Introduction:

The self-ionization of water (that is, the method in which the water ionizes to hydronium ions and hydroxide ions) takes place to a very limited extent. Whenever two molecules of water collide, there can be a transfer of a hydrogen ion from one molecule to the other. The products are a positively charged hydronium ion and a negatively charged hydroxide ion.

H₂O (l) + H₂O (l) ↔ H₃O⁺ (aq) + OH^- (aq)

We frequently utilize the simplified form of the reaction:

H₂O (l) ↔ H⁺ (aq) + OH^- (aq)

The equilibrium constant for the self-ionization of water is termed to as the ion-product for water and is represented the symbol K_w.

K_w = [H⁺][OH^-]

The ion-product of water (K_w) is the arithmetical product of the concentration of hydrogen ions and hydroxide ions. It will be noted that H₂O is not comprised in the ion-product expression as it is a pure liquid. The value of K_w is extremely small, in accordance by a reaction which favors the reactants. At 25°C, the experimentally determined value of K_w in the pure water is 1.0 × 10^-14.

K_w = [H⁺][OH^-] = 1.0 × 10^-14

In pure water, the concentrations of hydrogen and hydroxide ions are equal to one another. The pure water or any other aqueous solution in which this ratio holds is stated to be neutral. To determine the molarity of each ion, the square root of K_w is taken.

[H⁺] = [OH^-] = 1.0 × 10^-7 M

The acidic solution is a solution in which the concentration of hydrogen ions is more than the concentration of hydroxide ions. For illustration, hydrogen chloride ionizes to produce H⁺ and Cl^- ions on dissolving in water.

HCl (g) → H⁺ (aq) + Cl^- (aq)

This raises the concentration of H⁺ ions in the solution. According to LeChatelier's principle, the equilibrium is symbolized by:

H₂O (l) ↔ H⁺ (aq) + OH^- (aq)

is forced to the left, in the direction of the reactant. As an outcome the concentration of the hydroxide ion reduces.

A fundamental solution is a solution in which the concentration of hydroxide ions is more than the concentration of hydrogen ions. The solid potassium hydroxide dissociates in water to outcome potassium ions and hydroxide ions.

KOH (s) → K⁺ (aq) + OH^- (aq)

The increase in concentration of the OH^- ions causes the decrease in the concentration of H⁺ ions and the ion-product of [H⁺][OH^-] remains constant.

The pH Scale:

Deducing the acidity of a solution by employing the molarity of the hydrogen ion is cumbersome as the quantities are usually very small. Danish scientist Soren Sorenson (1868-1939) introduced an easier system for pointing out the concentration of H⁺ termed as the pH scale. The letters pH stand for the power of the hydrogen ion. The pH of a solution is the negative logarithm of hydrogen-ion concentration.

pH = -log [H⁺]

In pure water or a neutral solution the [H⁺] = 1.0 × 10^-7 M. Replacing to the pH expression:

pH = -log [1.0 × 10^-7] = -(-7.00) = 7.00

The pH of pure water or any neutral solution is therefore 7.00. For recording use, the numbers to the right of the decimal point in the pH value are the important figures. As 1.0 × 10^-7 consists of the two significant figures, the pH can be reported as 7.00.

The logarithmic scale condenses the range of acidity to numbers which are simple to use. Take a solution having [H⁺] = 1.0 × 10^-4 M. That is a hydrogen-ion concentration which is around 1000 times higher than the concentration in the pure water. The pH of such a solution is 4.00, a difference of just 3 pH units. It will be noted that whenever the [H⁺] is written in scientific notation and the coefficient is 1, the pH is simply the exponent with the sign changed. The pH of a solution having the [H⁺] = 1 × 10^-2 M is 2 and the pH of a solution having [H⁺] = 1 × 10^-10 M is 10.

As we are familiar that a solution having the [H⁺] higher than 1.0 × 10^-7 is acidic, as a solution having the [H⁺] lower than 1.0 × 10^-7 is basic. As a result, solutions whose pH is less than 7 are acidic, whereas those having a pH higher than 7 are basic. The Figure below illustrates the relationship, all along with several illustrations of different solutions.

Fig: pH Scale

Why does pure water have a pH of 7?

That question is in reality deceptive! However, pure water only consists of a pH of 7 at a particular temperature - the temperature at which the K_w value is 1.00 x 10^-14 mol² dm^-6.

This is how it comes in relation to:

To determine the pH you first require finding out the hydrogen ion concentration (or hydroxonium ion concentration - its similar thing). Then you transform it to pH. In pure water at room temperature the K_w value states you that:

[H⁺] [OH^-] = 1.00 x 10^-14

However in pure water, the hydrogen ion (that is, hydroxonium ion) concentration should be equivalent to the hydroxide ion concentration. For each and every hydrogen ion formed, there is a hydroxide ion formed too. That signifies that you can substitute the [OH^-] term in the K_w expression by the other [H⁺].

[H⁺]² = 1.00 x 10^-14

By taking the square root of each side provides:

[H⁺] = 1.00 x 10^-7 mol dm^-3

Converting that to pH:

pH = - log10 [H⁺]

pH = 7

That is where the well-known value of 7 comes from.

Common-ion effect:

Define:  Degree of ionization of the electrolyte is suppressed via the addition of a strong electrolyte having common ion. This effect is termed as the common ion effect.

In another word: The phenomenon of lowering the degree of ionization of a weak electrolyte via adding a solution of a strong electrolyte having a common ion is termed as common ion effect.

This theory comprise how the solubility of a salt changes whenever some ion which is common to both added substance and the salt in question. Let us take the equilibrium reaction comprising the saturated solution of silver chloride which is illustrated below.

AgCl (s) ↔ Ag⁺ (aq) + Cl^- (aq)

Let us add certain amount of silver nitrate (AgNO₃) to the above equilibrium system and observe what happens to the solubility of silver chloride.

Silver nitrate dissociates and generates silver ion (Ag⁺) and nitrate ion (NO₃^-) according to:

AgNO₃ (s) → Ag⁺ (aq) + NO₃^- (aq)

Here, Ag⁺ is the common ion to both AgNO₃ and AgCl. Adding AgNO₃ to the saturated solution of AgCl raises the concentration of the Ag⁺ ion thereby increasing the collisions between Ag⁺ ion and NO₃^- ion which outcomes in the formation of more solid AgCl.

According to the Le Chatelier's principle, it is a shift in equilibrium from right to left via forming more solid AgCl. The total result is the decrease in the solubility of AgCl as Ag⁺ ion and Cl^- ion are eliminated from the solution in the form of solid AgCl.

Thus, the effect of common ion is to reduce the solubility of the salt.

Whenever NaCl is added, the effects will still the same (that is, decrease in solubility of AgCl), as Cl^- ion is common to both the NaCl and AgCl.

When NaNO₃ is added, there is no effect on the solubility, as there is no common ion.

Application of Common Ion Effect: 

The knowledge of common ion effect is very helpful in analytical chemistry. This is frequently applied in the qualitative analysis.

An electrolyte is precipitated only if the concentration of its ions surpasses the solubility product (KSP). The precipitation is obtained only if the concentration of any one ion is raised. Therefore by adding a common ion, the solubility product can be raised.

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