Theoretical basics of Inorganic chemistry

Hydrolysis. The reaction of an anion or cation with water accompanied by cleavage of O–H bond is called hydrolysis. The term hydrolysis is derived from hydro, meaning water, and lysis, meaning breaking. It may be noted that in anionic hydrolysis shown in the solution becomes slightly basic (pH > 7) due to the generation of excess OH^– ions. In cationic hydrolysis there is excess of H⁺ ions which makes the solution slightly acidic (pH < 7) [24, p. 1].

 

Bronsted-Lowry concept of hydrolysis. HA and A– are conjugate acid-base pair

 

HA + H₂O ↔ H₃O⁺ + A^–

                                                      weak acid                     conjugate base

 

Since HA is a weak acid (poor proton donor), its conjugate base, A–, must be relatively strong (good proton acceptor). Owing to this fact, A^– ions tend to react with water by accepting proton from the latter to form HA molecule (anionic hydrolysis),

 

A^- + H₂O → HA + OH⁻

 

The presence of OH– ions makes the solution basic.

Similarly, BOH and B⁺ are a conjugate acid-base pair. Since BOH is a weak base, its conjugate acid, B⁺, would be relatively strong. Thus B⁺ would accept OH^– ions from water to form BOH molecules.

 

B⁺ + H₂O → BOH + H⁺

 

The presence of excess H⁺ ions makes the solution acidic [24, p. 2].

 

Hydrolysis of salts. Hydrolysis of salts disturbs the ionic equilibrium of water that causes change in concentration of hydrogen and hydroxyl ions and leads to acidic or alkaline reaction of a solution.

As it is known, water is a weak electrolyte, ionizing slightly and reversibly

 

HOH ↔ H⁺ + OH^-

 

If water is absolutely pure, the concentrations of hydrogen and hydroxyl ions are equal to each other and equal to 10^-7 mole/l at temperature of 25 °C. The product of these concentrations is named the ion product for water Kw

 

Kw = C(H⁺) ∙ C(OH^-) = 10^-14

 

To determine the acidity-alkalinity of a medium the hydrogen ion index pH and the hydroxyl ion index pOH are used

 

pH= - lg C(H⁺)

pOH = - lg C(OH^-)

 

It is easy to deduce from the above equations:

·         neutral solution pH = 7;

·         acidic solution pH < 7 (the higher acidity, the smaller it is);

·         alkaline solution pH > 7 (the higher alkalinity, the larger it is).

 

It is obvious that pH + pOH=14

There are various methods of measuring pH values. Qualitatively the reaction of a solution can be determined by means of special reagents called indicators, which change colour depending on concentration of hydrogen ion.

The most widely used indicators are litmus, phenolphthalein and methyl orange (table 8.1).

 

Table 8.1

Colours of indicators

Indicator	Reaction of solution
	acid	neutral	alkaline
Methyl orange	red	orange	yellow
Phenolphthalein	colourless	pale crimson	crimson
Litmus	red	violet	blue

 

Any salt can be obtained by neutralization of acids by bases. It is natural to assume, therefore, that solutions, at least those witch are products of complete displacement of hydrogen in acids by metals, must react neutral. However, this assumption holds only for salts of strong acids and strong bases. Salts of weak acids and strong bases, or vice versa, of strong acids and weak bases, do not react neutral when dissolved in water. As an example, a solution of zinc chloride ZnCl₂ reacts acid, indicating an excess of hydrogen ions as compared to hydroxyl ions C_(H+) > C_(OH-). 

In contrast to the previous example, a solution of potassium fluoride KF reacts alkaline which is characteristic of hydroxyl ion and the ratio of hydrogen and hydroxyl ions is C_(H+) < C_(OH-). 

The ionization theory attributes this phenomena to the interaction of water ions with the ions of dissolved salts, resulting the formation of an excess of hydrogen or hydroxyl ions. Although the concentration of hydrogen or hydroxyl ions is very low in water, they are in equilibrium with an immense number of unionized water molecules. When one of them is bound by the ions of the salt, the equilibrium is disturbed, causing new water molecules to ionize, which may lead to accumulation of considerable quantities of the other ion, making the solution react acid or alkaline.

Any reaction between the ions of a salt and the ions of water, attended with a change in the concentration of the latter, is called hydrolysis of the salt.

The chief cause of hydrolysis is the formation of slightly ionized substances (ions or molecules) [25, p. 3].

 

Acid-Base Neutralization. A solution is neutral when it contains equal concentrations of hydronium and hydroxide ions. When we mix solutions of an acid and a base, an acid-base neutralization reaction occurs. However, even if we mix stoichiometrically equivalent quantities, we may find that the resulting solution is not neutral. It could contain either an excess of hydronium ions or an excess of hydroxide ions because the nature of the salt formed determines whether the solution is acidic, neutral, or basic. The following four situations illustrate how solutions with various pH values can arise following a neutralization reaction using stoichiometrically equivalent quantities:

1.      A strong acid and a strong base, such as HCl(aq) and NaOH(aq) will react to form a neutral solution since the conjugate partners produced are of negligible strength (see figure 8.1):

 

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

 

Figure 8.1 The chart shows the relative strengths of conjugate acid-base pairs.

 

2.      A strong acid and a weak base yield a weakly acidic solution, not because of the strong acid involved, but because of the conjugate acid of the weak base.

3.      A weak acid and a strong base yield a weakly basic solution. A solution of a weak acid reacts with a solution of a strong base to form the conjugate base of the weak acid and the conjugate acid of the strong base. The conjugate acid of the strong base is a weaker acid than water and has no effect on the acidity of the resulting solution. However, the conjugate base of the weak acid is a weak base and ionizes slightly in water. This increases the amount of hydroxide ion in the solution produced in the reaction and renders it slightly basic.

4.     A weak acid plus a weak base can yield either an acidic, basic, or neutral solution. This is the most complex of the four types of reactions. When the conjugate acid and the conjugate base are of unequal strengths, the solution can be either acidic or basic, depending on the relative strengths of the two conjugates. Occasionally the weak acid and the weak base will have the same strength, so their respective conjugate base and acid will have the same strength, and the solution will be neutral. To predict whether a particular combination will be acidic, basic or neutral, tabulated K values of the conjugates must be compared [2, p. 804].

 

Salts of Weak Bases and Strong Acids. When we neutralize a weak base with a strong acid, the product is a salt containing the conjugate acid of the weak base. This conjugate acid is a weak acid. For example, ammonium chloride, NH₄Cl, is a salt formed by the reaction of the weak base ammonia with the strong acid HCl:

 

NH₃(aq) + HCl(aq) ⟶ NH₄Cl(aq)

 

A solution of this salt contains ammonium ions and chloride ions. The chloride ion has no effect on the acidity of the solution since HCl is a strong acid. Chloride is a very weak base and will not accept a proton to a measurable extent. However, the ammonium ion, the conjugate acid of ammonia, reacts with water and increases the hydronium ion concentration:

 

NH₄⁺(aq) + H₂O(l) → H₃O+(aq) + NH₃(aq)

 

The equilibrium equation for this reaction is simply the ionization constant. K_a, for the acid NH₄⁺:

 

[H₃O⁺][NH₃] / [NH₄⁺] = K_a

 

We will not find a value of K_a for the ammonium ion in Appendix H. However, it is not difficult to determine K_a for NH₄⁺ from the value of the ionization constant of water, K_w, and K_b (1), the ionization constant of its conjugate base, NH₃, using the following relationship [2, p. 806]:

 

K_w = K_a × K_b (1)

 

Salts of Weak bases and Strong acids. Some salts of weak bases and strong acids undergo cationic hydrolysis and yield slightly acidic solutions. Ammonium chloride is a typical example of this class of salts. It is the salt of a weak base. NH₄OH, and strong acid, HCl. It ionises in aqueous solution to form the cation, NH₄⁺.

 

NH₄OH ↔ NH₄⁺ + OH^-

       conjugate

    acid

 

NH₄⁺ is a Bronsted conjugate acid of the weak base NH₄OH. Therefore, it is a relatively strong acid. It accepts OH– ion from water (H₂O) and forms the unionised NH₄OH and H⁺ ion.

 

NH₄⁺ + H₂O ↔ NH₄OH + H⁺

 

The accumulation of H⁺ ions in solution makes it acidic.

The other examples of this type of salts are ferric chloride, aluminium chloride, and copper sulphate [24, p. 3].

 

Salts of Weak acids and Weak bases. The examples of this type of salts are ammonium acetate, ammonium cyanide and ammonium fluoride. Both the anion and the cation produced by ionisation of the salt undergo hydrolysis. The resulting solution is neutral, basic or acidic depending on the relative hydrolysis of the anions and the cations.

Ammonium acetate,CH3COONH4. It is the salt of weak acid, CH₃COOH, and weak base, NH₄OH. In aqueous solution it ionises to form the anion CH₃COO^– and the cation NH₄⁺. Since the acid and the base are both weak, their conjugate base (CH₃COO^–) and conjugate acid ( NH₄⁺) are relatively strong. They accept H⁺ and OH^– ions respectively from water and undergo considerable hydrolysis.

 

CH₃COO^- + H₂O ↔ CH₃COOH + OH^-    ...(1)

                                       conjugate base

 

NH₄⁺ + H₂O ↔ NH₄OH + H⁺     ...(2)

                                            conjugate acid

 

The overall hydrolysis may be represented as

 

CH₃COO^- + NH₄⁺ + H₂O ↔ CH₃COOH + NH₄OH

 

We have stated above that pH of the resulting solution will depend on the relative extent of anionic hydrolysis (1) and cationic hydrolysis (2). If both the ions react to the same extent (as shown for CH₃COONH₄), [OH^–] = [H⁺] and the solution is neutral. If the cation reacts to a greater extent, the solution is slightly acidic. If the anion is a little more reactive, the solution will be basic. Thus, a solution of CH₃COONH₄ is neutral, a solution of NH₄CN is slightly basic and a solution of NH₄F is slightly acidic [24, p. 3].

 

Equilibrium in a Solution of a Salt of a Weak Acid and a Weak Base. In a solution of a salt formed by the reaction of a weak acid and a weak base, to predict the pH, we must know both the K_a of the weak acid and the K_b of the weak base. If K_a > K_b, the solution is acidic, and if K_b > K_a, the solution is basic [2, p. 810].

 

The Ionization of Hydrated Metal Ions. If we measure the pH of the solutions of a variety of metal ions we will find that these ions act as weak acids when in solution. The aluminum ion is an example. When aluminum nitrate dissolves in water, the aluminum ion reacts with water to give a hydrated aluminum ion, Al(H₂O)₆³⁺, dissolved in bulk water. What this means is that the aluminum ion has the strongest interactions with the six closest water molecules (the so-called first solvation shell), even though it does interact with the other water molecules surrounding this Al(H₂O)₆³⁺ cluster as well:

 

Al(NO₃)₃ (s) + 6H₂O (l) ⟶ Al(H₂O)₆³⁺ (aq) + 3NO₃⁻ (aq)

 

We frequently see the formula of this ion simply as “Al³⁺ (aq)”, without explicitly noting the six water molecules that are the closest ones to the aluminum ion and just describing the ion as being solvated in water (hydrated). This is similar to the simplification of the formula of the hydronium ion, H₃O⁺ to H⁺. However, in this case, the hydrated aluminum ion is a weak acid (figure 8.2) and donates a proton to a water molecule. Thus, the hydration becomes important and we may use formulas that show the extent of hydration:

 

Al(H₂O)₆³⁺ (aq) + H₂O (l) → H₃O⁺ (aq) + Al(H₂O)₅(OH)²⁺ (aq)   Ka = 1.4 × 10⁻⁵

 

As with other polyprotic acids, the hydrated aluminum ion ionizes in stages, as shown by:

 

Al(H₂O)₆³⁺ (aq) + H₂O(l) → H₃O⁺(aq) + Al(H₂O)₅(OH)²⁺(aq)

Al(H₂O)₅(OH)²⁺ (aq) + H₂O(l) → H₃O⁺(aq) + Al(H₂O)₄(OH)²⁺(aq)

Al(H₂O)₄(OH)²⁺ (aq) + H₂O(l) → H₃O⁺(aq) + Al(H₂O)₃(OH)₃(aq)

 

Note that some of these aluminum species are exhibiting amphiprotic behavior, since they are acting as acids when they appear on the left side of the equilibrium expressions and as bases when they appear on the right side.

 

Figure 8.2 When an aluminum ion reacts with water, the hydrated aluminum ion becomes a weak acid.

 

However, the ionization of a cation carrying more than one charge is usually not extensive beyond the first stage. Additional examples of the first stage in the ionization of hydrated metal ions are:

 

Fe(H₂O)₆³⁺(aq) + H₂O(l) → H₃O⁺(aq) + Fe(H₂O)₅(OH)²⁺(aq)    Ka = 2.74

Cu(H₂O)₆²⁺(aq) + H₂O(l) → H₃O⁺(aq) + Cu(H₂O)₅(OH)⁺(aq)    Ka = ~6.3

Zn(H₂O)₄²⁺(aq) + H₂O(l) → H₃O+(aq) + Zn(H₂O)₃(OH)⁺(aq)    Ka = 9.6

 

The constants for the different stages of ionization are not known for many metal ions, so we cannot calculate the extent of their ionization. However, practically all hydrated metal ions other than those of the alkali metals ionize to give acidic solutions. Ionization increases as the charge of the metal ion increases or as the size of the metal ion decreases [2, p. 810].