Reversible Reactions and Chemical Equilibria

Reversible Reactions and Chemical Equilibria

In 1798, the chemist Claude Berthollet (1748–1822) accompanied a French military expedition to Egypt. While visiting the Natron Lakes, a series of salt water lakes carved from limestone, Berthollet made an observation that contributed to an im- portant discovery. Upon analyzing water from the Natron Lakes, Berthollet found large quantities of common salt, NaCl, and soda ash, Na₂CO₃, a result he found sur- prising. Why would Berthollet find this result surprising and how did it contribute

Berthollet “knew” that a reaction between Na₂CO₃ and CaCl₂ goes to comple- tion, forming NaCl and a precipitate of CaCO₃ as products.

Na₂CO₃ + CaCl₂ → 2NaCl + CaCO₃

Understanding this, Berthollet expected that large quantities of NaCl and Na₂CO₃ could not coexist in the presence of CaCO₃. Since the reaction goes to completion, adding a large quantity of CaCl₂ to a solution of Na₂CO₃ should produce NaCl and CaCO₃, leaving behind no unreacted Na₂CO₃. In fact, this result is what he ob- served in the laboratory. The evidence from Natron Lakes, where the coexistence of NaCl and Na₂CO₃ suggests that the reaction has not gone to completion, ran counter to Berthollet’s expectations. Berthollet’s important insight was recognizing that the chemistry occurring in the Natron Lakes is the reverse of what occurs in the laboratory.

CaCO₃ + 2NaCl → Na₂CO₃ + CaCl₂

Using this insight Berthollet reasoned that the reaction is reversible, and that the relative amounts of “reactants” and “products” determine the direction in which the reaction occurs, and the final composition of the reaction mixture. We recog- nize a reaction’s ability to move in both directions by using a double arrow when writing the reaction.

Na₂CO₃ + CaCl₂ < = = = = > 2NaCl + CaCO₃

Berthollet’s reasoning that reactions are reversible was an important step in understanding chemical reactivity. When we mix together solutions of Na₂CO₃ and CaCl₂, they react to produce NaCl and CaCO₃. If we monitor the mass of dissolved Ca2+ remaining and the mass of CaCO₃ produced as a function of time, the result will look something like the graph in Figure 6.2. At the start of the reaction the mass of dissolved Ca2+ decreases and the mass of CaCO₃ in- creases. Eventually, however, the reaction reaches a point after which no further changes occur in the amounts of these species. Such a condition is called a state of equilibrium.

Although a system at equilibrium appears static on a macroscopic level, it is important to remember that the forward and reverse reactions still occur. A reac- tion at equilibrium exists in a “steady state,” in which the rate at which any species forms equals the rate at which it is consumed.