pH Scale Reference

pH	Example
0	Battery acid
2	Lemon juice
4	Coffee
7	Pure water (neutral)
8	Baking soda
11	Ammonia
14	Lye (NaOH)

The logarithmic scale explained

pH is defined as −log₁₀[H⁺], where [H⁺] is the hydrogen ion concentration in moles per liter. Each unit change on the pH scale represents a 10-fold change in acidity:

pH 3 is 10× more acidic than pH 4
pH 3 is 100× more acidic than pH 5
pH 3 is 10,000× more acidic than pH 7 (neutral water)

This means small pH differences in, for example, ocean chemistry (the ocean has acidified from pH 8.2 to 8.1 since pre-industrial times) represent a 26% increase in ocean acidity - a significant change despite the small numerical difference.

pH in chemistry and biology

pH profoundly affects chemical reaction rates, enzyme activity, and cell function:

Human blood: maintained at pH 7.35–7.45. A drop below 7.35 (acidosis) or rise above 7.45 (alkalosis) is a medical emergency affecting enzyme and protein function.
Stomach acid: pH 1.5–3.5, necessary for pepsin activity and killing pathogens.
Soil pH: most crops grow best at pH 6–7. Acidic soil (below 5.5) limits phosphorus availability; alkaline soil (above 7.5) limits iron and manganese uptake.
Swimming pools: maintained at pH 7.2–7.8 for effective chlorine disinfection and swimmer comfort.

Buffers and pH stability

A buffer is a solution that resists pH changes when acid or base is added. Blood is buffered by the carbonate/bicarbonate system: CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻. The kidneys and lungs regulate this equilibrium to maintain blood pH within its narrow range. In the lab, common buffers include phosphate buffer saline (PBS) at pH 7.4 and acetate buffer at pH 4–5.