Glossary
pH
pH is a measure of how acidic a solution is — how many hydrogen ions it holds — on a scale that usually runs from 0 to 14, with 7 neutral.
pH measures the concentration of hydrogen ions in a solution. It is a negative logarithm, so the number falls as the concentration rises: pH 7 holds a certain amount, and every step down the scale holds ten times more. That is why the difference between pH 5 and pH 3 is a hundredfold, not a factor of two, and why a small pH change in the blood is a serious event. The scale runs 0 to 14 in ordinary water-based chemistry. Seven is neutral — pure water at body temperature; below 7 is acidic, and above is basic. The usual figures are worth holding: stomach acid around pH 2, urine around 6, blood held between 7.35 and 7.45, seawater about 8. A buffer is what keeps the blood in that narrow band, which is the reason the number can be tested as a range and not a single value.
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Source: NIST periodic table and reference data for the atomic masses and half-life figures, NIH PubChem for the substance data, and NIST Special Publication 811 for the unit conventions — the same sources the chemistry guide cites. All three are US federal government material, not subject to copyright in the United States. Every definition, example and calculation above was written for this course; no dictionary or textbook text is reproduced. (NIST and PubChem accessed 2026-10-26; definitions written 2026-10-27)