Does Limestone Dissolve in Water?

Short answer: barely on its own, but easily once carbon dioxide gets involved. Drop a piece of limestone into pure water and almost nothing happens. Add the carbon dioxide that is naturally dissolved in rain and groundwater, and the same rock slowly dissolves, enough, over thousands of years, to carve out entire cave systems. Here is why.

Limestone on its own is almost insoluble

Limestone is calcium carbonate, CaCO₃. In pure water it is classed as insoluble. A tiny amount does dissolve, as an equilibrium:

CaCO₃(s) ⇌ Ca²⁺(aq) + CO₃²⁻(aq)

but the position of this equilibrium lies far to the left. Only a few milligrams dissolve per litre, which is why a marble statue (also CaCO₃) keeps its shape in the rain over a human lifetime, though not over centuries.

How dissolved CO₂ changes everything

Rainwater is not pure. As it falls, and as it seeps through soil, it picks up carbon dioxide, and CO₂ in water makes a weak acid, carbonic acid:

CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq)

That weak acid attacks the carbonate. The overall reaction is:

CaCO₃(s) + CO₂(aq) + H₂O(l) → Ca²⁺(aq) + 2HCO₃⁻(aq)

which we usually write as the formation of soluble calcium hydrogencarbonate:

CaCO₃(s) + H₂CO₃(aq) → Ca(HCO₃)₂(aq)

Calcium hydrogencarbonate is soluble (it only exists in solution), so the limestone effectively dissolves. The trick is that the acid keeps removing carbonate ions, CO₃²⁻, by turning them into hydrogencarbonate, HCO₃⁻. By Le Chatelier’s principle, taking CO₃²⁻ out of that first equilibrium drags it to the right, so more and more CaCO₃ dissolves.

The reaction runs both ways (and that is the fun part)

Everything above is reversible. Lose the dissolved CO₂ (by warming the water, or by letting it stand in air) and the reaction reverses, depositing solid CaCO₃ again:

Ca(HCO₃)₂(aq) → CaCO₃(s) + CO₂(g) + H₂O(l)

That single reversible reaction explains a surprising amount of the real world:

• Hard water and limescale. Water that has trickled through limestone carries dissolved Ca(HCO₃)₂, this is “temporary” hard water. Boil it in a kettle and the reverse reaction deposits CaCO₃ as the white scale on the element.
• Stalactites and stalagmites. CO₂-rich water dissolves limestone above a cave, then drips through. In the open air of the cave it loses CO₂, deposits CaCO₃, and over ages builds the hanging stalactites and the rising stalagmites.
• Caves and sinkholes. The same slow dissolving, running for millennia, hollows out whole cave networks in limestone country.

The A-level link

This everyday question sits on top of two H2 topics.

• Group 2 chemistry. CaCO₃ is a Group 2 carbonate, and Group 2 carbonates are all essentially insoluble in water, which is exactly why “does it dissolve?” is an interesting question at all. The reverse reaction above is also a neat mirror of the thermal decomposition of Group 2 carbonates, CaCO₃ → CaO + CO₂, that you meet under thermal stability.
• Ionic equilibria and solubility. The dissolving of CaCO₃ is governed by its solubility product, Ksp. Adding acid removes CO₃²⁻ from the solution, and by Le Chatelier the dissolution shifts right. This is a textbook example of how pH affects the solubility of a sparingly soluble salt, a favourite exam idea.

So “does limestone dissolve in water?” is really a question about equilibrium: on its own, no; shift the equilibrium with a little acid, and yes.

Quick recap

• Pure water: CaCO₃ is essentially insoluble.
• Water plus CO₂: forms soluble Ca(HCO₃)₂, so it dissolves.
• Remove the CO₂: it comes back out as solid CaCO₃ (scale, stalactites).

Frequently asked questions

Does limestone dissolve in water? Barely in pure water (calcium carbonate is essentially insoluble), but readily in water containing dissolved CO₂, which turns it into soluble calcium hydrogencarbonate.

How does CO₂ make limestone dissolve? CO₂ forms carbonic acid, which reacts with CaCO₃ to give soluble Ca(HCO₃)₂: CaCO₃ + CO₂ + H₂O → Ca²⁺ + 2HCO₃⁻. Removing carbonate ions pulls the dissolution equilibrium to the right.

How do stalactites form? CO₂-rich water dissolves limestone underground; when it drips into a cave and loses CO₂, the reaction reverses and deposits solid CaCO₃, slowly building stalactites and stalagmites.

Practise this

See the Ionic Equilibria topic on ChemBank for Ksp and solubility questions, and the Periodic Table topic for Group 2 carbonate trends.

Book a trial lesson and turn everyday chemistry like this into exam marks.