Chemistry Chapter 21: Further Aspects of Equilibria

CHEMISTRY KEYWORDS:

Common Ion Effect: Reduction of the solubility of a dissolved salt by adding a compound that has an ion in common with the dissolved salt.

Partition Coefficient, K_pc: Ratio of the concentrations of a solute in 2 different immiscible solvents in contact with each other when equilibrium has been established.

 

CHEMISTRY DEFINITIONS:

Conjugate Pair (acid-base): An acid-base pair on each side of an acid-base equilibrium equation that are related to each other by the difference of a hydrogen ion.

Ionic Product of Water, K_w: Equilibrium constant for the ionisation of water.

Acid Dissociation Constant, K_a: Equilibrium constant for the dissociation of a weak acid.

Buffer Solution: A solution that minimises changes in pH when moderate amounts of acid or base are added.

Solubility Product, K_sp: Product of the concentrations of each ion in a saturated solution of a sparingly soluble salt at 298 K, raised to the power of their relative concentrations.

 

CHEMISTRY IMPORTANT NOTES:

Acids are proton donors, and bases are proton acceptors.

In a reaction at equilibrium, products are being converted to reactants at the same rate as reactants are being converted to products.

In a conjugate pair, the acid has 1 proton more than its conjugate base.

K_w is the ionic product of water. Its value at 298 K is 1.00 x 10^-14 mol²/ dm⁶.

The range of possible hydrogen ion concentrations in different solutions is very large. It can range from 10^-15 mol/dm³ to 10 mol/dm³. Danish chemist Soren Sorensen introduced the pH scale.

Even when there is an excess of OH^- ions, there is still a small concentration of H⁺ ions.

Monobasic acids contain only 1 replaceable hydrogen atom per molecule. Strong monobasic acids are completely ionised in solution.

Diluting the acid 10 times reduces the value of the H⁺ ion concentration to 1/10 and increases the pH by a value of 1.

Calculate the pH of a strong base solution:

·       Concentration of OH^- ions in solution.

·       Equilibrium expression for the ionisation of water.

·       K_w value.

K_a is the acid dissociation constant. At 298 K, the value of K_a for the dissociation of ethanoic acid is 1.74 x 10^-5 mol/dm³.

A high value of K_a indicates that the position of equilibrium lies to the right. The acid is almost completely ionised.

A low value of K_a indicates that the position of equilibrium lies to the left. The acid is only slightly ionised and exists mainly as HA molecules and comparatively few H⁺ and A^- ions.

Calculate K_a for a weak acid:

·       Concentration of the acid.

·       pH of the solution.

We ignore the concentration of hydrogen ions produced by the ionisation of the water molecules present in the solution, as the ionic product of water (1.00 x 10^-14 mol²/dm⁶) is negligible compared with the values for most weak acids.

We assume that ionisation of the weak acid is so small that the concentration of undissociated HA molecules present at equilibrium is approximately the same as that of the original acid.

Calculate the pH of a weak acid:

·       Concentration of the acid.

·       K_a value for the acid.

The value of pH calculated will not be significantly affected by these factors unless we require great accuracy. (above 2 assumptions)

A buffer solution is a solution in which the pH does not change significantly when small amounts of acids or alkalis are added. It is used to keep pH (almost) constant. It is either a weak acid and its conjugate base or a weak base and its conjugate acid, which minimises any change in pH when an acid or alkali is added. It plays an important part in many industrial processes, including electroplating, the manufacture of dyes and in the treatment of leather. It is also used to make sure that pH meters record the correct pH.

Buffer solution contains a relatively high concentration of both CH₃COOH and CH₃COO^-, also known as reserve supplies of acid (CH₃COOH) and its conjugate base (CH₃COO^-).

The pH of a buffer solution depends on the ratio of the concentration of the acid and the concentration of its conjugate base.

An increase in hydrogen ion concentration would greatly lower the pH of water, but when H⁺ ions are added to the buffer solution:

·       Addition of H⁺ ions shifts the position of equilibrium to the left because reactants combine with H⁺ ions to form more products, until equilibrium is re-established.

·       Large reserve supply of both acid and its conjugate base ensures that the concentration of the ions and molecules in the solution does not change significantly.

·       pH does not change significantly.

An increase in hydroxide ion concentration would greatly increase the pH of water, but when OH^- ions are added to the buffer solution:

·       Added OH^- ions combine with H⁺ ions to form water.

·       This reduces the H⁺ ion concentration. The position of equilibrium shifts to the right.

·       Product ionises to form more H⁺ and reactant ions until equilibrium is re-established.

·       Large reserve supply of both acid and its conjugate base ensures that the concentration of the ions and molecules in the solution does not change significantly.

·       pH does not change significantly.

In unpolluted regions of Earth, rainwater has a pH of 5.7, as carbon dioxide dissolves in rainwater to form a dilute solution of the weak acid carbonic acid, H₂CO₃ and its conjugate base, HCO₃^- acts as a buffer solution. It minimises changes in pH if very small amounts of acid or alkali are added to the rainwater.

In polluted regions of Earth, pH of rainwater falls to around 4. The rainwater can no longer act as a buffer because the concentrations of H₂CO₃ and HCO₃^- are not high enough to cope with the large amounts of acidic pollution involved.

No buffer solution can cope with the excessive addition of acids or alkalis. If very large amounts of acid or alkali are added, the pH will change significantly.

Buffer solutions that resist changes in pH in alkaline regions are usually a mixture of a weak base and its conjugate acid.

Calculate the pH of a buffer solution:

·       K_a of the weak acid.

·       Equilibrium concentration of the weak acid and its conjugate base (salt).

The cells in our body produce carbon dioxide as a product of aerobic respiration (oxidation of glucose to provide energy). Carbon dioxide combines with water in the blood to form a solution containing hydrogen ions. This is catalysed by the enzyme carbonic anhydrase. The production of H⁺ ions, if left unchecked, would lower the pH of the blood and cause acidosis. This may disrupt some body functions and eventually lead to coma.

If the H⁺ ion concentration increases:

·       Position of equilibrium shifts to the left.

·       H⁺ ions combine with HCO₃^- ions to form carbon dioxide and water until equilibrium is restored.

·       This reduces the concentration of hydrogen ions in the blood and helps keep the pH constant.

If the H⁺ ion concentration decreases:

·       Position of equilibrium shifts to the right.

·       Some carbon dioxide and water combine to form H⁺ and HCO₃^- ions until equilibrium is restored.

·       This increases the concentration of hydrogen ions in the blood and helps keep the pH constant.

Solubility is generally quoted as the number of grams or number of moles of compound needed to saturate 100 g or 1 kg of water at a given temperature. We say that a solution is saturated when no more solute dissolves in it.

The common ion effect is the reduction in the solubility of a dissolved salt achieved by adding a solution of a compound that has an ion in common with the dissolved salt. This often results in precipitation.

The addition of the common ion has reduced the solubility of the compound because its solubility product has been exceeded.

The solubility of an ionic compound in aqueous solution containing a common ion is less than its solubility in water.

Ammonia molecules are moving from the aqueous layer to the organic layer at the same rate as they are moving from the organic layer to the aqueous layer.

The partition coefficient is the equilibrium constant that relates the concentration of a solute partitioned between 2 immiscible solvents at a particular temperature.

K_pc depends on the relative solubilities of the solute in the 2 solvents used in the partitioning. Solubility depends on the strength of the intermolecular bonds between the solute and the solvent. This, in turn, depends on the polarity of the molecules of solute and solvent.

The greater the relative solubility in the mobile phase, the faster the rate of movement as the mobile phase passes over the stationary phase.

 

SUMMARY:

pH is a measure of the hydrogen ion concentration.

K_a is the dissociation constant for an acid. It is the equilibrium constant for the dissociation of a weak acid.

Acid strengths can be compared using pK_a values.

The ionic product for water, K_w = 1.00 x 10^-14 mol² dm^-6

A buffer solution is a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.

The pH of a buffer solution can be calculated by using the equilibrium concentrations of the weak acid and its conjugate base and the K_a value of the weak acid.

The solubility product, K_sp is the equilibrium expression showing the equilibrium concentrations of the ions in a saturated solution of a sparingly soluble salt, taking into account the relative number of each ion present.

The addition of a common ion to a saturated solution of a sparingly soluble salt causes precipitation of the sparingly soluble salt.

The partition coefficient, K_pc is the equilibrium constant which relates the concentration of a solute partitioned between 2 immiscible solvents at a particular temperature.