2 (a) Group 2 nitrates decompose when heated - CIE - A-Level Chemistry - Question 2 - 2019 - Paper 1

As the atomic number of Group 2 elements increases, the thermal stability of their nitrates generally increases. This can be explained by several factors:

1. Cationic Radius: The cationic radius increases down the group, leading to greater ionic size. Larger cations have a weaker ability to polarize the nitrate ion.

2. Polarization/Distortion Effects: Heavier Group 2 cations, such as Ba²⁺, are less polarizing than smaller cations, therefore they induce less distortion on the surrounding nitrate ions. As a result, this reduces the thermal stability of nitrates derived from the smaller cations.

3. Charge Density: Smaller cations, such as Ca²⁺, have a higher charge density than larger cations (like Ba²⁺), leading to increased polarization and therefore less thermal stability in their nitrates.

Calcium amide, Ca(NH₂)₂, will decompose more readily than barium amide, Ba(NH₂)₂. This is because the smaller Ca²⁺ ion has a higher charge density compared to the larger Ba²⁺ ion. This higher charge density increases the polarization of the amide ions, resulting in a lower thermal stability due to greater distortion of the NH₂⁻ ion.

The decomposition of barium amide can be represented by the following equation:

$2 \text{Ba(NH}_2\text{)}_2 \xrightarrow{\text{heat}} \text{Ba}_3\text{N}_2 + 4 \text{NH}_3$

The bond angle of NH₄⁻ is predicted to be approximately 109.5 degrees. This angle is a result of the tetrahedral geometry of the ion. In NH₄⁻, there are four bonding pairs of electrons and no lone pairs. According to the VSEPR theory, the structure will arrange itself to minimize repulsion between the electron pairs, leading to a tetrahedral shape with a bond angle of 109.5 degrees.