Given below are two statements :
One is labelled as Assertion (A) and the other is labelled as Reaso

Q: Why does thermal stability increase down the group?

As the cation size increases down the group, it becomes more comparable in size to the large carbonate anion. This leads to a more efficient and stable crystal lattice packing, requiring much more heat energy to break the bonds.

Q: Why does solubility decrease down the group?

Solubility depends on hydration energy overcoming lattice energy. Down the group, the hydration energy decreases much more rapidly than lattice energy because the cations get larger and their charge density drops. Thus, solubility decreases.

Q: Is Beryllium carbonate soluble in water?

Yes, compared to Barium carbonate, Beryllium carbonate is much more soluble because the small $Be^{2+}$ ion has a very high charge density, leading to a very high hydration energy that easily overcomes its lattice energy.

Q: Does this trend apply to other compounds like sulphates?

Absolutely! Sulphates ($SO_4^{2-}$) are also large polyatomic anions, so they follow the exact same trends for thermal stability and solubility as carbonates. Barium sulphate is also highly stable and insoluble.

Q: What happens when Barium carbonate is finally heated enough to decompose?

When heated to extreme temperatures, it decomposes into Barium oxide ($BaO$) and Carbon dioxide gas ($CO_2$). However, the temperature required is much higher than what is needed to decompose Beryllium or Magnesium carbonate.

Question

Given below are two statements :
One is labelled as Assertion (A) and the other is labelled as Reason (R).
Assertion (A) Barium carbonate is insoluble in water and is highly stable.
Reason (R) The thermal stability of the carbonates increases with increasing cationic size.
In the light of the above statements, choose the most appropriate answer from the options given below.

Options:

(A) Both (A) and (R) are true and (R) is the correct explanation of (A).

(B) (A) is true but (R) is false.

(C) Both (A) and (R) are true and (R) is not the true explanation of (A).

(D) (A) is false but (R) is true.

Sigma Percentile · Accepted Answer

Imagine you are an architect trying to build a highly stable, enduring structure. If you have massive, heavy boulders, would you use tiny pebbles to hold them together? Probably not. The structure would be wobbly and unstable. You would want large, sturdy blocks to match the boulders. This intuitive idea is the exact principle behind the thermal stability of alkaline earth metal carbonates!

### Analyzing the Setup

In inorganic chemistry, the carbonate ion ($CO_3^{2-}$) is considered a **large polyatomic anion**. When it forms ionic compounds with Group 2 metals (the alkaline earth metals), the stability of the resulting crystal lattice depends heavily on how well the cation and anion fit together.

As we move down Group 2 from Beryllium ($Be$) to Barium ($Ba$), the atomic and ionic radii increase due to the addition of new electron shells.

$$r_{Be^{2+}} < r_{Mg^{2+}} < r_{Ca^{2+}} < r_{Sr^{2+}} < r_{Ba^{2+}}$$

### The Master Equation of Stability

A fundamental rule dictates that **a large anion is best stabilized by a large cation**. Because the Barium ion ($Ba^{2+}$) is the largest in the group, it packs most efficiently with the large carbonate ion. This excellent packing leads to a highly stable crystal lattice.

Therefore, the thermal stability of the carbonates increases as we move down the group:

$$BeCO_3 < MgCO_3 < CaCO_3 < SrCO_3 < BaCO_3$$

Barium carbonate ($BaCO_3$) is so stable that it requires extremely high temperatures to decompose into Barium oxide and Carbon dioxide. This perfectly validates our Reason (R).

### The Solubility Paradox

Now, let's tackle the second part of the puzzle: solubility in water. For any ionic salt to dissolve, the energy released when water molecules surround the ions (the **Hydration Energy**) must be greater than the energy holding the crystal together (the **Lattice Energy**).

$$\Delta H_{	ext{hydration}} > \Delta H_{	ext{lattice}}$$

Here is the catch: Hydration energy is inversely proportional to the size of the ion. Small ions like $Be^{2+}$ have a very high charge density, so water molecules are strongly attracted to them, releasing a massive amount of hydration energy.

However, as we move down the group to $Ba^{2+}$, the ion becomes so large that its charge density drops. Consequently, its hydration energy decreases rapidly. In fact, the hydration energy decreases much faster than the lattice energy does.

Because the hydration energy for Barium carbonate is so low, it cannot overcome the lattice energy. Thus, **solubility decreases down the group**, making Barium carbonate practically insoluble in water. This perfectly validates our Assertion (A).

### Final Conclusion

Both the Assertion and the Reason are factually correct. More importantly, the Reason (increasing cationic size leading to higher thermal stability) is the exact scientific explanation for why Barium carbonate is so highly stable. Therefore, the correct choice is that both (A) and (R) are true, and (R) is the correct explanation of (A).

Select Answer:

Visualized Solution

$G r o u p 2 C a r b o na t es$

$T h er ma l S t abi l i t y P r in c i pl e$

$T r e n d o f T h er ma l S t abi l i t y$

$S o l u bi l i t y P r in c i pl e$

$T r e n d o f S o l u bi l i t y$

$C o n c l u s i o n$

$T h e W a y F or w a r d$

The Sigma Insight: Alkaline Metals

Analyzing the Setup

The Master Equation of Stability

The Solubility Paradox

Final Conclusion

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