Aqueous Solutions and Chemical

Equilibria

Most analytical techniques require the state of chemical equilibrium. At equilibrium, the rate of a forward process or reaction and that of

the reverse process are equal.

As water slowly seeps over the limestone surface (superficie calcarea) of the cave, calcium carbonate dissolves in the water according to the chemical equilibrium.

The flowing water becomes saturated with calcium carbonate. As carbon dioxide is swept away, the reverse reaction becomes favored, and limestone is deposited in formations whose shapes are governed by the path of the flowing water.

“Frozen Niagara” in Mammoth Cave National Park in Kentucky.

Stalactites and stalagmites are examples of similar formations found where water saturated with calcium carbonate drips from the ceiling to the floor of caves over eons.

Il carbonato di calcio è il sale di calcio dell'acido carbonico. Puro, a temperatura ambiente è un solido bianco poco solubile in acqua, cui impartisce una reazione lievemente basica: 100 grammi in un litro di acqua a 20 °C formano una sospensione il cui pH è circa 10! Perché?

CO32- + H2O HCO3

- + OH-

Come gli altri carbonati, subisce decomposizione per riscaldamento o per contatto con sostanze acide, liberando anidride carbonica. Su questa sua caratteristica si basa il metodo per la sua determinazione, detto analisi calcimetrica.

CaCO3 + HCl CO2 + H2O + CaCl2

Il carbonato di calcio è il maggiore componente del calcare sciolto nell'acqua e il principale responsabile della sua durezza. In combinazione con il riscaldamento dell'acqua per usi civili dà origine all'incrostazione calcarea.

In natura, il carbonato di calcio è il materiale che costituisce, in tutto o in parte, una grande varietà di tipi di rocce: il marmo, le rocce calcaree, il travertino. I minerali costituiti da carbonato di calcio sono l'aragonite e la calcite.

CaCO3 CO2 + CaO

In 1923, J. N. Brønsted in Denmark and J. M. Lowry in England proposed independently a theory of acid/base behavior that is especially useful in analytical chemistry. According to the Brønsted-Lowry theory, an acid is a proton donor, and a base is a proton acceptor. For a molecule to behave as an acid, it must encounter a proton acceptor (or base). Likewise, a molecule that can accept a proton behaves as a base if it encounters an acid.

Acids and Bases

An important feature of the Brønsted-Lowry concept is the idea that the product formed when an acid gives up a proton is a potential proton acceptor and is called the conjugate base of the parent acid. For example, when the species acid1 gives up a proton, the species base1 is formed, as shown by the reaction

Many solvents are proton donors or proton acceptors and can thus induce basic or acidic behavior in solutes dissolved in them.

An acid that has donated a proton becomes a conjugate base capable of accepting a proton to reform the original acid. Similarly, a base that has accepted a proton becomes a conjugate acid that can donate a proton to form the original base.

Water is the classic example of an amphiprotic solvent, that is, a solvent that can act either as an acid or as a base depending on the solute.

Amphiprotic Species

The simple amino acids are an important class of amphiprotic compounds that contain both a weak acid and a weak base functional group. When dissolved in water, an amino acid, such as glycine, undergoes a kind of internal acid/base reaction to produce a zwitterion—a species that has both a positive and a negative charge.

Other common amphiprotic solvents are methanol, ethanol, and anhydrous acetic acid. In methanol, for example, the equilibria are analogous to the water equilibria.

Amphiprotic solvents undergo self-ionization, or autoprotolysis, to form a pair of ionic species. Autoprotolysis is yet another example of acid/base behavior, as illustrated The extent to which water undergoes autoprotolysis at room temperature is slight.

The first two are strong acids because reaction with the solvent is sufficiently complete that no undissociated solute molecules are left in aqueous solution. The rest are weak acids, which react incompletely with water to give solutions containing significant quantities of both the parent acid and its conjugate base. Note that acids can be cationic, anionic, or electrically neutral. The same holds for bases. The acids in Figure become progressively weaker from top to bottom. Ammonium ion is an even weaker acid with only about 0.01% of this ion being dissociated into hydronium ions and ammonia molecules.

Perchlorate and chloride ions have no affinity for protons. The tendency of a solvent to accept or donate protons determines the strength of a solute acid or base dissolved in it. For example, perchloric and hydrochloric acids are strong acids in water. If anhydrous acetic acid, a weaker proton acceptor than water, is substituted as the solvent, neither of these acids undergoes complete dissociation. Instead, equilibria such as the following are established:

Perchloric acid is, however, about 5000 times stronger than hydrochloric acid in this solvent.

Acetic acid thus acts as a differentiating solvent toward the two acid revealing the inherent differences in their acidities. Water, on the other hand, is a leveling solvent for perchloric, hydrochloric, and nitric acids because all three are completely ionized in this solvent and show no differences in strength.

Types of Equilibrium Constants in Analytical Chemistry

Applying the Ion-Product Constant for Water

Aqueous solutions contain small concentrations of hydronium and hydroxide ions as a result of the dissociation reaction

How much is the water concentration?

Suppose we have 0.1 mol of HCl in 1 L of water. The presence of this acid will shift the equilibrium to the left. Originally, however, there was only 10-7 mol/L OH- to consume the added protons.

Using Solubility-Product Constants

Most, but not all, sparingly soluble salts are essentially completely dissociated in saturated aqueous solution. For example, when an excess of barium iodate is equilibrated with water, the dissociation process is adequately described by the equation

When we say that a sparingly soluble salt is completely dissociated, we do not imply that all of the salt dissolves. What we mean is that the very small amount that does go into solution dissociates completely.

What does it mean to say that “an excess of barium iodate is equilibrated with water”? It means that more solid barium iodate is added to a portion of water than would dissolve at the temperature of the experiment. Some solid BaIO3 is in contact with the solution.

In other words, the number of moles of Ba(IO3)2 divided by the volume of the solid Ba(IO3)2 is constant no matter how much excess solid is present.

The Effect of a Common Ion on the Solubility of a Precipitate?

The solubility of an ionic precipitate decreases when a soluble compound containing one of the ions of the precipitate is added to the solution. This behavior is called the common-ion effect.

If you do the exercise (at home) you can see that a 0.02 M excess of Ba2+ decreases the solubility of Ba(IO3)2 by a factor of about 5; this same excess of IO3

+ lowers the solubility by a factor of about 200.

The equation can frequently be simplified by making the additional assumption that dissociation does not appreciably decrease the molar concentration of HA.