Échelle de pH

Chemist Soren Sorenson developed the pH scale in 1909 to provide a more convenient way to quickly compare the acidity of different solutions. Since hydronium ion concentrations can range from 1 M to 1 × 10−14 M, expressing these values directly is impractical. The pH scale uses logarithms to compress this enormous range into a simple 0-14 scale.

Because pH is expressed on a logarithmic scale, a single-unit change corresponds to a 10-fold increase or decrease in hydronium ion concentration. For example, a solution with pH 3 contains ten times more hydronium ions than a solution with pH 4, and one hundred times more than a solution with pH 5. This logarithmic relationship makes pH a convenient way to express vast concentration differences.

The pH and pOH of any aqueous solution always sum to 14 at 25°C because the ion product of water (Kw) equals 1 × 10−14. Since pH is the negative logarithm of hydronium ion concentration and pOH is the negative logarithm of hydroxide ion concentration, this relationship allows you to calculate one value when the other is known. For instance, a solution with pOH of 10 has a pH of 4.

An acidic solution has a higher concentration of hydronium ions than hydroxide ions and a pH less than 7. A basic solution has a lower concentration of hydronium ions than hydroxide ions and a pH greater than 7. A neutral solution has equal concentrations of hydronium and hydroxide ions with a pH of 7. These distinctions apply at 25°C, where Kw equals 1 × 10−14.

The pH range of 0 to 14 is determined by the ion product of water (Kw), which equals 1 × 10−14 at 25°C. Taking the negative logarithm of this value yields 14, establishing the upper limit. The relationship pH + pOH = 14 ensures that aqueous solutions at standard temperature fall within this range, with pH values below 7 indicating acidity and above 7 indicating basicity.

Temperature changes alter the ion product of water (Kw), which shifts the pH values that define acidic, neutral, and basic solutions. At 80°C, pure water has a hydronium ion concentration of 4.9 × 10−7 M, giving pH and pOH values of 6.31, making neutral solutions have pH = 6.31 rather than 7. This distinction is important when studying processes like enzyme reactions in warm-blooded organisms at 36-40°C.

pH is calculated as the negative logarithm of the hydronium ion concentration: pH = −log[H3O+]. For example, a solution with 1 × 10−5 M hydronium concentration has a pH of 5. This formula can be rearranged to find hydronium concentration from pH, allowing chemists to understand weak acid solutions and dissociation constant calculations when analyzing solution properties.