Carbonic acid and its salts

Interaction of carbon dioxide with water

Syntheses involving carbon atoms are actively used in organic chemistry, as carbon is capable of bonding various atoms and molecules. A key inorganic compound is carbon(IV) oxide, known as carbon dioxide (CO2). This is a chemically active and stable substance. Biologically, it is indispensable: it serves as the primary carbon source for photosynthesis in plants, and in living cells, it is the end product of the oxidation of various organic compounds.

Upon contact with water, carbon dioxide demonstrates moderate solubility. The dissolution process involves a reversible reaction, the product of which is carbonic acid (H2CO3). It is important to understand that this is an extremely unstable compound exhibiting the properties of a weak acid. Under real-world conditions, no more than 0.2% of the total volume of dissolved gas converts into the acid form. It is precisely the presence of this acid that gives a characteristic slightly sour taste to water that is carbonated under pressure to produce sparkling water.

Carbonates and bicarbonates: preparation and solubility

The chemical properties of carbonic acid allow it to form two classes of salts: normal salts (carbonates) and acid salts (bicarbonates).

The formation of normal salts occurs when carbon dioxide is absorbed by solid alkalis or their solutions. For example, the reaction with sodium hydroxide yields classic sodium carbonate (Na2CO3). Obtaining acid salts requires special conditions: they are formed either by passing an excess of CO2 through a solution, or by the gradual, careful addition of a strong acid to a carbonate solution.

Solubility rules for these two classes differ significantly.

• Carbonates are mostly insoluble — salts of many divalent cations (calcium, magnesium, iron(II)) precipitate out. The only exceptions are readily soluble ammonium and alkali metal carbonates.
• Bicarbonates dissolve much better in water, and this rule also applies to acid salts of divalent metals.

An interesting paradox is observed with sodium salts: baking soda (NaHCO3) dissolves in water much worse than the normal salt (Na2CO3). This property is the basis of its production method: if carbon dioxide is passed through a concentrated solution of sodium carbonate, the bicarbonate precipitates out.

A natural illustration of the transition of salts from one state to another is karst — the process of the destruction of limestone formations with the formation of caves. Insoluble calcium carbonate under the action of water and dissolved carbon dioxide turns into soluble calcium bicarbonate and is gradually washed away.

Hydrolysis and acid-base equilibrium

In aqueous solutions, carbonic acid salts undergo hydrolysis, but the extent of this process strictly depends on the type of salt.

Normal salts (sodium carbonate) hydrolyze to a very significant degree. At a solution concentration of 0.1 mol/L, the degree of hydrolysis reaches 4%, and the hydrogen exponent (pH) rises to 11.7. Conclusion: in solution, this salt functions as a relatively strong base.

Bicarbonates are subject to hydrolysis to a much lesser extent. A solution of baking soda has a pH of about 8.5, meaning that basic properties are expressed more strongly in it than acidic ones. The main chemical feature of acid salts is their amphotericity. They are capable of reacting with both typical acids and strong alkalis (e.g., KOH).

The biological role of sodium bicarbonate is immense. It is a constant component of human blood, performing the function of a vital buffer system. Its main task is the rapid neutralization of aggressive acid vapors entering the body from the atmospheric air.

Behavior towards acids and heating

If a strong acid is added in excess to any carbonic acid salt, carbonic acid is immediately displaced and instantly decomposes into water and carbon dioxide. The evolution of CO2 bubbles serves as a reliable sign of the reaction even for water-insoluble carbonates upon the action of dilute solutions of strong acids.

The thermal stability of salts obeys strict regularities. The general rule states that the vast majority of carbonates decompose upon heating with the release of carbon(IV) oxide.

• Ammonium carbonate is distinguished by particularly low heat resistance.
• Only alkali metal carbonates (sodium, potassium) possess unique properties. They are capable of melting without chemical decomposition (the melting point of Na2CO3 is 851 °C).
• Limestone (CaCO3) undergoes thermal decomposition upon heating into lime (CaO) and carbon dioxide. When heated to 950 °C, the process proceeds almost completely in the forward direction. In the temperature range of 700–900 °C, the reaction is typically reversible, and the equilibrium can be shifted by changing the carbon dioxide pressure.

Briefly: answers to topical questions

Class and exam questions

What else to study on this topic

• Writing the hydrolysis equations for sodium and ammonium carbonate with the indication of the medium.
• Amphoteric properties of bicarbonates: reaction equations with HCl and KOH.
• The effect of carbon dioxide pressure on equilibrium shift during limestone calcination.
• Blood buffer systems and the mechanism of acid neutralization.

Other topics in "Chapter 12. Group IVA Elements"

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