acidsbases

Acids and bases are substances that influence proton transfer and electron-pair interactions in chemical reactions. Several definitions are used in chemistry. The Arrhenius model defines acids as substances that increase hydrogen ion concentration in water (forming hydronium, H3O+), and bases as substances that increase hydroxide, OH-. The Bronsted-Lowry definition broadens this to acids as proton donors and bases as proton acceptors, applicable in non-aqueous media as well. The Lewis definition extends further, describing acids as electron-pair acceptors and bases as electron-pair donors, encompassing reactions that do not involve protons.

Strength and measurement: Strong acids and bases dissociate completely in water, while weak ones dissociate only partially. The pH scale quantifies acidity or basicity; pH below 7 is acidic, above 7 is basic, and 7 is neutral at 25 °C. The relationship between pH and hydrogen ion activity is logarithmic. The autoionization of water, Kw = 1.0×10−14 at 25 °C, yields equal concentrations of H3O+ and OH− in pure water. Buffers resist pH change by containing a weak acid and its conjugate base or a weak base and its conjugate acid.

Conjugate pairs and reactions: Acid-base reactions produce conjugate acid-base pairs differing by one proton. The acidity or basicity of a species is often expressed by pKa or pKb values, guiding buffer design and titration. Organic acids (for example, acetic acid) and bases (such as ammonia) illustrate typical behavior in solution.

Applications: Concepts of acids and bases underpin many fields, including chemistry, biology, environmental science, medicine, and industry, influencing processes from digestion and metabolism to corrosion and wastewater treatment. In advanced contexts, Lewis acids and bases describe catalytic and coordination chemistry, with examples like metal cations and boron compounds acting as Lewis acids and amines or oxide ions as Lewis bases.