 carbonates decompose under normal laboratory temperatures, while Group 1 carbonates (such as \(\text{Na}_2\text{CO}_3\)) do not decompose under a standard Bunsen burner (with the exception of lithium carbonate, \(\text{Li}_2\text{CO}_3\)).

This difference is explained by two factors:

• Cation Charge: Group 2 cations have a \(2+\) charge, whereas Group 1 cations only have a \(1+\) charge. The higher charge exerts a stronger pull on the anion's electron cloud.
• Cation Size: Group 2 cations are smaller than their corresponding Group 1 counterparts in the same period (e.g. \(\text{Mg}^{2+}\) is smaller than \(\text{Na}^+\)).

Because Group 2 cations have a higher charge and smaller size, they possess a much higher charge density than Group 1 cations. Consequently, they polarise the carbonate anion far more effectively, making Group 2 carbonates much less thermally stable than Group 1 carbonates.