Acid rain occurs when pollutants like sulphur dioxide (SO₂) and nitrogen oxides (NOx) in the atmosphere react with water vapour to form sulphuric acid (H₂SO₄) and nitric acid (HNO₃), respectively. These acids dissolve in rainwater, decreasing its pH and making the rainwater more acidic than normal. Typical pH values of acid rain can be around 4.0 - 4.5, significantly lower than the normal pH of unpolluted rainwater, which is around 5.6. The environmental impacts of acid rain are extensive and damaging. It leads to the acidification of lakes and rivers, adversely affecting aquatic life by disrupting nutrient balances and increasing the solubility of toxic metals like aluminium. Acid rain also damages forests, particularly at higher elevations where the clouds and fog are more acidic than rain, affecting the soil's nutrient balance and damaging the leaves and barks of trees. Additionally, acid rain corrodes buildings, monuments, and infrastructure, particularly those made of limestone and marble, as the acid reacts with the calcium carbonate in these materials. This not only has cultural and historical implications but also economic impacts due to the increased need for maintenance and restoration.

pH and pOH are both measures of the concentration of ions in a solution but focus on different ions. pH is a measure of the concentration of hydrogen ions (H⁺), while pOH measures the concentration of hydroxide ions (OH⁻). The pH scale ranges from 0 to 14, with lower values indicating acidic solutions (high H⁺ concentration) and higher values indicating basic solutions (low H⁺ concentration). Similarly, the pOH scale also ranges from 0 to 14, with lower values indicating basic solutions (high OH⁻ concentration) and higher values indicating acidic solutions (low OH⁻ concentration). The pH and pOH are related through the equation pH + pOH = 14 (at 25°C). This relationship is a direct consequence of the water dissociation constant (Kw), which at 25°C is 1.0 × 10⁻¹⁴. This means that in any aqueous solution at 25°C, the product of the hydrogen ion concentration [H⁺] and hydroxide ion concentration [OH⁻] is always 1.0 × 10⁻¹⁴. Therefore, if the pH is known, the pOH can be easily calculated and vice versa.

A buffer solution is a special type of solution that resists changes in pH when small amounts of acid or base are added. It works by containing a weak acid and its conjugate base (or a weak base and its conjugate acid). When an acid (H⁺ ions) is added to the buffer, the conjugate base in the buffer reacts with these H⁺ ions to form the weak acid, minimising the change in pH. Conversely, when a base (OH⁻ ions) is added, the weak acid in the buffer donates H⁺ ions, which react with the OH⁻ ions to form water, again stabilising the pH. This ability to maintain a relatively constant pH is crucial in many biological systems, where enzymes and other biochemical processes require a specific pH range to function optimally. In industrial processes, buffers are used in fermentation, dyeing, and chemical manufacturing to maintain the desired pH, ensuring consistency and quality of the end products.

Temperature plays a significant role in determining the pH of water and aqueous solutions. In pure water, an increase in temperature causes a slight decrease in pH, meaning the water becomes more acidic. This is due to the enhanced ionisation of water at higher temperatures, leading to a greater concentration of hydrogen ions (H⁺). In aqueous solutions, the effect of temperature on pH can vary depending on the nature of the dissolved substances. For instance, the ionisation of weak acids and bases is also temperature-dependent. An increase in temperature typically increases the ionisation of weak acids and bases, resulting in a lower pH for acidic solutions and a higher pH for basic solutions. This temperature-dependent behaviour is crucial in chemical processes and biological systems, where slight variations in pH can have significant effects. For example, in enzymatic reactions within living organisms, optimal activity is often at a specific pH and temperature, and deviations from these conditions can lead to reduced efficiency or denaturation of the enzymes.

The solubility of carbon dioxide (CO₂) in water significantly affects the pH of natural water bodies like lakes, rivers, and oceans. When CO₂ dissolves in water, it forms carbonic acid (H₂CO₃), a weak acid, which slightly dissociates into hydrogen ions (H⁺) and bicarbonate ions (HCO₃⁻). This dissociation lowers the pH of the water, making it more acidic. The extent of this acidification depends on the concentration of dissolved CO₂, which can vary due to factors like atmospheric CO₂ levels, temperature, and biological activity. In natural water bodies, this process is part of the carbon cycle and plays a crucial role in regulating the pH of the ecosystem. However, increased levels of atmospheric CO₂ due to human activities (like burning fossil fuels) have led to increased dissolution of CO₂ in oceans, a phenomenon known as ocean acidification. This has serious implications for marine life, particularly for organisms with calcium carbonate structures, as the increased acidity can dissolve these structures.