Use scientific notation (e.g., 1e-7) for small values
pH = −log₁₀[H⁺]
pOH = −log₁₀[OH⁻]
pH + pOH = 14 (at 25°C)
Kw = [H⁺][OH⁻] = 10⁻¹⁴
Disclaimer
Results assume ideal aqueous solutions at standard conditions (25°C) unless otherwise specified. For accurate laboratory work, always verify with calibrated instruments.
pH is a measure of the acidity or basicity (alkalinity) of an aqueous solution. The term "pH" stands for "potential of hydrogen" or "power of hydrogen." It quantifies the concentration of hydrogen ions (H⁺) in a solution on a logarithmic scale ranging from 0 to 14. A pH of 7 is considered neutral (pure water at 25°C), values below 7 indicate acidic solutions, and values above 7 indicate basic or alkaline solutions.
The pH scale was developed by Danish chemist Søren Peder Lauritz Sørensen in 1909 while working at the Carlsberg Laboratory. Because pH is a logarithmic scale, each whole pH value below 7 is ten times more acidic than the next higher value. For example, a solution with pH 4 is ten times more acidic than one with pH 5, and 100 times more acidic than one with pH 6.
In any aqueous solution at 25°C, the product of the hydrogen ion concentration [H⁺] and hydroxide ion concentration [OH⁻] is constant: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴. This relationship, known as the ion product of water, means that as one concentration increases, the other must decrease proportionally. This is why acidic solutions have high [H⁺] and low [OH⁻], while basic solutions have the opposite.
Acidic Solutions (pH < 7)
Have higher [H⁺] than [OH⁻]. Examples include vinegar (pH ~2.5), lemon juice (pH ~2), and stomach acid (pH ~1.5-3.5). Strong acids like HCl completely dissociate in water.
Basic Solutions (pH > 7)
Have higher [OH⁻] than [H⁺]. Examples include baking soda solution (pH ~8.5), soap (pH ~9-10), and household ammonia (pH ~11-12). Strong bases like NaOH fully dissociate.
Acidic Substances
• Battery acid: pH 0
• Stomach acid: pH 1.5-3.5
• Lemon juice: pH 2.0
• Vinegar: pH 2.5
• Orange juice: pH 3.5
• Coffee: pH 5.0
• Milk: pH 6.5
Basic Substances
• Pure water: pH 7.0
• Blood: pH 7.35-7.45
• Sea water: pH 8.0
• Baking soda: pH 8.5
• Soap: pH 9-10
• Ammonia: pH 11-12
• Bleach: pH 12.5
pH measurement is crucial in many fields. In medicine, blood pH must be maintained between 7.35-7.45; deviations can indicate serious conditions like acidosis or alkalosis. In agriculture, soil pH affects nutrient availability for plants. In swimming pools, pH is maintained between 7.2-7.8 for comfort and chlorine effectiveness. Industrial processes, food production, and water treatment all rely on precise pH control for safety and quality.
Environmental scientists monitor the pH of water bodies to detect pollution and assess ecosystem health. Acid rain, with pH values as low as 4.2, can damage ecosystems and corrode buildings. Understanding and controlling pH is fundamental to chemistry, biology, environmental science, and many industrial applications.