Which among the following factors changes the value of ionic product of water? (a) change in temperature (b) addition of acid (c) addition of base (d) addition of either acid and base

Understanding the effects of temperature on the ionic product of water is crucial when exploring chemical equilibrium. The ionic product of water, denoted as \(K_w\), quantifies the amount of hydronium \(\text{H}_3\text{O}^+\) and hydroxide \(\text{OH}^-\) ions present in water. Specifically, \(K_w\) increases with higher temperature because this causes a greater degree of water dissociation into its ionic components. At room temperature, \(K_w\) is typically \(1.0 \times 10^{-14}\), but this value rises as the temperature increases.For instance, at elevated temperatures, there are more collisions between water molecules, which can lead to an increase in the rate and extent of dissociation. It is important to note that this behavior is an example of an endothermic process; energy absorbed from the surroundings facilitates the breaking of bonds within water molecules, leading to increased ionic concentrations.

When an acid is introduced to a water solution, it contributes additional hydronium ions \(\text{H}_3\text{O}^+\) to the system. Despite this, the ionic product of water, \(K_w\), remains unchanged because of the principle of chemical equilibrium. What happens is a shift in the position of equilibrium to counter the added hydronium ions by reducing the number of hydroxide ions \(\text{OH}^-\).

This balancing act—the decrease in hydroxide ions to match the increase in hydronium ions—ensures that the value of \(K_w\) is maintained. In a sense, the system absorbs the impact of the added acid through what is known as the Le Chatelier's principle. This principle posits that if a dynamic equilibrium is disturbed, the system will adjust itself to diminish the change and restore equilibrium.

Conversely, if a base is added to the water, it supplies extra hydroxide ions \(\text{OH}^-\), which could be predicted to increase the ionic product of water, \(K_w\). However, as with the addition of acid, Le Chatelier's principle comes into play, ensuring a balance is struck. The increase in hydroxide ions prompts a compensatory decrease in the concentration of hydronium ions \(\text{H}_3\text{O}^+\).

Through this self-regulatory mechanism, the product of the ion concentrations remains stable, keeping \(K_w\) constant despite the introduction of additional hydroxide ions. It's a delicate balance that demonstrates the resilience and stability of chemical systems under the governance of equilibrium laws.

Water dissociation is a fundamental concept that underpins the ionic product of water. It is the process by which water molecules \(H_2O\) separate into hydronium \(\text{H}_3\text{O}^+\) and hydroxide \(\text{OH}^-\) ions. This disassociation is an equilibrium reaction, meaning that at any given time in pure water, there are water molecules splitting apart and ions rejoining to form water molecules.

The reaction can be represented as: \([H_2O] \longleftrightarrow [\text{H}_3\text{O}^+] + [\text{OH}^-]\). At chemical equilibrium, the rate of dissociation equals the rate of recombination, resulting in a constant ionic product, \(K_w\). But it's not static—the ratio of dissociated ions can fluctuate in response to changes in environmental conditions such as temperature, although, within a closed system, the value of \(K_w\) remains constant unless the temperature changes.