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The pH scale, introduced by Søren Sørensen in 1909, measures the acidity or basicity of an aqueous solution. It is defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH = -log[H⁺]. At 25°C, a neutral solution has equal concentrations of hydrogen ions (H⁺) and hydroxide ions (OH⁻), each at 10⁻⁷ moles per liter, resulting in a pH of exactly 7. Pure water autoionizes to this extent, making it the reference point for neutrality. Solutions with pH below 7 are acidic (excess H⁺), while those above 7 are basic or alkaline (excess OH⁻). The scale is logarithmic, meaning each whole number change represents a tenfold difference in acidity. Temperature affects the ionization constant of water (Kw), so neutral pH shifts slightly at different temperatures—for example, at 100°C, neutral pH is approximately 6.1. Understanding pH is essential in chemistry, biology, medicine, and environmental science, governing processes from enzyme function to ocean acidification. The human body maintains strict pH regulation, with blood typically around 7.35-7.45.
