The pH of pure water is affected by temperature due to changes in the ion product of water (Kw). At 25°C, Kw is 1.0 × 10⁻¹⁴, and since the concentrations of hydrogen (H⁺) and hydroxide ions (OH⁻) are equal in pure water, the pH is 7. However, as temperature increases, Kw increases due to the enhanced dissociation of water molecules into hydrogen and hydroxide ions. This increased ionisation results in a higher concentration of both ions, but since the pH is dependent on the hydrogen ion concentration, the pH of pure water decreases (becomes more acidic) as temperature rises. Conversely, at lower temperatures, Kw decreases, leading to lower concentrations of these ions and a higher pH (more alkaline). This temperature dependency of pH is essential in processes where temperature control is critical, such as in biological systems and chemical reactions. It underscores the importance of considering environmental conditions when measuring and interpreting pH values, as the standard pH of 7 for neutrality is strictly true only at 25°C.

The pH scale, commonly ranged from 0 to 14, is based on the concentration of hydrogen ions in neutral water at 25°C, where the pH is 7 (neutral), and the concentrations of hydrogen (H⁺) and hydroxide ions (OH⁻) are both 1 × 10⁻⁷ M. However, in highly concentrated acidic or basic solutions, the hydrogen ion concentration can exceed 1 M (pH < 0) or the hydroxide ion concentration can exceed 1 M (pH > 14). This occurs because the pH scale is logarithmic and not limited to fixed endpoints. In extremely acidic solutions, a high concentration of hydrogen ions causes the pH to drop below 0. Conversely, in highly basic solutions, the low concentration of hydrogen ions (or high concentration of hydroxide ions) pushes the pH above 14. These extreme pH values are significant in industrial chemistry, where strong acids or bases are often used. They also challenge the traditional notion of the pH scale, highlighting its relative nature and dependence on specific conditions, particularly the solvent and temperature.

Yes, the pH of a solution can be negative, although this occurs under extreme conditions. The pH scale is commonly thought to range from 0 to 14, but these are not absolute limits. pH is defined as the negative logarithm of the hydrogen ion concentration. In highly acidic solutions, where the hydrogen ion concentration exceeds 1 Molar (M), the pH becomes negative. This situation arises with very strong acids, like concentrated sulfuric acid or hydrochloric acid, in high concentrations. For instance, a 1 M solution of a strong acid has a pH of 0, but a 10 M solution of the same acid would have a pH of -1. These extreme pH values are more theoretical and are rarely encountered in everyday situations. They are mainly of interest in industrial applications involving strong acids and in certain theoretical and experimental chemistry contexts. Understanding that pH can be negative helps in comprehending the logarithmic nature of the pH scale and the behaviour of highly acidic solutions.

The pKa value of an acid is a crucial indicator of its strength. It is the negative logarithm of the acid dissociation constant (Ka), which measures how readily an acid donates its proton to a base. A lower pKa value means a higher Ka, indicating a stronger acid that dissociates more completely in solution. This is because a lower pKa reflects a greater tendency of the acid to lose its proton. Understanding pKa values is vital in predicting the behaviour of acids in various chemical reactions, including synthesis and titration. In organic chemistry, the pKa values help predict the direction of proton transfer reactions, which are fundamental to understanding reaction mechanisms. In biochemistry, the pKa values of amino acids determine the ionisation state of proteins, affecting their structure and function. Therefore, pKa is not just a numerical value; it is a key concept that helps chemists understand and predict the behaviour of acidic substances in a wide range of chemical contexts.

The presence of a common ion in a solution of a weak acid affects its pH through the common ion effect. This phenomenon occurs when an ion that is a part of the weak acid's equilibrium (like the conjugate base) is added to the solution from another source. For instance, if acetate ions are added to an acetic acid solution, they shift the equilibrium position of the acetic acid dissociation reaction to the left (according to Le Chatelier's Principle), reducing the concentration of hydrogen ions and increasing the pH. This effect is crucial in understanding buffer solutions and their resistance to pH changes. The equilibrium expression for the weak acid is altered by the increased concentration of the common ion, resulting in a lower degree of ionisation of the acid. This principle is used in many practical applications, including controlling pH in industrial processes and biological systems, where maintaining a stable pH is essential.